Chemistry Course Book FALL 2017 - Fairview High School

Chemistry Course Book

FALL 2017

Units 1-6

1) Atomic Theory

2) Nomenclature/ Formula Writing

3) Chemical Reactions

4) Measurement/Math Review

5) Mole Concept

6) Solutions

Written, edited and compiled by

Brian Cox

Dan Albritton

With special thanks to Daniel Knowles

Semester 1 Table of Contents

Unit # Topic Pages

Unit 1 Study Guide 1

1 Notes 2-9

1 Assignment List 10

1 Problem Sets 11-34


2 Notes 36-46




3 Notes 61-68



Unit 4/5 Study Guide 79

4 Notes 80-82

5 Notes 83-89

4/5 Assignment List 90

4 Problem Sets 91- 96

5 Problem Sets 97 - 106

4/5 Unit 4/5 Review

Problem Set

107-110


6 Notes 112-116



Semester Exam Topics 126

Semester Exam

Required Review

127

Semester Exam

Extra Credit

128

Chemistry Success Guide

Prerequisites:

• Appropriate Math Preparation: B or Higher Average in Algebra I OR C or higher average in Algebra II

• Strong Work Ethic/ willingness to persevere when challenged

• Positive attitude/Active approach to learning

General Advice

• Pre-read (use assignment lists to preview new topics before class) and/or Preview Power points online

• Actively engage during class: active listening, appropriate note-taking and asking questions during class; make

best effort on problems during individual work time; actively discuss material during group time.

• Post-read/ notes follow-up – review material to follow up / rework HW problems

• Keep up with daily assignments/ ask questions / clear up question as they arise – don’t allow small issues to

become a major problem

Important Organizational Remainders

• Remember Homework packet is due at time of the unit test, clearly labeled and in order.

• Lab work is tested approximately each quarter; lab reports are due at the time of the unit test and may be used

during the lab test; regular unit tests are closed note. All unit tests are closed friend (individual work only).

Test Preparation Tips

• ≈ 3 days before test – review Unit Study Guide/Test Topics checklist to identify any topics you don’t understand;

• Review Text, notes, homework - ask a friend or instructor for help as needed

• Practice test strategy - 2 days before test- work unit review as a practice test (try to work problems under

testing conditions – use only materials that will be allowed during the test; work in a quiet environment

with a time limit.). It is critical to practice working problems!

• If necessary, write a practice test for yourself using the study test/topics list to guide selections of homework

problems. Put yourself in the mind of the instructor – what problems are likely to be represented on the test?

Work your practice test under actual testing conditions.

Testing Taking Tips

- Look over the test and decide how much time to allot to each section (multiple choice/free response)

- Multiple choice – eliminate as many choices as possible in order to be able to focus on fewer items.

- Time Management – General advice – do problems that are worth the most points first – generally free response.

- Collect easy points first; skip problems that are taking too long and come back to them later

Unit 1 ~ Atomic Theory ~ Test Topics

Scientific Method: Understand that the scientific method is a systematic approach to problem solving that utilizes

experimentation to gather information and test ideas. This method is not restricted to trained professional scientists. Most

people make use of some variation of the scientific method as part of their problem solving skills.

Chemical and Physical Changes: Understand the difference between the two and be able to classify a change as

physical or chemical. Key idea to remember: chemical changes alter the physical and chemical properties of a substance.

Elements, Compounds, and Mixtures

• Know the definitions of each: elements are pure substances composed of only one type of atom; compounds are

composed of two or more different elements chemically bonded together; mixtures contain different types of

atoms or compounds physically mixed together but NOT bonded chemically.

• Be able to classify substances as an element, compound, or mixture and draw/interpret atomic level picture of

each.

Phases of Matter: Know the fundamental characteristics of solids, liquids, and gases through an analysis of a heating or

cooling curve for water.

Dalton’s Atomic Theory

• Be able to use key ideas of atomic theory to explain how atomic theory accounts for the Law of Definite

Proportions (Constant Composition) and Law of Conservation of Mass.

• Know the two concepts from Dalton’s theory which are no longer believed to be true – 1) atoms are smallest

particle of matter (modern: atoms made up of protons, neutrons and electrons and 2) atoms of same element are

absolutely identical (masses can vary – concept of isotopes)

Structure of the Atom

• Thomson experiment – be able to describe the experimental design, results, and conclusions and apply

concepts to interpreting similar experiments.

• Rutherford experiment – Know this experiment well. Be able to describe the experimental design, results and

conclusions. Understand how Rutherford’s experiments disproved Thomson’s plum pudding model of the atom.

• Protons, Neutrons and Electrons

→ For each particle know: mass, charge, location in the atom (i.e, inside or outside nucleus)

→ Determine the number of protons, neutrons and electrons from atomic notation:

Example: 27Al +3

13

→ Know that the number of protons (atomic #) determines the chemical identity of an element.

• Isotopes – understand concept of atoms of the same element with same # of protons can have different numbers

of neutrons.

• Average Atomic Mass – Understand that the masses listed on the periodic table represent an average of the

masses of all of the isotopes of that element, taking into account the relative abundance or amount of each isotope.

• Ions: Understand ion = charged particle; + ions form when atoms lose electrons; - ions form when atoms gain

electrons.

Classifying Matter and the Changes It Undergoes

(Fill in using chemical and physical changes powerpoint on website)

Defn of Chemistry: “The study of ______________ and the ___________________ it undergoes”

Types of Changes Matter Can Undergo:

PHYSICAL CHANGE: a change in a substance that does _____ alter its fundamental properties (e.g. boiling

point, melting point, color, chemical reactivity) Examples:

- (e.g. boiling, freezing, melting)

- (e.g. salt or sugar dissolving in water)

- (e.g. glass shatters or a stick snaps into smaller pieces)

CHEMICAL CHANGE: a change in a substance that does alter its fundamental properties (b.p., m.p. color,

chemical reactivity)

Examples: 1)

2)

Note: _______ change and/or dramatic difference in appearance and/or very large energy changes usually

indicates chemical change

Importance of Chemical vs Physical Changes: Important in classifying types of matter e.g. compound vs. mixture.

Substances that contain more than one type of atom are classified as either a ______________________ (if the atoms can

only be separated by a chemical reaction) or a _____________________________ (if the atoms can only be separated by

a physical process.

Practice classifying chemical vs physical changes: HW 1-1

Classifying Matter:

➢ -substance composed of atoms of the _____________ type

➢ - a substance composed of two or more different types of elements ______________

____________ together

➢ - a substance composed of two or more different substances that can only be separated by

_______________ means

Atom:

Molecule:

Draw a picture showing the difference between atoms and molecules next to their names in the figure below:

Classifying Matter

ELEMENTS, COMPOUNDS AND MIXTURES

ELEMENT – substance composed of ______ type of atom

➢ ______________ be broken down into a simpler substance by either

__________________ or ___________ means

Examples:

Elements can be either monoatomic or diatomic:

7 diatomic elements: H2, N2,O2, F2, Br2, Cl2, I2

Atomic level representations of: (Use to represent atom)

Monoatomic Element Diatomic Element

COMPOUND – composed of ____ or more __________________ types of elements ________________________

_____________ together.

COMPOUNDS CAN ONLY BE BROKEN DOWN INTO THE ELEMENTS THAT MAKE UP THE SUBSTANCE BY

A ____________________ _________________________ .

Examples: 1) 2) 3)

Atomic Level picture of water:

MIXTURE- composed of ______ or more substances which can only be separated by a

_____________________________ _________________.

Most of the substances we see around us are mixtures.

Examples: 1) 2) 3)

Atomic Level Representation of different types of mixtures

Mixture of Elements:

Mixture of Elements and Compounds:

Mixture of Compounds:

ATOMIC THEORY

Theories are detailed explanations for observable events that are supported by large body of experimental evidence.

Atomic theory is one of the most successful scientific theories of all time based on its ability to explain observations and

predict the results of future experiments.

Key Ideas of Atomic Theory

1) All matter is composed of tiny particles called ________. Atoms are the smallest unit

of chemical combination.

(Later: atoms are made of ____________, ____________, and ______________ )

2) Atoms of the same _____________ are identical. Atoms of any one element are

different from those of any other element.

(Later: _____________– same chemical properties but different mass)

3) In chemical reactions, atoms are ________________ but the ________

____________and ________ of _________ present _____ ______ ____________.

Supporting evidence: Law of ____________________ of ___________

4) A ________________ is formed when atoms of two or more elements ____________

combined. The compound has the same _______ and ratio of ________.

Supporting evidence: Law of ______________ Proportions (_______________

Composition)

Experimental Evidence for Atomic Theory

Dalton’s atomic theory explains two fundamental laws of chemistry:

Law of Conservation of Mass – in a chemical reaction ________ ___ ____________________. In other

words, the sum of the masses of the reactants (_________ side of equation) equals the sum of the masses of the

products (_____________ side of the equation.)

Example: Mg + Br2 → MgBr2

24.3 g 159.8 g 184.1 g

Explanation using atomic theory:

1) Matter is composed of tiny particles called

2) In a chemical reaction atoms are _______________; the total # and type of atoms is constant.

Symbolic representation of concepts:

Use to represent Mg atoms and to represent Br atoms

Mg + Br2 → MgBr2

Notes on Coefficients and Subscripts

• Subscript: # below line; indicates # of atoms bonded together Br2 subscript

• Coefficient: # in front of formula; indicates number of units (atoms or molecules) present

Coefficient 2 Na

Examples: Br = 2 Br = 3 Br =

Br2 = 2 Br2 = 3 Br2 =

Law of Definite Proportions (or Constant Composition) – for a given compound the elements that make up the

compound and the ratio of their masses are always the same.

Example: Carbon dioxide is always composed of the same two elements, carbon and oxygen. Regardless of the total

mass of CO2, the ratio of the mass of oxygen ÷ mass of carbon is always 2.67 : 1

Explanation using atomic theory:

1) Matter is composed of _______.

2) 3) 2) In a compound the ________ of atoms and ______ of atoms is constant.

Structure of the Atom In the space below, sketch the Rutherford model of the atom. Label the nucleus and electron cloud. Be sure to make the

nucleus the appropriate relative size to the total volume of the atom. Fill in the missing information in the chart below:

Rutherford Model picture:

Atoms Are made of

Protons Neutrons Electrons

Location: Location: Location:

Charge: Charge: Charge:

Mass: 1 amu Mass: Mass:

Amu = Atomic Mass Unit = mass of 1 proton ; 1 amu = 1.66 x 10-24 g

If all atoms are made up of the same 3 particles, what makes one atom different from another?

Atomic Mass = +

Atomic # = # of ; determines chemical identity of atom

# of neutrons = atomic mass (p + n) atomic # (p)

Ion = _____________ atom

Ions form when atoms gain or lose _________________

+ IONS form when e- are _________ IONS form when e- are ________

NEUTRAL ATOM : # PROTONS # ELECTRONS + ION : # PROTONS # ELECTRONS - ION: # ELECTRONS # PROTONS

Example: 25 Mg +2

12

Protons: Atomic # (Lower left) =

Neutrons: Mass (top left) – atomic # =

Electrons: # protons (bottom left) – charge (top right) =

A second notation system is sued for neutral atoms (#protons = # of electrons):

He-4 is the same as 4 He The atomic # of 2 is unique to He and doesn’t need to be written.

2

Isotope- atoms with the same atomic ____________, but a different atomic __________. Isotopes have the

same # of ____________ but a different # of __________________.

charge

If no charge number is written charge is

assumed to = 0

SUMMARY: IONS vs ISOTOPES:

Isotopes: Chemically same, different mass: Same # of _____________, different # of ____________

Ions: Charged atoms: Same # of _________________, different # of ________________.

ISOTOPES IN NATURE

• A sample of the atoms of a naturally occurring element will contain a _________________ of all the different

___________ of that element.

• Example: Chlorine consists of two isotopes Cl-35 and Cl-37.

Average Atomic Mass of an Element on the Periodic Table

The atomic mass given on the periodic table represents the ___________ __________of all of the

____________of the element, taking into account the _________ ___________ of each isotope.

Note: Relative abundance = % of naturally occurring atoms of that isotope

Course Syllabus: Readings and Homework Problems

Note: Page # Refer to Zumdahl text; Problem Sets are in workbook (WB); Unit Reviews are extremely

important and will be worth a significant number of points.

HWK

#

Pages to Read Problems

1-1 Section 3.2 (p. 56-60) See below

1-2 Sec. 3.3,3.4, 3.5 (p. 60-65) Elements, Compounds and Mixtures Problem

Set

1-3 Section 3.1 (p.55) Atomic Theory description of phases of matter

1-4 Sec. 4.1,4.2,4.3 (p.73 -79) Atomic Theory Problem Set

1-5 Chemistry in Focus (p.83) Thomson Problem Set

1-6 Section 4.5 (p. 81 – 84) Rutherford Problem Set

1-7 Sections 4.6, 4.7 (pp. 84-88),

4.10 (p.96 – 100)

Structure of the Atom

1-8 Unit 1 Review Problem Set

Homework 1 - 1 (From Zumdahl p. 67 # 11, 12, 18 abcdejk)

11) Elemental mercury is a shiny, silver-colored, dense liquid that flows easily. Are these characteristics of mercury

physical or chemical properties?

12) If liquid elemental mercury is heated in oxygen, the volume of the shiny liquid decreases and a reddish orange solid

forms in its place. Do these characteristics represent a chemical change or a physical change?

Classify each of the events described below as chemical change (C) or physical change (P): 18a) A fireplace poker glows red when you heat it in the fire.

18b) A marshmallow gets black when toasted too long in a campfire.

18c) Hydrogen peroxide dental strips will make your teeth whiter.

18d) If you wash your jeans with chlorine bleach, they will fade.

18e) If you spill some nail polish remover on your skin it will evaporate quickly.

18j) A log of wood is chopped up with an axe into smaller pieces of wood.

18k) A log of wood is burned in the fire place.

Homework 1-2

Part 1 : Iron- sulfur Experiments

Initial observations: Iron filings:

Sulfur:

Experiment #1: Stir iron filings and sulfur together

Observations:

Is this an element, compound or mixture? Support your answer with two lines of experimental evidence.

Using to represent iron atoms and to represent sulfur atoms draw an atomic level picture of this experiment.

Experiment #2: Heat iron filings and sulfur together strongly in a flame; cool; examine mixture; check to see if iron can

be removed with magnet.

Observations:

Is this an element, compound or mixture? Support your answer with two lines of experimental evidence.

Using to represent iron atoms and to represent sulfur atoms draw an atomic level picture of this experiment.

Part 2: Interpret the following experiments

1) Filtration is a technique for separating a solid from a liquid. Filter paper allows water and molecules dissolved in

water to pass through while trapping solid particles. An orange colored liquid is separated into an orange solid and

a clear colorless liquid by filtration. Was the orange colored liquid an element, compound or mixture? Explain.

2) A clear, colorless liquid is heated on a hot plate. After 15 minutes the liquid has evaporated and a white solid

remains. Was the liquid an element, mixture or a compound?

3) Electrolysis of water

Initial observations: Water is clear and colorless; phenolthaphlein acid-base indicator is initially clear and

colorless. Salt is dissolved in water to increase electrical conductivity.

Final observations:

3A) Is electrolysis a chemical process or a physical process? Cite experimental observations to support your conclusion.

3B) Is water an element, compound or a mixture? How do you know?

3C) The volumes of the gases produced is in a ratio of 2:1. Which gas is in which tube? Label each gas in the diagram

above. Experimental Evidence which supports your conclusion as to which gas is present in which tube:

3D) Propose an experiment to test your hypothesis as to the identity of the gases. (Hint: Burning splint)

Part 3:

1) Classify the following as either an element (E), compound (C), or a mixture (M)

A) Air

B) Cu

C) Water

D) Salt Water

Part 4: How could you separate the following mixtures into its components?

A) Solid benzoic acid (soluble in hot water) from solid charcoal.

B) Two liquids, one with a boiling point of 100 oC and one with a boiling point of 78 oC.

Homework 1-3

Atomic Theory Description of Solids, Liquids, and Gases

CU Phet Simulation States of Matter is a useful reference for this assignment. phet.colorado.edu

Part 1: Describing the macroscopic (directly observable properties) of solids, liquids and gases

For each observation, indicate whether it describes solid (S), liquid (L) or gas(G). Many observations apply to more than

one phase

Examples:

1) Fixed or definite shape : S 2) No definite shape; takes shape of container : L, G

Observations

3) Definite volume 4) Expands to fill entire container

5) Easily compressible 6) Very difficult to compress

7) Phase you cannot move your hand through

8) Phase you can move your hand through but with noticeable resistance

9) Phase you can move your hand through very easily

Consider the 3 phases of water: solid (ice), liquid, gas (steam)

10) Which phase is coldest? Which is hottest?

Part 2: Atomic Level Descriptions – After watching the Phet Simulations on States of Matter complete the following:

Using draw atomic level representations of solids, liquid and gas

Hint: Watch spacing between particles and level of organization

Solid Liquid Gas

Questions: For each statement, indicate whether it best describes solid (s), liquid (l) or gas(g). A statement may apply to

more than one phase.

1) Very little empty space between atoms.

2) This phase is mostly empty space; very few atoms are present in the picture.

3) This phase has very strong attractive forces between atoms. Atoms cannot change their positions relative to other

atoms.

4) This phase has attractive forces that are strong enough to hold the group of atoms together, but allows for atoms to

slide past each other.

5) The attractive forces in this phase are extremely weak.

6) The atoms in this phase of matter are highly organized. There are often definite repeating patterns visible.

7) The atoms show completely chaotic and random motion. No patterns other than straight motion are visible.

8) The atoms in this phase show “pockets” of organization; In other words, some atoms seem detached or in random

motion, but little clusters of atoms form and then break apart.

9) In this phase of matter the atoms have extremely high kinetic energy; the atoms are in constant motion moving at

very high speeds.

10) In this phase atoms have very little kinetic energy (energy of motion). The atoms vibrate in place but do not

move past each other.

11) In this phase atoms have limited motion; the atoms are able to slide past each other.

Part 3: How does Atomic Theory explain the following observations?

1) Explain why gases can be easily compressed, but it is very difficult to compress a solid.

2) Why is it extremely difficult (not to mention usually painful) to put your hand through a solid object, somewhat

easy to move your hand through a liquid such as water (swimming!) and extremely easy to move your hand through

a gas (such as air).

Phase Transitions

1) In each of the 3 empty boxes above (Box A,B and C), indicate the appropriate phase of matter gas (G), Liquid

(L) or Solid (S)

2) Which 2 phases are in equilibrium (both phases present simultaneously) between points 2 and 3?

The temperature at which the transition between these two phases occurs is called the __________________

point.

3) Which 2 phases are in equilibrium (both phases present simultaneously) between points 4 and 5?

The temperature at which the transition between these two phases occurs is called the ___________________

point.

4) Compare the shape of the graph between points 4 and 5 vs. between points 5 + 6. Why are the shapes

different?

5) What changes in molecular behavior would be observed between points 4 and 5 vs between points 5 + 6?

6) Why is the heat energy required to transition from 2 to 3 lower than the heat energy required to transition

between points 4 to 5?

1

2

3

4

5

A

B

C

6

Homework 1-4 Dalton’s Atomic Theory

1) For each diagram below which is consistent with Dalton’s atomic theory write Correct. For each incorrect

picture identify which conceptual error is present:

I) atoms appear in the products that were not present in the reactant

II) atoms disappear from the product side that were originally present in the reactants

III) Reactant atoms are transformed into atoms of a different element

1A) H + H H H

1B) H + H He

1C) H H + O O H O H

1D) + O

2) Given the following: O2 + C → CO2

16.0 g ? g 22.0 g

A) How much C must have reacted? ____ g

B) Which Law (Conservation of Mass or Constant Composition) supports your answer?

C) How does Atomic Theory explain the above result?

All matter is composed of tiny particles called __________.

In a chemical reaction atoms the _______________ and type of atoms present on the reactant side must match the

number and _________ of atoms present on the product side. In other words, in a chemical reaction atoms rearrange

(change bonding partners) but the number and type of atoms are conserved. Draw a diagram using to represent

carbon atoms and to represent O atoms.

1E)

3) Given the following: 6 Mg + 2 N2 → 2 Mg3N2

? g + 56 g 201.8 g

A) How much Mg must have reacted? _____ g

B) Which Law (Conservation of Mass or Constant Composition) supports your answer?

C) Using both words and a diagram explain how Atomic Theory explains the above results.

4) Water is a clear, colorless that liquid that contains the elements hydrogen and oxygen. Hydrogen peroxide is a clear,

colorless liquid that contains the elements hydrogen and oxygen.

A) What happens when manganese dioxide, MnO2, a black powder, is added to water?

B) What happens when manganese dioxide, MnO2, a black powder, is added to hydrogen peroxide?

C) Are water and hydrogen peroxide the same substance? Explain.

D) How can atomic theory explain the above results? Draw a picture of a water molecules and a hydrogen peroxide

molecule to support your answer. Use H = and O =

5) Copper and chlorine form two different compounds. One is a light blue solid, while the other is a green solid.

According to Atomic Theory, what is the atomic level difference between the two compounds?

E) When water and hydrogen peroxides are chemically broken down into their constituent

elements the following results are obtained:

Water H: 2 g O: 16 g Hydrogen peroxide = H: 2 g O: 32 g

Which Law (Conservation of Mass or Constant Composition) explains these data?

Discovery of the Electron ~ HW 1-5

1) By the end of the 19th century, what observations had been made concerning electricity and electrical charges?

➢

➢

➢

Background

Late 19th century: Most scientists accept idea that matter is composed of atoms.

Unanswered Questions:

• Since matter can be neutral or positively or negatively charged and matter is made of atoms, what is the

relationship between atoms and electrical charges? How do atoms change when matter changes from neutral to

electrically charged?

• Are atoms truly the smallest particle of matter? Could there be particles smaller than atoms that make up atoms?

2) Description of Experiments with electricity using the cathode ray or Crooke’s tube:

If a sealed glass tube with most of the air pumped out (partial vacuum) containing a cathode (negative terminal

electrode) at one end and an anode (positive terminal) at the other end are connected to a high voltage source, a

beam of light (called a cathode ray beam) will move in a straight line from the cathode to the anode. (Note: The

cathode ray beam is invisible to the naked eye, but produces a green light when it strikes a fluorescent

background material). When a magnet is held near the cathode ray beam it curves. Sketch the cathode ray beam

path in the tube and show how the path changes when a magnet is held nearby.

Cathode Anode Cathode Anode

Cathode ray diagram Cathode ray diagram deflected by a magnet

Thomson ~ First Experiment

J.J. Thomson a British scientist, set out to determine the nature of the cathode ray. In the 19th century, scientists

discovered that electricity and magnetism are related forces. Since the cathode ray beam was deflected by a

magnet, Thomson hypothesized that the cathode ray beam might consist of electrically charged particles.

Thomson added a set charged plates above and below the beam. Sketch the result of the experiment below:

3) Based on the results of this experiment it can be concluded that cathode ray beam is composed of ______

charged particles. Supporting experimental evidence:

+

Thomson ~ Second Experiment

Unanswered Questions:

• What are the negatively charged cathode ray particles made of?

• How could you design an experiment to identify a particle that is too small to see?

Thomson hypothesized that the cathode ray particle might be a negatively charged atom. To test this hypothesis, Thomson

decided to compare the mass of the cathode ray particle (CRP) to the mass of hydrogen, the lightest known atom. To

understand the concept of Thomson’s 2nd experiment we will investigate the paths of moving spheres of plastic and steel.

4) Sketch the path of the different spheres as travel past the magnet below:

Held constant in experiment: Strength of magnetic field and velocity of spheres moving past magnet.

5) Which sphere was not deflected at all? Why did it not deflect?

6) Why did the other spheres deflect?

7) What determines the degree of deflection? Which spheres experience the greatest degree of deflection?

8) The actual Thomson experiment results indicated that the mass of the cathode ray particle was approximately

1/2000 of the mass of the hydrogen (the lightest known atom). What conclusion can be drawn from this

observation?

9) Further experiments by Thomson demonstrated that cathode ray particles are present in all elements. What is the

cathode ray particle called today?

= plastic sphere 100 g

= steel sphere 100 g

= steel sphere 10 g

Experimental Extensions and applications of Thomson Experiment:

10) If a stream of positively charged particles such as hydrogen ions move between two electrically charged plates,

will the beam bend towards the + or – plate? Explain.

11) What subatomic particle was discovered by experiments similar to experiments described in the previous

question?

12) Make a drawing showing the deflection of these beams when passed between a pair of charged plates. (Assume:

1) Velocity of are particles are equal 2) Strength of electric field in charged plates are identical).

A) Na+1 ions with a mass of 23 and K+1 ions with a mass of 39.

B) S-2 ions with a mass of 34 and O-2 ions with a mass of 16.

C) Cr+2 ions with a mass of 52 and Cr+5 ions with a mass of 52.

D) Cl-1 with a mass of 37 and K+1 with a mass of 37.

Discovery of the Nucleus – Rutherford Scattering Experiment ~ Homework 1-6

Background: Plum Pudding Model of the Atom

After Thomson’s discovery of the existence of ______________________ in atoms, he proposed the first model for the

internal structure of the atom. Thomson reasoned that since atoms were overall electrically __________________ and

contained negatively charged electrons, the atom contains an equal amount of __________________ charge. He

envisioned the atom consisting of negatively charged electrons embedded in a diffuse (____________ ______ ) cloud of

positive charge, much like a favorite English dessert that contains plums stuck in pudding.

One of the central components of the scientific method

is that a hypothesis must tested experimentally.

From 1909-1911, Ernest Rutherford and colleagues

set out to test the Plum Pudding Model.

Rutherford’s Challenge: How do you design an experiment to “see” inside an atom when atoms themselves are so small

that they cannot be seen directly?

Rutherford’s experiment involved firing a tiny probe particle at the atom to observe how atoms would affect the motion of

the probe. From the presence or absence of scattering in the pathway taken by the probe as they passed through the atom,

Rutherford hoped to determine if the particles that made up the atom were in fact spread out as Thomson predicted.

Analogy to Rutherford expt.

Sketch view of pattern of

bullets passing through

backside of tent.

Sketch Plum

Pudding

Model in Box

Label Components of Experiment (write Bold-Faced words in appropriate boxes above).

1) Lead Box containing radioactive element which emits alpha particles. The lead box contained a single

tiny opening in the front which allowed alpha particles to escape in only one fixed direction.

2) Alpha particle source (Radioactive element such as polonium or radium)

3) Alpha Particles

4) Visualization Screen - Fluorescent Circular Screen that emits flash of light when struck by alpha

particle.

5) Light Flashes – produced by alpha particles striking screen.

6) Thin Gold Foil

Target Atoms (Sheet of Gold Foil)

1) Why did Rutherford use a very thin sheet of gold foil as target atoms instead of a thick gold brick?

Characteristics of Alpha Particles (Probe particle) :

Alpha particles are released from certain radioactive substances such as radium, uranium or polonium atoms.

2) The relative sizes of an alpha particle and a gold atom are shown below. How does the size of an alpha particle

compare to the size of a gold atom? (Does the probe particle need to be bigger or smaller than the atom it is attempting

to pass through?)

Alpha particle

3) Look at the experiments below measuring the deflection of different types of radiation by an electric field. The alpha

particle path curves downward toward the negative plate. What kind of charge +, neutral or – does an alpha particle

have? What kind of charge would repel an alpha particle?

4) Alpha particles are invisible to the human eye. From previous work studying alpha particles, Rutherford knew that

alpha particles traveled in straight lines and that when an alpha particle collided with a screen covered with a chemical

called zinc sulfide a flash of light was observed.

A)

B) Inferring the trajectory based on position of light flashes: Imagine that alpha particles are fired at a dense

concentration of positive charges arranged in different shapes. The alpha particles in the experiments below are being

fired from left to right at a mystery shaped object surrounded by a circular visualization screen. Match each pattern of

light flashes on the circular screen with the different shaped objects. Sketch each shape in the appropriate circle

and draw straight lines coming from the lead box to show how some particles missed the object and some bounced

off.

Shape choices: A) B) Rectangle

Alpha particle source Visualization screen emits

flash of light where alpha

particle hits

Lead Box

Lead Box

Testing the Plum Pudding Model. What does the Plum Model Predict should happen when alpha particles

are fired at an atom?

Alpha Particle vs. Electron

5) At the time of Rutherford’s experiments, the only experimentally verified fact about the internal structure of the atom

was that it contained negatively charged electrons. One important aspect of the experimental design was to predict what

would happen in an encounter between an alpha particle and an electron.

A) From physics, it is known that when two objects collide, the outcome of the collisions can be understood in terms of

momentum which is calculated by the product of mass x velocity.

For example, what do you think would happen when an SUV driving at 65 mph collides head on with a mosquito

flying at 1.25 mph? The mass of the SUV is approximately a billion times greater than the mass of the mosquito. (Note:

pictures below are not to scale; the mosquito is near actual size, the truck would be much, much larger)

Which would be the outcome of this collision? (Choose the best answer).

A) The truck flies backward 500 feet from the force of the collision with mosquito.

B) Both the truck and the mosquito are instantly stopped dead in their tracks.

C) The mosquito splats onto the windshield. The truck keeps moving forward with no noticeable difference in speed.

6) The mass of an alpha particle is approximately 7500 times the mass of an electron. If both the alpha particle and

the electron are traveling at similar high velocities, will the electron be able to stop the alpha particle? Explain.

7) The fact that the alpha particle is positive and the electron is negative means they will attract each other, but a

reasonable prediction is that while the electron might stick to the alpha particle if they had a direct collision, the alpha

particle should not change its trajectory as the result of interacting with an electron. (Recall that both particles would be

moving very fast with little time near each other when passing by to attract each other.)

Also recall that in the Thomson model, the positive charge was described as diffuse or spreadout, so there was no

concentrated mass of positive charge that would be necessary in order to deflect the path of a positive alpha particle.

Analogy to alpha particle and concept of diffuse + charge: Why was the ship the Titanic able to pass through a fog

cloud made of water vapor but not through an iceberg?

8) Expected Results- In the space below draw the pattern (using arrowed lines) that would be predicted for the path of

alpha particles if the plum pudding model were correct:

9) Actual Results: How can you explain these actually observed results?

A) The vast majority of alpha particles passed straight through the gold foil with no significant deflection.

Explanation: Most of the atom must consist of

B) A very small, but statistically significant fraction of alpha particle (about 1 in 8000) were deflected off at

various angles. About 1 out of 100,000 alpha particles came flying straight back at the source.

In order to deflect alpha particles with high momentum (relatively heavy mass traveling at a high speed), the

alpha particles must have encountered a (heavy mass) of particles concentrated within the atom.

i) What kind of charge does an alpha particle have?

ii) What kind of charge must the concentrated mass of particles inside the atom have to deflect an alpha

particle?

10) Rutherford proposed that all atoms have a nucleus. A diagram of his interpretation of the experimental

results is shown below.

How does Rutherford explain the fact that a few of the alpha particles were deflected, but most were not

deflected?

11) The Rutherford model of the atom states that the atom contains a tiny, dense positively charged nucleus

surrounded by an electron cloud.

A) The model states that nucleus of the atom is positively charged. What is the experimental evidence

for this statement?

B) The nucleus of a hydrogen atom occupies approximately 0.000 000 000 0004 % of the total space of

the atom, while the “electron cloud” or empty space 99.999 999 999 999 6% of the total volume of the

atom. What is the experimental evidence that supports this statement?

10) Why was Rutherford so shocked to discover that all of the positive charge in the atom was concentrated in one

tiny spot, instead of having + and – charges all intermixed as Thomson’s model predicted? (Hint: What do you know

about the interactions between different charges?)

Summary of Experiment:

Conclusion #1: Most of the atom is

_____________ __________ .

Supporting Evidence: Only a small % of

alpha particles were deflected.

Conclusion #2: The contains a tiny

concentration of _____________ charge.

The deflection patterns of the alpha

particles were consistent with the

repulsion of ________ charges. The large

deflections could only be caused by a

concentration of ___________ charge.

The Structure of the Atom ~ Homework 1-7

1) Draw diagram of an atom that contains 2 protons, 1 neutrons and 2 electrons. Using the symbols place

each particle in its appropriate location within the atom.

2) For each atom give the overall charge and overall atomic mass.

# protons # neutrons # electrons Overall charge Atomic Mass (amu)

5 6 5

5 7 5

5 6 8

5 5 3

3) The chemical identity of an atom is determined by the number of ______________________.

What is the atomic #?

H: He Li Li Li Li+1

Protons: 1 2 3 3 3 3

Neutrons: 0 2 3 4 5 5

Electrons: 1 2 3 3 3 2

Atomic # = ; determines the chemical identity of the atom

4) An ion is an atom with a + or – electrical charge. How is an ion formed?

a) by either adding or subtracting protons from the atom

b) by either adding or subtracting neutrons from the atom

c) by either adding or subtracting electrons from the atom

d) all of the above

e) two of the above

nucleus

5) Given the following neutral atom:

Diagram how many would be present in each of the following:

A) +1 ion

B) -1 ion

C) +2 ion

D) -3 ion

6) Determine the # of protons, neutrons and electrons

11 B 27 Al +3 14 N

-3

5 13 7

Protons:

Neutrons:

Electrons:

7) An atom has an atomic number of 9 and an atomic mass of 19.

A) What is the chemical identity of the atom?

B) How protons and neutrons are present?

8) Write the notation for each carbon isotope in the vertical format and then determine # of protons, neutrons

and electrons. Assume all atoms are neutral (# protons = # of electrons).

C-12 → 12 C C-13: C C-14: C

6

Protons: 6

Neutrons: 6

Electrons: 6

9) Ion vs. Isotope

A) For each of the atoms below indicate the number of protons, neutrons and electrons.

i) 16 O -2 ii) 40 Ca +2 iii) 17 O

iv) 41 Ca

8 20 8 20

Protons:

Neutrons:

Electrons:

B) Circle any ions represented in # i – iv.

C) Which of the above represent isotopes of the same element?

10) Given: 2 different atoms X and Y.

Atom X has an atomic number of 9, an atomic mass of 19 and contains 10 electrons.

Atom Y has an atomic number of 10, an atomic mass of 19 and contains 10 electrons.

A) Is X an ion? Explain.

B) Is Y an ion? Explain.

C) Are X and Y isotopes of the same element? Explain.

11) Atomic Mass on the Periodic Table

• The mass of boron on the periodic table is listed as 10.8 amu, although there are no atoms of boron that

actually weigh 10.8 amu. Where does the number 10.8 come from?

• Hint #1: Boron has two isotopes, B-10 and B-11. How would you calculate the average of 10 and 11?

• Hint #2: How would you calculate the average of 10,10,11,11,11,11,11,11,11,11?

12) Lithium has 2 common isotopes, Li-6 and Li -7. The weighted average mass of lithium is given on the periodic table

as 6.94 amu. Which statement is most likely to be correct? Explain

A) There is exactly the same amount of Li-6 and Li-7 is nature (e.g. 50% for each).

B) There is slightly more Li -6 than Li-7.

C) There is slightly more Li-7 than Li-6.

D) There a large % of Li-6 in nature, but only a small % of Li-7.

E) There a large % of Li-7 in nature, but only a small % of Li-6

13) The average mass of a carbon atom is 12.011. Assuming you could pick up one carbon atom, what is the change that

you would randomly get one with a mass of 12.011? Explain.

a) 0% b) 0.011% c) about 12% d) 12.011% e) greater than 50% f) none of the above

HW 1-8 Unit 1 – Scientific Method and Atomic Theory Unit Review

Chemical/Physical Changes- classify each of the following changes as chemical or physical.

1) Meat blackens when cooked to long on a barbeque.

2) Water boils when heated strongly.

3) Salt dissolves in water.

Elements/ Compounds/ Mixtures (From Zumdahl Textbook p. 68, #25,26,38,39)

4) Magnesium filings are silver colored and sulfur is a yellow powder. Suppose a teaspoon of magnesium filings and

teaspoon of powdered sulfur are stirred. Would this constitute an element, compound or a mixture?

5) Suppose the magnesium filings and sulfur in the previous question are heated strongly forming a white crystalline

solid. Would this constitute an element, compound or a mixture?

6) Pure substance, X is melted, and the liquid place an electrolysis apparatus. When electric current is passed

through the liquid, a brown solid forms in one chamber and a white solid forms in the other chamber. Is substance

X an element, compound or a mixture?

7) If a piece white blackboard chalk is heated strongly in a flame, the mass of the chalk will decrease, and eventually

the chalk will crumble into a fine white dust. Does this change suggest that the chalk is composed of an element

or a compound?

Solids/Liquids/Gases (#8-11 from Zumdahl Textbook p. 67, #5,8,10 ; p. 104, #11)

8) Compare and contrast the ease with which molecules are able to move relative to each other in the three states of

matter.

9) How is the rigidity of a sample of matter affected by the strength of the forces among the particles in the sample?

10) In a sample of gaseous substance, more than 99% of the overall volume of the sample is empty space. How is this

fact reflected in the properties of gaseous substance, compared with the properties of a liquid or solid substance?

11) Sketch an atomic level diagram of a solid, liquid and gas phase of a substance.

Dalton’s Atomic Theory (Zumdahl Textbook 106, #16)

12) Correct each of the following misstatements from Dalton’s Atomic Theory.

a) Elements are made of tiny particles called molecules.

b) The atoms of a given element are exactly the same as the atoms of a different element.

c) A given compound may vary in the relative number and types of atoms depending on the source of the

compound.

d) A chemical reaction may involve the gain or loss of atoms as it takes place.

13) When electric current is passed through a water, hydrogen and oxygen gas is are produced. Is this a chemical or

a physical change? Why is the hydrogen gas volume twice that of the oxygen gas volume?

14) Hydrogen and oxygen can combine in two different ways to form H2O and H2O2. Based on our in class

demonstration (HW 1-4, #4) are these two compounds the same or different? Explain.

15) When phosphorous is heated in the presence of high concentrations of chlorine gas, two different products are

obtained. One compound is a clear, colorless fuming liquid while the other is a pale yellow-white solid.

According to Dalton’s Atomic Theory and the Law of Constant Composition, what is the difference

between the two phosphorous-chlorine compounds?

Conservation of Mass

16A) If 56 g of Fe reacts with 32 g of S how many grams of FeS will be formed?

16B) How would Atomic Theory support your answer? Draw an atomic level picture using

to represent Fe atoms and to represent S atoms.

17)A stick of wood weighs 10.0 g. The wood is burned in air. After burning, the blackened wood weighs 6.2 g. How

would you explain this result? Does this result violate the Law of Conservation of Mass? Explain.

18) A piece of steel wool (essentially iron) weighs 1.12 g. The steel wool is burned in air. The blacked metal weighs

1.60 g. How would you explain this result? Does this result violate the Law of Conservation of Mass? Draw an

atomic level picture to support your answer.

Structure of the Atom/ Thomson/ Rutherford Experiments (Zumdahl Textbook p.106. 22)

19) Indicate whether each of the following statements true or false. If false, correct the statement so that becomes true.

A) Rutherford’s bombardment experiments with metal foil suggested that the alpha particles were being deflected by

coming near a large, negative charged nucleus.

B) The proton and electron have similar masses but opposite charges.

C) Most atoms also contain neutrons, which have the same mass as protons but have no charge.

D) Protons and electrons are found inside the nucleus and neutrons travel outside of the nucleus.

20) In 4 separate experiments, a beam of protons (H+ ions), a beam of neutrons, a beam of electrons and a beam of

positrons are passed between two electrically charged plates. A positron is the anti-particle of an electron,

meaning that it has the exact same mass, but opposite electrical charge. The results of the experiment are

summarized in a single diagram below. Which letter path is produced by each particle? How do you know?

21) Rutherford Experiment

A) Since alpha particles are invisible to the naked eye, how did Rutherford track the trajectory of the alpha particles in the

gold foil experiment?

B) What result did Rutherford expect if the Thomson plum pudding model was correct?

C) If the black dots in the center represents the nucleus of the atom, which of the following models is most consistent with

the results of Rutherford’s scattering experiment? What experimental evidence supports your answer?

D) What is the charge on the nucleus? What experimental evidence supports your answer?

+ plate

- plate

A

B

C D

E) In addition to bombarding gold foil with alpha particles, Rutherford also repeated his experiments with many other

metal foils. How would the results of an experiment in which aluminum foil was bombarded with alpha particles be both

alike and slightly different than the results of the gold foil experiment? (Hint: How is the structure of a gold atom very

similar to the structure of an aluminum atom? How are the atoms different, i.e. what makes one atom gold and another

atom aluminum?)

22) Given the following ions:

70 +3 16 -2

Ga O

31 8

e-: e-:

n: n:

p: p:

a) Indicate the # of protons, neutrons and electrons present for each.

b) Did each atom gain or lose electrons to form its ion? How many electrons were gained or lost?

Ga+3: O-2:

23) Which of the following represent different isotopes of the same elements? Circle any atoms that are ions.

A) H-3 and H-1 B) H and H+1 C) H-3 and C-14 D) 35 Cl -1 37 Cl -1

17 17

Note in 23B assume H and H+1 both have a mass of 1 amu.

24) There are two known isotopes of Cl, Cl-35 and Cl-37. The mass on the periodic table for Cl is given as 35.5.

A) What % of Cl atoms in nature have a mass of 35.5?

B) What is the average of 35 and 37?

C) What is the average of 35, 35, 35, 37?

D) Where does the number 35.5 on the periodic table come from?

E) Challenge question: Chlorine gas as a chemical formula of Cl2. If the mass of Cl2 molecules are measured in

nature it is found that some Cl2 molecules weigh 70 amu, some weigh 72 amu and some weigh 74 amu. How

would explain these three different masses for Cl2 molecules?

UNIT 2 ~ Nomenclature and Formula Writing ~ Test Topics

Resources provided: Periodic table (unmarked – does NOT include colored regions for classes or charges)

and list of element names correctly spelled. List of polyatomic ions (excluding 7 from memorize list – see

below).

To be Memorized (will be extensively used in Unit 3):

1) 7 Diatomic Elements H2, N2,O2,F2, Cl2, Br2, I2 (remember 7 pattern from periodic table)

2) names and charges of 7 common polyatomic ions

Acetate: C2H3O2-1 or CH3COO-1 Ammonium: NH4

+1

Carbonate: CO3-2 Hydroxide: OH-1 Nitrate: NO3

-1

Sulfate: SO4-2 Phosphate: PO4

-3

3) Prefixes for class III:

1 = mono 2 = di 3 = tri 4 = tetra 5 = penta 6= hexa

7 = hepta 8 = octa 9 = nona 10 = deca

1) Given a periodic table with element names, be able to use element locations to:

a) Determine whether an element is a metal (left of staircase) or nonmetal (right of staircase).

b) Determine what type of bond, ionic or covalent will likely form when two elements bond.

Ionic: metal + nonmetal; Covalent: nonmetal + nonmetal

c) Determine the number of valence electrons for a given element. Group number at the top of each column

indicates the number of valence electrons.

d) Determine what charged ion of an element will typically form if is an ionic compound. Key : Remember that

the noble gases have a completely filled outer energy level level of electrons. Other atoms will gain or lose

electrons in order to obtain a completely filled valence level.

2) Be able to interpret chemical formulas by number of atoms present.

3) Be able to describe ionic and covalent bonds. Given a formula of a compound, be able to determine whether ions

are present and how many are present.

4) Nomenclature and formula writing

a) Be able to determine the type or class of rules to use in naming or writing a chemical formula by locating the

first element in the name/formula in the appropriate region on the periodic table.

b) Given the formula of a compound, be able to write the name of the compound; given the name of a

compound, be able to write the formula.

(Or HOFBrINCl)

Unit 2 Notes: Nomenclature and Formula Writing

Introduction - Review of Periodic Table, Atomic Structure and Bonding

Nomenclature: System of ________ compounds Formula: ______ and ____ of atoms in compound

Practice recognizing classes of named compounds:

Class 1: Class 2: Class 3:

1) 1) 1)

2) 2) 2)

3) 3) 3)

Arrangement of Elements in the Periodic Table

• Elements are arranged in order according to atomic # which = # of ___________________ .

• Periodic Table is organized into 7 rows or ____________________ and 18 vertical columns called

__________________ or _______________________.

• Most important principle of organization of elements:

• Example: Group 8A (Family 18), Noble Gases: All elements of this family are

________________ _________________ ____________.

• Elements that exist in pairs of atom are called diatomic.7 Diatomic Elements to memorize for unit test (we will

need this in unit 3):

1) _____ 2) ______ 3) _______ 4) ________ 5) _______ 6) _______ 7) _________

On periodic table arrangement 6 of the 7 elements forms a “7” pattern starting with element #7, N2.

Acronym: “HOFBrINCl”

Metal and Nonmetals

• Metals – shiny solids (Hg only liquid metal), which are shiny (reflect light off of polished surfaces), can be easily

shaped and are excellent conductors of heat and electricity.

• Metals are separated from nonmetals on the periodic table by a staircase. Metals are to the ___________ of the

staircase, and nonmetals are to the _____________ of the staircase.

• In chemical reactions with nonmetals, metal tend to ______ electrons to form ___ ions

• Nonmetals – any element that does not exhibit metallic properties. In chemical reactions with nonmetals,

nonmetals share electrons; in reaction with metals, nonmetals ________ electrons to form _____ ions.

Importance of Valence Electron Energy Levels

• Electrons are arranged inside the atom in ___________________ ___________________.

• The first energy level can hold a maximum of ______ electrons. When the first level is filled electrons are

added to the second level out from the nucleus.

Diagram of filled first level:

• Levels 2-7 (ignoring transition metals) can hold a maximum of _______ electrons.

Diagram of filled first and second levels:

The outermost electrons are called __________________________ __________________.

Diagram of Li with valence level electron labeled:

Lewis Dot Structures: Uses dots to represent valence electrons. In the space above next to your previous diagram

(don’t cross out your electron level diagram like power point slide) draw the symbol for Li with a single dot to represent

the valence electron.

Valence Electrons and the Periodic Table

Note each = valence electron

• The __________ # at the top of groups 1A, 2A, 3A, 4A,5A,6A and 7A and 8A indicates the number of

______________________ (outermost) electrons.

Valence electrons and chemical reactions

• The # of valence electrons determines __________________________ _______________.

• Atoms with filled valence level will ___________ chemically react.

• A completely filled valence level = chemically _________________ or ________________.

• Atoms WITHOUT a filled a valence level WILL ______________ ________ in order to obtain a

_______________ ________ valence electron level.

Example: H vs He vs Li

Helium

• Helium has a ________________ __________ first level ( _____ valence electrons)

Diagram for He:

• Helium will __________ burn when mixed with oxygen and a spark.

• Filled valence level =

Practice:

= F = Na

nb

Hydrogen

• Hydrogen has only _____ valence electron. Diagram of H energy level:

• Hydrogen needs to __________ ______ valence electron to _______ its outer energy level.

• Hydrogen is ________________ _______________. It _______________ when mixed with ____________ and

a spark.

• Hydrogen reacts with oxygen to obtain a filled valence level like He.

Lewis Dot for H: Lewis Dot for Water:

Lewis Dot for He:

Note similarity valence electron configuration of H in water compared to He

• Lithium burns brightly in oxygen. Lithium wants to lose ____ valence electron to ______ its valence electron

level.

Lewis Dot for Li: Lewis Dot of LiO:

Note: Li+ ion has same electron configuration as He.

• Atoms without filled valence electron levels will _____________, ________________, or

___________________ valence electrons to reach stability, a full valence electron level

Neutral Atom

Predicting Atomic Charges in Compounds Containing Metals Bonded to Nonmetals

Metals (M) want to __________ all valence electrons and will form + ions.

Nonmetals (X) want to _______ electrons up to a total of _____.

M → X →

M → X →

M → X →

Gain of e- Loss of e-

General Categories of Chemical Bonds

Chemical Bond – ___________ ____________ between atoms or ions that binds them together as a unit.

Atoms bound together to order to become more stable by obtaining a ________________

___________ _______________ electron level.

Ionic Bonding – (Attractive force between Ions; forms between a ___________ and a ____________)

• A complete ___________________ of one or more electrons from a __________________ to a

nonmetal.

• The + metals are attracted to the – nonmetal ions by opposite electrical charges.

Example: e- is transferred

Na + Cl → [Na] [Cl] Dots represent valence e-

Practice Example : MgF2 , (Mg = metal, F = nonmetal):

F Mg F →

Covalent Bonding (Co = together/sharing; formed between ______________ and ___________)

• The _________________ of one or more ____________ of electrons between two nonmetal atoms.

• Example: 2 e- are shared

H + H → H H or H H

Chemical Nomenclature

Chemical Nomenclature – System for naming chemical compounds

Binary salts – compounds consisting of two elements

Ternary salt- contains three elements (metal + polyatomic ion)

Polyatomic ion – two elements covalently bonded together with an overall charge

3 Sets of Rules or Classes of Naming Rules

Class or Type I - Ionic Bond between metal with a predictable + charge and a negative nonmetal or

polyatomic ion ; name cations, anion ends in “ide”

Examples: NaCl – sodium chloride ; MgCl2, magnesium chloride

Class or Type II- Ionic Bond with variable oxidation state transition metal a negative nonmetal or polyatomic

ion; Name cation, use Roman numerals used to specify charge of transition metal, anion ends in “ide”.

Example: CuCl2 – copper(II) chloride

Class or Type III – Covalent bonds between Two nonmetals; use prefixes to indicate number of atoms

Example: CO2 – carbon dioxide; N2O4 – dinitrogen tetroxide

Class or Type is determined by location of FIRST element on periodic table.

Class or Type I - ________ bond between a ___________ with a _______________ + charge and a

negative nonmetal or polyatomic ion.

Rules for Naming Class I:

Name metal cation (positive + ion) first using element name; Name nonmetal anion (negative – ion 2nd);

change element name of anion to ends in “ide”

Examples: NaCl ; MgCl2

Practice - Name the following compounds:

1) AlCl3

2) ZnBr2

3) Ga2O3

Rules for Writing Formulas:

1) Use location on periodic table to find charge for each element.

2) Sum of charges must total zero.

3) If + charge ≠ charge, use “Criss-Cross” technique

4) Final subscripts must be simplest ratio.

Example #1: (+ charge on cation = - charge on anion)

What is the formula of potassium bromide?

K = group 1 = lose valence e- to fill outer level → ion

Br = Group 7 = gain valence e- to fill outer level → ion

K+1 + Br-1 → KBr

Example #2: (+ charge on cation ≠ - charge on anion)

What is the formula of calcium nitride?

Charges: Ca N Criss-Cross move charge # (ignore + or – ) to subscript on opposite ion

Final Answer: Ca N

Practice- Write the formulas of the following compounds:

1) Magnesium sulfide

2) Aluminum oxide

3) Potassium sulfide

Polyatomic ions - two or more different atoms covalently bonded together with overall charge.

These groups of atoms will bond to other ions to form compounds.

Memorize the following

Name: Formula & Charge Atomic level picture

1)

2)

3)

4)

5)

6)

7)

Concept Check: How many IONS are present in each of the following?

(Hints for a-c each ion is enclosed in [ ];

A) [K]+[OH]- B) [Na]+[CO3-2 ] [Na]+ C) [Na]+ [PO4]-3 [Na]+

[Na]+

For D- I, circle the polyatomic ion and give the total # of ions (from both monoatomic ions such as Na+1

or Ca+2 and polyatomic ions).

D) NaOH E) Na2SO4 F) Ca3(PO4)2

G) NH4OH H) (NH4)2S I) Al(NO3)3

Naming for Ternary salts (compounds containing a polyatomic ion)

Simple Naming Rule: Give each ion its name

Example problems: KNO3 - potassium nitrate Na2CO3 – sodium carbonate

Practice: Name the following:

1) NaOH _____________________________________________

2) Mg(C2H3O2)2 ____________________________________________________________________

3) NH4Cl ____________________________________________________________________

Writing Formulas:

-Find charges of monoatomic ions from periodic table

- Look up or memorize charges of polyatomic ions

- Sum of charges must total zero; use criss-cross if necessary

- If more than one polyatomic ion use ( )

Note: most polyatomic ions (exceptions include hydroxide, peroxide and cyanide) end in either “ate” or

“ite” while elements end in “ide”.

Examples: Ammonium hydroxide = →

Calcium phosphate = →

Practice Formulas:

1) calcium nitrate

3) gallium carbonate

2) cesium hydroxide

Notes: Class or Type II compounds -

Transition metals can form ___________ oxidation states (+) charges, resulting in more than one

possible formula between transition metals and negative ions.

Rules for Naming Class II Compounds

• Metal cation first, negative ion second

• ___ ___________ of metal in written in Roman Numerals in ( )

• For neutral compound: sum of + = sum of –

• Final answer must be simplest ratio

• Example #1: CuCl:

• Example #2: CuCl2:

Practice: Name the following compounds:

1) FeCl3

2) CuSO4

3) Co3(PO4)2

Class II compounds: Writing Formulas • Charge on transition metal will be given in ( )

• Use periodic table or polyatomic ion list to determine negative charge.

• Sum of charges must be zero.

• If + charge ≠ - charge, use “Criss-Cross” technique

• Final subscripts must be simplest ratio.

Examples: 1) Iron(II) oxide

2) Iron(III) oxide

Practice: Write the formulas

1) Tin(IV) oxide

2) Nickel(II) phosphate

3) Manganese(II) nitrate

CLASS 3 – COVALENT BONDS BETWEEN TWO NONMETALS

Rules for Naming Class 3 Compounds

• First element in formula is named first. (First element is generally further to left on periodic table. If two

nonmetals are in the same column element further done is generally written first.)

• Second Element always ends in “ide”.

• __________ are used to denote # of atoms present.

• First element prefix only if more than 1 atom present (i.e. no mono prefix on first element); Second element

always has prefix.

Practice: Name the following 1) BF3

2) CCl4

3) O2F2

Writing Formulas Class III Compounds

• Translate prefixes exactly – DO NOT SIMPLIFY

• Important note: No “criss-crosses” because NO IONS

Explanation for Not simplifying formula

• Ionic compounds such as NaCl exist in lattice consisting of trillions of ions; Formula of ionic

compound is simplest ratio; Covalent compounds consist of small molecules true formula is the

one written.

Practice Formula Writing Class III compounds

1) tetraphosphorous hexoxide

2) Xenon tetrafluoride

3) Sulfur trioxide

Examples: Name the following

1) CO

2) CO2

3) N2O4

UNIT 2 ASSIGNMENT LIST

2-1 Sections 4.8, 4.9 (p.88-96) Intro to periodic table

2-2 Section 12-6, (352-53) Writing Lewis Dot Structures

2-3 Section 4.10 (p.96 – 100) Intro to Ionic and Covalent Bonding

2-4 Practice test student written See directions below

2-5 Section 4.11(p.100 – 103)

Sec. 5.2, 5.2 (p. 114 – 117)

Nomenclature Problem Set Class I

2-6 Section 5.5 (p. 127 – 130) Nomenclature Problem Set Polyatomic Ions

2-7 Section 5.2 (p. 117 – 121) Nomenclature Problem Set Class II

2-8 Section 5.3 (p. 122-124) Nomenclature Problem Set Class III

2-9 Section 5.4 (p.124-126) Unit 2 Review Problem

Directions for HW 2-4: Student written practice test to be worked with a partner

• Minimum requirements:

• Write questions on one side paper, answer key on the back. List the name of the test writer and the test

taker on the front of the paper.

• 2 practice writing formula from names and 2 practice writing names from formulas each for both classes

I and III. (Total: 4 practice naming and 4 practice formulas)

• 3 practice writing formula from names and 3 practice writing names from formulas for class II each and

problems from class I or II that include polyatomic ions. (Total 6 practice naming and 6 practice

formula)

• Problems must be mixed up in random order and contain no identifying labels (e.g. class I problem)

included.

• In separate section or in additional problems make sure that all 7 polyatomic ions (formula, name,

charge) are quizzed.

• In separate section quiz all 7 diatomic elements.

• In separate section quiz all class 3 prefixes.

• Exchange practice quiz with a friend before unit test; Take practice quiz using a periodic table only; then

grade quiz using answers on back or have friend grade paper.

HW Set 2-1- Intro to Periodic Table

(# 1-4 taken from Zumdahl, p. 108, #43-46)

1) What property determines the numerical order of elements on the periodic table?

2) In which direction on the periodic table, horizontal or vertical elements with similar chemical properties

aligned? What are the families of elements with similar chemical properties called?

3) List the characteristic physical properties that distinguish the metallic elements from the nonmetallic elements.

4) Where are the metallic elements found on the periodic table? Are there more metallic elements or nonmetallic

elements?

5) What is the difference between core or inner electrons and valence electrons? Which type of electron core or

valence is most important in determining chemical behavior?

6) The electron energy levels of Na, Cl and Ne are given below. Using dots to represent valence electrons, draw a

Lewis dot structure for each element. Example: H

Homework 2-2- Introduction Lewis Dot Structures

1) Using a dot to represent a valence electron, draw in the correct number of valence electrons around the symbol for

each in the periodic table below. Example: H

2) How can the number of valence electrons be determined from the periodic table?

3) The electron levels of Ar and K are given below.

A) Using dots to represent valence electrons, draw a Lewis dot structure for each element.

B) Predict which element Ar or K would be unreactive (inert). Explain your reasoning.

IA

IIA IIIA

IVA

VA

VIA

VIIIA

Note Group III is also commonly labeled as family 13, Group IVA is labeled as family 14, Group VA is labeled

as family 15 , Group VIA is labeled as family 16, Group VIIA is labeled as family 17, Group VIIIA is labeled as

family 18 on most periodic tables today.

VII

A

4) For each of the following elements indicate whether the atom would gain or lose electrons to obtain a completely

filled energy level, the number of electrons gained or lost, and the resulting charge on the ion. (element # is given in [

] to make easier to locate element on PT.)

Example #1 : F [#9] Gain 1 electron Charge: -1

Example #2: Mg [#12] Lose 2 e- +2

a) Na [#11]

b) Ca [#20]

c) Cl [#17]

d) O [#8]

e) S [#16]

f) N [#7]

g) P [#15]

h) K [#19]

i) Al [#13]

5) For each element in question #4, write M above its symbol if it is a metal and NM above its symbol if it is a

nonmetal. Ex. #1: F = NM Ex. #2 Mg = M

6) Complete the following statement (choose the word that best fills in the blank):

Metals (gain/lose) ________ electrons to form ( +/-) _____ ions.

Nonmetals (gain/lose) ________ electrons to form ( +/-) _____ ions.

Homework 2-3 – Introduction to Ionic and Covalent Bonding

1) Interpreting formulas – counting # of atoms: Indicate the # of atoms present for each:

a) C2H4 b) Mg(CN)2 c) Na3(PO4)3 d) Fe2(SO4)3

C: Mg: Na: Fe:

H: C: P: S:

N: O: O:

2) Label each picture below as Ionic or Covalent.

Choose the word that best completes the blank.

In an Ionic Bond a Metal (transfers electron(s) to / shares electron pairs with ) _______________________________ a

nonmetal.

In a Covalent Bonds a Nonmetal (transfers electron(s) to / shares electron pairs with )

_______________________________ a nonmetal

3) For each of the following pairs of elements:

i) Label metals with M and nonmetals with NM

ii) Predict the type of bond Ionic (I) or Covalent (C) would most likely form.

Example: Li + F : Li (M) + F (NM) = I

A) C + S = ______ B) K + S = ______

C) Mg + Cl = ______ D) Fe + O = ______

E) B + F = ______ F) N + H = ______

G) S + F = ______ H) Cu + Br = _______

4) For each of the following elements diagram how electrons would be transferred from the metal to the

nonmetal and draw the final correct ionic lewis dot structure.

Example:

A) LiF B) MgO C)AlN

D) Na2O E) Al2S3

5) For each of the following compounds:

• Label metals with M and nonmetals with NM

• Predict the type of bond Ionic (I) or Covalent (C) would most likely form.

• Indicate how totals ions are present

Elements M / NM I or C # of ions

LiF M + NM I 2

SO3 NM + NM C 0

CS2

K2S

MgCl2

Fe2O3

BF3

NH3

SF6

CuBr2

Homework 2–5: Class or Type I

For each of the following chemical formula or name, locate the first element to determine the Type or Class of Rules to

use in naming. Write the Type/Class in the blank provided.

SnBr4 ____ AlH3 ____ FeO ____ OF2 ____

XeF6 ____ Ca3N2 ____ CBr4 ____ MnF2 ____

lithium oxide ____ diboron hexachloride ____ copper(I) hydroxide____

barium nitrate ____ cobalt(II) sulfate ____ sulfur dioxide ____

Name or write the formula for the following Type/Class 1 chemicals:

Formula Name

Li2O Lithium oxide

NaI

CaF2

Al2S3

CaBr2

SrO

AgCl

CsI

RaCl2

Ag2S

MgI2

Name Ions Formula

Lithium oxide Li+1 O-2 Li2O

Aluminum iodide

Silver oxide

Potassium nitride

Calcium phosphide

Magnesium fluoride

Sodium sulfide

Barium hydride

Zinc oxide

Cadmium bromide

Lithium fluoride

Class or Type I - Ionic Bond between

metal predictable + charge and a

negative nonmetal or polyatomic ion ;

name cations, anion ends in “ide”

Examples: NaCl – sodium chloride ;

MgCl2, magnesium chloride

Homework 2–6: Polyatomic ions

Polyatomic ions are “many” atoms bonded together with a charge. You must memorize the following 7 polyatomic ions.

Fill in the formula and charges for each ion.

Ammonium = ________ Acetate = ________ Carbonate = ________

Hydroxide = ________ Nitrate = ________ sulfate = ________ Phosphate = ________

All other polyatomics are provided to you on the unit test. Name or write the formula for the following Type/Class 1

chemicals that have polyatomic ions:

Formula Name

Li2CO3 Lithium carbonate

NaNO3

Ca(OH)2

Al2(SO4)3

NH4NO3

AgNO2

KClO4

H2O2

LiOH

NaHCO3

KMnO4

Name Ions Formula

Lithium carbonate Li+1 CO3-2 Li2CO3

Aluminum hydroxide

Silver phosphate

Potassium dichromate

Calcium perchlorate

Magnesium acetate

Sodium nitrate

Barium sulfite

Zinc sulfate

Cadmium nitrite

Hydrogen cyanide

Homework 2 – 7: Class or Type II

Type or Class 2 Rules uses Roman Numerals in ( ) to indicate the charge on the metal, since the transition metals can

have variable oxidation states or charges.

Name or write the formula for the following Type/Class 2 compounds:

Formula Name

Fe2O3 Iron(III) oxide

Cu(OH)2

Co2(SO4)3

CuNO3

SnBr2

MnO2

PbCO3

CrCl3

Nb2O5

Ti(NO3)4

Mo2O5

SbN

Ni2(SO3)3

PtO

HgCl2

Name Ions Formula

Manganese(II) hydroxide Mn+2 OH-1 Mn(OH) 2

Lead(IV) phosphate

Platinum(IV) chloride

Rust: Iron(III) oxide

Copper(II) acetate

Iron(II) nitrate

Cobalt(III) sulfite

Cobalt(III) sulfate

Iron(III) nitrite

Mercury(I) oxide

Mercury(II) oxide

Vanadium(V) fluoride

Chromium(III) carbonate

Mercury(I) sulfide

Examples:

CuCl is copper(I) chloride

CuCl2 is copper(II) chloride

Homework 2 – 8: Class or Type III

Type or Class 3 Rules uses prefixes to indicate the number of atoms, since the nonmetals share valence electrons (no

charges!). First element only needs a prefix for 2 or more atoms while the second element ALWAYS has a prefix.

Example: NO - nitrogen monoxide N2O3 - dinitrogen trioxide

Concept Check: For each of the following indicate how many total ions are present:

A) Class I - AlCl3 B) Class II - Cu(NO3)2 C) Class III – N2O4

Name or write the formula for the following Type/Class 3 chemicals:

Formula Name

ClF

OF2

B2H6

PCl3

ClBr

ClF3

N2O5

P4O10

Cl2O

Name Formula

phosphorus pentachloride

tetraphosphorus hexoxide

sulfur hexafluoride

sulfur trioxide

sulfur dioxide

bromine pentafluoride

iodine monobromide

xenon tetrafluoride

phosphorus trihydride

HW 2-9 ~ Unit 2 Nomenclature Review Problem Set

Periodic Table Overview

1) List the 7 diatomic elements.

2) Identify each of the following as a metal, nonmetal

A) I B) Cs C) Cu

3) For each of the following give the charge on the most common ion.

A) Cl B) Mg C) Al D) O E) Na F) Zn G) Ag

4) Which family of elements is chemically inert(unreactive)? What is special about the valence electron

configuration of these elements that makes these elements unreactive?

Ionic vs Covalent Bonding

5) Ionic bonds are formed between a ________________ and a ___________________. Ionic bonding involves the

transfer of electrons: the metal _____ electrons to form ____ ions and the nonmetal ________ to form ____ ions.

For example, Na will lose _____ valence electron to form a _____ ion while Cl will _______________ this one

electron to form a ____ ion. The + and – charges will attract each other.

6) Covalent Bond are formed when two _______________________ share a pair of electrons.

7) List the type of bonding ionic (I) or covalent (C) that would form between the following pairs of elements.

A) Mg and Cl B) N and O C) Cu and Br

D) C and O E) Pb and I

Rules for Naming and Writing Formulas, Class I, II and III

8) For each of the elements pairs from the previous problem, identify the type of naming rules that would be used

classes I, II, or III. (Hint: Remember to locate first element in pair on periodic table).

Type/Class I Naming and Formulas

9) These rules involve “predictable” metals forming ionic bonds with nonmetals. What do we mean by saying the

charges on the metals are predictable?

10) Name the following: a) NaCl

b) CaBr2

11) Write formulas for the following:

a) silver oxide b) beryllium chloride

Polyatomic ions

12) Polyatomic ions are groups of atoms (2 or more) covalently bonded together with an overall charge. Write the

formula and charge for each of the 7 polyatomic ions you must memorize for the test.

Ammonium , Acetate , carbonate

Hydroxide , nitrate , sulfate phosphate

13) For each of the following: i) circle the polyatomic ion ii) list the ions including charge and iii) indicate the # of

each ion present iv) Give the total # of atoms of each element

A) LiNO3 B) Ca3(PO4)2 C) Co2(CO3) 3

Ions: Li+1, NO3-1

# of each ion: 1, 1

Total Atoms: Li: 1 ; N: 1 ; O: 3

Type/ Class II Naming and Formulas

B) The rules involve transition metals forming ionic bonds with nonmetals. What does it mean to say that charge on

transition metals is “unpredictable” or that transition metals have “variable oxidation states?”

C) Why do we will call Ba(NO3)2 barium nitrate but we call Fe(NO3)2, iron(II) nitrate? What does the roman

numeral after the iron represent?

Type/Class III Naming and Formulas

D) The rules involve nometals forming covalent bonds with other nonmetals. Since the elements are sharing valence

electrons, there are no charges like types 1 and 2. Instead we use ________________ to designate the number of

atoms in the formula.

E) Why is NO called nitrogen monoxide, but N2O is called dinitrogen monoxide (ie why is there no prefix for N in

NO?)

F) Explain why the type I and II compounds like calcium oxide and and lead(IV) oxide, require finding charges,

criss-crossing and then writing the final formula in simplest terms, whereas the class III compound, dinitrogen

tetraoxide you simply translate the prefixes (di = 2, tetra = 4) with no charges, no criss-crosses and you do not

simplify the final answer?

Practice Formula Writing:

1) calcium nitrate 2) calcium nitrite

3) calcium nitride 4) lead(II) oxide

5) lead(IV) oxide 6) phosphorous trichloride

7) iron(II) carbonate 8) zinc nitride

9) cobalt(III) sulfate 10) sodium acetate

11) tetraphosphorus hexoxide 12) ammonium chloride

Practice Naming:

1) KOH 2) NH4NO2

3) CaH2 4) KMnO4

5) Na2Cr2O7 6) XeF2

7) CuCl2 8) NI3

9) CrSO4 10) Cr2(SO4)3

11) B2O3 12) Cu3(PO4)2

Unit 3 ~ Chemical Reactions ~ Test Topics

Writing Chemical Equations

• Write and interpret symbols used in chemical equations (element symbols, compound symbols, phases of matter,

etc.)

Conservation of Mass and Balancing Chemical Equations

• Law of Conservation of Mass – Understand that mass is conserved because matter is made up of atoms that are

neither created nor destroyed in chemical reactions (simply rearranged).

• Be able to balance chemical equations be changing coefficients so that there are the same number and type of

atoms on both the reactant and product side of an equation.

• Using symbolic diagrams to represent atoms, illustrate the concept that atoms are conserved in chemical reactions.

Understand the difference between coefficients and subscripts.

Classifying Chemical Reactions

• Understand the following classifications for chemical reactions:

1) Synthesis 2) Decomposition 3) Combustion of a hydrocarbon

4) Single Replacement 5) Double Replacement 6) Acid-Base

• Given a chemical equation, be able to classify the type of reaction taking place.

• Given the reactants, be able to predict the products of the reaction.

• Understand the concept of activity series in a single replacement reaction. Given a chart of activity series rules, be

able to predict whether a particular single replacement reaction will take place or not.

• Given solubility rules, be able to identify the solid for a precipitation reaction.

• Given the reactants of a reaction, be able to identify the solid for a precipitation reaction.

Acid/Base Reactions

• Know the Bronsted-Lowry definitions of acids and bases; Acid – H+ donor; Base H+ - acceptor

• Be able to identify which substance is the acid and which is the base in a reaction. Be able to recognize:

Acids: HCl = hydrochloric acid HNO3 = nitric acid

CH3COOH = acetic acid or HC2H3O2 H2O = water*

Bases: NH3 = ammonia KOH = potassium hydroxide

NaOH = sodium hydroxide H2O = water *

• * Water is an amphoteric substance, one that can act as either an acid or a base.

• Properties of acids and bases:

Taste: Acids sour, Bases bitter Feel: Acids sting, Bases are slippery

Reactions with metals: Acids react with active metals, bases generally do not.

Electrical conductivity: Both acids and bases are excellent conductors

Indicators: Acid in litmus = RED Base in litmus = BLUE

Acid in phenolphthalein = COLORLESS Base in phenolphthalein = PINK

Representing Chemical Reactions with Equations Reaction of sodium metal and water (Phenolphthalein added)

Initial observations: Final observations:

Phenolphthalein indicator = colorless in acidic and neutral solution, magenta in basic solution.

Interpreting observations to determine products:

Neutral metals do not dissolve in water, but + charged metals do dissolve. What is the charge on sodium ion?

Basic solutions contain a high concentration of hydroxide ion:

Bubbles indicate the presence of a _____. Keeping in mind the Law of Conservation of Mass, what are the 3

different elements on the reactant side? Which two of these products are gases in their elemental

form?

The actual product gas will burn explosively. Which gas is the product?

Is the gas monoatomic or diatomic in its elemental form?

Write the final equation for reaction of sodium (monoatomic or diatomic?) and water below:

Common symbols in chemical reactions:

rxn = reaction NR = no reaction

→ yields or produces → heat added to initiate rxn

(s) = ; ↓ = precipitate (solid product) formed

(l) = (aq) = (dissolved in water)

(g) = ; ↑ =

∆

Balancing Chemical Equations

All chemical reactions need to obey the Law of ___________________________________

Defn of Balanced Equation:

Importance: Balanced chemical equations are used to provide chemical recipes (stoichiometry)

Key idea: Chemical equations can only be balanced by changing ____________________________

Never by changing _____________________________. Final coefficients must be simplest, whole #’s.

Remainder: Coefficients are numbers in ___________ of a formula that indicate how many of that unit (element or

compound) are present. Subscripts are numbers within formula that indicate # of atoms.

3 Br2 ←

Using to represent Br atoms, draw pictures to represent the following: Br Br2 2 Br

2 Br2 3 Br2

Example: Balancing by using coefficients not subscripts:

Given the reaction H2 + O2 → H2O

Can you balance by changing the subscript of H2O to H2O2? Explain.

Single Replacement Reactions: Element + Compound → Compound + Element “Like Replaces Like”

Metals replace _________ and halogens replace ____________.

→

Where MA, MB are metals, C is a nonmetal or negative polyatomic ion.

→

Where M is a metal, X and Y are halogens (F, Cl, Br or I)

Example #1 (metal replaces metal)

Demo: Al(s) + CuCl2 (aq) →

Initial Observations: Al CuCl2

The blue color in the aqueous CuCl2 solution is caused by

Final Observations:

Al (a metal) has replaced _____, (a metal).

Final rxn: Al(s) + CuCl2 (aq) →

Example #2: Halogen replaces halogen : X2 + MY → MX + Y

2

Where M is a metal, X and Y are halogens (F, Cl, Br or I)

KI(aq)

+ Cl2 (aq)

→

Which is the element KI or Cl2?

Is the element a metal or a halogen?

Which is also a halogen, K or I?

Halogen replaces halogen. Recall that elemental form (0 charge, element not part of compound) of

halogens are diatomic.

http://www.mcgrawhill.ca/school/applets/sf10/unit1/single_replacement.html

http://www.mcgrawhill.ca/school/applets/sf10/unit1/single_replacement.html

Activity Series Some processes are spontaneous in one direction but not the other.

Example: Rocks can fall off a cliff to the ground but are not able to fly up off the ground back onto the

cliff without outside intervention.

Single replacement rxns will run in one direction but not the other.

Example: Mg + 2 HCl →

But MgCl2 + H2 →

The activity series chart can be used to predict whether a single replacement reaction will occur.

Using the Activity series chart

Is the element in reaction a metal or a halogen? (To decide which chart column to use)

What is the metal or halogen in the compound?

If the element is higher on the activity chart than the metal or halogen in the compound the reaction will

occur.

Example #1: Will the reaction Ag + HCl → spontaneously occur?

Step #1: Element is =

Step #2: Metal in compound is always first; element in formula =

Step #3: On metal activity chart, is element, Ag, higher than H?

Example #2: Will reaction the reaction F2 + KCl → spontaneously occur?

Step #1: Element is =

Step #2: Halogen in compound is always second; element in formula =

Step #3: On halogen activity chart, is element, F, higher than Cl?

Double Replacement/ Precipitation Reactions

Reaction between two ionic (composed of + and – ions) compounds dissolved in water. Upon mixing both sets of positive

and negative ions switch partners. In most cases, a solid product called a precipitate is formed in the reaction.

AX + BY →

Compound + Compound → Compound + Compound

Where A and B = + charged metal ions ; X and Y are negatively charged nonmetals or negative polyatomic

ions

Important note: + ions are always written 1st in formulas and – ions are always written 2nd.

Key foundation concept: Ionic compounds dissociate into separate _______ when dissolved in

___________.

Example: NaCl (s)

+ H2O

(l)→

Draw Diagram of Na+ and Cl- ions surrounded by water molecules:

Fill in the ions that would be formed in each beaker when aqueous solutions of NaCl and AgNO3 are mixed:

Demo Reaction: CuCl2 (aq)

+ Na2

CO3 (aq)

→

Step 1: List ions

To determine charge on Cu:

Concept: Draw pictures to represent the dissociation of the ions:

CuCl2

Na2

CO3

Note that subscripts for Cl and Na were ignored because they indicate # of different ions but the subscript within the

CO3

-2

was included because it is part of the polyatomic ion.

Step 2: Combine + and – ions from different compounds

Step 3: Formulas of Compounds

Step 4: Balance the equation

Step 5: Use solubility rules chart to determine phases of products and write final balanced equation

including phases.

(aq)

(s) ppt.

Notes: Acid-Base Reactions

Hydrogen, H, consists of a ______________ and an ________________.

Diagram of H atom:

If an electron is removed from the H atom, only the ___________ remains.

Proton symbol: ______

Bronsted-Lowry Definitions of Acids and Bases:

Acid: H+ ___________________ Base: H+ ____________________

Examples of Acids: vinegar, Orange Juice, wine, aspirin

Examples of Bases: coffee, cigarettes, tonic water, baking soda

Properties of Acids Properties of Bases

Taste:

Feel:

React with active metals?

Solutions conduct electricity?

Acid- Base Indicators

Indicators are chemicals that turn _____________________ __________ in acidic or basic solutions

(different pH’s).

AciD in litmus = _________, base in litmus = _________

Acid in phenolphthalein = ______________ Base in phenolphthalein = ______________

Need to be able to identify the following chemicals in an equation as acids or bases:

Acids Bases

Hydrochloric acid = ______ Ammonia (covalent, no ions) =

Nitric acid = ________ Sodium hydroxide (ionic) = Ions: +

Acetic acid = ________ Potassium hydroxide (ionic) = Ions: +

Water = _______ Water = _______

Recognizing Patterns: Concept of Acid-Base Conjugate Pairs:

When an acid loses an H+ it forms it is conjugate base. HA → A- + H+ ; HA and A- are conjugate pairs.

When a base accepts an H+ it forms a conjugate acid. B- + H+ → HB; B- and HB are conjugate pairs.

Conjugate acid/base pairs always differ by an H+.

Example: HCl + H2O → Cl- + H3O+

Conjugate Acid – Base Pairs: Write the conjugate for each of the following

HNO3 C2H3O2H HCl

H2O NH3 OH-

Acid-Base Reactions

Example #1: NH3 + HCl →

Water is amphoteric meaning it can act as an __________ or a __________.

Example #2:

NH3 + H2O →

When water is combined with a base it acts an _______.

Example #3:

HNO3 + H2O →

When water is combined with an acid it acts an _______.

Example #4: Ionic bases – must remember to dissociate ions:

NaOH(aq) + HCl(aq) →

Unit 3 Assignments

Reading in Zumdahl text Homework

3-1 Sec. 6.1, 6.2 (p. 143 – 149)

Section 6.3 (p. 149 – 156)

Balancing Equations Problem set

3-2 Create your own practice test;

exchange with a partner

Min. Requirements: 2 eachproblems form syn,

decom, comb., S.R., D.R. ; 3 balancing; 4 acid-base

3-3 Sections 7.5 ( p. 181-183) 7.7 (p.

188-191)

Reaction Problem Set #1-14

3-4 Single Replacement Rxn

3-5 Activity Series

3-6 Section 7.2 (p. 166 – 175) Double Replacement Precipitation rxns

3-7 Practice Quiz

3-8 Section 7.4 (p. 177 – 180) Acid- Base Reactions

3-9 Unit 3 Review Problem Set (WB)

HW 3-1

Level 1 Balancing Problems: Example: K + Br2 → KBr

Atomic Level Picture: K = Br =

Atom Count: K: Br: → K: Br:

Level 1 Practice: Balance the following reactions:

1) _____ Mg + ______ O2 → ______ MgO

2) _____ Na3N → _____ Na + ______N2

3) _____ Al + ______ MgO → ______ Mg + _____ Al2O3

Level 2: Helpful hint: Least Common Multiple:

Example: ______ KClO3 → ______ KCl + ______ O2

K: Cl: O: → K: Cl: O

Problem: 2 O on right will not divide evenly into 3 O on left. What is the least common multiple of 2 and 3?

4) ____ Al + _____O2 → _________ Al2O3

5) ______ Al + _____ HCl → _________ AlCl 3 + _______ H2

Level 3: Most complex problems have elements that appear in multiple terms on each side and/or issues with

odd / even.

Hints: Balance elements that appear in only 1 term on each side first.

Remember that the subscript for an element term on one side is an even # and is odd in one term on the

other side, doubling odd #’s is the easiest way to convert odd #’s to even #’s.

Example #1: _____ CH4 + ____ O2 → _____ CO2 + _____ H2O

Start with _____ or _______ because those elements appear in only 1 term on both sides.

Save ______ for last because it appears in multiple terms and would require filling in 3 coefficients

simultaneously which is more complex.

Practice:

6) _____ C6H12O6 → _____ C2H6O + _____ CO2

7) _____ C5H12 + ______ O2 → ______CO2 + _____ H2O

8) _____ C2H5OH + _____ O2 → ______ CO2 + _____ H2O

Odd-even:

Example #1: ______ H2O2 → ______ H2O + _____ O2

Example #2: _______ C4H10 + _____ O2 → ______ CO2 + ____ H2O

9) ______ C8H18 + _____ O2 → ______ CO2 + ____ H2O

10) ______ NH3 + ______ NO → _______ N2 + _____ H2O

Challenge: Cu + HNO3 → Cu(NO3)2 + NO + H2O

Reactions Problem Set

Homework 3-3 : Synthesis, Decomposition, Combustion of a Hydrocarbon Predict the products and balance the reactions for each type listed below:

Synthesis: Demo: Initial observations: Al: I: Final Observations:

Al(s) + I2(s) →

1) Na + O2 → 2) Al + F2 →

3) Sr + O2 → 4) K+ S8 →

5) Na + N2 →

Decomposition: Demo: Initial observations: NI3 Final Observations:

NI3 (s) →

6) AlCl3 → 7) HgO →

8) H2O → 9) Al2O3 →

10) FeBr3 →

Combustion

Demo: “Methane Mamba”: Observations:

CH4(g) + O2(g)

11) C2H4O2 (l) + O2 (g)

12) C6H12O6 (g) + O2 (g)

13) C2H2 (g) + O2 (g)

14) C2H6 (g) + O2 (g)

Homework 3-4 Single Replacement: Element + Compound → Compound + Element

Metal replaces Metal : MAC + MB → MBC + MA , where MA and MB = metals; C = - ion

Halogen replaces Halogen: X2 + MY → MX + Y2 , where X and Y are halogen, ; M = + metal ion

Predict the products of the following reactions and balance the final equation:

1) HCl + Zn → 2) ZnCl2 + Mg →

3) Cl2 + CaBr2 → 4) Al + SnCl2 →

5) F2 + NaCl → 6) Br2 + KI →

Homework 3-5 Single Replacement Reactions and the Activity Series Use the Activity series

Chart below to predict whether a reaction will occur. If it does not, write NR, if it does write the products and

balance the final equation.

1) Ni + MgSO4 → 2) Mg + NiSO4 →

3) Cl2 + KI → 4) I2 + KCl →

5) Cu + HCl → 6) Ca + HCl →

Homework: 3-6 Double Replacement (Precipitation Reactions) For each reaction: list the ions, write the products and use the solubility rules chart to determine phases aqueous (aq) or

solid (s). Place a box or a circle around the product(s) that form a precipitate (solid) or ↓ symbol.

22) HCl (aq) + AgNO3 (aq)

Ions:

23) FeCl3 (aq) + NaOH (aq)

Ions:

24) CaCl2 (aq) + Na3PO4 (aq)

Ions:

25) Na2SO4 (aq) + BaCl2 (aq)

Ions:

26) AlCl3 (aq) + Pb(NO3)2 (ag)

Ions:

HW 3-7 - Practice Quiz– Predicting Reaction Products

Part 1: Predicting Reaction Products (#1-8) Directions: For each of the reactions below:

∙ Write the Reaction type ( SYN = synthesis, DECOMP = decomposition, COMB. = combustion, S.R. = single

replacement, D.R. = double replacement) in the space provided.

Write the correct formulas for the products. (You do not need to balance the final equation with coefficients – balancing

will be evaluated separately in Part 2.) In all cases a reaction occurs (no NR). Phases of matter are NOT necessary.

1) Reaction Type: Al + F2 →

2) Reaction Type: HCl + Zn →

3) Reaction Type: C2H6 + O2 →

4) Reaction Type: FeBr3 →

5) Reaction Type: FeCl3 + NaOH →

6) Reaction Type: KCl + F2 →

7) Reaction Type: AlCl3 + Pb(NO3)2 →

8) Reaction Type: C3H8 + O2 →

PART 2: BALANCE THE FOLLOWING EQUATIONS BY FILLING IN THE MISSING COEFICIENTS. (If you leave

a ________ blank the coefficient will be understood to be 1.) Use the top for work, final answer fill in the blanks.

1) Al + HCl → . H2 + AlCl3

Final Answer: ______ Al + _______ HCl → _______ H2 + _______ AlCl3

2) C4H10 + O2 → . H2O + CO2

Final Answer: ______ C4H10 + _______ O2 → _______ H2O + _______ CO2

3) NH3 + O2 → . H2O + NO2

Final Answer: ______ NH3 + _______ O2 → _______ H2O + _______ NO2

Homework: 3-8 Acid/Base Reactions

Acids are ______ _______________; Bases are _______ _____________.

Conjugate acid-base pairs differ by an _____.

For each of the following acids or bases write its conjugate:

1) HCl : 2) OH- :

3) HNO3 : 4) C2H3O2-:

NaOH and KOH are referred to as ionic bases meaning that they dissociate in + and – ions in when dissolved in water.

Write the dissociation equation for each compound.

NaOH(aq) → KOH(aq) →

Write the products for the following reactions:

Part 1: NH3 and H2O covalent bases (don’t dissociate).

1) NH3(aq) + HNO3(aq) 2) C2H3O2H(aq) + NH3(aq)

3) HCl + H2O → 4) C2H3O2H(aq) + H2O →

Part 2: Ionic bases (dissociate into ions)

5) HCl(aq) + NaOH(aq) 6) NaOH(aq) + C2H3O2H(aq)

7) HNO3(aq) + NaOH(aq) 8) KOH(aq) + HNO3(aq)

9) C2H3O2H(aq) + KOH(aq) 10) HCl(aq) + KOH(aq)

HW 3-9 ~ Unit 3 Reactions ~ Reactions Review Problem Set Writing Chemical Equations

1) Chemical equations include symbols that represent elements, compounds, phases of matter and other important

information. Provide the appropriate symbol for each of the following

a) Solid b) Liquid c) gas d) aqueous

e) yields or produces f) heat added g) solid ppt produced

h) gas released

2) What is the difference between a liquid and aqueous, H2O(l) and NaCl(aq) ?

3) Balance the following equations:

A) Ga + O2 → Ga2O3

B) C3H8O + O2 → CO2 + H2O

C) C10H22 + O2 → CO2 + H2O

Reaction Patterns:

1) Which reaction type always produces CO2 and H2O as the two products?

2) This reaction involves the rapid burning of a hydrocarbon and the element ________________.

3) Diatomic or Not? Consider the following single replacement reaction:

Zn + 2 HCl → ZnCl2 + H2

In this reaction both Cl and H have subscripts of 2, for different reasons.

a) Which element has a subscript of 2 because it represents a diatomic form (isolated element, no charge)?

b) Which element has a subscript of 2 because it has a charge of -1 in a compound and is bonding with a metal with

a +2 charge? (i.e. subscript comes from the cross-cross)

4) For each of the following identify the pattern type (synthesis, decomposition, combustion of a hydrocarbon, single

replacement, double replacement)

4a) C → A + B 4b) A + B → C

4c) AX + BY → AY + BX 4d) CxHy + O2 → CO2 + H2O

4e) MAC + MB → MBC + MA

5) Use the activity series chart to predict whether following reactions will take place. If the reaction does occur,

write the products.

A) Mg + AgNO3 →

B) I2 + NaF →

6) Use the solubility rules to identify the solid product in the following reactions. Circle the solid product (precipitate).

A) Cu(NO3)2 (aq) + 2 NaOH(aq) → Cu(OH)2 + 2 NaNO3

B) AgNO3 (aq) + Na2CO3(aq) → Ag2CO3 + 2 NaNO3

7) Reaction Practice: For each of the following reactions, identify the type, predict the products and then balance the

equation.

A) Type: Ca3N2 →

B) Type: C4H10 + O2 →

C) Type: Ca + HCl →

D) Type: Mg + N2 →

E) Type: NiCl2 + KOH →

F) Type: F2 + KCl →

G) Type: Al + O2 →

H) Type: HgO →

I) Type: Na3PO4 + Co(NO3)2 →

J) Type: C5H12 + O2 →

Properties of Acids and Bases:

1) Indicators: What color are the following solutions in the indicator

Indicator Acidic Solution Basic Soluion

Litmus

Phenolpthalein

2) _________________ react with active metals.

3) True or False: Acids taste bitter, Bases taste sour.

4) Which of the following aqueous solutions would conduct electricity due to the presence of high concentrations of

ions?

A) HNO3 B) pure, neutral H2O C) KOH D)Na3PO4

5) Acids always act as H+ ____________________, while bases act as H+ ____________

6) Conjugate acid-base pairs such as HCl and Cl- always differ by an _____.

7) Predict the products of the following acid – base reactions:

A) HCl + NaOH →

B) HNO3 + NH3 →

C) KOH + C2H3O2H →

D) HCl + H2O →

Unit 4/5 Mathematical Tools for Chemistry and the Mole Concept

Units, the Metric System, and Dimensional Analysis

• Memorize the following commonly used SI prefixes in the table in the box on Homework 4-1.

• Be able to convert from one set of units to another using Dimensional Analysis (the Factor Label Method). YOU

MUST SHOW UNIT CANCELLATION when solving problems for this unit.

Significant Figures (Zumdahl, pp. 23-28)

• Be able to define and explain importance of significant figures and uncertainty.

• Be able to record a measurement to the correct number of significant figures and uncertainty.

• Be able to determine the number of significant figures in a recorded measurement.

• Be able to apply rules for addition/subtraction (fewest decimal places) and multiplication//division (fewest

significant figures in a calculation.

Scientific Notation

• Be able to convert numbers between decimal form and scientific notation (and vice versa)

Mole Concept

• Know the definition of a mole ( 1 mole = 6.022 x 1023 particles), why it is such a large number, and

how it is used to relate number of particles to the mass of particles.

• Understand mole concept utilizing concept of counting by weighing. Atoms, molecules and ions are too

small and numerous to be counted with the naked eye.

• Understand that average atomic mass on the periodic table can represent two things:

1) The mass of 1 atom measured in atomic mass units (amu)

2) The mass of 1 mole (6.022 x 1023 particles), why it is such a large number, and how it used (to relate number of

particles to the mass of particles).

• Be able to interconvert between grams ↔ mole ↔ # of particles (atoms, molecules, etc.) using the “mole bridge.”

Note : this schematic will NOT be provided on the exam!

6.022 x 1023 Mass on Periodic Table

% Composition, Empirical and Molecular Formulas

Know the definitions of:

% composition by mass = mass of the element divided by mass of compound x 100%

Empirical formula = the simplest whole ratio of atoms in a compound

Molecular formula = the actual formula of a covalently bonded compound

Note: Subscripts in a chemical formula may represent # of atoms or # of moles but NOT a ratio of masses.

• Be able to calculate percent composition given the masses of the individual elements or calculate the masses of

the elements given % composition

• Be able to calculate the empirical formula of a compound from the elemental analysis (% composition by mass)

using the four steps discussed in class: 1) Determine masses of each element 2) Convert grams to moles

3) Divide by the smallest # of moles 4) If needed, multiply by smallest integer to get whole # ratio

• Be able to calculate the molecular formula from the molar mass and the empirical formula.

Molecular formula = n x empirical formula or

Atoms, Ions or

Molecules (particles) MOLES GRAMS

n = molecular formula mass

empirical formula mass

SIGNIFICANT FIGURE NOTES

Counting #’s – known exactly; no _____________________.

Measured #’s –_______________ involve an _________________ with an uncertainty in the LAST measured digit.

Uncertainty –_____________ of __________________ _ ____________

Example: 13.76 g + .01 g ; True value lies within range 13.7__ ↔ 13.7__ g

Definition of Significant Digits:

- Digits in a measurement that are __________________ given the ______________ of the measuring device

- All of the places in a measurement that are ____________ plus ____ estimated place.

Importance of Significant Figures - The conclusions that you can draw from data cannot ___________ the accuracy your

measuring device can actually measure.

Rules for Recording Significant Digits:

1) Digital Electronic Device (e.g. digital scales): record all of the numbers ________ as they appear on the

screen.

Example: Screen reads: 1000.00 g

Record: ; Uncertainty = + g; True value is 999.99, 1000.00 or 1000.01

Incorrect 1000 g; Implies uncertainty is + g; True value would be 999 , 1000 , or 1001 g

2) Nondigital Electronic Device (rulers, graduated cylinders)

DETERMINE THE ______________ MARKED UNIT

ESTIMATE _______ PLACE TO THE ____________ OF THE SMALLEST MARKED UNIT

EXAMPLE: Smallest marked unit = .1 cm → Estimate to nearest .01

Graduated cylinder – Read at _____- ____________ from the ____________ of the _________________.

3 4 5 centimeters

a b c

RULES FOR INTERPRETING THE NUMBER OF SIGNIFICANT FIGURES IN A MEASUREMENT

When is a significant digit not Significant?

When it is a - !

- It is sometimes necessary to insert a zero to locate a decimal even though a place has not been accurately measured.

Example: Newspaper Headline: 500,000 ATTEND FREE CONCERT IN CENTRAL PARK

In reality, this # is an _____________ the exact # of people who attended is .

Can’t report # as 5_ _ _ _ _ _ Can’t report # as 5 (very different than ½ million!)

Convention: Use zero’s to take help locate decimals even though we haven’t actually measured those places.

COUNTING #’S – numbers whose values are known with no .

Example:

NONZERO DIGIT’S: are significant.

Examples:

LEADING ZERO’S: – zero’s in of all nonzero digits (i.e. to the left of all nonzero digits)

LEADING ZERO’S are ____________ SIGNIFICANT.

Example:

CAPTIVE ZERO’s : ZERO’s nonzero digits.

CAPTIVE ZERO’s ARE _____________ SIGNIFICANT.

Example:

TRAILING ZERO’S : Zero’s at the of a number (i.e. to the right of all nonzero digits)

TRAILING ZERO’S ARE SIGNIFICANT UNLESS MARKED BY A .

Examples:

CONVERSION FACTORS – → ____________

Conversion are exact definitions so you can choose the number of significant figures. Choose # of significant figures in

conversion factor to at least match # of significant figures in the given.

In conversion problems: final SF in calculation match =

Example: 6.0 in ( 2 sig figures) → convert to feet → final answer 2 sig figures

CALCULATIONS INVOLVING SIGNIFICANT FIGURES

CALCULATIONS USING MEASURED VALUES CAN BE NO MORE ACCURATE THAN THE

_____________ ____________________ MEASUREMENT

POINT OF CONFUSION: Definition of “least accurate” is different from addition and subtraction compared to

multiplication and division

Least accurate: Addition and Subtraction =

Least accurate: Multiplication and Division =

RULES ADDITION/SUBTRACTION

Round to - past decimal

Example: 3.05 cm ← places past decimal

+ 12.1 cm ← places past decimal

114 cm ← places past decimal

129. 15 → cm (Rounded to least accurate, 0 places past decimal)

RULES MULTIPLICATION/DIVISION

Round final answer to same # of sig figures as factor with fewest SF

Example 3.05 cm ← SF

x 12.1 cm ← SF

110 cm ← SF

4059.55 → cm3 ( Rounded to least accurate, SF)

Notes ~ The Mole Concept

Chemist’s Definition of a Mole:

____________________ particles (atoms, molecules, ions or formula units) of a substance.

How Many Particles Are Present? ▪ In chemistry, it is important to know the # of atoms (or molecules) that you have.

▪ Chemical _____________, like ______, give us ratios of atoms: ___ H atoms bonded to ___ O.

▪ Chemical _____________: ____ H2 + O2 ___ H2O give us ________ of reacting particles.

▪ Both formulas and equations are needed to calculate __________ for reactions.

BIG Problem

▪ Atoms are too _________ to _____ with the naked eye!

▪ STM image of Si on 0.09 nm scale. (0.09 x 10 m = 9 x 10 )

▪ Counting would take a _____________!

Solution to Counting Problem ~ The Mole Concept

The mole connects the __________ of particles present with a property we can measure in the lab: ___________

(Counting by _______________ and Counting by _______________)

↔ ↔

Can’t see Can Measure

6.022 x 1023 grams from PT

The Mole is a Number

▪ The mole represents a # of particles.

▪ A pair = ____ A dozen = ____ A mole = ____________________

▪ This number is often referred to as “Avogadro’s Number”

The Mole is a BIG Number

▪ 6.022 x 1023 represents an incomprehensively large number!

▪ If a mole of marbles were laid on the surface of the earth, how much of the earth would be covered?

▪ Cover entire earth to a height of ______ miles!

Concept of Counting by Grouping

A box of atoms contains 12 atoms.

Count the atoms below by drawing a box around each set of 12 atoms. How many total sets are present? How

many total atoms are present?

Counting by Grouping – to avoid working with extremely large numbers, we count by grouping

into units containing 6.022 x 1023 atoms – called a mole.

Examples:

▪ Instead of 6.022 x 1023 , write ____ mole.

▪ Instead of 3.011 x 1023 , write ______ (6.022 x 1023) or _____ mole.



Counting by Weighing – determining the # of marbles in a bag

How can you account for small variations in the masses of individual marbles?

How could you determine the number of marbles in the bag using a scale?

Developing the Definition of a Mole

• The definition of a mole was chosen to make it easy to determine the mass of 1 mole of atoms using the periodic

table.

What is the pattern in the following data?

Element(atomic #) # of atoms # of moles Mass in grams He (#2) 6.022 x 1023 1.000 4.00 g

Li (#3) 6.022 x 1023 1.000 6.93 g

B (#5) 6.022 x 1023 1.000

10.81 g

C (#12) 6.022 x 1023 1.000 12.01 g

Cu (#29) 6.022 x 1023 1.000 63.55 g

Au (#79) 6.022 x 1023 1.000 196.97 g

Key to Applying Mole Concept:

Mass of 1 mole of atoms ( atoms) of an element = __________ mass

from periodic table expressed in __________.

Chemical Formulas

What information is in a chemical formula?

• A ratio of __________. In 1 H2O molecules there are _____ atoms of H and ____ atom of O.

• A ratio of __________. In 1 mole of H2O there are _____ moles of H and ____ moles of O.

Formulas are NOT ratio of grams. In 1 mole of H2O, there are NOT 2 g of H for every g of O.

How do chemists determine the formula of a compound?

Step 2: Determination of empirical formula

Empirical Formula = _______________ ratio of atoms in formula

IONIC COMPOUND (+ metal ion/-nonmetal or polyatomic ion)

Empirical Formula = _____________________ formulas

Ions bond together in a lattice – giant “molecule” ;

Actual formula would be too difficult to write,

e.g. Na1023Cl10

23, so formula is written as simplest ratio, ________.

COVALENT COMPOUND (2 or

more nonmetals sharing pairs of

electrons exist as molecules.

Empirical Formula (simplest ratio)

may or may not be true ratio

True Formula is called ___________

Formula

Example :

H2O empirical = molecular

Benzene empirical ≠ molecular

Empirical= Molecular =

MOLECULAR FORMULA DETERMINATION – requires empirical formula + ________ _______

Molar Mass Determined by Experiment (in problems, must be provided).

Step 1: Elemental Analysis

• Identify _____________ present

• Determine ____ ________________ by _________

for each element.

• Determine _____

______________by ______ for

each element

Empirical Formula vs Molecular Formulas

Empirical Formula = ________________________________ ratio of atoms in a compound.

Molecular Formula = ________________________________ ratio of atoms in a compound.

Ionic compound → Empirical Formula is actual formula

Molecular compound (covalent bonds) → Molecular formula (actual formula) is usually different than empirical formula.

Determining empirical formula is still key step in determining molecular formula.

Examples of Empirical and Molecular formulas for covalently bonded compounds:

Carbon dioxide: Empirical = CO2 Molecular = CO2

Benzene: Empirical = CH Molecular = C6H6

Glucose: Empirical = Molecular = C6H12O6

% Composition by Mass

Definition of % : part out of _____; To calculate % ____________ x 100%

Mass % of each element must add up to ______

Example #1 ~ % from mass data~ What is the % by mass of compound containing 28.0 g Fe and 8.0 g O?

% Fe =

%O =

Example # 2: ~ % from a chemical formula~ What is the % by mass of each element in water, H2O?

Step 1 – calculate grams of each element and molar mass

H2:

O:

H2O:

%H = %O =

Work HW 5-4, #1.

Definition of Mass % = Mass of x 100%

Total Mass of Compound

Determining Empirical Formula from Elemental Analysis Data

Example #1: Elemental Analysis (% by mass) for a compound is: Al: 32.13% F: 67.87%

Determine the empirical formula:

Recall Empirical formula =

Step 1: Find masses of elements. Since data is in mass % assume 100 g sample; then can replace % with g.

Step 2: Convert to moles.

Al: 32.13 g Al

F: 67.87 g F

Step 3: Divide by the smallest # of moles:

Al: F:

Final Answer: Al F Work HW 5-4, #2

Example #2: An oxide of aluminum is formed by the reaction of 4.151 g of Al with 3.692 g of oxygen.

Calculate the empirical formula of this compound.

Step 1: Find masses of elements. (In this problem masses are given; If % are given replace % by g)

Step 2: Convert to moles.

Al: 4.151 g Al

O: 3.692 g O

Step 3: Divide by the smallest # of moles:

Al:

O:

We cannot write an empirical formula with 1.5 moles (atoms) of oxygen. There is no such thing as AlO1.5. Atoms do not

exist as fractions! We need one more step.

Step 4: Multiply the numbers from step 3 by the smallest integer that will convert them to whole numbers

These whole #’s are the subscripts in the empirical formula:

1.500 x 2 = ________ atoms of O

1.000 x 2 = ________ atoms of Al

Additional Hints for Step 4: Converting decimals to whole #’s . In our problems, after dividing by smallest # of moles if

decimals are still present they will represent in halves, thirds or quarters:

Examples: Consider moles elements X and Y after dividing by smallest # of moles:

Moles X Moles Y Multiply by Final Answer

1.00 2.00 N/A XY2

1.25 1.00 4 X5Y4

1.33 1.00

1.50 1.00

1.67 1.00

1.75 1.00

1.99 1.00

1.89 1.00

Determining the Molecular Formula of Covalently Bonded Compounds

Molecular formula = n x (empirical formula) ; where n = small whole #

Molar Mass = n x empirical formula mass → n = molar mass

Empirical formula mass

Example: Napthalene, the active ingredient in mothballs, has an empirical formula of C5H4. In a separate experiment,

the molar mass is determined to be 128.17 g/mole: What is the molecular formula of this compound?

Step 1: Calculate empirical formula mass: C5:

H4:

Step 2: n = molar mass

Empirical formula mass

Step 3: Molecular Formula = empirical formula x n =

Al O

Work HW 5-4, #3 and 4

Work HW 5-5

Unit 4 and 5 Assignments

HWK # Pages to Read Problems

4-1 Sec. 2.1, 2.2, 2.3 (p.15 – 22) Scientific and Metric Conversions Prob. Set

(workbook)

4-2 Section 2.6 (p. 28 – 33) Dimensional Analysis Problem Set (workbook)

4-3 Spooks/Gorks + Double Decker Problem Set

4-4 Sections 2.4, 2.5 (p.22 – 28) Significant Figures Problem Set

5-1 Sec. 8.1, 8.2 (p. 203 – 208) Mole Concept Problem set (workbook)

5-2 Sec. 8.3, 8.4 (p.208 – 218) Molar Mass (workbook)

5-3 Mole Bridge (workbook)

5-4 Sections 8.5, 8.6,8.7

(p.218 -227)

Mass % and Empirical formulas :

Self-Checks: 8.7 (p. 220), 8.8 (p. 225)

8.9 (p.226), 8.10 (p. 227)

5-5 Section 8.8 (p. 227 – 229) Self-check 8.11 (p. 228); Page 234, #78,80.81

Unit

Review

4/5

Unit 4 and 5 Review Problem Set

HW 4-1 Scientific Notation and Metric Conversions Problem Set

Scientific Notation: Method for ______________________________very small or very large numbers.

Scientific Notation Format: N x 10x , where 1 ≤ N < 10, x = whole number

Examples: Very Large (# of atoms in a 12 g sample of carbon) : 602 000 000 000 000 000 000 000 = 6.02 x 10 atoms

Very Small ( carbon- carbon bond length) 0.000 000 000 154 m = 1.54 x 10 m

Write the following numbers in scientific notation.

1) 810,000 2) 0.000 000 000 13

3) 102 000 000 000 000 000 4) 0.000 000 723

Convert the following numbers from scientific notation to standard notation:

5) 3.01 x 104 6) 1.2 x 10-2

Calculator Practice: Use EXP or EE button for Scientific Notation:

Example: 6.02 x 1023 → 6.02 EE23 or EXP23 Screen will read 6.02 23

7) 3.01 x 6.02 x1023 = 8) 5.89 x 1022 ÷ 6.02 x 1023 =

Fill in missing exponents in the box below. Use this table to solve metric conversion problems in this unit.

MEMORIZE THIS TABLE FOR UNIT TEST

1 km= 10 m k = kilo

1 m = 10 dm d = deci

1 m = 10 cm c = centi

1 m = 10 mm m = milli

1 m = 10 µm µ = micro

1 m = 10 nm n = nano

Note: Prefixes apply to all SI units. Example: 1000 mL = 1 L or 103 g = 1 kg

FACTOR LABEL CONVERSION PROBLEM SET 4 - 2

Factor Label Method Key Concepts:

➢ Method for converting between units, e.g. cm → m or g → kg

➢ Treats units as algebraic quantities:

(7)(7) = 7 ; (x)(x) = x ; (cm)(cm) = cm

7 = x = cm =

7 x cm

➢ Conversation factors written as ratios

Example: Given: 1 in = 2.54 cm → Possible ratios: 1 in or 2.54 cm

2.54 cm 1 in

➢ Start by writing given, write ( _________), choose ratio to cancel unwanted units.

Example: What is the height in cm of a person that is 68 inches tall?

Correct: Incorrect:

For each of the expressions below calculate the final answer and express in the correct units.

( )

( )( )( )

Given the conversion factors below, fill in the missing numbers and units. Calculate the final answer and express it in the

correct units.

Given: 1 mole Ne = 20.18 g Ne 1 mole Ne = 6.022 x 1023 atoms

( )( )

3.011 x 1020 atoms Ag atoms

6.022 x 1023 atoms Ag atoms

1 mole Ag atoms

=

46.0 g Na 1 mole Na

22.99 g Na

32.0 g O2 2

4 mole Na

1 mole O2

1 mole O2

1)

2)

=

60.54 g Ne 1 mole Ne 6.021 x 1023 atoms Ne

atoms = 3)

Double Decker Problems – have units in both the numerator and denominator

Use the conversion factors below for questions #4 and 5: Given: 1 km = 0.62 miles ; 60 s = 1 min ; 60 min = 1 hr

On Germany’s superhighway, the Autobahn, there are sections of roads with no speed limits. The speed of a car one of

these sections was measured at 5.56 x 10-2 km/s. Convert this speed into units of miles/hour (mi/hr). Be sure to show

unit cancellation!

( )( )( ) It is well know that light travels faster than sound. For example, thunder is always observed after a lighting flash, never

before. By what factor is the speed of light faster than the speed of sound?

A) The speed of light is 3.00 x 105 km/s. Convert this speed into miles/hour.

( )( )( )

(

Given:

1 km = 103 m k = kilo 1 m = 101 dm d = deci

1 m = 102 cm c = centi 1 m = 103 mm m = milli

1 m = 106 µm µ = micro 1 m = 109 nm n = nano

Perform each of the following unit cancellations by showing unit cancellation:

6) 235 mL → ? L (1 L = 103 mL)

( )

7) 656 nm → ? m (Draw your own ( ) and fill in.)

656 nm

4)

5.56 x 10-2 km/s

0.62 miles

1 min 1 hr

=

5)

3.00 x 105 km/s

s =

B) Divide your answer to part A by the speed of sound which is 759 mi/hr. Your answer represents the factor by which

light travels faster than speed. What happens to the units when you divide?

235 mL =

Part 2: Two or more step conversions (Hint: convert through base unit; the base unit for length is the meter)

8) 8.43 cm → ? µm

( )( )

9) 295 mm → ? km (Draw your own ( ) and fill in.)

295 mm

10) The distance between stars in space are so great that astronomers often use a unit called a parsec to express the

distance. For example, the distance between our Sun its closest star, Barnard’s Star is 1.8 parsecs. To understand why

this distance is not commonly expressed in m, calculate how many meters away Barnard’s Star is from our Sun.

Express your answer in correct scientific notation. (Which is simpler number to work with 1.8 parsecs or your final

answer to this problem?)

Given: 1 parsec = 3.3 light years ; 1 light year = 9.46 x 1012 km

Add in your own ( ).

1.8 parsecs

(

8.43 cm =

1 m

1 m

Homework 4-3 ~ Challenge Factor Label Problems

Directions: If you understand the factor label method you should be able to solve problems with units that are completely

unfamiliar to you. Use the relationships listed below to solve the practice problems. Part of the challenge here is to

identify which of the provided conversion factors are relevant and which are not necessary to solve the problem. UNIT

CANCELLATION MUST BE SHOWN TO EARN CREDIT.

Sporks and Gorks Conversion factors:

4 gorks = 9 orks 5.5 x 105 phrods = 4.4 x 10-4 clods

0.3 phrods = 6.02 x 1023 7 sporks = 2 tlorks

4.9 x 1012 kvorks = 7.7 x 10-7 norks 3 sporks = 10 orks

1.2 x 103 kvorks = 420 tlorks 17 clods = 0.07 tlorks

1) 38 gorks = ? sporks

2) 699 tlorks = ? phrods

3) Challenge Problem: Banana slugs are small, slimy creatures that typically inhibit the moist rain forests in the Pacific

Northwest and northern California. Banana slugs are not very fast; they move with an estimated sprint velocity of 6

inches/hour. Imagine that a banana slug is kidnapped by terrorists and taken to the anatomy dissecting classroom at

Fairview. The banana slug, realizing that it is in mortal danger and unwilling to donate its body to science, uses its martial

arts skills to fight its way to freedom. It escapes out the west facing door near the chemistry labs and begins its spring to

freedom, the Oregon border 1000 miles away. Assuming the banana slug continues to operate at full throttle at all

times, is not affected by the altitude, and does not stop for a latte, will it be able to reach home within one year?

Given: 12 in = 1 ft ; 5280 ft = 1 mile ; 24 hr = 1 day; 365 d = 1 yr

6 in/ hr → mi/ yr

Homework 4-4 Significant Figures Problem Set

Summary Sig. Fig. Rules: Count # → infinite, don’t follow rules; Nonzero digits → ALWAYS

Leading Zero’s (0.07 g) → NEVER ; Captive zero’s (708 g) → ALWAYS ;

Trailing zero’s no decimal (400 g) → NO Trailing zero’s with decimal (400. g) → YES

Unit Conversion Factors – match given: 50.0 m (3 SF) → 0.0500 km (3 SF)

1) Determine the number significant figures in each of the following numbers:

1A) 5.432 g _______ 1B) 40.319 g ________ 1C) 3 pencils _______ 1D) 0.189 lb _______

1E) 300 kg _______ 1F) 300. kg ________ 1F) 0.000235 g _______ 1H) 2500.0 cm _______

1I) 0.002300 mg _______ 1J) 3.450 x 103 m _______

Summary Sig. Fig. Rules Calculations: Add/Sub = fewest places past decimal ; Mult/Div = few S.F

2) Perform each of the following calculations. In the first blank write the unrounded calculator answer. On the second

blank, write the final answer rounded to the correct number of significant digits including correct units!

Unrounded (calc) Rounded Units

2a) 3.482 cm + 8.51 cm + 16.324 cm 28.316 28.32 cm

2b) 48.0032 g + 9.17 g + 65.4321 g __________ __________ ____

2c) 80.4 cm – 16.532 cm __________ __________ ____

2d) 106.5 cm – 30. cm __________ __________ ____

2e) 48.2 cm x 1.6 cm x 2.12 cm __________ __________ ____

2f) 64.35 cm3/ 8.149 cm __________ __________ ____

2g) 0.057 mL x (760 mm)(273 K) __________ __________ ____

(740 mm)(250 K)

3A) 3.0 x 108 m/s x (2 x 106 s) = __________ __________ ____

3B) 1.4 x 10-2 m x (3.25 x10-6) = __________ __________ ____

3C) 6.023 x 1014 mm2/ 5.81 x 1012 mm = __________ __________ ____

Homework 5-1 ~ Mole Concept

1) Give 2 example applications of why knowing the # of particles (atoms or molecules) is important in chemistry.

Why can’t we simply count atoms by looking under an electron microscope?

2) Concept of Counting by Grouping – Fill in the following table:

Group Number of objects

1 Pair of Aces 2

2 Pairs of Aces

1 Dozen Doughnuts 12

3 Dozen Doughnuts

Dozen Doughnuts 60

1.5 Dozen Doughnuts

3) Concept of using moles for Counting by Grouping- Fill in the missing information in the following

table:

(Hints: Does the identity of the element change the relationship between number of moles and the number of

atoms? What patterns do you see in the numbers?)

Element Number of Atoms “Groups of 6.022 x 1023” Number of Moles

He 6.022 x 1023 1(6.022 x 1023) 1.000

Li 6.022 x 1023 1(6.022 x 1023)

Ne 1(6.022 x 1023) 1.000

B 3.011 x 1023 = ½( 6.022 x 1023) 0.5000

B 1.204 x 1024 = 2(6.022 x 1023) 2.000

B 1.806 x 1024 3(6.022 x 1023) 3.000

Cu 3.011 x 1023

Cu 1.204 x 1024 =

Cu 1.806 x 1024

Zn ½( 6.022 x 1023)

Zn 2(6.022 x 1023)

Zn 3(6.022 x 1023)

Ag 0.5000

Ag 2.000

Ag 3.000

4) Why do chemists count atoms in groups of moles instead of the actual number of atoms?

Given: 3.011 x 1024 atoms of iron, Fe.

A) Calculate the number of moles of Fe represented by 3.011 x 1024 atoms by dividing 3.011 x 1024 by 6.022 x

1023.

3.011 x 1024 atoms

B) Which is an easier number to work with: 3.011 x 1024 atoms or your answer to 4A (the number of moles)?

C) Why do chemists prefer to represent the number of atoms present in groups of moles instead of using the

actual number of atoms present?

5) Concept of Counting by Weighing

A) Calculating average mass of particles: What is the average mass of the following ten atoms (mass

units are amu):

11 11 11 11 11 11 11 11 10 10

B) A black box contains an unknown of atoms (you cannot see inside the box). If the average

mass of 1 atom is 10.8 amu and the total mass of the contents in the box is 1.08 x 1010 amu, how

many atoms are in the box?

Number of atoms (N) = Total Mass of Atoms ÷ Average Mass of Atoms

6)Foundations of the Mole Concept – How many C atoms are present in 12.01 g of carbon (atomic mass in

grams)?

Average Mass of 1 Carbon atom = 12.01 amu

Individual atoms massed in amu; large # of atoms are massed in grams

Average Mass 1 C atom in grams = 12.01 amu x 1.661 x 10-24 g = 1.995 x 10-23 g

Number of atoms (N) = Total Mass of Atoms ÷ Average Mass of Atoms

N = 12.01 g ÷ 1.995 x 10-23 g/atom = atoms

1 amu

( mole (

atoms

7) What is mass of 1.000 moles (6.022 x 1023) atoms of each of the following elements to the nearest .01 g?

Element Number of moles Number of Atoms Mass in Grams

Helium (#2) 1.000 6.022 x 1023

Lithium (#3) 1.000 6.022 x 1023

Boron (#5) 1.000 6.022 x 1023

Neon (#10) 1.000 6.022 x 1023

Sulfur (#16) 1.000 6.022 x 1023

8) Use the periodic table and the mole concept to identify the elements based on the information given

below:

A) 1 mole (6.022 x 1023 atoms) of this element weighs 26.98 g. What is the element?

B) 1 mole (6.022 x 1023 atoms) of this element weighs 30.97 g. What is the element?

9) You are asked to count out exactly 6.022 x 1023 atoms of Be (element #4). Describe how you would

obtain the correct number atoms since the individual atoms are too small to be seen and too numerous to

be counted out individually.

10) Connecting Concepts of Counting by Weighing and Counting by Grouping

In the pictures below each atom is represented by and has a mass of 10 units.

A) A box of atoms contains 12 atoms. What is the total mass of 1 box?

(assume mass of box can ignored, just count atoms)

1 box

B) The total mass of a set of boxes weighs 600 mass units.

How many boxes are present?

How many atoms are present?

C) The total mass of a set of boxes weighs 1260 mass units.

How many boxes are present?

How many atoms are present?

11) 1 mole (6.022 x 1023) of carbon atoms weighs 12 g. However, in the lab you are not restricted to working with only

1 mole. You can work with or more less than 1 mole. For each of the following what would be the mass in grams for

each quantity?

1 mole 0.25 moles 0.33 moles 0.50 moles 0.75 moles

12 g g g g g

2 moles 2.5 moles 3 moles

g g g

12) Complete the following table:

moles Number of atoms Element Mass in grams

1.000 6.022 x 1023 He

0.5000 3.011 x 1023 He

2.000 1.204 x 1024 He

3.000 1.806 x 1024 He

1.0000 Ne

0.5000 Ne

2.000 Ne

3.000 Ne

3.011 x 1023 20.04

2.0000 48.62

13) Which of the following contains the same number of atoms as 12.01 g of carbon? Explain.

A) 13.88 g Li B) 12.0 g He C) 5.4 g of B D) 20.18 g Ne

Explanation:

Homework 5-2- Molar Mass of compounds/molecules 1) Molar Mass of molecules:

Helium is a monoatomic element: 6.022 x 1023 atoms (1 mole) weighs _____ g.

The mass of 6.022 x 1023 (1 mole) of Bromine atoms weighs 79.90 g.

Elemental Bromine is a diatomic molecule with a formula of Br2.

1 mole of 6.022 x 1023 Bromine molecules (Br2) weighs 159.80 g. Where does the number 159.80 g

come from?

2) The molar mass of a substance is calculated by __________________ the atomic weights of the

component atoms.

3 ) If 1 mole of H = 1.01 g and 1 mole O weighs 16.00 g ;

How much does 1 mole of H2O weigh?

4) Complete the following table describing water molecules.

# of H2O molecules Total # H atoms Total #O atoms

1 2 1

2

100

6.022 x 1023

1 mole moles mole

Since 6.022 x 1023 molecules = 1 mole, subscripts for each element in a formula also represent the _______ of

that element present in 1 mole of the compound.

5) Complete the following table

moles Number of molecules Formula Mass in grams

6.022 x 1023 H2 2.02

6.022 x 1023 N2

1 O2

1 CO2

6) Calculate the molar mass to 2 places past the decimal for each of the following compounds:

A) NH4Br B) Mg(NO2)2 C) Ca3(PO4)2

N: Mg: Ca:

H: N: P:

Br: O: O:

7)There are 3 different known gaseous compounds containing nitrogen, N and oxygen, O. N has an atomic mass of 14.0

g/mole; O atomic mass = 16.0 g/mole. (All molar masses are units of grams/mole)

Compound X: Molar Mass = 30.0 g/m Compound Y: Molar Mass = 44.0 g/m

Compound Z: Molar Mass = 46.0 g/mole

Use the information above to determine the chemical formula of each compound. Clearly explain your reasoning.

Homework 5-3 – Mole Bridge Conversions Use the mole bridge and dimensional analysis to solve the following. Round the final answer to match same

number of SF as the original given.

Example: What is the mass in grams of 9.34 moles of iron, Fe?

( ) = g Fe (3 SF)

Part 1: moles → grams

1) What is the mass in grams of 0.521 moles of Ni (atomic #28)?

( )

2) What is mass in grams of 2.49 moles of SO3?

Part 2: grams → moles

3) 26.2 g Au (atomic #79) → moles?

4) What is the number of moles in 1.271 grams of ethanol, C2H5OH?

Part 3: moles → atoms or molecules

5) 0.0015 moles of Cu (atomic #29) → ? # of atoms

6) How many molecules are present in 1.25 x 10-2 moles of lead acetate, Pb(CH3CO2)2 ?

Ans: 1) 30.6 g 2) 199 g 3) 0.133 mol 4) 0.02759 moles 5) 9.0 x 1020 6) 7.53 x 1021

Part 4: atoms or molecules → moles

7) What is the number of moles in 3.312 x 1024 atoms Se?

# particles

(atoms/molecules MOLES Mass in grams 6.022 x 1023 Gram from PT

0.521 mole Ni =

9.34 mol Fe

1 mole Fe

55.8 g Fe

mole Fe

Part 5: grams → moles → atoms or molecules

8) How many atoms are present in 104.49 g of Bi (atomic #83)?

9) How many molecules of CO are present in 6.37 g of carbon monoxide, CO?

Part 6: atoms or molecules → moles → grams

10) What is the mass in grams of 1.804 x 1024 atoms of Zr (atomic #40)

11) What is the mass in grams of 4.03 x 1024 molecules of benzene, C6H6?

Part 7: # of molecules → # atoms

12A) How many C atoms and H atoms are present in 1 molecule of methane, CH4?

12B) How many C atoms and H atoms are present in 10 molecules of methane, CH4?

12C) How many C atoms and H atoms are present in 1000 molecules of methane, CH4?

12D) How many C atoms and H atoms are present in 6.022 x 1023 molecules of methane, CH4?

12E) How many C atoms and H atoms are present in 7.089 x 1024 molecules of methane, CH4?

Ans: 7) 5.500 moles 8) 3.1100 x 1023 atoms 9) 1.37 x 1023 molecules 10) 273.3 g 11) 523 g

12A) C:1 H:4 12B) C:10 H:40 12C) C:1000 H:4000

12D) C: 6.022 x 1023 H:4 (6.022 x 1023) = 2.409 x 1024

12E) C: 7.089 x 1024 H:4 (7.089 x 1024) = 2.836 x 1025

Homework 5-4 - % Composition by Mass and Empirical Formula

1) Homework Self-Check 8.7, p. 220 – Antibiotic Penicillin has a formula of C14H20N2SO4. Determine mass % of each

element.

2) Self-check 8.8 from HW on p. 225. It is observed that 0.6884 g of lead (Pb, atomic #82) combines with 0.2356 g of

chlorine (Cl, atomic #17) to form a compound. Calculate the empirical formula of the compound.

3) Self check 8.9 from HW p.226: Sevin, a commercial insecticide used to protect cotton, vegetable and fruit crops from

insect damage is made from carbamic acid. Elemental analysis of carbamic acid:

0.8007 g of C, 0.9333 g of N, 0.2016 g H, and 2.133 g of O. Determine the empirical formula:

4) Self check 8.10 from HW ,p.227: The most common form of nylon is 63.68% C, 12.38 % N, 9.80% H and 14.14% O.

Calculate the empirical formula of Nylon.

ANS: 1) C: 53.81%, H: 6.453% , N: 8.969%, S: 10.27%, O:20.49% 2) PbCl2 3) CNH3O2 4) C6NH11O

Homework 5-5 - Molecular Formula

1) p. 234, #78: A compound with empirical formula of CH was found to have a molar mass of 78 g/mole. Determine the

molecular formula.

2) p. 234, #80: A compound with empirical formula C2H5O was found in a separate experiment to have a molar mass of

90 g/mole. Determine the molecular formula.

3) p.234, #81: Elemental analysis: C: 42.87% H: 3.598 % O: 28.55% N: 25.00% Determine the empirical formula.

In a separate experiment the molar mass was determined to be 168.13 g/mole. Determine the molecular formula:

4) Self Check 8.11 p. 228: A gasoline additive has the following elemental analysis:

71.65% Cl 24.27% C 4.07%H Determine the empirical formula.

In a separate experiment, the molar mass was determined to be 98.96 g/mole. Determine the molecular formula.

Answers: 1) C6H6 2) C4H10O2 3) E.F = C2H2ON ; M.F. = C6H6O3N3 4) E.F. = CH2Cl; M.F. = C2H4Cl2

Unit 4 Review

Scientific Notation

1) For A and B, rewrite each scientific notion number in standard form. For C and D express each number in scientific

notation:

1A) 1.995 x 102 1B) 1.995 x 10-4

1C) 0.0005219 1D) 93,000,000

Dimensional Analysis- perform the following conversions, clearly showing unit cancellation. Match SF of final answer to

SF of given.

1) Years ago in England, land was measured in units such as fardells, nooks, yards and kides. Given the following: 2

fardells = 1 nooke 4 nookes = 1 yard 4 yards = 1 kide

How many kides are there in 50.0 fardells?

2) 250. mL = ? L 4) 3.00 x 108 m/s = ? km/hr

Recording Significant Figures from Measuring Device

5) Report the volume measure to the correct # of SF (units are mL).

6) Report the length measurement to the correct # of SF (units = cm).

7) Significant Figure Rules Summary – Fill in the blank describing the SF rules

Answer Choices: ALWAYS ; NEVER; SOMETIMES; INFINITE ; YES: SIGNIFICANT ; NO: NOT SIGNIFICANT

COUNTING #’S and Conversion factors – NONZERO DIGIT’S: ZERO’S:

LEADING :

CAPTIVE:

TRAILING :

Decimal present?

Significant Figures Rules in Measurements:

8) How many significant figures are present in each of the following?

A) 5 pencils B) 234 cm C) 702 g D) 0.020 g

E) 300 cars F) 300 mL G) 3.0 x 102 mL

H) 300. mL I) 0.01010 m J) 120.0 mL

9) Perform the following calculations: Report your final answers to the correct number of significant figures:

A) 323.74 mL B) 603.00 cm C) 490 cm3 D) 97.581 g

- 20.0 mL x 50.0 cm ÷ 7.00 cm + 2.22 g

Unit 5 Review Problems

1) 63.5 g of copper contains approximately the same number of atoms as:

A) 13.8 g Li B) 4.5 g Be C) 24.0 g C D) 12.0 g He E) 39.9 g Ar

2) The subscripts in a chemical formula represent the number of atoms present. Another way to view the subscripts

is that they are numbers of moles of the atoms.

In 1 molecule of water there are _____ atoms of H and ______atom of O.

In 1 mole of water there are ____ moles of H atoms and ______ mole of O atoms.

3) What is the mass of 5.00 moles of Ca3(PO4)2?

4) How many moles are present in 14.9 g of KCl?

.

5) Dimethlynitrosamine (CH3)2N2O is a carcinogenic (cancer-causing) substance that may be formed in foods,

beverages, or gastric juices from the reaction of nitrite ion (used as a food preservative) with other substances.

A) Calculate the molar mass of dimethylnitrosamine, (CH3)2N2O.

B) What is the mass (in grams) of 1.00 x 106 molecules of (CH3)2N2O?

C) How many molecules of (CH3)2N2O are present in 2.50 g of dimethylnitrosamine?

D) How many N atoms are present in 2.50 g of (CH3)2N2O?

E) How many H atoms are present in 2.50 g of (CH3)2N2O?

Answers: A) 74.6 g/mole B) 1.23 x 10-16 g C) 2.03 x 1022 molecules (CH3)2N2O

D) 4.06 x 1022 N atoms E) 1.21 x 1023 molecules (CH3)2N2O

You must show work to earn credit. No work = No credit

6) Calculate the % by mass of each element in C5H10O.

7) A compound contains only the elements N and O. A total sample mass of 23.0 g contains 7.00 g of N. Calculate

the % by mass of each N and O.

8) A compound contains the following % by mass: Ca: 28.03% oxygen: 22.38% chlorine: 49.59% Determine

the empirical formula of the compound.

9) A compound has the following elemental analysis: Cu: 66.75% P: 10.84% O: 22.41%. Determine the empirical

formula of the compound.

10) A compound has a molar mass of 70 g/mole. The simplest (empirical) formula of the compound in CH2. What is

its molecular formula?

11A) A compound is analyzed and found to contain 43.64% phosphorous and 56.36% oxygen by mass. Calculate the

empirical formula of the compound.

11B) The compound from 11A has a molar mass of 283.88 g/mole. What is compound’s molecular formula?

12)A compound contains 47.08% carbon, 6.59% hydrogen, and 46.33% chlorine by mass; the molar mass of the

compound is 153 g/mole. What are the empirical and molecular formulas of the compound?

Answers to selected problems: You must show work to earn credit. No work = no credit!

9) Cu3PO4 10) C5H10 11A) P2O5 11B) P4O10 12) Empirical formula = C3H5Cl Molecular Formula: C6H10Cl2

Unit 6 ~ Solution Chemistry ~Test Topics Density

Memorize density equation: D = m/V and be able to solve for any variable from the equation given the other two.

(Example: Solve for D given m and V, solve for m given D and V, etc.)

- Understand that density is a physical property of a substance which may be used to identify a substance or make

predictions of behavior (e.g help predict if a solid substance will float or sink in a particular liquid).

- Graphical analysis: Determine the density of if a substance from the slope of a graph of Mass (Y-axis) and

Volume (X-axis).

Solutions Summary:

Solute – component present in smaller amount

Solvent – component present in larger amount (typically water)

Solution – homogenous mixture of solute and solvent

Equations:

Molarity = Moles of Solute

Dilution Equation: MCVC= MDVD

Where MC = Molarity of Concentrated Solution VC = Volume of Concentrated Solution

MD = Molarity of Dilute Solution VD= Volume of Dilute Solution Note: MC > MD

Quick Guide to Common Problems:

1) Calculating Molarity given grams solute and mL solution:

Convert g → moles and mL → L (Recall 1000 mL = 1 L); divide moles/L

2) Recipe for Preparing Molar Solution from a solid:

1) Calculate moles using moles = Molarity x Volume in Liters

2) Calculate grams by multiplying molar mass (PT) x moles (calculated in part 1)

3) Directions:

Dissolve ___ g of ____ in a small volume of water a _____ mL volumetric flask; Add water up to

mark.

3)Preparing a Molar Solution by Dilution

1) Use dilution equation MCVC= MDVD to find volume of concentrated solution needed, VC .

2) Directions: Dilute ____ mL (VC) of ______ M stock solution with a enough water to prepare

____ mL of solution.

Liters of solution

Density Notes

Introduction to Concept • Which is heavier, 1 lb of lead or 1 lb of feathers?

• Which occupies more volume, 1 lb of lead or 1 lb of feathers?

• Why is the answer to the first question the different from the answer to the second question?

Density Equation

Units:

Does size matter?

The density of the aluminum cylinder on the left is 2.7 g/mL. What would be the density of two aluminum cylinders?

m = 2.7 g v = 1.0 mL m = v=

A) 1.4 g/ mL B) 2.7 g/mL C) 5.4 g/mL D) Can’t determine

Calculating density using correct significant figures

Calculate the density of a piece of wood that has a mass of 5.0 g and a volume of 210 mL to the appropriate

number of SF.

Solution Chemistry ~ Notes ☺

Definition of a Solution

A solution is a ___________________________________________________________.

Solution

Salt Water dissolved in water

Soda Water dissolved in water

Relationships

o Solute: component present in ______________ amount

o Solvent: component present in ______________ amount

The solute is dissolves in the solvent to form the solution.

SOLUTION = ____________________ + ______________________

Important note: Unless otherwise specified the solvent in our course is always ________.

Importance of Solutions:

• In nature, it is very rare to find pure solids, liquids or gases; most substances are mixtures.

• Most elements and compounds of interest are __________ at RT; however rxn rate between solids is very

_________ due to low _________ _____ that limit contacts between reacting particles.

• In the lab, most chemical reactions are carried out by dissolving __________ in ____________ solvents to

increase # of _______________ and hence rxn rates.

**********************************************************************************************

Molarity ~ Notes

• For solutions, chemists use the idea of ______________________ to describe the amount of solute in a given amount

of solution.

• Concentrations are always RATIOS:

• Compare the concentrations of the following three containers (dots represent solute):

100 mL 200 mL 400 mL

• “Molarity” is a common method of expressing the concentration of a solute dissolved in a solution. Molarity,

abbreviated by M, is defined as:

Question: What volume of water would be required to produce 1 L of salt water?

A) Less than 1 L B) exactly 1 L C) More than 1 L

Explanation:

• Sample Molarity Calculation: What is the molarity (M) of a 250. mL solution containing 9.45 g of CsBr?

Write definition of Molarity: M = moles solute

L solution

Convert to appropriate units:

Solve for M:

Fill in the missing information:

Grams of KBr Moles of KBr Volume of Solution Molarity of Solution

119 g 1.00 moles 1000. mL 1.00 M

0.500 moles 500. mL

238 g 2000. mL

Notes on Preparation of Molar Solutions

Method 1: Preparation a Desired Volume of a Molar Solution from a Solid

Calculate moles of solute needed.

Calculate grams of solute needed.

Weigh solute with scale, and transfer to volumetric flask calibrated to prepare desired volume.

Rinse container with water to make sure that solute is transferred.

Dissolve solute in small volume of water.

Fill up to mark with water.

Example: Give detailed directions for the preparation of 250. mL of a 0.100 M solution of CuCl2 from solid CuCl2.

Molarity = moles of solute

Step 1: Find moles using equation: moles =

Recall 1000 mL = 1 L

Step 2: Calculate grams using moles from step 1 and molar mass.

Step 3: Directions:

Dissolve _____ g of _______ in enough water to prepare _______ mL of solution.

Liters of solution

Method 2: Preparation of a Molar Solution by Dilution a Concentrated Stock Solution.

Concept of Dilution: Same # of moles dissolved in a larger volume

Dilution Equation: MCVC= MDVD

Where MC = Molarity of Concentrated Solution VC = Volume of Concentrated Solution

MD = Molarity of Dilute Solution VD= Volume of Dilute Solution

Note: MC > MD

Example Problem: Describe how you would prepare 250. mL of a 0.100 M CuCl2 solution, starting with a

0.500 M CuCl2 stock solution.

MC = VC = MD = VD=

Starting with Dilution Equation: MCVC= MDVD

Solve for VC

VC =

Recipe: Dilute ____ mL of the _____ M stock solution with enough water to prepare _______ mL of solution.

___ mL of ___ M stock solution

ADD

WATER

Unit 6 ~ Solutions ~ Homework Assignments

6-1 Section 2.8 (p. 41-44) Density Problem Set (workbook)

6-2 Sec. 15.1, 15.2 (p. 451 – 455) Page 476 #32, 34ac, 36ac, p.480,# 115,119

6-3 Section 15.4 (p. 457 – 462) Molar Solutions Problem Set

6-4 Section 15.5 (p. 462 – 465) Page 480, #124, 125

6-5 Unit 6 Review Problem Set

HW 6-2 p. 476, #32 To prepare 500. mL of 1.02 M sugar solution, which of the following would you need?

A) 500. mL of water and 1.02 mol of sugar

B) 1.02 mol of sugar and enough water to make a total volume of 500. mL

C) 500. g of water and 1.02 mol of sugar.

D) 0.51 mol of sugar and enough water to make the total volume 500. mL

p. 476, #34- Calculate the molarity from mol solute and volume of solution.

34a) 0.50 mol KBr; 250 mL 34c) 0.50 mol KBr; 750 mL

p. 480, #115 Concentrated hydrochloric acid is made by pumping hydrogen chloride gas into distilled water. If

concentrated HCl contains 439 g of HCl per liter, what is the molarity?

p. 476, #36 – Calculate molarity from moles of solute and volume of solution

36a) 1.25 g KNO3; 115 mL 36b) 1.25 mg KNO3; 1.15 mL

p.480, #119 – If 10. g of AgNO3 is available, what volume of 0.25 M AgNO3 can be prepared?

HW 6-4 p.480 #124 When 50. mL of 5.4 M NaCl is diluted to a final volume of 300. mL, what is the concentration of

the diluted solution?

p. 480, #125 When 10. L of water is added to 3.0 L of 6.0 M H2SO4 what is the molarity of the resulting solution? Assume

the volumes are additive.

Homework 6-1 Density Problems

1) Determine the density of Sponge Bob Square Pants to the correct of SF using the volume displacement method

screen reads: 7.50 g

2) A graph of mass vs volume data is given below for a gummy bear.

A) Calculate the slope of the line from the data.

B) What is the density of the gummy bear? Explain your answer (Hint: What is the formula for slope? What is the

formula of density? What values are on each axis?)

m=

Vfinal =

-Vinitial =

∆V =

Al Zn Fe Cu Pb

3) All of the cylinders above have the same mass. What do you notice about the relationship between density and volume

of different metals with the same mass?

Same mass, smaller volume → (higher or lower) density

4) Eureka Problem- Archimedes and the Gold Crown

Archimedes was a famous Greek scientist who lived on the island of what is now Sicily, Italy in the 3rd century BC. According to

legend, Archimedes was asked by the ruler of the city to determine whether or not the king had been cheated by his goldsmith. The

goldsmith had been given a particular mass of gold and asked to make a crown for the king. The king suspected that the goldsmith

had replaced some of the gold with silver and kept the extra gold for himself. The king asked Archimedes to determine whether the

crown was in fact pure gold without ruining the crown by cutting it open. The key to solving the mystery is the fact that gold and

silver have different densities. According to contemporary accounts, Archimedes was sitting in the bath thinking about the problem

when he realized when he was submerged that his body displaced a volume water equal to the volume of his body; he then leaped out

of the bath and ran naked through the streets shouting “Eureka (I have found it!)

Calculate the volume occupied by 500. g (3 SF) each metal (gold and silver). Will 500. g of pure gold and pure

silver displace the same volume of water when placed in a bath?

Given: Density of Gold = 19.3 g/mL ; Density of Silver = 10. 5 g/mL

5a) Density of liquid water vs ice: Use the diagram of liquid water/ice to explain the following observation: Liquid

water has a density of 1.0 g/ mL while solid water (ice) has a density of 0.91 g/mL.

5b) Overheard comment on a group hike: “ The best way to have cold water on a hike on a hot day is to freeze the water in your

bottle the night before. The ice will then slowly thaw to cold water during the hike. The only problem is that freezing the water

makes it heavier to carry.” Does freezing water increase its mass? Explain.

5c) The real concern about freezing water bottles is that filling the bottle completely to the top, sealing the bottle and then

freezing it can cause problems. What could happen and why?

6a)Diet soda and regular soda have very different densities. Regular soda sinks in water while diet soda floats. If the

density of water is 1.0 g/mL what can you say about the density of regular vs diet soda?

Density of soda > density of soda

6b) Regular soda gets its sweet taste from sugar, while diet sodas use artificial sweeteners like Nutrasweet. The difference

in density is primarily due to the fact that in the concentration (# of molecules per can volume) of sugar molecules vs

artificial sweetener molecules required to produce the desired level of sweet taste. Which picture could represent sugar

dissolved in water in regular soda vs. Nutrasweet dissolved in water in diet?

7) Given the following data:

7a) Which phases tends to have the lowest densities and which phase tends to have the highest densities? Highest

density phase: Lowest density phase:

7b) Propose an atomic level observation to account for the differences in density in different phases (support answer with

atomic level pictures). = atoms

Higher density = Lower Density =

Gases: He, Air,

Liquids: gasoline, kerosene, benzene, water, carbon

tetrachloride

Solids: Magnesium, salt, aluminum, iron, copper, silver, lead,

uranium, gold

Homework 6-3 ~ Preparation of Molar Solutions Problem Set Part 1: Preparation of a Molar solution from a solid solute.

For each aqueous solution, calculate the grams of solute required and provide detailed directions for the preparation of the

solution.

1) 250. mL of 0.300 M KMnO4 from solid KMnO4.

Calculation:

Detailed Directions: Dissolve _____ g of _______ in enough water to prepare _______ mL of solution. Diagram steps

below:

2) 500. mL of 0.250 M KI from solid KI.

Calculation:


below:

3) 2.00 L of 4.00 M C6H12O6 from solid C6H12O6.

Calculation:


below:

Part 2: Preparation of a Molar Solution by Dilution of a Concentrated Stock Solution.

Calculate the volume of concentrated stock solution needed to prepared the desired concentration and provide detailed

directions from the preparation of the solution.

4) 500. mL of 6.00 M from 12.4 M (concentrated) HCl stock solution.

Calculations:



5) 100. mL of 1.00 M H2SO4 from 18.0 M (concentrated) H2SO4.

Calculations:



6) 25.0 mL of 3.00 M HNO3 from 15.0 M HNO3.

Calculations:



HW 6-5 Unit 6 Review ~ Density and Solutions

1) Use the graph below to determine the density of hydrogen peroxide.

Molarity and Dilution

2) In an aqueous solution of hydrochloric acid, HCl, identify the solute, solvent and solution.

3) How many moles would be present in 400. mL of 2.00 M solution?

4) What volume of a 2.00 M KBr solution would contain 11.9 g of KBr?

5) A solution contains 3.5 moles dissolved in 100. mL. How many moles would be present if the solution were

diluted to a new total volume of 1000. mL? What would the new concentration be after dilution?

6) If you wish to prepare 1.0 L of a 3.00 M solution of aqueous KBr solution, which of the following statements

about the volume of water needed to prepare the solution are correct?

A) Less than 1.0 L of water would be needed

B) Exactly 1.0 L of water would be needed

C) More than 1.0 L of water would be needed

7) Application of Preparing Molar solutions concepts~ Some experiments require solutions with extremely low

concentrations. These are extremely difficult to prepare directly because of the very small number of grams

present. Example: What is the best method to prepare 10.0 mL of a 0.00500 M NaCl solution?

A) Approach #1: Calculate the grams of solute required and provide detailed directions (next page) for the

preparation of 10.0 mL of a 0.00500 M of NaCl solution from solid NaCl .

Dissolve _____ g of _______ in enough water to prepare _______ mL of solution.

A) Our scale reads to the nearest 0.01 g. Based the grams calculated in part a, why would be it in be difficult in

practice to make this solution directly from the solid?

B) Approach #2: An alternative strategy would be to prepare 10.0 mL of 0. 500 M NaCl which would require 0.29

g of NaCl – an amount which is measureable on our scale. We could then prepare our desired solution by diluting

our stock solution. Calculate the volume of 0.500 M NaCl stock solution needed and provide detailed

directions for the preparation of the 10.0 mL of 0.00500 M NaCl solution.



8) Concept of Serial Dilution – sometimes the concentrations needed are so small that dilutions need to be done in a

series.

Solution 1

A) First Dilution: What would be the concentration of solution of obtained by diluting 1 mL of 1 M NaCl

stock solution into a total solution volume of 10. mL? Fill in your answer in the box above solution 2.

B) Second Dilution: If we repeated our experiment by taking 1.0 mL of the concentration you prepared in

part A (Solution 2 is new stock for 2nd dilution) into a total solution volume of 10. mL, what would the

final concentration of the new solution be? Fill in final answer in the box above solution 3.

C) Continue the process to prepare solutions 4 and 5 and fill in answers in boxes provided.

1 M Stock

Solution

M M M M

2 3

4 5

Chemistry Final Exam Study Guide ~ Semester 1

General Information

• Final Exam will be worth 200 pts and cover all topics from semester 1 with equal emphasis for each

unit. Exam is essentially a giant retest of core ideas from term.

• Format: multiple choice and Free Response

• Exam is CLOSED BOOK, CLOSED NOTE, CLOSED FRIEND!

• Provided materials: Periodic table, activity series chart, solubility rules chart, polyatomic ions list (minus

7 you had to learn) + calculator (you may NOT use your own calculator).

Highlighted topics:

Unit 1 ~ Atomic Theory

Topics: Scientific Method, chemical and physical changes, Atomic level interpretation of behavior of solids,

liquids and gases, elements, compounds and mixtures, Dalton’s Atomic Theory, Thomson and Rutherford

Experiments, Protons, Neutrons and Electrons, Ions and Isotopes

Unit 2 ~ Nomenclature

Topics: Organization of Periodic Table, Ionic and covalent bonding, chemical formulas and # of atoms in

formula, Class 1,2,3 naming rules and writing formulas

Unit 3 ~ Reactions

Topics: Components of a Reaction: coefficients, subscripts phases; Balancing; 6 Reaction Patterns: synthesis,

decomposition, combustion of a hydrocarbon, single replacement (including using activity series charts), double

replacement (including using solubility rules charts).

Unit 4 ~ Math Foundations

Topics~ Unit Conversions, Scientific Notation, Recording measurements with correct significant figures (SF)

and units; SF rules for recorded measurement; SF rules in calculations; Density

Unit 5 ~ Mole Concept

Topics: Definition of a Mole/ Mole Concept, Molar Masses, Mole Bridge (calculations from particles to moles

to grams, etc). % by mass calculations, Empirical Formula, Molecular Formula

Unit 6 ~ Solutions

Topics: density, concept of concentration, Molarity and dilution; Preparation of Molar Solutions from solid and

by dilution of a concentrated stock solution.

Required Practice Problems for First Semester Final Exam

A small, but useful set of sample problems. This does NOT cover every possible exam topic. Look back and rework the

unit review problem sets from the semester. USE YOUR OWN PAPER. NOT ENOUGH SPACE PROVIDED.

Unit 1~ Atomic Theory

1) Rutherford’s experiment invalidated Thomson’s Plum Pudding model of the atom.

a) In Rutherford’s atom, where is most of the mass located?

b) What two subatomic particles are located inside the nucleus?

c) Why did most of the positive alpha particles pass straight through the gold atoms in the foil?

2) How many protons, neutrons and electrons are present in the following ions?

25 +2 35 -1

Mg e-: Cl e-:

12 n: n:

p: p:

Unit 2: Nomenclature

3a) How many valence electrons does Na have? What charge does Na ion typically have in a compound?

3b) If neutral Na transfer an electron to neutral Cl, what type of bond has been formed?

3c) If two nonmetals such as C and O form a bond electrons are shared. What is this bond type called?

4) Name the following compounds:

a) MgCl2 b) CuCl2 c) CO2

5) Write formulas for the following compounds:

a) zinc chloride b) iron(II) oxide c) iron(III) oxide d) sulfur trioxide e) chromium(III) phosphate

Unit 3 ~ Reactions: State the type of reactions, predict the products, and balance.

6a) Al + Br2 → 6b) CH4 + O2 → 6c) HgO →

6d) Mg + HCl → 6e) Cu(NO3)2 + Na3PO4 →

Acid-Base Rxns: 6f) NaOH + HCl → 6g) HNO3 + NH3

Unit 4~ Math

7) Prefixes: How many mL are in 1 L?

8) Use the factor label method to convert the speed of light 3.0 x108 m/s into mi/hr.

9) What is the sum of 20.23 g + 1.1g calculated to the correct number of significant figures?

10) What is the product of 30.0 m x 70 m calculated to the correct number of significant figures?

Mole ~ Unit 5

11) Calculate the number of atoms in 18.02 g of Be:

12) What is the molar mass of Cu(NO3)2?

13) If cobalt heated in the presence of sulfur, a compound is produced that contains 55.06 % Co by mass. Calculate the

empirical formula of the compound.

14) A compound has a molar mass of 70 g/mole and an empirical formula of CH2. What is the molecular formula of the

compound?

Extra Credit Opportunity for Fall Final

For each of the Units, you may earn extra credit points by doing the following:

• WRITE OR TYPE A 1 PAGE SUMMARY OF KEY IDEAS OF UNIT → 1 pt./Unit

Include examples of solved problems as appropriate; No credit for handing in the notes/materials already provided

in this course guide. Must be your own summary of key ideas. INDIVIDUAL WORK MAY NOT USE A

GOOGLE DOC.

• SOLVE ALL OF THE PROBLEMS FOR UNIT GIVEN BELOW → 2 pts.

Honest effort, work shown and all problems from a given unit completed to earn credit.

Units 1-5 (Unit 6 not included since we just finished this unit): Total pts possible = 15 (5 summaries + 10 problems).

Do not have to complete all units to earn points, however each subsection of work must be complete. For

example, 1 pt earned for summarizing Unit 1, 2 points earned completing problems in Unit 2 = 3 total point.

Fractions of completed summaries or problems will not be counted.

Problems posted on website.

EXTRA CREDIT DUE AT TIME OF FINAL EXAM

ALL WRITTEN WORK SUBMITTED IS INDIVIDUAL - you may discuss

problems with a friend, but final work submitted is your own. Each student must

have made an honest individual attempt on all problems.

Chemistry Course Book FALL 2017 - Fairview High School

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