Matter and Change Chapter 1. Matter pure substances – compounds and elements anything that has mass and takes up space mixtures – two or more pure substances.

Matter and Change

Chapter 1

Matter

pure substances – compounds and elements

anything that has mass and takes up space

mixtures – two or more pure substances mixed in the same container…not bonded

Solid, Liquid, or Gas?

• Solid

• Liquid

• Gas

SHAPE

• definite shape

• indefinite shape

• indefinite shape

VOLUME

• definite volume

• definite volume

• indefinite volume

Physical Properties• physical property – a characteristic that

can be observed or measured without changing the identity of the substance

melting pointstate of matter

masscolor

Physical Changes• physical change – a change in a

substance that does not involve a change in the identity of the substance

cutting breaking

tearingdissolving

boiling

All phase changes are physical changes.

Chemical Properties• chemical property – a characteristic that

can be observed or measured with a change in the identity of the substance

flammability

reacts with oxygen

reacts with an acid

Chemical Changes• chemical change – a change in a

substance that does involve a change in the identity of the substance

color change

…four good indications of a chemical change.

gas released (often with an odor)

energy change (light, heat, …) precipitate

formed

Precipitate

• precipitate – a solid formed from two aqueous solutions during a chemical reaction

aqueous – dissolved in water

Homogeneous Mixturesthe same throughout…each

sample contains the same ratio of ingredients

Heterogeneous Mixtures

different throughout…each sample contains a different

ratio of ingredients

Separation of MixturesPHYSICAL changes only

by hand

density

evaporation

magnetism

filtration chromatography

Separation of elemental Fe

This doesn’t work with iron that is bound into a compound, only with elemental Fe.

magnetism

Distillation Apparatus

Homogeneous Mixture, Heterogeneous Mixture or Pure Substance?

Observations vs. Interpretations

• observations – the facts…what you see or measureExample: The solution turned cloudy white.

The test tube felt warm to the touch.

• interpretations – your opinion of what you seeExample: From the previous observations, a

chemical change has occurred.

Qualitative vs. QuantitativeObservations

• Qualitative observations – ones that do not involve numbers

• Quantitative observations – ones that DO use numbers (quantity)

Qualitative vs. QuantitativeWhich one is it?

1. A white precipitate was formed.

2. The reaction produced 40.0 g of water.

3. The test tube felt warm.

4. The temperature rose 25 degrees.

1. qualitative

2. quantitative

3. qualitative

4. quantitative

Extensive vs. Intensive Properties

• extensive – DOES depend on the amount of matter presentExample: mass, volume, amount of energy

within a substance

• intensive – does NOT depend on the amount of matter presentExamples: color, melting point, density, luster

If mass and volume are extensive properties, why is density an intensive property?

Measurements and CalculationsPages 28-65

Chapter 2

Prefix Symbol Exponential

tera T 1012

giga G 109

mega M 106

kilo k 103

hecto h 102

deka D or da 101

----- ----- 100

deci d 10-1

centi c 10-2

milli m 10-3

micro 10-6

nano n 10-9

pico p 10-12

Quantity QuantitySymbol

Unit Name Unitabbreviation

Length l meter m

Mass m kilogram kg

Time t second s

Temperature T kelvin K

Amount of a Substance n mole mol

Electric Current I ampere A

Luminous Intensity Iv candela cd

Size of Units

1L=1dm3

1cm3 = 1mL

Density

MD =

V

Density

Mass

Volume

Density is an INTENSIVE property… it does NOT depend on the amount of matter you have.

1. What is the density of a block of marble that has the dimensions 5.00cm length, 4.00cm width, and 15.5cm

height and has a mass of 853 g?

3

MD =

V853 g

D = 5.00 cm x 4.00 cm x 15.5 cm

gD = 2.75

cm

2. Diamond has a density of 3.26 g/cm3. What is the mass of a diamond that has a volume of 0.351 cm3?

3 3

3

MD =

Vg M

3.26 = cm 0.351 cm

g3.26

cm

30.351 cm

M

M = 1.14 g

3. What is the volume of a sample of liquid mercury that has a mass of 76.2 g, given that the density of mercury is

13.6 g/mL?

MD =

Vg 76.2 g

13.6 = mL V

76.2 gV =

g13.6

mLV = 5.60 mL

water

1.00 g gDensity = = 1.00

1.00 mL mL

Significant Figures

• digit or figure – 0,1,2,3,4,5,6,7,8,9

• significant digit or figure – a digit that helps you to understand the details of the entire number given to you

Significant Figures

5005 m

0.0045 L

4.500 g

100 cm3

100. cm3

0.04020 g

Scientific Notation• Move the decimal until one non-zero digit

appears to the left of the decimal.

• Be sure the power of the ten reflects the direction of that move.

• Keep the same number of significant figures in the scientific notation as your original.

50050 m

0.00450 L

= 5.005 x 104 m

= 4.50 x 10-3 L

Accuracy(getting all

measurements right)

Precision

(getting all measurements

the same)

Percent Error

What is the percent error for a mass measurement of 17.7 g, given that the correct value is 21.2 g?

17.7 g - 21.2 g %E = x 100%

21.2 g

experimental accepted

accepted

Value - Value Percentage error = x 100%

Value

%E = - 17%

In order to prepare for the chemistry exam on Wednesday, you should study for at least 1.0 hour the night before the exam. How many milliseconds is this?

3.6 x 106 ms

The summer Olympics showed us that Usain Bolt could run 10.36 meters per second. How fast is this in miles per hour?

(3.281 ft = 1.000 meters)

23.18 miles/hr

The volume of the Powerade® that I drank yesterday was 946 mL. What is that volume in dm3?

.946 dm3 or 9.46 x 10-1 dm3

Chapter 3: Atoms

The Building Blocks of Matter

An atom is the smallest particle of an element that retains the chemical properties of that

element.

The Early Atom

• As early as 400 B.C., Democritus called nature’s basic particle the “atomon” based on the Greek word meaning “indivisible”.

• Aristotle succeeded Democritus and did not believe in atoms. Instead, he thought that all matter was continuous. It was his theory that was accepted for the next 2000 years. (Read page 43 of your textbook.)

Basic Laws of Matter

• Law of Conservation of Mass- mass is neither created nor destroyed during ordinary chemical reactions or physical changes.

CH4 + 2O2 → 2H2O + CO2

16g + 64g → 36g + 44g

Antoine Lavoisier

stated this about 1785

Basic Laws of Matter

• Law of Definite Proportions – no matter how much salt you have, it is always 39.34% Na and 60.66% Cl by mass.

Example: Sodium chloride always contains 39.34% Na and 60.66% Cl by mass.

2NaCl → 2Na + Cl2100g → 39.34g + 60.66g116.88g → ? + ?

Joseph Louis Proust

stated this in 1794.

Basic Laws of Matter

• Law of Multiple Proportions- Two or more elements can combine to form different compounds in whole-number ratios.

Example

John Dalton proposed this

in 1803.

Dalton’s Atomic Theory

• In 1808, Dalton proposed a theory to summarize and explain the laws of conservation of mass, definite proportions, & multiple proportions.

I was a school teacher at the

age of 12!

Dalton’s Atomic Theory

1. All matter is composed of extremely small particles called atoms.

2. Atoms of a given element are identical in size, mass, and other properties.**

3. Atoms cannot be subdivided, created, or destroyed.**

4. Atoms of different elements combine in simple whole-number ratios to form chemical compounds.

5. In chemical reactions, atoms are combined, separated, or rearranged.

**Today, we know these parts to have flaws.

John Dalton - 1808

Flaws of Dalton’s Theory…

2. Atoms of a given element are identical in size, mass, and other properties.

3. Atoms cannot be subdivided, created, or destroyed.

Isotopes – atoms with the same number of protons but a different number of neutrons

Subatomic particles – electrons, protons, neutrons, and more

The Atom

• Atom - the smallest particle of an element that retains the chemical properties of that element.

CARBON

Subatomic Particles

• Protons- positively charged particles found in the nucleus of an atom.

• Neutrons- neutral particles found in the nucleus of an atom.

• Electrons- negatively charged particles found in the electron cloud.

Joseph John Thomson

• In 1897 the English physicist Joseph John Thomson was able to measure the ratio of charge of the cathode ray particles to their mass.

• He found that the ratio was always the same regardless of the metal used to make the cathode or the nature of the gas inside the cathode ray tube.

• Thomson concluded that cathode rays were composed of identical, negatively charged particles called electrons.

Cathode Ray Tube Experiment Accomplishments

• Proved that the atom was divisible and that all atoms contain electrons.

• This contradicted Dalton’s Atomic Theory.

• This allowed a new model of the atom.

Discovery of the Nucleus

• In 1911, Ernest Rutherford performed a Gold Foil Experiment.

• He and his colleagues bombarded a thin piece of gold foil with fast moving, positively charged alpha particles.

Alpha Particles

• Alpha () particles are Helium-4 nuclei.

• This means they are two protons and two neutrons (with no electrons).

• Thus, they are positive.

4 +22 He

Gold Foil Experiment

Gold Foil Experiment

Gold Foil Experiment

• The volume of the nucleus was very small compared to the volume of the atom.

• Therefore, most of the atom was composed of empty space. Niels Bohr later found that this empty space was where the electrons were located.

Bohr’s Model of The Atom

Atomic Number

atomic number (Z) - the number of protons in the nucleus of each atom of a given element. – (Henry Moseley)

The number of p+ identifies the element.

Atomic Number increases from left to right on the periodic table.

Electrons

The number of electrons in a neutral atom is equal to the number of protons in that atom.

e- = p+

•Electrons can be lost or gained.

• When electrons are lost or gained, ions are formed.

Metals vs. Nonmetals

• Metals form cations.

Na Na+ + 1e-

• Nonmetals form anions.

Cl + 1 e- Cl-

Mass Number

• mass number (A)- the number of p+ & no in the nucleus of an atom.

# of neutrons = mass number – atomic number

Why aren’t electrons included when determining the mass number of the atom?

Isotopes

isotope- two or more atoms having the same atomic number (same #p+) , but different mass numbers (due to different #no).

Isotope Notation

Nuclear Notation

Hyphen NotationUses the elements symbol followed by a hyphen & the mass

number.

C-12

How many protons, neutrons & electrons are there in the following?

Cl-38 35Cl-1

Br-80 32S-2

N-14 56Fe+3

Changes in the Nucleus

Nuclear Reaction- changes that occur in the atom’s nucleus.

• Nuclear reactions can change the composition of an atom’s nucleus permanently.

Types of Radioactive Decay

Alpha Radiation ()- stream of high energy alpha particles.• Consists of 2 protons & 2 neutrons making it identical to

a He-4 nucleus.• Alpha particles can be represented by:

• Most alpha particles are able to travel only a few centimeters through air and are easily stopped by clothing etc.

4 4 +2 42 2 2 He He

Alpha Decay

239 235 494 92 2Pu U + He

234 230 492 90 2U Th + He

parent daughter

Types of Radioactive Decay

Beta Radiation () – consists of a stream of high speed electrons. These electrons are not electrons that are in motion around the atom’s nucleus.

• Beta particles can be represented by:

• Can penetrate through clothing and damage skin.

0 -1 0 01 1 1 e e

Beta Decay

6 6 02 3 -1He Li + 24 24 011 12 -1Na Mg +

parent daughter

Checking for Understanding

21084 Po 146C

alpha decay

beta decay

4 2062 82He + Pb

14 07 -1N +

22688 Ra

The Mole

mole (mol)- SI Unit for the amount of a substance that contains as many particles as there are atoms in exactly 12g of carbon-12.

• A unit of counting, like the dozen.

Avogadro’s Number

Avogadro’s Number - the number of particles in exactly one mole of a pure substance.

1 mole = 6.0221415 X 1023

1 mol = 6.02 x 1023

Amedeo Avogadro

Atomic Mass

atomic mass - the mass of one mole of an atom

• Atomic mass is expressed in atomic mass units (amu) or (u) or g/mol.

• Can be found on the periodic table.

• All atomic masses are based on the atomic mass of carbon-12 being 12 amu.

Molar Mass

molar mass - the mass of one mole of a pure substance.

• Molar mass is written in units of amu or g/mol.

Atomic mass vs. Molar mass

• atomic mass - the mass of one mole of an atom.

• molar mass - the mass of one mole of a pure substance.

Atomic Mass vs. Molar Mass

Example Atomic Mass Molar Mass

Na

Ag

C

O

22.99 g/mol 22.99 g/mol

107.87 g/mol 107.87 g/mol

12.01 g/mol 12.01 g/mol

16.00 g/mol 16.00 g/mol

Molar Mass of Compounds

Compound Molar Mass

H2O

C6H12O6

NaCl

Cl2

(NH4)3PO4

CuSO4·5H2O

18.02 g/mol

180.18 g/mol

58.44 g/mol

70.90 g/mol

149.12 g/mol

249.72 g/mol

The Mole Bridge

Cheer

• I say grams, you say molar mass.

grams – molar mass

grams – molar mass

Grams to Moles

Converting grams to moles: divide by molar mass.

1. How many moles of Ca are in 5.00g of Ca?1 mol Ca

5.00g Ca x = 40.08 g Ca

0.125 mol Ca

Moles to Grams

Converting from moles to grams: multiply by molar mass

1. What is the mass in grams of 2.25 moles of Fe?

55.85 g Fe2.25 mol Fe x =

1 mole Fe126 g Fe

Types of Particles

• Atoms – C, Cu, He

• Molecules – O2, C12H22O11, CO2 (all nonmetals in the formula)

• Formula units – NaCl, CaCl2, Mg(NO3)2 (includes a metal in the formula)

1 mole = 6.02 x 1023 particles

Particles to Moles

Converting particles to moles: divide by Avogadro's Number.

1. How many moles of Pb are in 1.50 X 1025 atoms of Pb? 2.49 x 101 moles Pb

Moles to Particles

Converting moles to atoms: multiply by Avogadro's Number.

1. How many molecules of NO are in 0.87 moles of NO?

5.2 x 1023 molecules NO

Grams to Moles to Particles

Example: How many molecules of N2 are in 57.1g of N2?

257.1 g N x 2

2

1 mol N

28.02 g N

232

2

6.02 x 10 molecules Nx

1 mol N

242= 1.23 x 10 molecules N

Particles to Moles to Grams

Example: How many grams of NaF are in 7.89 X 1024 formula units of NaF?

247.89 x 10 f.un. NaF x 23

1 mol NaF

6.02 x 10 f.un. NaF41.99 g NaF

x 1 mol NaF

= 550. g NaF

Chapter 4

Arrangement of Electrons in Atoms

The emission of light is fundamentally related to the behavior of electrons.

The Concept of Energy

• Energy – the ability (capacity) to do work

• Units of Energy: calorie (cal), Joule (J)

• Work is done when an object moves some distance in response to a force (pushing or pulling).

work = force x distance

Forms of Energy

• Mechanical energy

• Chemical energy

• Solar energy

• Radiant or Electromagnetic energy

• Electrical energy

• Nuclear energy

• Thermal energy

• Sound energy

• Magnetic energy

Where Does Energy Go?

Law of Conservation of Energy – energy cannot be created nor destroyed

Energy is continually transferred from one thing to another, never disappearing.

Types of Energy

• kinetic energy – the energy that objects have because they are moving

KE – “energy of motion”

• potential energy – the energy that is available for doing work at some later time

PE – “energy of position”

Energy Units

• What units are used for energy? joule = J (SI unit) calorie = cal (non SI unit) Calorie = Cal (food calorie =

1000 cal)

1.00 cal = 4.184 J

James Prescott Joule

Example: How many calories of energy is 1067 joules?

1.00 calorie1067 J x 255. calories

4.184 J

Measurable Properties of Light

• Wavelength () – the distance between corresponding points on adjacent waves

Units: meters, nanometers, etc.

1m = 1x109nm

Measurable Properties of Light

• Frequency ()- number of waves that pass a given point in a specific time (usually one second)

Unit: Hertz (Hz) -11Hz = or s

s

Types of Electromagnetic Radiation

• Gamma rays

• X-rays

• Ultraviolet

• Visible

• Infrared

• Microwaves

• Radio waves

Memorize these in order. Know E, , and order.

High energy, E

High frequency,

High wavelength,

•Violet

•Indigo

•Blue

•Green

•Yellow

•Orange

•Red

Relationship between and

c =

c = speed of light in m/s

= wavelength in m

= frequency in 1/s

What is the constant, h?

E = energy in J

h = constant in J.s

= frequency in 1/s or Hz

-34Planck's constant, h = 6.626 x 10 J s

E h

Combining the Equations you get…

c

E = h

=

Check for Understanding

13. What would the energy be for light with a frequency of 5.68 x 1012 1/s?

14. What is the energy of light that has a wavelength of 7.89 x 10-11 m?

Check for Understanding

-34 12

-21

13. E = hv

1 E = (6.626 x 10 J s)(5.68 x 10 )

s

E = 3.76 x 10 J

-34 8

-11

-15

hc14. E =

m(6.626 x 10 J s)(3.00 x 10 )

s E = 7.89 x 10 m

E = 2.52 x 10 J

Photoelectric Effect

• Photoelectric effect- the emission of electrons from a metal when light shines on the metal.

• The photoelectric effect does not occur when the light’s frequency is below a certain amount regardless of the intensity of the light.

Absorption vs. Emission

• Absorption – when energy is “taken in” by electrons

• Emission – when energy is “given off” by electrons

Energy States of an Electron

• Ground State- the lowest energy state of an atom. (stable)

• Excited state- state in which an atom’s potential energy is increased from that of the ground state. (unstable).

Spectroscopy

• Spectroscope – used to separate light into a spectrum by wavelength so it can be examined.

• Line-Emission Spectrum – produced when an electron jumps from a higher energy level to a lower energy level.

Acts as an atomic fingerprint.

How to use a Spectroscope

Discovery of theLine-Emission Spectrum

• Hydrogen atoms were excited by passing a high-voltage current through hydrogen gas causing the gas to glow a lavender color.

• When viewed with a spectroscope (diffraction grating or prism) the lavender light separated into four narrow lines of different color.

Hydrogen Line Spectrum

• Each of the lines seen in the hydrogen spectrum is a result of light at a different wavelength.

• Since light of a particular wavelength has a definite frequency and a definite energy, the lines of the hydrogen spectrum must be a result of the emission of photons with specific energies.

Other Emission Spectrum

helium

carbon

The Bohr Model

Bohr Model of the Atom

• The fact that hydrogen atoms only released photons of specific frequencies indicated that differences between the atom’s energy levels were fixed.

• The puzzle behind the hydrogen-atom spectra was solved in 1913 by Niels Bohr.

Bohr Model of the Atom

• The orbits are separated from one another by empty space where the electrons cannot exist.

• The energy of the electrons becomes higher as they get farther away from the nucleus.

• Electrons can move from one energy level to the next by gaining or losing a finite amount of energy.

Chapter 4 – Part II

Arrangement of Electrons in Atoms

The emission of light is fundamentally related to the behavior of electrons.

Heisenberg Uncertainty Principle

• In 1927, Werner Heisenberg stated that it is impossible to simultaneously determine both the position and velocity of an electron or any other particle. This became known as the Heisenberg Uncertainty Principle.

• Quantum theory was more widely accepted after this proposal.

Quantum Model of the Atom

The dots represent 1 electron and the region in which you are likely to find it

90% of the time.

Energy Levels

• Quantum numbers describe the location of electrons in an atom.

Principal Quantum Number, n

• principal quantum number, n – indicates the main energy level occupied by the electron

• n = 1→7

(notice how many periods are

on the periodic table…7 !)

(Does this sound like the Bohr model?)

Energy Sublevels

• Every energy level can be broken down into sublevels.

• Each sublevel has an energy that is slightly different.

• Not the traditional step ladder anymore, but still you can’t have energies in between the rungs of our new ladder.

No more

Bohr model

Orbitals

• orbital – a 3-D region in space within an energy level in which the electron is most likely to be found (90% of the time).

An orbital can hold a maximum of 2 e- !

Pauli Exclusion Principle

Sublevels in the Quantum Model

s orbital

2 electrons

1 orientation

Sub Levels in the Quantum Model

p orbitals: px, py, pz

6 electrons total

3 orientations

Sub Levels in the Quantum Model

d orbitals: dxy, dxy, dyz, dx2-y

2, dz2

10 electrons total

5 orientations

Sub Levels in the Quantum Model

f orbitals

14 electrons total

7 orientations

Summary of orbitals

energy

levels sublevels

n = 1 1s

n = 2 2s, 2p

n = 3 3s, 3p, 3d

n = 4 4s, 4p, 4d, 4f

n = 5 5s, 5p, 5d, 5f

n = 6 6s, 6p, 6d

n = 7 7s, 7p

Orbital Diagrams

1s

2s

2p

3s

4s

3p

3d

Aufbau Principle – fill in the lowest possible energy orbital

Pauli Exclusion Principle – a maximum of 2 e- may occupy one orbital, both with opposite spins

Aufbau Principle – fill in the lowest possible energy orbital

Pauli Exclusion Principle – a maximum of 2 e- may occupy one orbital, both with opposites spins

Aufbau Principle – fill in the lowest possible energy orbital

Hund’s Rule – within equal energy orbitals, the e- are distributed to have the maxiumum unpaired e- possible

Valence Eletrons

valence electrons – electrons in the highest energy level

Br: 1s22s22p63s23p64s23d104p5

1st E level 2nd E level 3rd E level 4th E level

4th is the highest Energy level, so there are 7 valence electrons

Chapter 5: Periodic Trends

Who told the elements where to go?

MENDELEEV!

(Father of the periodic table)

History of the Periodic TableEarly 1800’s

J.W. Dobereiner organized elements into triads.

Triads- groups of three elements with similar properties. The middle element of the triad was thought to be an approximate average of the properties of the first & third element.

History of the Periodic TableMid 1800’s

J.A.R. Newlands- developed the law of octaves in 1865.

Law of octaves- stated that if elements were arranged by increasing atomic mass, the properties of the eighth element were similar to properties of the first, the ninth like those of the second, the tenth like those of the third.

History of the Periodic TableMid 1800’s

In 1869, Dmitri Mendeleev published the first periodic table.

• Mendeleev arranged the elements horizontally by increasing atomic mass and placed elements in groups (vertically) based on similar properties.

History of the Periodic TableEarly 1900’s

In 1913, Henry Moseley developed the Modern Periodic Table. He determined the atomic numbers of many elements and then used those values to arrange the periodic table in rows by increasing atomic number and in columns by similar properties.

Metals, Nonmetals, & Metalloids

Metals• Make up most of the periodic table.• Ductile- can be drawn into wire• Malleable- can be hammered into thin sheets.• Lustrous- shiny• Good Conductors of heat & electricity• Located to the left of the step ladder on the

periodic table.

Metals, Nonmetals, & Metalloids

Nonmetals

• Brittle-break when hammered.

• Poor conductors of heat & electricity.

• Lack luster

• Located to the right of the step ladder on the periodic table.

Metals, Nonmetals, & Metalloids

Metalloids

• Semimetals.

• Properties of both metals & nonmetals.

• Located along the step ladder on the periodic table.

• Examples: B, Si, Ge, As, Sb, Te

Metals, Nonmetals, & Metalloids

Alkali Metals- Group 1

• Most reactive group of metals.• Usually found in combined

form as a salt due to their high reactivity.

• Combine vigorously with nonmetals especially groups 16 & 17.

• React readily with water.• Soft and silvery appearance.

Alkaline Earth Metals- Group 2

• Found in the earth’s crust but not in the elemental form due to their high reactivity. They are usually found in rock structures.

• 2nd most reactive group of metals.• More dense than group 1.• Shiny silvery-white color.

Transition Metals- Groups 3-12

• Compose the d-block.• Typical metallic

properties.• Good conductors.• Lustrous.• Produce colored ions.

Main Group Elements Groups 13-18

• Compose the p-block.• Properties of elements

vary greatly.• Contains all of the

nonmetals & metalloids as well as some metals.

Halogens -Group 17

• Most reactive groups of non-metals.• React vigorously with metals (especially groups

1 & 2) to produce salts.• Fluorine is a poisonous pale yellow gas,

chlorine is a poisonous pale green gas, bromine is a toxic and caustic brown volatile liquid, and iodine is a shiny black solid which easily sublimes to form a violet vapor on heating.

• Found in nature in the combined state.

Noble Gases- Group 18

• Least reactive of all elements.• Often called inert gases.• All are gases.• The noble gases are all found in minute

quantities in the atmosphere, and are isolated by fractional distillation of liquid air. Helium can be obtained from natural gas wells where it has accumulated as a result of radioactive decay.

Inner Transition Metals- Periods 6 & 7

• Compose the f-block.• Fill in Between Groups 3 & 4 on the

Periodic Table.

Lanthanides (Period 6)- Rare Earth Metals

• Shiny reactive metals

Actinides (Period 7)• Unstable & radioactive metals.• Most are laboratory made.

Reactivity of Metals Trend

Period Trend- Metals increase in activity from right to left on the periodic table.

• The alkali metals are the most reactive group of metals.

Group Trend- Metals increase in reactivity from top to bottom with a group.

• Ra is the most reactive alkaline earth metal.

Reactivity of Nonmetals Trend

Period Trend- Nonmetals increase in activity from left to right on the periodic table with the exception of the noble gases.

• The halogens are the most reactive group of nonmetals.

Group Trend- Nonmetals increase in reactivity from bottom to top with a group.

• F is the most reactive halogen.

Reactivity Trends

What is the most reactive metal on the periodic table? Explain.

Circle the most reactive nonmetal in each row:

1. Te Po S

2. Br I Cl

Atomic Radii

Atomic Radius- ½ the distance between the nuclei of two identical atoms that are bonded together.

• Measure from nucleus to nucleus and divide by two.• The larger the atomic radius the larger the atom.

Atomic Radii Trend

Atomic Radii increases as you travel down & left on the periodic table!

Atomic Radii Trend

Atomic Radii Trends

1. Circle the atom with the largest atomic radii.a. I F Cl

b. O B N

c. S Sb Sr

2. Put the following elements in order of increasing atomic radii.a. Na, P, Rb

3. Which element has the largest atomic radii?

Ionic Radii

Cation- a positive ion formed from the loss of electrons. Ex. Ca+2

• Metals typically form cations.• Cations are smaller than their parent atoms.

Example: Na > Na+

Ionic Radii

Anion- a negative ion formed from the gain of electrons. Ex. Cl-1

• Nonmetals typically form anions.• Anions are larger than their parent atoms.

Example: Cl-1 > Cl

Ionic Radii Trends

1. Circle the larger ion.

a. Cl-1 Cl

b. Na Na+1

c. N N-3

d. Fe+2 Fe+3

e. C+4 C-4

Ionization Energy

Ionization Energy- the energy required to remove an electron from an atom.

• Nonmetals have higher ionization energies than metals.

• The lower the ionization energy of an atom the easier it is to remove an electron.

Ionization Energy Trend

General Group Trend (representative elements)- Ionization energy increases as you travel up a group in the periodic table.

Reason: As you travel up a group there are less principle energy levels, therefore, electrons are located closer to the nucleus. It is harder to remove electrons that are closer to the nucleus to due the nuclear charge attraction.

Ionization Energy Trend

Ionization energy increases as you travel up & right on the periodic table!

Ionization Energy Trends

1. Circle the atom with the highest ionization energy.

a. F O B

b. As N P

c. F C Li

Electron Affinity

Electron Affinity- energy change that occurs when an electron is acquired by a neutral atom.

• Nonmetals have a more negative (LOWER) electron affinity than metals.

• Metals typically have a more positive (HIGHER) electron affinity than nonmetals.

• Noble gases have an electron affinity between metals & nonmetals.

Electron Affinity Trends

1. Circle the atom with the highest electron affinity.

a. N S Ar

b. O Ne Ca

c. Cu Ar Kr

Electronegativity

Electronegativity- measure of the ability of an atom in a chemical compound to attract electrons.

• The higher the electronegativity the greater the attraction for electrons.

• The electronegativity scale was determined by Linus Pauling.

• The electronegativity scale ranges from 0 to 4.• Since some noble gases do not form compounds

they have electronegativity values of 0.• Used to determine the polarity of an bond.

Electronegativity Trend

Electronegativity increases as you travel up & right on the periodic table!

Electronegativity Trends

1. Circle the atom with the highest electronegativity.

a. F S Sn

b. Li N O

Chemical Bonding

Chapter 6

Pages 174-213

15-1 The Attachment Between Atoms

atoms combine to form

ionic bonds covalent bonds(M + NM) (NM + NM)

chemical bond – a mutual electrical attraction between the nuclei and valence electrons of two atoms that binds the atoms together

Ionic Bonding• ionic bond – when electrons are taken by one

atom from another atom

metal and a nonmetal

NaCl

cation and anion

(The charges are “hidden” to make a neutral compound.)

The simplest ratio of the packed ions is called:

The Formula Unit Ex: NaCl

Ions

• cations (+)• anions (-)

• monatomic ions – ions formed from one atom Examples: Na+ or O-2

• polyatomic ions - ions formed from two or more atoms bonded together

Examples: NH4+ or SO4

-2

Ionic Compounds

• solid at room temperature• high melting points (thus are usually solid at RT)• formula unit represents the lowest ratio of ions

that combine to form a neutral compound• most are crystalline solids• when dissolved in water, the ionic compounds

will break up into ions (dissociate)• the solutions of ionic compounds will conduct

electricity (electrolytes)

Covalent Bondingcovalent bond – when electrons are shared

between two atoms– two nonmetals – No ions formed! (no electrons are taken)

Pure Covalent

Ionic

There is another type of bond, not

purely covalent and not

purely ionic.

Polar Covalent

Nonpolar Covalent

Electronegativity

• electronegativity – a measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound

The difference in electronegativity values for two atoms will indicate whether the two atoms form an ionic bond or a polar or nonpolar covalent bond.

Electronegativity Differences

• 0.0 to 0.3 nonpolar covalent

• 0.3 to 1.7 polar covalent

• 1.7 and up ionic

Learn these values…Although the general rule is a metal and nonmetal

will form an ionic bond and two nonmetals will form a covalent bond.

Ionic, Polar Covalent, or Nonpolar Covalent?

What kind of bond would each pair form?

1. N and S

2. S and C

3. Mg and Cl

4. C and F

5. Ba and O

VSEPR Theory

Valence

Shell

Electron

Pair

Repulsion Theory

Repulsion between the sets of valence-level electrons

surrounding an atom causes these sets to be oriented as far

apart as possible.

Regions of Electron Density

What is a Region of electron density?

• Single bond (2e- connecting 2 atoms)

• Double bond(4e- connecting 2 atoms)

• Triple bond(6e- connecting 2 atoms)

• Lone pair (unbonded pair) (2e- alone on an atom)

LINEAR180o sp hybridization2 Regions of Electron Density2 Bonds

bonded pair of electrons

bonded pair of electrons

TRIGONAL PLANAR120o sp2 hybridization3 Regions of Electron Density3 Bonds

3 bonded pairs of electrons

TETRAHEDRAL109.5o sp3 hybridization

4 Regions of Electron Density4 Bonds

4 bonded pairs of electrons

TRIGONAL PYRAMIDAL107o sp3 hybridization

4 Regions of Electron Density3 Bonds & 1 Lone Pair

3 bonded pairs of electrons

1 lone pair of electrons

BENT105o sp3 hybridization

4 Regions of Electron Density2 Bonds & 2 Lone Pairs

All of these have 4 regions of electron density and sp3 hybridization

(although the number of bonded pairs is different)

TRIGONAL BIPYRAMIDAL120o & 90o

5 Regions of Electron Density5 Bonds

OCTAHEDRAL90o

6 Regions of Electron Density6 Bonds

SF6

ICl4-

SQUARE PLANAR90o

6 Regions of Electron Density4 Bonds & 2 Lone Pairs

Hybridization

Regions of e- Density Hybridization

2 sp

3 sp2

4 sp3

Chapter 6: Chemical Bonding

Part II – Polarity

What do the differences in electronegativity indicate between

two atoms?

the bond type that they will form

–Ionic (difference of 1.7+)

–Polar covalent (difference of 0.3-1.7)

–Nonpolar covalent (difference of 0.0-0.3)

What are the three types of bonds?

Nonpolar Covalent Bond

• Nonpolar covalent bond – a covalent bond in which the electrons are shared equally

Example: Cl2

Both chlorines have the same electronegativity.

Polar Covalent Bond

• Polar covalent bond – a covalent bond in which the electrons are shared unequally

Example: HCl

Which end represents the hydrogen end & which end represents the chlorine end?

H Cl

How do you know this answer? Cl has a higher electronegativity

Dipole

• dipole – partial negative or partial positive charge formed during unequal sharing of electrons (in polar bonds only)

• + -

• The direction of a dipole is from the dipole’s positive pole to its negative pole.

Nonpolar vs. Polar Covalent Bond

The dipole is drawn over the HCl molecule. The arrow points towards the more electronegative element, Cl.

There is no dipole drawn over the H2 molecule…

WHY?

There is not an uneven distribution of electrons.

Nonpolar vs. Polar Covalent Bond

Compare the electron clouds for the H2 and HCl molecules. Why are they different?

The electrons are not shared equally within the HCl molecule.

Do we usually draw the electron clouds with the Lewis structures?

NO

Dipole

Example: PCl3 has three dipoles because there are three polar bonds.

P

Cl Cl Cl

What is the more electronegative element, P or Cl?

So, how would the dipoles be drawn?

These three dipoles are drawn beside the bonds,

pointing towards the more electronegative

element, Cl.

Each dipole represents a polar covalent bond.

You Try It!

• Draw the Lewis Structures on your white board for the following molecules (Aha!…they must all be covalent, then!) and LABEL ALL THE DIPOLES along each bond.

1. CO2

2. HBr

3. NH3

Draw them on your white board and only when you are

ready to check the Lewis structures AND dipoles, go on

to the next slide.

Answers

Did you make CO2 linear?

Did you make NH3 trigonal pyramidal?

Did you show the dipoles pointing to the more electronegative element?

Dipoles within Molecules

• There are dipoles created along bonds, however, a molecule can exhibit overall polarity.

An overall dipole is created.

Here is the dipole along the first H-O bond…

Here is the dipole along the second H-O bond…

An overall dipole is created.

Two individual dipoles along each bond cause the molecule to have a greater electron cloud towards the oxygen end leaving the other end of the molecule more positive. This causes a partial negative end towards the oxygen and the a partial positive end towards the hydrogens (you can only have two “ends” to a small molecule.

The bond is polar and so is the overall molecule.

When you have a two atom molecule, it is easy to tell if it is a polar molecule.

partially negative end

partially positive end

Polarity of a Molecule

• Polarity of a molecule depends upon two things– the polarity of the bonds– the shape of the molecule

• All molecules with only nonpolar bonds are nonpolar molecules.

• Molecules with polar bonds, may or may not be polar molecules, depending on the shape.

carbon monoxide, CO

YES

Is the molecule polar? YES…a two atom molecule is polar if the bond is polar.

Is the bond polar?

ammonia, NH3

YES

Is the molecule polar?

YES…the molecule is not symmetrical and does have a partial positive end and a partial negative end. You can draw on overall dipole.

Are the bonds polar?

methane, CH4

YES

Is the molecule polar?

NO…the molecule is symmetrical and does NOT have a partial positive end and a partial negative end. (The center and the outside does not count as an “end”.)

Are the bonds polar?

Polar or Nonpolar Molecule?

Rule #1: All molecules with nonpolar bonds are always nonpolar molecules.

Rule #2: Molecules with polar bonds, will be polar only if they are nonsymmetrical (like NH3).

You Try It! #2

Draw the following molecules, including the dipoles along each bond.

Determine whether they are polar or nonpolar molecules.

• CH3Cl

• SF6

• O2

• H2S

• KCl

Draw them on your white board. Click to

see the answers.

Chapter 7 – Part IFormula Writing and Naming

May I please have a glass of

Be sure to know how to name…

• Ions– polyatomic and monotomic, cations and anions

• Ionic compounds– ones with and without Roman numerals

• Molecular/Covalent compounds– use prefixes for these (mon-, di-, tri-, tetra-, …)

• Acids

Chapter 7 – Part IIMath with Formulas

Percent Composition• What is the percent composition of oxygen

in carbon dioxide?

Percent composition is determined by mass. You need to know the total mass of the element for the bottom of your fraction and the mass of the part of the compound in question for the top of the fraction.

2 x 16.00 g%O = x 100.% = 72.7% O

44.01 gmolar mass of carbon dioxide

molar mass of oxygen

multiplied by 2 because there are 2 oxygen in the formula

Empirical Formulas• What is the empirical formula for glucose,

C6H12O6?

An empirical formula is just the lowest ratio of the molecular formula. It is not representative of the real molecule, but it just shows the ratio of atoms in that molecule.

C6H12O6 has a ratio of 6:12:6 which can be reduced, so the answer is

empirical formula is CH2O (1:2:1 ratio)

Determining Molecular Formulas from Empirical Formulas

• What is the molecular formula for a compound that weighs 110.g/mol and has an empirical formula of C3H3O?

The molecular formula must weigh some multiple of the empirical formula. First determine the mass of the empirical formula.

C3H3O weighs 55 g/mol and the real compound weighs 110. g/mol, or twice that much.

So, the real formula is 2x C3H3O or C6H6O2