Overview of the Basics CHAPTER 1-3 Review Chemistry: The Molecular Nature of Matter, 6 th edition By Jesperson , Brady, & Hyslop

Overview of the Basics

CHAPTER 1-3 Review

Chemistry: The Molecular Nature of Matter, 6th editionBy Jesperson, Brady, & Hyslop

2

CHAPTER 1-3 Review

Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

Learning Objectives Scientific Method Matter: definition, elements, compounds, mixtures, changes/properties Atomic Theory

Law of definite proportions Law of conservation of mass

Chemical formulas Chemical equations

Balancing Measurements: units, conversions, uncertainty Significant Figures Density Subatomic particles Atomic #, mass #, atomic weights Periodic Table Ionic Compounds, hydrates, molecular compounds Basic nomenclature


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Chapter 1 Scientific Method

1. Make observations/collect data• Empirical fact• Something we see, hear, taste, feel, or smell• Something we can measure • Organize data so we can see relationships

2. Law or Scientific Law • Usually an equation• Based on results of many experiments• Only states what happens• Does not explain why they happen

3. Hypothesis• Mental picture that explains observed laws • Tentative explanation of data• Make predictions• Leads to further tests• Go to laboratory and perform experiments

4. Theory• Tested explanation of how nature behaves • Devise further tests• Depending on results, may have to modify

theory• Can never prove theory is absolutely correct

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Chapter 1 Elements


• Substances that can’t be decomposed into simpler materials by chemical reactions

• Substances composed of only one type of atom• Simplest forms of matter that we can work with

directly• More complex substances composed of elements

in various combinations

diamond = carbon gold sulfur


5

Chapter 1 Elements


6

Chapter 1 Classification of Matter

http://ridenourmhs.wikispaces.com/ESUnit2

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Chapter 2 Chemical vs Physical Properties


Physical properties Can be observed without changing chemical makeup of substance

Chemical properties

Solids: Fixed shape and volumeParticles are close together

Liquids: Fixed volume, but take container shapeParticles are close together

Gases: Expand to fill entire containerParticles separated by lots of space

• Chemical change or reaction that substance undergoes

• Chemicals interact to form entirely differentsubstances with different chemical and physical properties

• Describe behavior of matter that leads to formation of new substance

• “Reactivity" of substancee.g. Iron rusting

– Iron interacts with oxygen to form a new substance.


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Chapter 1 Atomic Theory

Developed by John Dalton to explain Law of Conservation of Mass & Law of Definite Proportions

1. Matter consists of tiny particles called atoms.2. Atoms are indestructible.

• In chemical reactions, atoms rearrange but do not break apart.

3. In any sample of a pure element, all atoms are identical in mass and other properties.

4. Atoms of different elements differ in mass and other properties.

5. In a given compound, constituent atoms are always present in same fixed numerical ratio.


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Chapter 1 Law of Definite Proportions

Atoms react as Whole particles.

When two elements form more than one compound, different masses of one element that combine with same mass of other element are always in ratio of small whole numbers.

e.g. Fool’s gold, pyrite, iron(III) sulfide Mass ratio always 1.00 g of iron to 0.574 g of sulfur

e.g. WaterMass ratio always: 8 g O to 1 g H


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Chapter 1 Law of Conservation of Mass

sulfur sulfur dioxide trioxideMass S 32.06 g 32.06 g

Mass O 32.00 g 48.00 g


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Chapter 1 Molecules and Chemical Formulas

Atoms combine to form more complex substances = Molecules

Chemical Formulas: • Specify composition of substance• Chemical symbols represent atoms of elements present• Subscripts:

– Given after chemical symbol– Represents relative numbers of each type of atom

Example: Fe2O3 : iron and oxygen in 2:3 ratio


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Chapter 1 Hydrates

• Crystals that contain water moleculese.g. Plaster: CaSO4∙2H2O calcium sulfate dihydrate

– Water is not tightly held• Dehydration

– Removal of water by heating– Remaining solid is anhydrous (without water)

Blue =

CuSO4 •5H2O

White = CuSO4


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Chapter 1 Depicting Molecules

H C H

H

H

CH4

methane


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Chapter 1 Chemical Equations

• Use chemical symbols and formulas to represent reactants and products. – Reactants on left hand side– Products on right hand side– Arrow () means “reacts to yield”

e.g. CH4 + 2O2 CO2 + 2H2O – Coefficients

• Numbers in front of formulas• Indicate how many of each type of molecule

reacted or formed– Equation reads “methane and oxygen react to

yield carbon dioxide and water”


15

Chapter 1 Conservation of Mass in Reactions

• Mass can neither be created nor destroyed• This means that there are the same number of each type

of atom in reactants and in products of reaction– If number of atoms same, then mass also same

CH4 + 2O2 CO2 + 2H2O 4H + 4O + C = 4H + 4O + C

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Chapter 1 Balanced Chemical Equations


Ex. 2C4H10 + 13O2 8CO2 + 10H2O

4 C and 10 H per molecule 2 O per

molecule

2 H and 1 O per

molecule

1 C and 2 O per

moleculeEx. 2C4H10 + 13O2 8CO2 + 10H2O

2 molecules of C4H10

13 molecules of O2

10 molecules of C4H108

molecules of CO2

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Chapter 2 Intensive vs Extensive Properties


Intensive properties – Independent of sample

size– Used to identify

substancese.g. Color

Density Boiling point Melting point Chemical reactivity

Extensive properties – Depend on sample sizee.g. volume and mass

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Chapter 2 Measurements


1. Measurements involve comparison– Always measure relative to reference

e.g. Foot, meter, kilogram– Measurement = number + unit

e.g. Distance between 2 points = 25• What unit? inches, feet, yards, miles• Meaningless without units

2. Measurements are inexact– Measuring involves estimation– Always have uncertainty– The observer and instrument have

inherent physical limitations

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Chapter 2 International System of Units


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21



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Chapter 2 Decimal Multipliers


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Chapter 2


10001

1. Distance (d )Centimeter (cm)

1 cm = 10–2 m = 0.01 m Millimeter (mm)

1 mm = 10–3 m = 0.001 m 2. Volume (V)

1 L = 1000 mL 1 mL = 1 cm3

3. Mass (m)1 g = 0.1000 kg =

g

4. Temperature (T)273.15 K = 0°C

4 Common Lab Measurements

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Chapter 2


Uncertainty in Measurements

Example: Consider two Celsius thermometers• Left thermometer has markings every 1˚C

– T between 24 °C and 25 °C– About 3/10 of way between marks– Can estimate to 0.1 °C = uncertainty– T = 24.3 0.1 °C

• Right thermometer has markings every 0.1 °C– T reading between 24.3 °C and 24.4 °C– Can estimate 0.01 °C – T = 24.32 0.01 °C

Measurements all inexactLimitations of reading instrument

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Chapter 2 Significant Figures


1. All non-zero numbers are significant. e.g. 3.456

has 4 sig. figs.2. Zeros between non-zero numbers

are significant. e.g. 20,089 or 2.0089 × 104 has 5 sig. figs

3. Trailing zeros always count as significant if number has decimal point e.g. 500. or 5.00 × 102

has 3 sig. figs

Scientific convention: All digits in measurement up to and including first estimated digit are significant.

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Chapter 2 Significant Figures


4. Final zeros on number without decimal point are NOT significant

e.g. 104,956,000 or 1.04956 × 108 has 6 sig. figs.

5. Final zeros to right of decimal point are significante.g. 3.00 has 3 sig. figs.

6. Leading zeros, to left of first nonzero digit, are never counted as significant

e.g. 0.00012 or 1.2 × 10–4 has 2 sig. figs.

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Chapter 2 Significant Figures: Rounding


1. If digit to be dropped is greater than 5, last remaining digit is rounded up.e.g. 3.677 is rounded up to 3.68

2. If number to be dropped is less than 5, last remaining digit stays the same.e.g. 6.632 is rounded to 6.63

3. If number to be dropped is exactly 5, then if digit to left of 5 is a. Even, it remains the same.e.g. 6.65 is rounded to 6.6b. Odd, it rounds up.e.g. 6.35 is rounded to 6.4

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Chapter 2 Significant Figures: Calculations


Multiplication and Division• Number of significant figures in answer = number of

significant figures in least precise measurement

e.g. 10.54 × 31.4 × 16.987

4 sig. figs. × 3 sig. figs. × 5 sig. figs. = 3 sig. figs.

Addition and Subtraction• Answer has same number

of decimal places as quantity with fewest number of decimal places.

12.9753 319.5+ 4.398

4 decimal places1 decimal place3 decimal places1 decimal place336.9

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Chapter 2 Significant Figures: Exact Numbers


• Numbers that come from definitions– 12 in. = 1 ft– 60 s = 1 min

• Numbers that come from direct count– Number of people in small room

• Have no uncertainty• Assume they have infinite number of significant figures. • Do not affect number of significant figures in

multiplication or division

30

Chapter 2 Scientific Notation


• Clearest way to present number of significant figures unambiguously– Report number between 1 and 10 followed by

correct power of 10 – Indicates only significant digits

e.g. 75,000 people attend a concert– If a rough estimate

• Uncertainty 1000 people• 7.5 × 104

– If number estimated from aerial photograph• Uncertainty 100 people• 7.50 × 104

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Chapter 2 Accuracy & Precision


Accuracy– How close measurement is to true or

accepted true value• Measuring device must be calibrated

with standard reference to give correct value

Precision– How well set of repeated

measurements of same quantity agree with each other

– More significant figures equals more precise measurement

32

Chapter 2 Dimensional Analysis


• Also called the Factor Label Method• Not all calculations use specific equation• Use units (dimensions) to analyze problemConversion Factor • Fraction formed from valid equality or

equivalence between units• Used to switch from one system of

measurement and units to another

Given Quantity

Desired Quantity

Conversion Factor

× =

33

Chapter 2 Dimensional Analysis


Example: Convert 0.097 m to mm.• Relationship is 1 mm = 1 × 10–3 m• Can make two conversion factors

• Since going from m to mm use one on left.m 101

mm 13 mm 1

m 101 3

m 101mm 1m 097.0 3

= 97 cm

34

Chapter 2 Density


• Ratio of object’s mass to its volume

• Intensive property (size independent)– Determined by taking ratio of two extensive properties (size

dependent)– Frequently ratio of two size dependent properties leads to size

independent property

• Density useful to transfer between mass and volume of substance

• Density decreases slightly as temperature increases• Units: g/mL or g/cm3

volumemassdensity

35

Chapter 3


Rutherford Nuclear Atom

Discovery of the NucleusRutherford

Alpha scattering expt

Discovery of Protons1918 Rutherford

Mass spectrometer

Discovery of electron mass and charge

Millikan Oil Drop expt

Discovery of the Electron1897 ThomsonCathode ray tube expt

Discovery of Neutron:1932 Chadwick

Discovery of Subatomic Particles in the late 1800’s and early 1900s

36

Chapter 3 Properties of Subatomic Particles


Particle Mass (g) Electrical Charge Symbol

Electron

9.10939 10–28 –1

Proton 1.67264 10–24 +1

Neutron 1.67495 10–24 0

e01

p11

11 ,H

n10

Three kinds of subatomic particles of principal interest to chemists

Nucleus (protons + neutrons)

Electrons

37

Chapter 3 Atomic Notation


Atomic number (Z) = Number of protons that atom has in nucleusIsotopes = Atoms of same element with different

masses– Same number of protons ( )– Different number of neutrons ( )

Isotope Mass number (A)– A = (number of protons)+(number of neutrons) =

Z + N– For charge neutrality, number of electrons and

protons must be equalAtomic Symbols = Summarize information about

subatomic particles– Every isotope defined by two numbers Z and A

Ex. What is the atomic symbol for helium? He has 2 e–, 2 n and 2 p Z = 2, A = 4

p11

n01

He42

38

Chapter 3 Isotopes


• Most elements are mixtures of two or more stable isotopes

• Each isotope has slightly different mass• Chemically, isotopes have virtually identical chemical

properties• Relative proportions of different isotopes are

essentially constant• Isotopes distinguished by mass number (A):

e.g. – Three isotopes of hydrogen (H)– Four isotopes of iron (Fe)

39

Chapter 3 Carbon-12 Atomic Mass Scale


• Need uniform mass scale for atomsAtomic mass units (symbol u)

– Based on carbon:• 1 atom of carbon-12 = 12 u (exactly)• 1 u = 1/12 mass 1 atom of carbon-12 (exactly)

Why was 12C selected?– Common– Most abundant isotope of carbon – All atomic masses of all other elements ~ whole

numbers– Lightest element, H, has mass ~1 u

40

Chapter 3 Calculating Atomic Mass


• Generally, elements are mixtures of isotopese.g. HydrogenIsotope Mass % Abundance 1H 1.007825 u 99.985 2H 2.0140 u 0.015How do we define atomic mass?

– Average of masses of all stable isotopes of given elementHow do we calculate average atomic mass?

– Weighted average– Use isotopic abundances and isotopic masses

41

Chapter 3 Periodic Table


– Summarizes periodic properties of elementsEarly Versions of Periodic Tables

– Arranged by increasing atomic mass– Mendeleev (Russian) and Meyer (German) in 1869– Noted repeating (periodic) properties

Modern Periodic Table – Arranged by increasing atomic number (Z ):– Rows called periods– Columns called groups or families

• Identified by numbers• 1 – 18 standard international• 1A – 8A longer columns and 1B – 8B shorter columns

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43



44

Chapter 3 Periodic Table Groups


1A 2A B B 7A 8A

Alkali Metals Alkaline Earth

Metals

Transition Metals

Lanthanide & Actinide

Halogens Nobel Gases

Very reactive

Metals except for H

+1 ions

React with Oxygen to form compounds that dissolve into alkaline solutions in water

Reactive

+2 ions

Oxygen compounds are strongly alkaline

Many are not water soluble

Metals

Form ions with several different charges (oxidation states)

Tend to form +2 and +3 ions

Lanthanides 58 – 71

Actinides 90 – 103

Actinides are radioactive

Reactive

Form diatomic molecules in elemental state

-1 ions

Salts with alkali metals

Inert

Heavier elements have limited reactivity

Do not form ions

Monoatomic gases

45

Chapter 3 Metals, Nonmetals, and Metalloids


46

Chapter 3 Metals, Nonmetals, and Metalloids


Metals Nonmetals Metalloids

• Metallic luster, malleable, ductile, hardness variable

• Conduct heat and electricity

• Solids at room temperature with the exception of Hg

• Chemical reactivity varies greatly: Au, Pt unreactive while Na, K very reactive

• Brittle

• Insulators, non-conductors of electricity and heat

• Chemical reactivity varies

• Exist mostly as compounds rather then pure elements

• Many are gases, some are solids at room temp, only Br2 is a liquid.

• Metallic shine but brittle

• Semiconductors: conduct electricity but not as well as metals: examples are silicon and germanium

47

Chapter 3 Ions and Ionic Compounds


Ions– Transfer of one or more electrons from one

atom to another – Form electrically charged particles

Ionic compound – Compound composed of ions– Formed from metal and nonmetal– Infinite array of alternating Na+ and Cl– ions

Formula unit – Smallest neutral unit of ionic compound– Smallest whole-number ratio of ions

48



Metal + Non-metal ionic compound 2Na(s) + Cl2(g) 2NaCl(s)

Na+ + ClNa + Cl NaCl(s)

e

Michael Watson Richard Megna/Fundamental Photographs Richard Megna/Fundamental Photographs

Anions = Negatively charged ions Cations = Positively charged ions

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Electrical conductivity requires charge movementIonic compounds:

– Do not conduct electricity in solid state– Do conduct electricity in liquid and aqueous states where ions are free

to move Molecular compounds:

– Do not conduct electricity in any state– Molecules are comprised of uncharged particles

50



Negative (–) charge on anion = number of spaces you have to move to

right to get to noble gas

51

Chapter 3 Rules for Writing Ionic Formulas


1. Cation given first in formula2. Subscripts in formula must produce electrically neutral

formula unit3. Subscripts must be smallest whole numbers possible

– Divide by 2 if all subscripts are even– May have to repeat several times

4. Charges on ions not included in finished formula unit of substance

– If no subscript, then 1 implied

52

Chapter 3 Determining Ionic Formulas


“Criss-cross” rule– Make magnitude of charge on one ion into subscript for

other – When doing this, make sure that subscripts are reduced

to lowest whole number.

Ex. What is the formula of ionic compound formed between aluminum and oxygen ions?

Al2O3Al3+ O2–

53

Chapter 3 Transition Metal and Post-Transition Metal Ions


54

Chapter 3 PolyatomicIons


Example: What is the formula of the ionic compound formed between ammonium and phosphate ions?

Ammonium = NH4+

Phosphate = PO43–

(NH4)+ (PO4)3–

(NH4)3PO4

55

Chapter 3 Nomenclature


56

Chapter 3 Nomenclature: Ionic Compounds


Cations: – Metal that forms only one positive ion

• Cation name = English name for metal– Na+ sodium – Ca2+ calcium

– Metal that forms more than one positive ion– Use Stock System

• Cation name = English name followed by numerical value of charge written as Roman numeral in parentheses (no spaces)

• Transition metal– Cr2+ chromium(II) Cr3+ chromium(III)

57

Chapter 3 Nomenclature: Ionic Compounds


Anions:– Monatomic anions named by adding

“–ide” suffix to stem name for element

– Polyatomic ions use names in Table 3.5

58

Chapter 3 Nomenclature: Hydrates


• Ionic compounds – Crystals contain water molecules – Fixed proportions relative to ionic substance

• Naming– Name ionic compound– Give number of water molecules in formula

using Greek prefixes

mono- = 1 hexa- = 6di- = 2 hepta- = 7tri- = 3 octa- = 8tetra- = 4 nona- = 9penta- = 5 deca- = 10

59

Chapter 3 Molecular Compounds


Molecules– Electrically neutral particle – Consists of two or more atoms

Chemical bonds– Attractions that hold atoms together in molecules– Arise from sharing electrons between two atoms– Group of atoms that make up molecule behave

as single particleMolecular formulas

– Describe composition of molecule– Specify number of each type of atom present

60

Chapter 3 Nonmetal Hydrides


Nonmetal hydrides– Molecule containing nonmetal + hydrogen– Number of hydrogens that combine with nonmetal =

number of spaces from nonmetal to noble gas in periodic table

N O F Ne

61

Chapter 3 Organic Compound Formulas


Molecular formula – Indicates number of each type of atom in moleculee.g. C2H6 for ethane or C3H8 for propane– Order of atoms

• Carbon Hydrogen Other atoms alphabeticallye.g. sucrose is C12H22O11

Emphasize alcohol – write OH group last– C2H5OH

Structural formula– Indicate how carbon atoms are connected– Ethane = CH3CH3

– Propane = CH3CH2CH3

62

Chapter 3 Nomenclature: Molecular Compounds


• Goal is a name that translates clearly into molecular formulaNaming Binary Molecular Compounds

– Which two elements present?– How many of each?

Format:– First element in formula

• Use English name– Second element

• Use stem and append suffix –ide – Use Greek number prefixes to specify how many atoms of

each element

63

Chapter 3 Nomenclature: Binary Molecules


1. hydrogen chloride

2. phosphorous pentachloride

3. triselenium dinitride

• Mono always omitted on first element• Often omitted on second element unless more than one

combination of same two elementse.g. Carbon monoxide CO Carbon dioxide CO2

• When prefix ends in vowel similar to start of element name, drop prefix vowel

1 H 1 Cl HCl

1 P 5Cl PCl5

3 Se 2N Se3N2

64

Chapter 3 Nomenclature: Exceptions for Binary Molecules


Binary compounds of nonmetals + hydrogen– No prefixes to be used– Get number of hydrogens for each nonmetal from

periodic table– Hydrogen sulfide = H2S– Hydrogen telluride = H2Te

Molecules with Common Names– Some molecules have names that predate IUPAC

systematic names– Water H2O ▪ Sucrose C12H22O11

– Ammonia NH3 ▪ Phosphine PH3

Overview of the Basics CHAPTER 1-3 Review Chemistry: The Molecular Nature of Matter, 6 th edition By Jesperson , Brady, & Hyslop

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