1

What is

chemistry?

Chemistry

2

What is Chemistry?CuCl2 + NH3 What happens?

CuCl2soluble

Cu(OH)2(s)insoluble

Cu(NH3)4+2

soluble

3

Molecular

Interpretation

Cu+2 +2OH- Cu(OH)2

This course: understand & predict

Cu(OH)2 + 4NH3

Cu(NH3)4+2 + 2OH-

NH3 + H2O NH4+ + OH-

4

Chemistry

•the study of matter

•and the changes

matter undergoes

Focus:

•Underlying principles

•Molecular point-of-view

5

Classification of Matter

?? H2, cement, air, NaCl

Homo-geneous

Hetero-geneous

Mixture

Cmpd Element

Substance

Matter

“solution”

Physically

Separate

Chem. Separate.

Give examples of physical separation & chemical separation

Physical vs. Chemical Change

6

Physical change: no new substance is formed. e.g. melt, bend, cut, boil

Chemical change: new substance formed (always involves a chemical reaction).

e.g. Fe(s) + S(s) FeS(s)

7

Units of Measurement

Quantity SI Unit Others

Length m cm

Mass kg g

Time s

Temperature K oC

Metric Prefixes (know them!)

8

giga- G E9mega- M E6kilo- k E3deci- d E-1centi- c E-2milli- m E-3nano- n E-9pico- p E-12

How many grams is 14 pg?

9

Derived Units

How many cm3 in 1 m3 ?

For example: volume

10

Derived Units

Density = mass/volume

What is the mass of 17.4 mL

of ethanol (D = 0.798 g/mL)?

1 g H2O @ 4oC = 1 cm3 = 1mL

11

Temperature Scales

Kelvin

0

273

373

(Size of

degree is

the same)

Celsius

Motion

stops -273

Water freezes

Water boils 100

0

12

Temperature Scales

K = oC + 273.15

Try it:

1. 72oC = ? K

2. What is room temperature in oC?

3. If T changes by 37oC, what is

the change in Kelvin?

13

Scientific NotationAvogadro’s Number:

602,200,000,000,000,000,000,000

6.022 x 1023 particles

Mass of H atom:

0.00000000000000000000000166 g

1.66 x 10-24 g

14

Using Scientific Notation

N. x 10n

Add/subtract: n must be the same

Add 7.4 x 103 and 2.1 x 102

Multiply (divide): add (subtract) n’s

Divide 8.0 x 102 by 6.2 x 104

15

Significant Figures

•All non-zeros are sig.1.234 ?

•Zeros between non-zeros are sig.40,501 ?

How wide is this room?

16

Significant Figures

Zeros to left are NOT sig.0.000349 ?

Zeros to right are sig. if there is a decimal point.0.09030 ? 400 ?

17

Sig Figs in Calculations

Add/subtract: answer has same number of digits to right of decimal as the number with the fewest digits to right of decimal.

89.321 + 9.2 =

18

Sig Figs in Calculations

Multiplication/division: answer has the same number of sig figs as the number with least number of sig figs.

2.8 x 4.5039 =

19

Pure Numbers

How many sig figs in:

Five people?

2.54 cm = 1 in

20

Conversion Factors

5 apples = $1.00

5 apples$1.00

= 1

In Wegman’s, apples cost 5 for a dollar.

21

Conversion Factor

5 apples = 1 dollar

1 dollar = 100 pennies

10-2 m = 1 cm (or 1 m = 100 cm )

1 mole = 6.02 x 1023 atoms

1 g H2O = 1 mL H2O

22

Conversion Factors: Try it.

Convert radius of Ca atom (0.197 nm) to cm. (Hint: go through the base unit.)

Adult has 5.2 L of blood. What is volume in m3 ?

What is 65 miles per hour (mph) in m/s ? (1609 m = 1 mi)

23

Atomic Theory

Democritus vs. Dalton

500 BC 1808 AD

Pure

thought

Observation

+ Reason

Matter is not continuous!! Weird.

24

Dalton’s Atomic Theory

Elements composed of atoms

Atoms of element are identical*

Compounds: simple integers of

atoms

Chem Rxn: rearrangement, not

creation, of atoms

25

Dalton’s Theory

Law of Definite Proportions:

(mass of C and O in CO2)

Law of Multiple Proportions:

(CO and CO2)

Based on observations:

26

Subatomic Particles

The Electron

27

The Electron

Symbol: e-

Charge = -1 = -1.6 x 10-19C

Mass = 9.09 x 10-28 g

28

The Proton

Ernest

Rutherford

1910

29

The Proton

Charge = +1

Mass = 1.67 x 10–24 g

30

The Neutron

Mass ratio of He/H problem

Chadwick 1932

(nuclear chemistry)

Charge = 0

Mass = 1.67 x 10–24 g

(similar to proton)

31

Relative Proportions

Protons &

Neutrons

Electron cloud

32

Atomic Number & Mass Number

Mass number

(#p + #n)

Atomic number

(#p)

XA

Z

H1

1H

2

1H

3

1

33

Practice

How many p, n, e- in:

Hg ?200

80

-2 ion of Oxygen-16 ?

Hg-200 or

34

Periodic Table

Group/Family vs. Period

Metals/Nonmetals/Metalloids

Key Groups

1 or 1A: Alkali Metals

2 or 2A: Alkaline Earth Metals

17 or 7A: Halogens

18 or 8A: Noble Gases

35

Atoms, Molecules, Ions

Water Ammonia Methane

36

Systematic Names: Why?

Common Name Formula

Water H2O

Lime CaO

Lye NaOH

Potash K2CO3

Laughing Gas N2O

Baking Soda NaHCO3

Muriatic Acid HCl

37

Chemical Formulas

Molecular formula: Exact

numbers of atoms of each

element in a molecule.

Cl

H

CChloroform

CHCl3

38

Chemical FormulasEmpirical: smallest whole number

ratio of atoms in a compound.

Glucose: C6H12O6 CH2O

Hydrogen Peroxide: H2O2 HO

39

Molecular Compounds

Molecular cmpds: bonded nonmetals

Many molecular compounds

are binary (only 2 elements)

ending in ---ide.

HCl = hydrogen chloride

Organic compounds have a special naming system.

40

Molecular Compounds

Name is:

prefix-atom-prefix-atom-idemono 1

di 2

tri 3

tetra 4

penta 5

hexa 6

N2O dinitrogen monoxide

PCl3phosphorus trichloride

41

Molecular Compounds

SiCl4

AlCl3

H2O NH3 H2S

Try it!

42

Ionic Compounds

Ionic Compounds: usually

a positive metal ion +

negative nonmetal ion.

Binary: NaCl = sodium chloride

Ternary: BaCO3 = barium carbonate

43

IonsIon: Charged atom or group of

atoms (monatomic or

polyatomic; Cl- vs CO3-2)

chlorine atom: 17 p & 17 e-

chloride ion: 17 p & 18 e-

Cation: +positive Anion: -negative

44

Ionic Formulas

Sum of charges is zero.

e.g. Aluminum Oxide

Al2O3

“criss-cross”

2(+3) + 3(-2) = 0

Al3+ O2-

45

Ionic FormulasIonic compounds always use

empirical formulas.

Sodium Chloride: NaCl

Polyatomic Ions(know them!)

46

Memorize the ‘ate’ Ions

47

ate:C2H3O2

- CO32- C2O4

2-

ClO3- CrO4

2- Cr2O72-

MnO4- NO3

- PO43-

SCN- SO42- S2O3

2-

Rule: if the non-oxygen atom is in an odd Group, charge is -1 (except PO4

3-)

Figure Out the ‘hydrogen ---ate’ ions

48

CO32-

HCO3-

SO42-

HSO4-

Adding a hydrogen decreases the charge by 1.

Figure Out ‘ite’ ions

49

ClO4- perchlorate 1 more O

ClO3- know the ‘ate’

ClO2- chlorite 1 less O

ClO- hypochlorite 2 less O

Try it: name IO2-

Memorize Some Others

50

H3O+ Hg2

2+ NH4+

CN- O22- OH-

Have fun!

51

Ionic Compounds

Some metals form

more than one cation:

Fe+2 Fe+3

Classical Ferrous Ferric

Stock Iron(II) Iron(III)

FeCl3 = Iron(III) chloride

Common Metals with more than one charge (know them!)

52

Cu+ copper(I) Sn2+ tin(II)

Cu2+ copper(II) Sn4+ tin(IV)

Fe2+ iron(II) Hg22+ mercury(I)

Fe3+ iron(III) Hg2+ mercury(II)

Pb2+ lead(II) Au+ gold(I)

Pb4+ lead(IV) Au3+ gold(III)

53

Practice

NH4ClO3 calcium phosphate

PbO

cupric sulfite

iron(III) carbonate

54

Naming Acids

Acid: substance yielding

H+ (proton) in water.

HCl = hydrogen chloride

HCl(aq) = hydrochloric acid

55

Naming Acids

56

Binary Acids

Anion Example

_ide Chloride Cl-

Acid Example

hydro_ic acid hydrochloric acid

HCl

57

Oxyacids

Anion Example

_ite chlorite ClO2-

Acid Example

_ous acid chlorous acid

HClO2

58

Oxyacids

Anion Example

_ate chlorate ClO3-

Acid Example

_ic acid chloric acid

HClO3

59

Naming Acids

Anion Acidchloride HCl hydrochloric acid

hypochlorite HClO hypochlorous acid

chlorite HClO2 chlorous acid

chlorate HClO3 chloric acid

perchlorate HClO4 perchloric acid

60

Acids: Try It !!!

H2SO4

HF

HNO3

HCN

H2SO3

61

Have Some Fun

Read about naming bases,

and hydrates on your own.

Quiz on compound names and formulas.

62

Mass Relationships in Chemical Reactions

The Mole

63

Masses (not weights)

• proton = 1.673 x 10-24 g

• neutron = 1.675 x 10-24 g

• electron = 9.109 x 10-28 g

64

Atomic Mass Unit

1 amu 1/12 mass of C-12 atom

Mass of C-12 atom 12.00… amu

Mass of H 1 amu mass 1 p

Mass of O 16 amu

65

Average Atomic Mass

Periodic Table lists

weighted-average

atomic mass of the

naturally occurring

isotopes.

Determining Mass and Abundance of Isotopes.

66

Mass Spectrometry or Mass Spectroscopy

ion detector

vacuum pump

accel. field

ionizing e- beam

injection

Mass Spectrometry

67

1. Substance injected into vacuum tube.2. Particles are ionized by e- beam.3. Ionized particles are accelerated in elec. field.4. Particles are deflected by magnetic field

that can be varied in strength.5. Deflection depends on particle mass-to-

charge ratio, particle speed, and magnetic field strength.

6. Particles enter ion detector.

Strontium Example

68

Relative peak heights:• Sr-84 0.68• Sr-86 12.00• Sr-87 8.47• Sr-88 100.00

Total 121.15

Strontium Weighted Average

69

.68 121.15

x 84.0 amu = 0.47 amu

12.00121.15

x 86.0 amu = 8.52 amu

8.47121.15

x 87.0 amu = 6.08 amu

100.00121.15

x 88.0 amu = 72.64 amu

87.71 amu

Fractionalabundance

70

Average Atomic Mass

e.g. Chlorine has 2 isotopes:

75.77% Cl-35 @34.969 amu

24.23% Cl-37 @36.966 amu

(.7577)(34.969 amu) +

(.2423)(36.966 amu) = 35.453 amu

71

Average Atomic MassTRY IT: Of every 100 atoms of

Cu, 69.09 are Cu-63 (62.93 amu)

and the rest are Cu-65 (64.93 amu).

Calculate the mass of Cu as given

on the Periodic Table.

72

Molecular Mass

The sum of atomic masses in amu

Molecular mass of HNO3:

H: 1 x 1.008 amu = 1.008 amu

N: 1 x 14.01 amu = 14.01 amu

O: 3 x 16.00 amu = 48.00 amu

= 63.02 amu

TRY IT: What is mass of N2O4?

73

The Mole

Mole:

number of C atoms in

exactly 12 g of C-12.

“chemist’s dozen”

74

Avogadro’s Number

1 mole =

6.022 x 1023 particles

75

Avogadro’s Number

602,204,500,000,000,000,000,000

# of grains of sand in the world!

# of stars in the universe!

76

The Mole

1 mole C atoms = 6.02 x 1023 atoms

1 mole CO2 molecules = 6.02 x 1023

molecules

1 mole pencils = 6.02 x 1023 pencils

1 mole Br- ions = 6.02 x 1023 ions

77

Molar Mass

Molar Mass

mass (in g) of one mole

e.g. 1 mol C atoms

= 6.02 x 1023 atoms C

= 12.0 g C

78

(6 x 12.0 g) + (12 x 1.0 g) + (6 x 16.0 g)

= 180.0 g

Mass of one mole of glucose,

C6H12O6

Molar Mass

79

Molar Mass: Other Terms

Gram atomic mass:

gam of Fe is 55.9 g

Gram molecular mass:

gmm of CO2 is 44.0 g

Gram formula mass:

gfm of NaCl is 58.4 g

80

Molar Mass

Mass (g) # Particles

Molar mass

g/mol

6.02 x 1023

part./molMoles

Conversion Factors

81

Mole Practice

1. What is mass in g of one C-12 atom?

2. How many grams in one amu?(Hint: use carbon as an example.)

3. What is the molar mass of C6H8O6 ?

4. How many H atoms in 25.6 g urea?

(NH2)2CO

82

Percent Composition

% by mass of each element in a cmpd

=n x molar mass of element

molar mass of cmpdX 100

where n = element subscript

CO2

83

Percent Composition

What is % of O in CO2?

=n x molar mass of element

molar mass of cmpdX 100%O

=2 x 16.00

44.01X 100

= 72.71%

84

Percent Composition

What is the mass of Cu (in kg)

in 3.71 x 103 kg of

chalcopyrite (CuFeS2)?

TRY IT !!

85

Percent Composition

If % composition is known,

a compound’s empirical formula

can be determined.

86

Empirical Formula

What is empirical formula of compound:

24.75% K 34.77% Mn 40.51% O?

Divide by smallest number KMnO4

assume 100 g

24.75 g 34.77 g 40.51g

.6330 mol .6330 mol 2.532 mol

go to moles

87

Empirical Formula

Sometimes it is not so easy!

What is emp. formula of Vitamin C?

C = 40.92%

H = 4.58%

O = 54.50%

Try it.

88

Empirical Formulas of C-H-O Compounds

Heat

O2

H2O CO2absorber absorber

Sample

89

Example

If combustion of an 11.5 g sample

of ethanol produced 22.0 g of CO2and 13.5 g of H2O, what is the

empirical formula of ethanol?

Try It

90

If combustion of a 21.4 g sample

of an unknown C-H-O compound

produced 29.5 g of CO2 and 24.1 g

of H2O, what is the empirical

formula of the compound?

91

Molecular Formulas

If the empirical formula and

approximate molar mass are

known, molecular formula can

be determined.

92

Molecular Formulas

Problem: A compound contains

only 1.52 g N and 3.47 g O and

has molar mass 95. What is it’s

molecular formula?

93

Molecular Formulas

N: 1.52 g 0.108 mol

O: 3.47 g 0.217 molNO2

NO2 has empirical molar mass of

46.02 g compared to molar mass 95 g.

Molar mass

Emp.mol.mass= 95 g

46.02 g N2O4

94

Molecular Formulas

Problem: 200 g of a compound

contains 77.34 g C, 32.44 g H and

90.22 g N and has molar mass .

What is it’s molecular formula?

TRY IT !!

95

Chemical Equations

Chemical Reaction:

Atoms & molecules

react to form

new substances.

A chemical reaction is represented

by a chemical equation.

96

Evidence of Reactions

• Macroscopic evidence of chemical

reactions:

Heat Light

Gas Solid (precipitate)

Color change

• Chemical change involves creation of

new substances with new properties,

usually accompanied by a noticeable

energy change.

97

Chemical Equations

H2 + O2 H2O

But mass must be conserved!

Hydrogen burns to form water.

98

Chemical Equations

Balanced: 2H2 + O2 2H2O

Also show state:

2H2 (g) + O2 (g) 2H2O (l)

2H2 O2 2H2O

+

99

Balancing Equations

Identify reactants & products

Write correct formulas first

Change only coefficients

2H2 + 1O2 2H2O

Double check

100

Balancing Equations

Potassium chlorate decomposes

upon heating to yield oxygen

and potassium chloride:

KClO3 KCl + O232 2

101

Balancing Equations

Complete combustion of

Ethane (C2H6) yields carbon

dioxide and water.

Try it.

102

Types of Reactions

• Combination

• Decomposition

• Single Replacement

• Double Replacement

• Combustion

103

Combination

2 Br2+ 2 NaBrNa

104

Decomposition

CaCO3 CaO + CO2

105

Single Replacement

Cu AgNO32 Cu(NO3)2 2+ Ag

Cu + ZnCl2 No Reaction

+

106

Double Replacement

BaCl2 + H2SO4 2BaSO4+ HCl

3CrCl3(aq) + NaOH(aq)

3Cr(OH)3(s) + NaCl(aq)

Soluble Ionic Compounds

107

Ionic compounds that are soluble in

water actually exist as separated ions.

CrCl3(aq) is actually:

Cr+3(aq) + 3Cl-(aq)

Net Ionic Equations

108

CrCl3(aq) + 3NaOH(aq)

Written as a net ionic equation is:

Cr(OH)3(s) + 3NaCl(aq)

Cr3+(aq) + 3OH-(aq) Cr(OH)3(s)

Cl- and Na+ are “spectator ions”

109

Combustion

C2H5OH + 3O2 2CO2+3H2O

110

Stoichiometry

The quantitative study of reactants

and products in a chemical reaction.

Interpret coefficients in chemical

equation as molecules or moles.

2H2 + O2 2H2O

2 mol H2 + 1 mol O2 2 mol H2O

111

Stoichiometry

2H2 + O2 2H2O

2 mol H2 1 mol O2 2 mol H2O

means stoichiometrically

equivalent to

112

Stoichiometry

2Li + 2H2O 2LiOH + H2

How many moles of H2 are

obtained from complete

reaction of 6.23 moles of Li ?

6.23 mol Li1 mol H2 = 3.12 mol H2x2 mol Li

113

Stoichiometry

2Li + 2H2O 2LiOH + H2

How many g H2 from 81 g Li ?

81 g Li 1 mol Li6.9 g Li

1 mol H22 mol Li

2.0 g H21 mol H2

= 12 g H2

x x

x

114

Stoichiometry

You try it!A key reaction in smog formation is:

2NO + O2 2NO21. How many g NO2 formed by

complete reaction of 1.44 g NO ?

2. How many molecules of NO2formed from 0.00052 mg of NO ?

115

LimitingReagents

How many cheese sandwiches

can be made from:

25 slices of cheese

12 tomatoes

7 heads of lettuce

4 slices of bread

Bread is the “limiting reagent”

116

Burn S in F2: S + 3F2 SF6

What if 4 mol S added to 20 mol F2 ?

Since 1 mol S 3 mol F2

Then 4 mol S 12 mol F2

Thus S is limiting and F2 is excess.

LimitingReagents

117

Limiting Reagents

Urea is made by reacting

ammonia and carbon dioxide:

2NH3 + CO2 (NH2)2CO + H2O

If 641 g NH3 reacts with 1140 g CO2•Which chemical is limiting ?

•How much urea is formed ?

•How many g of excess reagent is left ?

118

2NH3 + CO2 (NH2)2CO + H2O

641 g 1140 g

Start with NH3 & convert to urea:

641 g NH31 mol NH317.0 g NH3 2 mol NH3

1 mol urea

1 mol urea60.1 g urea = 1130 g urea

x x

x

119

2NH3 + CO2 (NH2)2CO + H2O

641 g 1140 g

Start with CO2 & convert to urea:

1140 g CO21 mol CO244.0 g CO2 1 mol CO2

1 mol urea

1 mol urea60.1 g urea = 1560 g urea

x x

x

1130 g

120

2NH3 + CO2 (NH2)2CO + H2O

641 g 1140 g

1130 g

1560 g

Since NH3 yields less urea, it is

the limiting reagent, and 1130 g

urea will be produced.

CO2 is the excess reagent.

121

To calculate the excess amount of

CO2, start with limiting reagent, NH3.

641 g NH3

1 mol NH317.0 g NH3 2 mol NH3

1 mol CO2

1 mol CO2

44.0 g CO2 = 829 g CO2

2NH3 + CO2 (NH2)2CO + H2O

1140 g – 829 g = 311g excess CO2

x x

x

122

LimitingReagents

You try one!

Al reacts with iron(III) oxide highly

exothermically. If 124 g Al reacts with

601 g iron(III) oxide:

•Write a balanced equation.

•How much aluminum oxide is formed ?

•How much excess reagent is left ?

123

Reaction Yield

from limiting reagent calculation

from experiment

% Yield =Actual Yield

Theoretical Yieldx 100%

124

Reaction YieldAmmonia is produced by the Haber

process by reacting nitrogen gas and

hydrogen gas:

N2 + 3H2 2NH3

What is the % yield if 254 g NH3

are produced by reacting 246 g of N2

with 62 g of H2 ?

125

N2 + 3H2 2NH3246 g 62 g 254 g (actual)

1. Determine limiting reagent

246 g N228.0 g N2

1 mol N2

1 mol N2

3 mol H2

2.02 g H2

1 mol H2

= 53.2 g H2

Thus N2 is the limiting reagent.

x x

x

126

2. Calculate theoretical yield of NH3

246 g N228.0 g N2

1 mol N2

1 mol N2

2 mol NH3

17.0 g NH3

1 mol NH3= 299 g NH3

limiting

N2 + 3H2 2NH3246 g 62 g 254 g (actual)

x x

x

127

3. Calculate % yield of NH3

vs.

299 g theo.

=254 g 299 g

x 100%

= 85.0 % yield

N2 + 3H2 2NH3246 g 62 g 254 g (actual)

% Yield =Actual Yield

Theoretical Yieldx 100%

128

TRY IT !!

Methanol is made by reacting hydrogen

gas and carbon monoxide gas.

Write balanced equation.

Calculate the % yield if 3.57 x 104 g of

CH3OH is produced by reacting 68.5 kg

carbon monoxide and 8.60 kg hydrogen.