A.P. Chemistry Unit 1

1
A.P. ChemistryUnit 1

• Chapters 1, 2, 3

2
Why take Chemistry?
3
Purpose of AP Chem at CG?
4
Introductory terms

• Table groups

5
Definition Review Try to fill in

• 1. Chemistry
• 2. Chemical Property
• 3. Physical Property

6

• 4. Intensive property
• 5. Extensive property

7
Terms to recognize

• 6. Element
• 7. ATOM
• 8. Compound
• 9. Molecule
• 10. Formula Unit

8
TERMS

• 11. Pure substances
• 12. Chemical reaction
• 13. Physical change
• 14. Mixture
• 15. Homogeneous
• 16. Heterogeneous
• 17. Chemical symbol
• 18. Chemical formula

9
More terms

• 19. Chemical Equation
• 20. Reactants
• 21. Products
• 22. Coefficients
• 23. Subscripts
• 24. Matter
• 25. Law of conservation of energy
• 26. Law of conservation of matter

10
27. Name Steps in the Scientific Method

1. Asking a question
2. Forming hypotheses
3. Researching previously found information
4. Designing experiments
5. Conducting experiments / collecting data
6. Determining variables dependent and independent
7. Organizing and analyzing data
8. Stating conclusions
9. Considering sources of error
10. Communicating results
11. Planning future experiments

11
Classification of Matter
12
Atoms Molecules Elements Compounds?
1. Aluminum foil 6. oxygen gas 2. Carbon
Dioxide 7. sodium chloride 3. Zinc
8. water 4. Graphite
9. chlorophyll 5. Helium 10.
nitrogen
13
1. E, A 6. E, M 2. C, M 7. C, ?M 3.
E, A 8. C, M 4. E, A
9. C, M 5. E, A 10.
E, M
14
Solid, Liquid, Gas(room temperature)

1. Ammonia
2. Gasoline
3. Graphite
4. H2O
5. Shaving cream
6. Aluminum
7. Ice cream
8. Helium
9. Bromine
10. Sugar

15
Solid, Liquid, Gas(room temperature)

1. G
2. L
3. S
4. L
5. L G ?
6. S
7. L
8. G
9. L
10. S

16
Pure Substance or Mixture?

1. Water
2. Hydrogen
3. Salt
4. Tea
5. Sodium
6. Sugar
7. Iron oxide
8. concrete

1. Raisin cookie
2. Gatorade

17
Pure Substance or Mixture?

1. P
2. P
3. P
4. M
5. P
6. P
7. P
8. M

1. M
2. M

18
Chemical or Physical Change?

1. Dog is groomed
2. Child gets taller
3. Gas forms when Baking soda is mixed with vinegar
4. Pencil is sharpened
5. Paper burns
6. Leaves turn color
7. Ice melts

1. Sugar dissolves in water
2. Cookie bakes in oven
3. Cake mix is combined with water.

19
Chemical or Physical Change?

1. P
2. C
3. C
4. P
5. C
6. C
7. P

1. P
2. C
3. P (?)

20
Metric fundamental units

• Kilogram
• Meter
• Second
• mole
• Kelvin
• Coulomb

21
Metric units for?....

1. Length
2. mass
3. area
4. volume
5. Density
6. Weight and Force

• 7. Energy
• 8. particles
• 9. Pressure
• 10. Current
• 11. Potential
• 12. Power

22
fundamentals?....

1. yes
2. yes
3. yes
4. yes
5. yes
6. Kgm/s2
7. Kgm2/s2

• 8. yes
• 9. Kgm/s2
• 10. C/s
• 11. Kgm2/Cs2
• 12. Kgm2/s3
• 13. yes

23
Metric units?

• Fundamental?
• Can you convert?

• Deci
• Centi
• Milli
• Kilo
• Nano
• Giga
• Micro
• Mega

24
AP Chemistry

• Chapter 2
• Atoms,
• Molecules
• and Ions

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John Dalton

• English school teacher
• 1766-1844
• Author of the Modern Atomic Theory

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• loved studying the weather
• saw the applications for chemistry in his ideas
  about the atmosphere.
• Proposed Atomic Theory 1803
• Dalton's theory was presented in New System of
  Chemical Philosophy (1808-1827).

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John Dalton

• Was colorblind
• Daltonism
• His eyes were used to prove it is a brain disorder

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Postulates of Daltons atomic theory

• Matter is made of atoms which stay the same
  during a chemical change.

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An element is a substance made of one type of
atom, each of which has the same properties.
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A compound is matter made of two or more elements
combined in fixed proportions.
32
A chemical reaction involves rearrangement of
atoms into new substances, but no loss or gain of
any atoms.
33
Law of definite proportions

• Molecules of the same compound are all the same.

34
The same elements can make many different
compounds.
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Atomic Structure
36
Modern Theory Says

• Atoms are made of
• Protons
• neutrons
• electrons
• a small dense nucleus of protons and neutrons and
  surrounding electrons.

37
Famous experiments leading to this view of the
atom

• Joseph John Thomson
• The Cathode Ray Tube Experiment (Cambridge
  1897)
• discoverer of the electron, Nobel Prize Winner.

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The Cathode Ray Tube and Thomsons Plum Pudding
atom
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Robert Millikan

• The Oil Drop Experiment (USA, 1909)
• measured charge on an electron
• calculated the mass of electrons.

40
Millikans Experiment

• Used calculation to determine the charge on each
  suspended droplet
• All were multiples of 1.6x10-19 Coulomb

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Ernest Rutherford

• Studied Gold Foil Experiments by Geiger and
  Marsden (1911)
• atoms are mainly empty space with a small,
  massive, dense, positively charged nucleus.
• Also discovered the proton.

42
Gold Foil Experiment Video?

• Led to the idea of a nuclear atom
• Led to the idea that atoms are mainly made of
  empty space.

43
Modern Atomic theory

• the atom consists of 40 fundamental particles.
• The electron is a quark, but the proton and the
  neutron are not.

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ISOTOPES

• Atoms of the same element that have different
  masses (different of neutrons in the nucleus).

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ISOTOPES TO KNOW

• The three isotopes of hydrogen
• Protium H-1
• Deuterium H-2 Water made with this is called
  heavy water.
• Tritium H-3 and is radioactive!

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AZ symbols

• A mass , nuclear particles, PN
• Z atomic , nuclear charge, P
• Neutrons ?
• X element symbol

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Write an AZ nuclide symbol for

• strontium 90
• silicon 30
• radon 226

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Atomic number vs mass number

• Protons nuclear charge
• vs
• protons neutrons number of nuclear particles

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• Atomic weight vs atomic number
• Weight of protons, neutrons, electrons
• Vs
• Proton number nuclear charge

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Atomic particles
51
Atomic Weights

• All relative to the Carbon-12 isotope
• Carbon-12 is the mass standard
• One mole of carbon 12 12 grams

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Atomic Weights

• Represent average mass of isotopes and their
  percent composition in nature.

53
Measuring atomic weight

• Units are amu
• Atomic mass units
• One mole of amu 1 gram

54
Measuring the weight of atoms

• By Mass spectrometry.
• An unknown is compared to a known sample ( the
  standard).
• Particles are accelerated through a gas and bent
  by a magnetic field.
• The curvature of their pathway is measured and
  mass is calculated.
• FBvq.

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Mass Spec and examples

• See worksheet

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Example

• Given data for Chromium, determine its average
  atomic weight.
• Isotope Mass Frac. Abundance
• Cr-50 49.9461 0.0435
• Cr-52 51.9405 0.8379
• Cr-53 52.9407 0.0950
• Cr-54 53.9389 0.0236

51.9959 u
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• Coppers two isotopes are mass numbers 63 and
  65 What percent abundance is each if the
  average atomic mass is 63.5?
• No calculator!

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• What is the mass in grams of a mole of titanium
  atoms?

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• What is the mass of one atom of Ca-40?
• Answer in g and amu

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History of the periodic table
61
Dmitri Mendeleev (Russia)

• Wrote periodic law
• Chart based on atomic weights.

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John Newlands, Great Britain

• law of octaves.
• ridiculed b/c of inconsistencies

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Julius Lothar Meyer, Germany

• periodic law around the same time as Mendeleev.
• not credited as Mendeleev b/c ?? Predictions
  undiscovered..

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Henry Mosely

• 1887-1915
• Studied with Rutherford,
• measured nuclear charge / atomic number of
  elements.

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Henry Mosely

• Reordered periodic table by at. and it is better

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Henry Mosely

• Volunteered for service in WWI. Was a signal
  officer for the British Army and killed in action
  at Gallipoli in 1915.

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The Periodic Table

• Periods
• 7 horizontal Rows
• Families / Groups 18 vertical columns

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Sketch and label
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Binary

• Made of two elements.
• Examples?

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Molecular

• Made of molecules
• Nonmetal atoms making compounds
• Covalent shared electrons.

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Ionic

• Made with a metal ion or ammonium and an anion..

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Ionic or Molecular?

• CuBr2
• H2O
• Li2O
• NH4I
• RbClO

• Na2CO3
• C4H10
• MgSO4
• Al2 (SO4)3
• KF

73
Making ions

• Cations vs anions
• Only changes in electrons

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Formula unit vs molecule

• Ionic vs not ionic
• Low ratio of ions vs formula of particle

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Ionic charges

• Metals vs nonmetals
• Metals make cations
• Nonmetals make anions

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Ion names

• Metals vs nonmetals
• The ide ending for nonmetals

77
Nomenclature practice time

• Use the ion sheets
• Use the flowchart
• Follow the rules
• Practice!

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Name these Compounds

1. HClO
2. HF
3. H2O2
4. PbCrO4
5. LiC2H3O2
6. CO

• CdI2
• N2O5
• CuSO4
• SrBr2
• H3PO4
• Ca(NO3)2

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Name these Compounds

1. Hypochlorous acid
2. Hydrofluoric acid
3. Dihydrogen dioxide
4. Lead II chromate
5. Lithium acetate
6. Carbon monoxide

• Cadmium iodide
• Dinitrogen pentoxide
• Copper II sulfate
• Strontium bromide
• Phosphoric acid
• Calcium nitrate

80
Write Formulas for these Compounds

1. Cobalt III hydroxide
2. Barium phosphate
3. Magnesium chloride
4. Aluminum iodide
5. Sodium oxide
6. Perchloric acid

1. Nitrous acid
2. Oxalic acid
3. Hydrobromic acid
4. Diphosphorous pentoxide
5. Dinitrogen monosulfide
6. Silver carbonate

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Write Formulas for these Compounds

1. Co(OH)3
2. Ba3(PO4)2
3. MgCl2
4. AlI3
5. Na2O
6. HClO4

1. HNO2
2. H2C2O4
3. HBr
4. P2O5
5. N2S
6. Ag2CO3

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Organic Chemistry Intro

• carbon based with hydrogen and oxygen mainly.
• Hydrocarbons C and H only
• Alkanes Hydrocarbons with only singly bonded
  carbon atoms.
• 1st rule C bonds 4 times
• Alcohols Have an -OH functional group
• Isomers Molecules with the same formula but
  different structures
• R general symbol for a carbon chain

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ORGANIC NOMENCLATURE

• If R __Carbon, it is called
• 1 Carbon meth(yl)
• 2 Carbons eth(yl)
• 3 Carbons prop(yl)
• 4 Carbons but(yl)
• 5 Carbons pent(yl)
• 6 Carbons hex(yl)
• 7 Carbons hept(yl)
• 8 Carbons oct(yl)
• 9 Carbons non(yl)
• 10 Carbons dec(yl)

84
Draw and decide if isomers exist.

• Propane
• 1-butanol
• pentane

85
Name each. If they have any isomers, name one of
them too.

• CH3CH2CH2CHOHCH2CH2CH2CH2CH3
• CH3CH2CH2CH2CH2CH3
• CH3OH

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STOICHIOMETRY

• Using a balanced equation to make theoretical
  predictions.
• Beqs show COM!

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Simple reactions

• Combination (synthesis)
• Smaller reactants make a more complex product.

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Simple Reactions

• Decomposition
• A reactant forms simpler products

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Classify a-e

• Combination
• a c
• Decompostion
• b d e

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Combustion

• Hydrocarbon and oxygen gas reactants
• CO2 and H2O products

91
Combustion write reactions

• Octane
• 2-Hexanol
• 3-Heptanol

92
weight Units

• Molecular or formula weight u or amu
• Molar mass g/mol

93
Formula weight ormolecular weight of

• A. carbon dioxide
• 44.0 amu
• B. water
• 18.0 amu
• C. oxygen gas
• 32.0 amu
• D. Table salt
• 58.5 amu

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Molar Mass of

• A. carbon dioxide
• 44.0 g/mol
• B. water
• 18.0 g/mol
• C. oxygen gas
• 32.0 g/mol
• D. Table salt
• 58.5 g/mol

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Mole

• 6.02x1023
• Used to count particles
• A mole of miniature marshmallows would cover the
  USA to a depth of 600 miles.

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Mole Relationships

• 1 mole 6.02x1023 at or mc or fu
• 1 gram 6.02x1023 amu
• 1 mole S at wt of any formula (g)
• 1 mole gas (STP) 22.4 L

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Practice Mole Calculations

• 1. 0.0365 g
• 2. 1.0x1024 atoms
• 3. 23.6 mol
• 4. 1.1x10-21 amu
• 250. Liters
• 42.1C, 6.4H, 51.5O
• 51.9 N

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Example Problem

• Determine the empirical and molecular formula of
  a compound found by combustion to contain 39.9
  Carbon, 6.7 hydrogen and 53.4 oxygen. The
  molecular weight of the compound is 120 amu.
• Empirical CH2O
• Molecular C4H8O4

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Problem Solving

• Pretend to have 100 grams . Or.
• If grams are given, use them!
• Change grams to moles for each element.
• Look at mole ratios to work out lowest whole
  number subscripts.
• Use known molar mass to find molecular formula
  with integer multiplier.

100

• A sample of a compound is found to contain 17.5 g
  Na, 39.7 g Cr and 42.8 g O. What is its
  empirical formula?
• Na2Cr2O7

101

• Sorbic acid is added to food as a mold inhibitor.
  Its composition is 64.3 C, 7.20 H , and the
  rest oxygen. Its molecular weight is 112 u.
  What is the molecular formula for sorbic acid?
• C6H8O2

102
Challenge Problem, AP levelAn organic acid
contains only C, H and O. A 12.72 mg sample of
the acid is completely burned in oxygen. It
yields 18.63 mg of carbon dioxide and 7.62 mg of
water. What is the mass percentage of each
element in the organic acid? Whats the
empirical formula?
103

• Review Write an equation for the formation of
  carbon dioxide from its elements.
• C O2 ----gt CO2
• How many grams carbon are needed to produce 150.
  grams of carbon dioxide?
• 40.9 grams

104
Problem Solving

• Stoichiometry uses the balanced equation ratios.
• Balanced equation coefficients are about particle
  to particle ratios.
• Coefficients mean moles or atoms/molecules/formula
  units
• Or (by thinking of Avogadros Hypothesis)
  volumes(any units) for gaseous substances.

AMEDEO AVOGADRO
105
Stoichiometry

• Write an equation for the reaction of sodium in
  water.
• 2Na 2H2O ---gt 2Na 2OH- H2
• How many molecules of hydrogen gas are produced
  by the reaction of 0.25 grams of sodium METAL?
• 3.3x1021 molecules

106
Stoichiometry

• Write the equation for the dehydration of ethyl
  alcohol and butanoic acid into ethyl butyrate, an
  ester.
• C2H5OH C3H7COOH ---gt C3H7COOC2H5 H2O
• How many grams of water can be made from
  8.22x1023 molecules of ethyl alcohol (ethanol)?
• 24.6 grams

Ethyl butyrate is the odor of pineapples.
107
Limiting Reagents and Percent Yield

• Write an equation for the synthesis of aluminum
  chloride.
• 2Al 3Cl2 ---gt 2AlCl3
• If 3.00 g Al react with 13.0 g Cl2, how much
  AlCl3 can be produced?

108
Use an IRF box
2 Al 3Cl2 ------gt 2AlCl3
13 g 0.183 moles 0 moles


3 g 0.111 moles
I Initial moles R reacted moles F final moles
109

• Determine the limiting reagent.
• Use initial moles compared to how many are
  required for each reaction.
• Low number limits the process.

110
Use an IRF box
2 Al 3Cl2 ------gt 2AlCl3
13 g 0.183 moles 0 moles
Limiting reagent!!

3 g 0.111 moles
I Initial moles R reacted moles F final moles
111
Use an IRF box
2 Al 3Cl2 ------gt 2AlCl3
13 g 0.183 moles 0 moles
All 0.111 moles used
0 moles left over
3 g 0.111 moles
I Initial moles R reacted moles F final moles
112
Fill in R Row

• Mole ratios in R row must match the reaction
  coefficient ratios.
• The next coefficient divided by LR coefficient,
  multiplied by limiting moles ---gt R moles

113
Use an IRF box
2 Al 3Cl2 ------gt 2AlCl3
13 g 0.183 moles 0 moles
0.111 moles 0.167 moles 0.111 moles

3 g 0.111 moles
I Initial moles R reacted moles F final moles
114
Fill in F row

• Subtract for reactants, add for products
• Once the box is filled in with moles, any
  question can be answered.

115
Use an IRF box
2 Al 3Cl2 ------gt 2AlCl3
13 g 0.183 moles 0 moles
0.111 moles 0.167 moles 0.111 moles
0 moles 0.016 moles 0.111 moles
3 g 0.111 moles
I Initial moles R reacted moles F final moles
116
Now solve the problem!

• If 3.00 g Al react with 13.0 g Cl2, how much
  AlCl3 can be produced?
• 14.8 grams
• If 12.0 grams is recovered, what is the percent
  yield?
• 81.0

117
Problem Solving

• Thinking of the balanced equation as a recipe
  might help.
• Determine how many times the reaction recipe
  can be carried out with each amount of moles.
• Reactant that can make the fewest batches is
  the limiting reagent.

118
Limiting Reagent and Percent Yield

• Write an equation for the synthesis of lithium
  hydroxide from lithium oxide and water.
• Li2O H2O ---gt 2 LiOH
• If 42.0 grams lithium oxide react with 20.0 grams
  water, how much LiOH can be produced?
• 53.1 grams
• If 45.0 grams are obtained, what is the percent
  yield of the experiment?
• 84.7

Lithium metal
119
Limiting Reagents and Yield

• Write an equation for the preparation of
  hydrocyanic acid and water from ammonia, methane
  and oxygen.
• 2NH3 3O2 2CH4 --gt 2HCN 6H2O
• How many grams of HCN can be obtained from the
  reaction of 25.0 grams ammonia, 75.0 grams oxygen
  and 25.0 grams methane?
• 39.7 grams
• If 11.0 grams HCN is obtained, what is the
  percent yield of the reaction process?
• 27.7

120
Thats the end of Ch. 3!

• Mole day is coming SOON!
• Make the moleata!
• Talk to all other classes about donations for
  mole day?

121
Mole Day

• CG Chemistry

122
T-Shirt Ideas
123
T-shirt Ideas
124
Build the Moleata!
125
A.P. Chemistry

• Chapter 4
• Chemical Reactions An Introduction

126
I. Ionic Solutions

• Deionized water vs. tap water, bath water, lake
  water, ocean water?
• Ions!

127
Svante Arrhenius
128

• Svante August Arrhenius b. Uppsala, Sweden,
  February 19, 1859
• son of Svante Gustaf and Carolina
• educated at the Cathedral school Showed an
  aptitude in mathematics and physics.
• 1876 University of Uppsala mathematics,
  chemistry and physics.
• 1881 Stockholms Academy of Sciences.

129

• twice married - in 1894 to Sofia Rudbeck, (one
  son) and 1905 to Maria Johansson (one son and two
  daughters)
• died at Stockholm, October 2, 1927, and is buried
  at Uppsala.

130
Ionic solutions are .....

• electrolytic
• Capable of conducting electricity
• Many ionic solids are electrolytic in water.
• the ions electrolytes (particles that conduct
  electricity)

131
What ions are found in a solution of...

• KOH?
• CaCl2?
• (NH4) 2SO4?
• Write equations for what these compounds do in
  water.

132

• KOH (s) --gt K (aq) OH- (aq)
• CaCl2 (s) --gt Ca2 (aq) 2Cl- (aq)
• (NH4) 2SO4 (s) ---gt 2NH4 (aq) SO42- (aq)

133
Nonelectrolytic substances

• Make no ions (electrolytes) in solution.
• Examples
• sucrose, C12H22O11
• methanol, CH3OH
• urea, NH2CONH2
• antifreeze, HOC2H4OH
• All are molecular

134

• Dissolve but do not make ions
• C12H22O11 (s)----gtC12H22O11 (aq)
• CH3OH(l) -----gtCH3OH (aq)
• NH2CONH2 (s)-----gtNH2CONH2 (aq)
• HOC2H4OH (l)-----gtHOC2H4OH (aq)

135
Strong and Weak Electrolytes

• refers to degree () of ionization of solute.
• Acids and bases are described as strong or weak.

136
Strong vs. Weak Acids

• Strong acids ionize 100
• Weak acids ionize only partially

137
Students must Know the 6 strong acids

• nitric
• perchloric
• sulfuric
• hydrochloric
• hydrobromic
• Hydroiodic
• Write equations showing what each strong acid
  does in water.

138
The Strong Acids in Water

• Nitric HNO3 ---gt H NO3-
• Perchloric HClO4 ---gt H ClO4-
• Sulfuric H2SO4 ---gt 2H SO42-
• Hydrochloric HCl ---gt H Cl-
• Hydrobromic HBr ---gt H Br-
• Hydroiodic HI ----gt H I-

139
Dilute vs. Concentrated Solutions

• refers to the amount dissolved per volume of
  solution.
• dilute solutions small amount dissolved.
• Concentrated more dissolved

140
Writing Chemical Equations

• can be done
• A. molecularly show the whole mixture.
• B. Ionically
• complete ionic indicates any electrolytes in mix
• Net ionic only shows species that changed.
• AP CHEM requires net ionic equation writing.

141
I. Precipitation Reactions

• two aqueous solutions are mixed and one of the
  products is insoluble.
• Aqueous Dissolved in Water.
• Precipitate insoluble species.
• Spectator Ions in the mixture but do not take
  part in a reaction.

142

• SOLUBLE OR NOT?
• Know Solubility rules of ionic compounds.....
  page 136.
• Knowing the solubility song / chart helps.

143
Examples Write the net ionic reactions!

• 1. Potassium Chloride and Silver nitrate react
  in aqueous solution.
• Net ionic Cl- Ag ----gt AgCl
• 2. Ammonium sulfate and calcium chloride react
  in aqueous solution.
• Net ionic Ca2 SO42- ----gt CaSO4
• 3. Sodium carbonate and copper II sulfate react
  in aqueous solution.
• CO32- Cu2 -----gt CuCO3

144
More Practice. What will happen when the
following mix?

• NiCl2 and Na3PO4
• NaCl and Fe(NO3)2
• Al2(SO4)3 and KOH
• Pb(C2H3O2)2 and NH4Cl

145
Combustion burns in O2

• hydrocarbons
• Products always CO2 and H2O(l)
• exothermic
• Heat makes H2O vaporize.

146
Equations to Balance

• Recognize some organic alkanes and alkenes,
  alkynes and alcohols.
• Octane
• 2-Hexene
• 1-butyne
• 3-Heptanol

147
Acid-Base Basics

• Taste, feel ?
• pH ranges ?
• Ions they make in solution ?
• Household examples of each
• Page 139

148
The Strong Acids in Water

• Nitric HNO3 ---gt H NO3-
• Perchloric HClO4 ---gt H ClO4-
• Sulfuric H2SO4 ---gt 2H SO42-
• Hydrochloric HCl ---gt H Cl-
• Hydrobromic HBr ---gt H Br-
• Hydroiodic HI ----gt H I-

149
The Strong Bases in water

• LiOH LiOH ---gt Li OH-
• NaOH NaOH ---gt Na OH-
• KOH KOH ---gt K OH-
• Ca(OH)2 Ca(OH)2 ---gt Ca2 2OH-
• Sr(OH)2 Sr(OH)2 ---gt Sr2 2OH- Ba(OH)2
  Ba(OH)2 ---gt Ba2 2OH-

150
Acid and Base Definitions

• Arrhenius
• Acids contain hydrogen and make hydrogen ions in
  water. Bases contain hydroxide and make hydroxide
  ions in water.
• Bronsted and Lowry acids and bases Proton donors
  and proton acceptors.

151
Indicators to know

• Indicators are molecules that change colors at
  different pH levels.
• phenolphthalein
• Litmus
• Methyl red
• Others?

152
Reactions of Acids/Bases

• Learn to Write Net ionic equations for many
  examples!

153
Neutralization reactions

• acid base ? water salt.
• salt metal cation combined with an anion (often
  soluble in water)
• General Equation
• HA MOH -----gt H2O MA
• make examplesto net ionic

154
Carbonates and Acids

• Some Salts acids ? gases.
• Carbonates Acids ----gt carbon dioxide, water
  and a salt.
• General Equation
• MCO3 HA ----gt H2O CO2 MA
• Write examples.. To net ionic.
• Demo and test for gas Acetic acid and sodium
  (bi)carbonate

155
Sulfites and Acids

• Sulfites Acids -----gt sulfur dioxide, water
  and a salt.
• General
• MSO3 HA -------gt H2O SO2 MA
• examples Write net reactions.
• SO2 gas. Stinky, irritating

156
Sulfides and Acids

• A sulfide reacts with an acid to produce hydrogen
  sulfide gas and a salt.
• General
• MS HA --gt H2S MA
• examples
• test for gas?

157
One Base Reaction

• Decompostition of Hydroxides
• MOH? MO H2O
• Examples..

158
Oxidation and Reduction

• Aka single displacement, synthesis or
  decomposition reactions.
• electrons exchanged / atoms change oxidation
  states (charges).

159
Define Oxidation and Reduction

• Oxidation the loss of electrons
• Reduction The gain of electrons
• LEO says GER

160
Rules for deciding Oxidation State

• Elements alone have an ox. state of zero.
• H is always 1, unless it is with a metal as a
  hydride.
• O is always 2- unless it is in a (rare) peroxide
  H2O2, Li2O2, K2O2, or Na2O2.
• All common /main group metals keep their periodic
  pattern charge.
• Other semi and nonmetal elements oxidation
  states are determined last.
• Sum of ox. States charge on species.

161
Tell each elements Oxidation State

1. NH4Cl
2. Cl2O
3. N2O
4. NO
5. NO2
6. P2O5
7. KMnO4
8. Fe2(SO4)3
9. Na2C2O4

1. H3PO4
2. KNO3
3. Ca(NO2)2
4. BrO2 -
5. BrO3-
6. BrO4 -
7. CH4

162
Types of redox reactions

• Synthesis / Combination reactions
• Decomposition Reactions
• single replacement reactions
• Combustion reactions
• Other complex reactions

163
What is oxidized, whats reduced?

• Copper nitrate solution reacts with zinc metal to
  make aqueous zinc nitrate and metallic copper.
• Lithium metal reacts with a cobalt II chloride
  solution to make metallic cobalt and aqueous
  lithium chloride.

164
What is oxidized, whats reduced?

• Ca O2 ---gt CaO
• HgCl2 ----gt Hg Cl2

165
Will redox occur? Using the Activity Series

• See the AP pages for the reduction potential list.

166
What is reduced and what is oxidized? What are
Ox. And Red. Agents? Write 1/2 reax.

• Examples
• iron nail in copper sulfate.
• Aluminum foil in tin II chloride solution.
• Copper wire is placed in silver nitrate solution.

167
CH. 19, section 1

• Balancing complex redox reactions in acid/base
  environments

168
Last Topic Solution Chem/StoichVolumetric
Analysis
169
Measuring Concentration of Solutions

• can be done in several different ways,
  including....
• Molarity
• moles dissolved per liter of solution.

170
Other Concentration Definitions

• Molality moles dissolved per kilogram of
  solvent
• Mass percentage mass of solute compared to mass
  of solution.
• Mole fraction moles dissolved compared to moles
  of total solution particles.

171
Molarity Equation

• M n
• V
• M Molarity
• n moles dissolved solute
• V volume of solution in liters

172
Molality Equation

• m n
• kg
• m molality
• n moles dissolved solute
• kg mass of solvent in kilograms.

173
Mass Percent Equation

• Mass of solute___
• Mass of solution
• Express concentration as a percentage.
• Any units for mass will do.

174
Mole Fraction Equation

• Moles of solute ________________
• Moles of solute moles of solvent
• Express result as a decimal number

175
Practice Molarity Problems

• What is the molarity of a solution containing
  34.2 grams of sucrose in 2.00 liters of solution?
• 0.0500 M
• What volume of 0.65 M HCl holds 3.0 grams of HCl?
• 0.126 L or 126 mL

176
Molarity and Dilution Problems

• What mass of silver nitrate must be added to a
  flask to make 500. mL of 0.025 M solution?
• 2.1 g
• How many mL of 4.00 M acetic acid are needed to
  make 500. mL of 1.00 M solution?
• 125 mL

177
More Problems!

• How many mL of 8.4 M KNO3 are needed to make 3.00
  Liters of 2.5 M solution?
• 890 mL
• Suppose 200.mL of water are added to 400.mL of
  1.20 M HNO3. What is the molarity of the
  resulting mixture?
• 0.800 M

178
Stoichiometry with Molarity Quantitative
Analysis

• A 1.000 L sample of polluted water was analyzed
  for lead II ion by adding excess sodium sulfate
  to it. The mass of lead II sulfate precipitating
  was 220.0 mg What is the mass of lead in the
  water?
• 150.3 mg
• What would be the concentration of lead in the
  water?
• 0.000725 M

179
Solution Stoich

• A flask contains water mixed with some HCl. The
  solution is titrated with 0.225 M KOH until a pH
  of 7 is reached. 15.20 mL of the KOH solution
  are needed. What is the mass of the HCl in the
  flask?
• 0.125 grams

180
Solution concentrations

• If 35.0 grams of potassium nitrate are dissolved
  in 55.0 grams of water, the solution has a
  density of 1.108 g/mL. Determine the molarity,
  molality, mass, and mole fraction concentration
  values of the solution.

181
Learn to balance complex redox reactions
separate note page

• Include Ch. 19 section 1 problems with the
  chapter 4 problem set.
• Booknotes not required for 19.1

182
Chemical or Physical Property? 1-7

1. Its a liquid
2. The pH is 12
3. It burns in air
4. It tastes sweet
5. It is green
6. It weighs 5 lbs.
7. It bubbles in acids

183
Chemical or Physical Property? 1-7

1. P
2. P
3. C
4. P (?)
5. P
6. P
7. C

184
Intensive vs. Extensive Properties?

1. Its mass is 50 g.
2. It dissolves in oil.
3. Its density is 1.5g/ml
4. It is 6 inches long
5. It conducts electricity
6. It is acidic
7. It is at room temperature.

185
Intensive vs. Extensive Properties?

1. E
2. I
3. I
4. E
5. I
6. I
7. E

186
Homogeneous or Heterogeneous?

1. Concrete
2. Jello
3. Muddy water
4. Diamond
5. Hair
6. Children in a class
7. Tossed salad
8. milk

187
Homogeneous or Heterogeneous?

1. He
2. Ho
3. He
4. Ho
5. He
6. He
7. He
8. Ho

188
Element, Compound or Mixture?

1. C
2. E
3. C
4. M
5. E
6. C
7. C
8. M

1. M
2. M

189
Name these elements

• Sb
• As
• Ni
• Fe
• Zr
• Ra
• Au

8. Na 9. Sr 10. Ag 11. Ba 12. P 13. F 14. Mg

190
Name these elements

• antimony
• arsenic
• nickel
• iron
• zirconium
• radium
• gold

8. sodium 9. strontium 10. silver 11. barium 12.
phosphorus 13. fluorine 14. magnesium
191
Write symbols for these elements

1. Aluminum
2. Tin
3. Rubidium
4. Argon
5. Helium
6. Neon
7. uranium

8. lead 9. potassium 10. calcium 11. zinc 12.
chlorine 13. copper 14. tungsten
192
Write symbols for these elements
8. Pb 9. K 10. Ca 11. Zn 12. Cl 13. Cu 14. W
1.Al 2. Sn 3. Rb 4. Ar 5. He 6. Ne 7. U
193
Solubility Song

• Sing
• Make a Chart
• What is the chemistry of a soluble ionic
  compound?

194
Solubility Quiz Soluble or not?

1. Iron II hydroxide
2. Potassium phosphate
3. Barium nitrate
4. Strontium sulfate
5. Calcium chloride
6. Silver acetate

195
Solubility Quiz Soluble or not?
Name two solutions with soluble salts that would
combine to form the precipitates in 1, 4 and 6

1. not
2. sol
3. sol
4. not
5. sol
6. not

196
Ions to Know

• Thoughts on patterns

197
What is the common ionic charge for each element?

1. calcium
2. argon
3. potassium
4. nitrogen
5. chlorine
6. aluminum
7. oxygen

198
What is the common ionic charge for each element?

1. 2
2. 0
3. 1
4. 3-
5. 1-
6. 3
7. 2-

199
Name these ions

• SO42-
• NO3-
• PO33-
• C2H3O2-
• NH4
• S2O32-
• C2O42-

200
Name these ions

• Sulfate
• Nitrate
• Phosphate
• acetate
• ammonium
• thiosulfate
• oxalate

201
Write formulas for these ions

• Carbonate
• Nitrite
• sodium
• iodite
• sulfite
• sulfide
• bromate

202
Write formulas for these ions

• CO32-
• NO2-
• Na
• IO2-
• SO32-
• S2-
• BrO3-

203
What is the charge on these ions?

• PO3
• N
• AsO4
• ClO2
• S2O8
• NH4
• IO4

204
Something Newin Chapter 1

• Antoine LaVoisier
• 1743-1794
• French Chemist
• father of modern chemistry

205

• At age 28 married 13-year-old Marie-Anne who
  translated from English for him and illustrated
  his books she was well educated in chemistry
  herself.
• burned P and S in air, and proved the products
  weighed more than the reactants but the weight
  gained was lost from the air.
• Thus established the Law of Conservation of Mass.

206
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207

• 1778 demonstrated the "air" responsible for
  combustion named this portion of air oxygen and
  the other part of air azote (Greek for no
  life).
• discovered that hydrogen combined with oxygen to
  produce water.

208
Antoinne LaVoisier
209

• 1787 invented the system of nomenclature still
  used today.
• 1789 published the first modern chemical
  textbook, with his theories
• a clear Law of Conservation of Mass
• There is no such thing as phlogiston
• a list of elements, including oxygen, nitrogen,
  hydrogen, phosphorus, mercury, zinc, and
  sulfur.. but also included light and caloric

210
Lavoisier "I have tried...to arrive at the
truth by linking up facts to suppress as much as
possible the use of reasoning, which is often an
unreliable instrument which deceives us, in order
to follow as much as possible the torch of
observation and of experiment."
211
Phlogiston Theory

• Ancient Greeks thought there were four substances
  in the world E,A, F, W
• In the 1600s Johann Becher added to the list
  Phlogiston is a 5th element its in any
  substance that burns!

212

• LaVoisier
• worked as a tax collector
• beheaded during the French revolution for using
  public money to fund his research.

213
LaVoisier beheaded in France.
214
A Word about ENERGY

• What is it?
• Law of Conservation of Energy

215
Energypart of chemical reactions

• Energy is required to break bonds.
• Energy is released as bonds form.
• This is true in all physical and chemical changes.

216
Reactions and Energy

• Endothermic
• Exothermic
• More about NRG in later chapters.

217
9. What is the Ag in your alloy? Compare
the percent (what is your error) you got to the
theoretical value for the Ag in dimes made
before 1950.Suggest an error to account for the
difference. Choose one with the correct
direction.
218
Results of the atomic theory

• It yields definitions of
• Elements
• Compounds
• Chemical Reactions

219
Daltons Postulates lead to two laws

• Conservation of Mass
• Multiple ProportionsIf two elements form more
  than one compound, the mass ratios of one of the
  elements in one compound to the same element in
  the other compound is always in a small whole
  number ratio.
• Think of benzene and methane

220
Mendeleev

• Born in Siberia 1834
• Youngest of 14 children.
• Hated everything in school except science.
• Father died when he was 2, mother favored him as
  a student and child, she died after he got
  admitted to university at age 15.

221
Mendeleev

• In 1855 was told he had two years to live,
  probably had tuberculosis.
• Worked as a professor of chemistry at St.
  Petersburg, Russia
• Organized known elements according to their
  properties and thus discovered the periodic law.

222
Mendeleev

• Meyer also discovered a periodic law, but
  Mendeleev published first.
• Was a talented and popular public speaker
• Married Feozva, had two children they did not
  get along. Divorced her and married Anna, with
  whom he had four children.
• The Czar looked the other way on his bigamy

223
Mendeleev

• His periodic law was most accepted after it was
  shown his predictions of the existence of other
  elements were correct.
• Eka-silicon and two others were discovered.
• Died in 1907 at the age of 73.

224
Molecular vs. Empirical formula

• Molecular formula is the real formula for a
  compound
• Empirical formula is the lowest ratio of elements
  in the compound.
• Example
• ethylene glycol is C2H6O2 (molecular)
• empirical formula is CH3O (lowest ratio)

225
Structural Formula

• Arrangement of the atoms in a formula to show
  what shape, function or type of molecule it is.
• H-O-H CH3COOH
• Many organic molecules are frequently written
  structurally.

226
Organic Carbon Chain Classes

• Alkanes
• Alkenes
• Alkynes

227
The End of Chapter 12 notes

• Time really matters.
• T-shirt? Submit ideas Asap
• Mole Day volunteers needed make a moleata, run a
  contest.

228
A.P. Chemistry Chapter 3

• Calculations with Chemical Formulas and Equations

229
What is a Mole?