AP Chemistry Unit 1

1
AP ChemistryUnit 1

• Review
• Chapters 1-4

2
The Atom
3
The Atom

• What are the 3 particles in an atom, where are
  they, and what is their approximate mass in amu?

4
Protons and Neutrons

• Mass of proton 1.673 x 10-27 kg, 1.007 amu
• Mass of neutron 1.675 x 10-27 kg, 1.009 amu
• Mass of electron 9.109 x 10-31 kg, 0.0005486
  amu
• What is the conversion factor?
• 1 amu 1.660 x 10-27 kg

5
Types of Matter
6
States of Matter

• 3 basic states, 2 advanced
• What are they?

7
States of Matter

• Solid lowest energy common state, usually most
  tightly packed
• Liquid more energy than solid, fluid, particles
  slide by each other
• Gas highest energy common state, particles far
  apart, fluid
• Plasma most common state, high energy ionized
  gas
• Super solid exists at lt 0.2 K, crystal where
  atoms condense to a single point

8
Properties of Matter
9
Extensive and Intensive Properties

• Extensive properties
• What and how much of something there is
• Intensive properties
• What something is

10
Physical Properties

• Observations
• Color, size, hardness

11
Chemical Properties

• Descriptions of how something can change
• Either looked up or experimented

12
Scientific Measures and Units
13
Fundamental Units
14
Group Discussions

• Why use Kelvin?

"There is nothing new to be discovered in physics
now. All that remains is more and more precise
measurement." -Lord William Thomson Kelvin 1900
15
Celsius vs Centigrade

• 1741 Anders Celsius made thermometers and set
  100 Celsius as the freezing point of water and 0
  Celsius as the boiling point
• 1887 International Commission of Weights and
  Measures flips the scale and makes this
  Centigrade scale the standard

16
What is a Mole?

• Group Activity
• How many moles would fit in the room if atoms
  were the size of a grain of sand?
• Form a hypothesis and write a protocol
• Find out!

17
What is a Mole?

• Discussion
• Were you close?
• Why or why not?

18
SI Prefixes
19
Significant Figures

• All certain digits plus one
• Certain digits are what is graduated on the
  instrument
• Last digit always assumed uncertain
• 1 more than what is graduated, an estimate

20
Rules for Sig Figs
21
Rounding in Science
22
Processing Data
23
Conversions and Unknowns

• Cross Multiplication
• Uses known relationship to find unknown, relys on
  only 2 units
• Dimensional analysis
• Uses known relationships to cancel units out and
  convert

24
Practice

• Convert 24 minutes into seconds
• Find volume of 3.5 mol of an ideal gas at STP
• Convert miles to cm (Hint Appendix 6)
• Convert angstroms to in
• Convert 55 mph to Tm/s

25
Uncertainty

• All measurements have uncertainty
• Liquids
• Measure from the miniscus
• Rulers
• Dont measure from the end

26
Uncertainty

• Last digit always estimated digit
• One place more than what is graded

27
Precision and Accuracy

• Precision
• how close a set of measurements are to each other
• Accuracy
• how close a measurement is to the actual value

28
Precision

• Calculate by finding average among readings and
  then how far each is from that
• Most commonly shown in standard deviation

29
Standard Deviation

• Appendix A1 Section 5 pg A12
• Find the average
• Standard Deviation is calculated using
• sv((S(xy-xavg)2)/(n-1))
• Where s is the standard deviation, xy is an
  individual measurement, xavg is the average
  measurement, and n is the total number of
  measurements.

30
Standard Deviation Practice

• Find the standard deviation in the following set
• 2 cm, 4.1 cm, 3 cm, and 2 cm

31
Percentage Error

• error (Experimental-Accepted)/Accepted 100
• Tells how close experimental data is to the
  actual value
• Negative if accepted value is above
• Positive if accepted value is below

32
Practice

• Experimental 3.45 cm Actual 3.91 cm
• Experimental 95 s Actual 89 s
• Experimental 21 kg Actual 20.5 kg
• A sample of 1.0 mole of water is found to have a
  mass of 17.998 g. What is the percent error?

33
Some Important Laws
34
Famous Dead Guys

• Antoine Lavoisier
• French 1743-1794
• Created modern chemistry
• Made Law of Conservation of Mass
• Much work with oxygen
• Beheaded during revolution

35
Law of Conservation of Mass

• The total mass cannot change, matter is neither
  created or destroyed.
• REMEMBER THIS WHEN BALANCING EQUATIONS

36
Law of Definite Proportions

• Also called the Law of Constant Composition
• A pure compound always has a constant formula

37
Law of Multiple Proportions

• Elements always react and form compounds in whole
  number ratios

38
Famous Dead Guys

• John Dalton
• English 1808
• Teacher
• Explained Law of Conservation Of Mass
• First true Atomic Theory

39
Daltons Atomic Theory

• All matter is composed of small particles called
  atoms
• Atoms of the same element are identical and
  different from those of other elements
• Atoms cannot be divided, created, or destroyed
• Atoms combine in whole number ratios to form
  compounds
• In reactions atoms combine, separate, and
  rearrange

40
First Atomic Model

• Greeks knew that matter was made of smaller pieces

41
Atomic Models
1808
1897
1911
1913
42
Famous Dead Guys

• Ernest Rutherford, Hans Geiger, Ernest Marsden
• University of Manchester 1911
• Fired a particles at a sheet of gold foil
• Discovered nucleus
• Known as Rutherfords Gold Foil Experiment

43
Rutherfords Gold Foil Experiment

• Expected particles to pass through easily
• Large deflections, some reflections even
• Nucleus must be small compared to the atom
• What did they find about the charge of the
  nucleus?

44
Nomenclature
45
Binary Ionic

• Binary ionic compounds are compounds with only 2
  elements held together by an ionic bond

46
Binary Ionic Compounds Type 1

• Cation (positive) named first, then anion
  (negative)
• Monatomic cations named by the element.
• Monatomic anion is the base of the element with
  the end replace with -ide

Pg 62
47
Examples

• NaCl
• Ca3P2
• LiI

Sodium Chloride
Calcium Phosphide
Lithium Iodide
48
Binary Ionic Compounds Type 2

• Same rules as Type 1
• Typically transition metals
• Name must include oxidation state on the cation
  in parentheses and Roman Numerals
• Fe3 is Iron (III)

49
Examples

• Fe2O3
• CuF
• PbS2
• Hg2H2

Iron (III) Oxide
Copper (I) Fluoride
Lead (IV) Sulfide
Mercury (I) Hydride
50
Polyatomic Ions
51
Examples

• H2O2
• KCN
• Co2(SO3)3
• H2O or HOH
• H3PO4

Hydrogen Peroxide
Potassium Cyanide
Cobalt (III) Sulfite
Hydrogen Hydroxide
Hydrogen Phosphate (Phosphoric acid)
52
Binary Covalent Compounds Type 3

• 2 nonmetals
• The first element in the formula is named first
  using the whole name
• The second element is named as an anion
• Prefixes are given to tell how many atoms are
  present
• Mono- isnt used for the first element

53
Prefixes

• 1 Mono-
• 2 Di-
• 3 Tri-
• 4 Tetra-
• 5 Penta-
• 6 Hexa-
• 7 Hepta-
• 8 Octa-
• 9 Nona-
• 10 Deca-

54
Practice

• NO
• CCl4
• PCl3
• P4O8

Nitrogen Monoxide
Carbon Tetrachloride
Phosphorous Trichloride
Tetraphosphorous Octaoxide
55
Acids

• Dissociate to produce H3O ions
• If the anion doesnt contain oxygen
• Add hydro to the beginning and replace ate with
  ic in the anion name
• If the anion is an oxyanion
• If the anion ends in ate, it is changed to ic
• If the anion ends in ite, it is changed to ous

56
Examples

• HNO2
• HNO3
• HCl
• HF
• C2H4O2

Nitrous Acid
Nitric Acid
Hydrochloric Acid
Hydrofluoric Acid
Acetic Acid
57
Isotopes
58
Isotopes

• Atoms of an element with non-standard masses
• So what has to be different?

59
Isotopes

• Isotopes have the same number of protons and
  electrons with differing neutrons
• Identified by mass number

60
Isotopic Notation

• The mass number is in the upper left before the
  symbol of the element
• Carbon-12 is 12C

61
Isotopes of Hydrogen
Hydrogen is the only element with specific names
for the isotopes
62
Practice

• Write the isotopic notation for the following
• Protons 3 Neutrons 5
• Protons 7 Neutrons 7
• Protons 108 Neutrons 170
• Mass 195 Neutrons 112
• Mass 19 Neutrons 5
• Mass 78 Neutrons 37

8Li 14N 278Hs 195Bi 19Si 78Nb
63
Mass Number and Atomic Number

• Atomic - lower left, symbolized by Z, protons
• Mass - upper right, symbolized by A, neutrons
  protons

64
Atomic Mass Units

• Make masses manageable
• Relative mass
• 1 amu 1/12 mass of 12C 1.660540 x 10-27 kg
• Protons and neutrons both 1 amu
• Protons 1.007276 amu
• Neutrons 1.008665 amu

65
Practice

• A proton is 1,837 times larger than an electron
  and has a mass of 1.007276 amu. Determine the
  mass of an electron in amu.
• 0.0005483266 amu

66
Why Atomic Masses Arent Always Whole Numbers

• Atomic mass of an element is a weighted average
  of all natural isotopes
• Sum of percentage x atomic mass for each isotope

67
Carbon

• 98.93 Carbon-12
• 1.07 Carbon-13
• 0.989312 amu 0.010713 amu 12.0107 amu

68
Oxygen

• 99.757 Oxygen-16
• 0.038 Oxygen-17
• 0.205 Oxygen-18
• Determine atomic mass for Oxygen
• 0.99757160.00038170.0020518
• 16.004 amu

69
Molar Mass

• Mass of 1 mole of a pure substance
• Atomic mass, but in grams instead of amu

70
Practice

• What is the mass of 1 mole of neon?
• What is the mass of 4.5 moles of lead?
• What is the mass of 2 moles of methane gas (CH4)?
• How many moles are in 110 g of manganese?

71
Stoichiometry!
72
Empirical Formulas

• Smallest ratio of atoms, determined by percent
  composition

73
Molecular Formula

• Gives the actual formula for a compound
• Combines empirical formula with molar mass of a
  compound

74
Percent Composition

• Helps determine empirical formula
• Take total mass divided by molar mass of each
  component element
• Ex. HCl is 2.7 hydrogen by mass and 97.3
  chlorine

75
Practice

• What is the empirical formula for a compound with
  total atomic mass of 18 amu and is 11.1 hydrogen
  and 88.9 oxygen?
• A compound is 58.5 g/mol and 39.3 sodium. What
  is the other element?
• A compound is 46 g/mol, 52.1 carbon, 34.8
  oxygen, and 13.1 hydrogen. What is the formula?

76
Chemical Reactions

• Things reacting and combining are called
  reactants
• Things being made are products
• Gives ratio of reactants and products with ratio
  of coefficients

77
Example

• Combustion of ethanol
• C2H6O 3 O2 ? 3 H2O 2 CO2

78
Symbols and Abbreviations

• Solid (s)
• Liquid (l)
• Gas (g)
• Aqueous (aq)
• Heat ?
• Light ?
• Proton 1H
• Neutron 1n
• Electron e-

79
Balancing Reactions
80
Basic Rules

• Make sure there are equal numbers of all elements
  on both sides of equation
• Make sure any charges are balanced on both sides
  of equation, add electrons if necessary
• Only change coefficients, not individual formulas

81
Example

• H2 O2 ? H2O
• 2 H2 O2 ? 2 H2O

82
Practice

• KMnO4 H2SO4 ? HMnO4 K2SO4
• 2 KMnO4 H2SO4 ? 2 HMnO4 K2SO4
• C2H5OH O2 ? CO2 H2O
• C2H5OH 3 O2 ? 2 CO2 3 H2O
• N2 H2 ? NH3
• N2 3 H2 ? 2 NH3

83
Using Reactions as Recipes

• Reaction shows basic ratios
• Use this to determine unknowns based on known
  values

84
Practice

• The combustion of 2-butene yields 24 g of water,
  what mass of oxygen was consumed?

85
Limiting Reactants

• A reactant that decides how much product can be
  made
• Determined by finding moles of all reactants
• The least amount of reactant is the most amount
  of product that can be formed

86
Example

• 53 g of methane are reacted with 55.8 g of water.
  What mass of carbon monoxide is produced and what
  is the limiting reactant?
• CH4 H2O ? 3 H2 CO
• 3.3 moles of methane and 3.1 moles of water, so
  water is the limiting reactant.
• 3.1 moles of CO is 87 g

87
Another Example

• 41 g of methane are reacted with excess water.
  What mass of hydrogen is produced?
• CH4 H2O ? 3 H2 CO
• 2.5 moles of methane
• 7.5 moles of hydrogen
• 15 g of hydrogen

88
Reaction Yields

• Theoretical yield maximum possible yield in
  ideal conditions given the amount of reactants
• Percent yield what percent of the actual yield
  was actually obtained, actual/theoretical

89
Practice

• What percent yield was obtained in a reaction of
  hydrogen with oxygen to make 78 g of water when 5
  moles of oxygen was reacted with excess hydrogen?
• 2 H2 O2 ? 2 H2O
• Theoretical yield 90 g
• Percent yield 87

90
Chemical Reactions

• Chemistry to the MAX!

91
Solvents

• A liquid or gas that dissolves another compound,
  element, or mixture
• Thing being dissolved is called the solute

92
Solvents

• Polar unequal charge distribution causing a
  positive and negative region or regions
• Nonpolar No unequal charge distribution
• Memory Key Like dissolves like
• Polar solvents dissolve polar solutes
• Nonpolar solvents dissolve nonpolar solutes

93
Water as Solvent
-

94
Water as Solvent

• Water is a polar molecule
• When something is dissolved in water it is
  hydrated
• Bent or seesaw structure created a positive and
  negative region
• Explains hydrophobic and hydrophilic lipids and
  other molecules

95
Water as Solvent

• It can dissolve polar compounds
• Anions or cations

96
Partial Charges

• Regions of uneven distribution, positive or
  negative
• d or d-

97
Solubility

• How well a solute is dissolved in a given solvent
• Depends on magnitude of intermolecular forces
  (bonds) in relation to the magnitude of the
  electronegativities of the solvent and solute

98
Solubility and Electronegativities

• If the difference in electronegativities between
  the 2

99
Electrolytes
100
Electrolytes

• A substance (typically ionic) that will conduct
  electricity when in water

101
Electrolytes

• Water (H2O) does not conduct electricity
• Strong electrolytes conduct electricity well
• Weak electrolytes do not conduct electricity well
• Nonelectrolytes do not conduct electricity at all

102
Famous Dead Guy

• Svante Arrhenius
• 1859-1927 Swedish
• Arrhenius Equation relates reaction rates to
  temperature (well get to it all too soon)
• Discovered conductivity in water came from
  hydrated ions
• Wasnt accepted for almost 20 years

103
Strong Electrolytes

• Compounds that completely ionize
• Conduct electricity very well
• Soluble salts, strong acids, strong bases

104
Strong Electrolytes

• Soluble salts
• NaCl dissolves into almost 100 Na and Cl-
• Strong acids and bases
• Dissolve to almost 100 H and conjugate based or
  OH- and conjugate acid respectively
• HCl, NaOH

105
Weak Elecrolytes

• Do not ionize well
• Typically stronger bonds than strong electrolytes
• Weak acids and weak bases

106
Nonelectrolytes

• Do dissolve in water, but dont ionize at all
• Molecules stay entirely in tact
• Ethanol or sugar

107
Molarity

• Concentration
• Measure of amounts of solutions
• Units of Molar (M)
• M moles of solute/liters of solvent

108
Example

• 2.5 moles of NaCl is dissolved in 0.75 L of
  water. What is the concentration?
• M 2.5 mol/0.75 L
• 3.3 M

109
Practice

• 0.2340 moles of HCl are mixed with 275 mL water.
  What is the concentration?
• 50 g of NaCl are dissolved in 50 g of water, what
  is the concentration?
• How many moles of H2SO4 are in 100 mL of a 6 M
  solution?
• How many grams of water are needed to make 10. mL
  of 10. M NaOH?

110
Dilution

• Decreasing the concentration by adding more
  solvent
• Only increases the volume of solvent, does not
  affect total amount of solute present
• Make sure moles of solute are equal before and
  after
• M1V1 M2V2

111
Example

• A 12 M HCl sample needs to be diluted to make 100
  mL of 1.0 M HCl. How much of the 12 M sample is
  needed?
• 100 mL 1.0 M 12 M x mL
• x 8.3 mL of 12 M HCl
• Units of volume dont matter, as long as theyre
  the same

112
Practice

• How much 5 M NaOH needs to be used to make 3 L of
  2.7 M NaOH?
• How much water should be added to 12 mL of 2 M KI
  to dilute it to 0.5 M?
• 2.34 L of water are added to 0.841 L of an
  unknown concentration acid. The final
  concentration was 1.00 M, what was the initial
  concentration?

113
Precipitation Reactions

• A reaction that creates one or more insoluble
  products
• Insoluble product is called the precipitate

114
Salts

• Generally a replacement of the acidic hydrogen on
  an acid with a metal
• NaCl
• BaNO3

115
General Solubility Rules for Salts in Water

• Most nitrate salts are soluble
• Most sulfate salts are soluble
• Most alkali or ammonium salts are soluble
• Cl-, Br-, and I- salts are soluble unless with
  Ag, Pb2, and Hg22
• Hydroxide salts are only slightly soluble
• Sulfate, chromate, carbonate, and phosphate salts
  are only slightly soluble

116
Reaction Equations
117
Molecular Equations

• Shows reactants and products as whole molecules
• HCl(aq) NaOH(aq) ? H2O(g) NaCl(aq)

118
Complete Ionic Equation

• Shows dissolved ions as ions
• H (aq) Cl-(aq) Na (aq) OH-(aq) ?
• H2O(g) Na (aq) Cl-(aq)

119
Net Ionic Equation

• Removes ions or molecules on both sides of the
  equation, these are called spectators
• H (aq) OH-(aq) ? H2O(g)

120
Titrations
121
Titrations

• Also called volumetric analysis
• Uses an acid or base of known concentration added
  to a base or acid of unknown concentration in a
  neutralization reaction to determine the unknown
  concentration
• Uses the same equation as dilutions

122
Titrations

• Titrant solution with known concentration added
  to the analyte
• Analyte the solution with unknown concentration
• Titrant delivered from a burette into analyte in
  a beaker or Erlenmeyer flask

123
Titrations

• The point where all the analyte has reacted with
  the added volume of titrant is called the
  equivalence point
• The endpoint is where no more titrant should be
  added, ideally the same as the equivalence point
• The endpoint is marked by a change in an indicator

124
Indicators
125
Titration Curve
Equivalence Point
126
Equivalence Point

• Can be found exactly by taking the first
  derivative of the titration curve

127
Simplifying Neutralization Reactions

• Replace acids with HnA
• Replace bases with B(OH)n
• N represents how many acidic hydrogens or
  hydroxides are present in the reaction
• HCl NaOH ? NaCl H2O
• HA BOH ? AB H2O

128
Normality

• When polyprotic acids or polybasic bases are used
  it can be easier to keep straight
• Units of N (Normal)
• Multiply molarity by how many acidic hydrogens or
  hydroxides to get normality

129
Redox
130
Redox

• Oxidation-Reduction reactions
• Deals with the transfer of electrons
• Therefore not all reactions are redox
• Oxidation states used to determine transfers

131
Oxidation States

• Common oxidation states determined from group on
  the periodic table
• Group 1 usually 1
• Group 2 usually 2
• Group 16 usually -2
• Group 17 usually -1
• and refer to charge as ions

132
Oxidation State Rules
133
Oxidation State Rules

• Always written as charge magnitude (/-) followed
  by number
• If compound is overall neutral the oxidation
  states must be equal
• Make sure to take quantity of atoms into account

134
Solving Redox Equations

• Always fill in F, H, and O first
• Make sure to multiply states by atoms
• Make sure charges equal sums of oxidation states
• Never assign Group 1 or 2 negatives or 16 or 17
  positive unless you have a good reason
• Make sure elemental states always are 0

135
Practice

• H2O
• I2
• C2H3O2-
• CrO42-
• HPO42-
• C13H17NO2

136
Redox

• Oxidation an increase in oxidation state,
  losing an electron
• Reduction a decrease in oxidation state,
  gaining an electron

137
Redox

• Oxidizing agent a compound that oxidizes
  another compound, gets reduced, gains an electron
• Reducing agent a compound that reduces another
  compound, gets oxidized, loses an electron

138
The Easy Way to Solve Redox

• Write balanced equation
• Write known oxidation states below each atom (H,
  O, polyatomic ions, elementals)
• Under that write the total oxidation state by
  multiplying the state by the number of atoms
• Determine unknown oxidation states by charge on
  the molecule
• Make sure the sum of the oxidation states equals
  any charge on the molecule

139
Practice

• List oxidation states. Which are oxidized, which
  are reduced? What are the oxidizing and reducing
  agents?
• CH4 2O2 ? CO2 2 H2O
• AgNO3 HCN ? AgCN HNO3
• Ag(S2O3)23- e- ? Ag 2 S2O32-
• Fe2 MnO4- ? Mn2 2 O2 Fe3 4 e-

140
End of Unit

• Questions?

141
Acid-Base Chemistry
142
Acid Definitions

• Lewis (Arrhenius) acids generate H ions in
  solution
• Brønsted-Lowry acids are proton donors

143
Base Definitions

• Lewis (Arrhenius) bases generate OH- ions in
  solution
• Brønsted-Lowry bases are proton acceptors

144
Discussion

• Compare and contrast Brønsted-Lowry and Lewis
  acids and bases. Are they the same?

145
Strong Acids/Bases

• Completely dissociate (ionize) in solution

146
Strong Acids
147
Strong Bases
148
Neutralization

• The reaction of an acid with a base to produce
  water
• When just enough acid and base have been mixed so
  that 100 of each has reacted it is said to be
  neutral

149
Tips for Neutralization Reactions

• Moles of acid (H) must be equal to moles of base
  (OH-).
• Watch for polyprotic (multiple acidic hydrogens)
  acids.
• If moles are unequal the solution will not be
  neutral.

150
pH and pOH
151
pH

• Goes from 1-14
• Water is 7
• pH -logH
• means concentration as molarity

152
pH

• Strong acids dissociate completely, which means
  they have a higher H and a lower pH

153
pOH

• Goes from 1-14
• Water is 7
• pOH -logOH-
• means concentration as molarity

154
pOH

• Strong bases dissociate completely, which means
  they have a higher OH- and a lower pOH

155
pH and pOH

• High pH is the same as a low pOH and vice-versa
• pOH can be found using 14-pHpOH and vice-versa

156
Practice

• What is the pH of a 0.3 M solution of HCl?
• What is the pH of a 0.003 M solution of H2SO4?
• What is the pOH of 5.0 g NaOH in 1.0 L of water?