Section 1'1 1'3

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Lesson 1 Section 1.1 1.3

• The Scientific Method
• ? Hypothesis
• ?? explains observations
• ?? is falsifiable
• ? A law
• ?? is founded upon thorough observational
  experience
• ?? is based upon a mathematical treatment
• ?? has predictive capability
• ?? is generally irrefutable
• ?? examples
• ??? Newtons laws of motion
• ??? the laws of thermodynamics

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• The Scientific Method (continued)
• Theory
• ?? designed to explain an hypothesis
• ?? sometimes motivated to explain anomalous
• experimental data
• ?? is sometimes stand-alone
• ?? is based upon a model
• ??? conceptual model
• ??? mathematical model (has predictive
  capability)
• ?? can never be absolutely proven

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Hypothesis
Theory
Confirm/revise
Test
Test
Confirm/revise
Observations
Experiments
Experiments
Test
Confirm/revise
Law
4

• States of matter
• Solid
• Liquid
• Gas

Classification of matter
Matter
Pure substance
Mixture
Element
Compound
Heterogeneous
Homogeneous
5
Lesson 2 Section 1.4 1.8

• Properties of matter
• Physical properties characteristics associated
  with appearance,
• texture, color, state, etc.
• ? Chemical properties characteristics associated
  with reactivity

6

• Energy
• The capacity to do work. Work is the ability to
  move an object.
• Total energy
• ?? Kinetic energy energy of motion
• ?? Potential energy the energy an object is
  capable of using by
• virtue of its position in an external
  force field
• Internal energy
• The law of conservation of energy energy is
  neither created nor
• destroyed.

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• Fundamental units m, g, s, K, mol, C, cd
• Derived units a combination of two or more
  fundamental units
• Area has units of length2
• ?? Example km2
• Volume has units of length3
• ?? Example cm3
• Density

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Temperature conversions
Note The Kelvin scale is the thermodynamic scale.
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• Categories of properties
• An extensive property depends upon the amount.
• ?? Examples
• An intensive property does not depend upon the
  amount.
• ?? Examples

11

• Rules for significant figures
• All non-zero numbers are significant.
• ?? Example 134.2 4 sig. figs.
• Zeros between non-zero numbers are significant.
• ?? Example 9087 4 sig. figs.
• Place-holding zeros are not significant.
• ?? Examples 0.0032 2 sig. figs.
• 0.00320 3 sig. figs.

12

• Rules for significant figures (continued)
• Multiplication division the answer has the
  same number of
• sig. figs. as the number that has the fewest
  sig. figs.
• ?? Example 45.47 0.650 29.6
• Addition subtraction the answer has the same
  number of
• decimal places as the number that has the
  fewest number of
• decimal places.
• ?? Examples 54.13 0.073 54.20
• 54.13 0.003 54.13

13

• Rules for significant figures (continued)
• Round to the correct number of sig. figs. after
  you
• have done all the calculations.
• Exact numbers and exact conversions do not limit
  the
• number of sig. figs.

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• Precision and accuracy
• Accuracy how close a measured answer is to the
  actual
• answer.
• Precision how close a series of measurements
  are to one
• another.
• Systematic errors lead to poor accuracy.
• Excessive random error leads to poor precision.

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The target
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• Units and dimensional analysis
• Always include units in your analysis.
• Conversion factor a ratio of units that is
  designed to
• convert from one unit to another.
• ?? Example
• Picket fence method convert 25 mi/hr to m/s.
  Use the
• following conversion 1 km 0.6214 mi

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• Problem-solving tips
• Identify what is given.
• Identify what you must find.
• Devise a strategy that connects what you have to
  you need
• to find. Dont forget units!
• If the problem requires calculations, do all the
  algebra first,
• then do the arithmetic (number crunching).
• Check your answer for reasonableness.

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Lesson 3 Section 2.1 ? 2.4

• Some fundamental laws of chemistry
• the law of conservation of matter matter is
  neither created nor
• destroyed.
• the law of definite proportions all samples of
  a given compound
• have the same proportions of its constituent
  elements.
• the law of multiple proportions when two
  elements (A B)
• form two different compounds, the masses of B
  that combine
• with 1 gram of A can be expressed as a ratio of
  small whole
• numbers.

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• Atomic theory
• Every element is composed of particles called
  atoms.
• All atoms of the same element have the same mass
  (this one
• is not quite true!).
• Atoms combine in whole number ratios to form
  compounds.
• Atoms of one element cannot change into atoms of
  another
• element (this debunks the science of
  alchemy).

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• The structure of the atom
• the electron
• ?? very light particles that comprise part of
  the atom
• ?? possesses a negative charge

- plate
Electron beam
plate
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• The structure of an atom (continued)
• J. J. Thomson determined the charge-to-mass
  ratio of an
• electron (-1.76 x 108 C/g)
• R. Millikan determined the charge of an electron
  (-1.60 x 10-19
• C)
• mass of an electron

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• Structure of an atom (continued)
• Rutherfords experiment with ? particles
  bombarding gold
• foil showed that atoms were mostly comprised of
  space.
• Rutherfords nuclear theory
• ?? Most of the atoms mass is contained within
  a small core
• called a nucleus.
• ?? Most of the atoms volume is empty space.
• ?? There are as many electrons outside the
  nucleus as there
• are positively charged particles (called
  protons) in the
• nucleus.
• Studies by J. Chadwick showed there are neutral
  particles
• (called neutrons) in the nucleus.

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• Structure of an atom (continued)
• Protons and neutrons make up the nucleus of the
  atom, and the
• electrons circle around the nucleus.

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Lesson 4 Section 2.5 ? 2.8

• Elements
• The element is determined by the number of
  protons (p).
• Atoms of a given element may have different
  numbers of
• neutrons (n).
• Isotope symbol atomic number
• mass number
• X element symbol

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• Elements (continued)
• Isotope symbol (continued)
• Shorthand notation
• Example isotopes of chlorine
• Chlorine-35
• Chlorine-37

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• Elements (continued)
• Ions formed from the gain or loss of electrons
• The Periodic Table discuss the periodic table

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• Atomic mass
• MA is the average atomic mass
• fn is the fraction of isotope n
• mn is the mass of isotope n

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• Atomic mass (continued)
• Example Mg has three naturally occurring
  isotopes with
• amus of 23.99 (78.99), 24.99 (10.00), and
  25.98 (11.01)
• Calculate the atomic mass of Mg.

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• Molar mass and Avogadros number
• 1 mole 6.0221421 1023 particles
• Note this is the number of atoms in exactly 12
  g of
• Atomic and molar masses are in g/mole

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Lesson 5 Section 3.1 ? 3.5

• Types of chemical bonds
• Ionic bonds bonds between cations and anions
• ?? The cation is generally a metal ion or a
  polyatomic
• cation and the anion is generally a
  nonmetal ion or a
• polyatomic anion.
• ?? Examples KBr, MnCl2, NH4Cl
• Covalent bonds bonds that are formed by the
  sharing of
• electrons
• ?? Covalent bonds generally form between two
  nonmetals.
• ?? Examples NO, H2O, PCl5, Br2

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• Types of chemical formulas
• Empirical formula gives the relative number of
  atoms of each
• element in the compound.
• Molecular formula gives the actual number of
  atoms of each
• element in the compound.
• Structural formula shows the actual covalent
  bonds in the
• compound.
• Molecular models model kits

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• Example hydrogen peroxide
• HO empirical formula
• H2O2 molecular formula
• H ? O ? O ? H structural formula

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Classification of elements and compounds
Pure substances
Elements
Compounds
Atomic
Molecular
Molecular
Ionic
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• Formation of ions from the elements
• Group 1A elements form cations with a 1 charge.
• Group 2A elements (alkaline earth metals) form
  cations with
• a 2 charge.
• Transitions metals always form cations, but the
  charges vary.
• Exceptions Zn always forms Zn2 and Ag always
  forms Ag.
• Group 6A nonmetal elements form anions with a ?2
  charge.
• Group 7A elements (halogens) form anions with a
  ?1 charge.
• All charges for the elements in the formula must
  add to zero.

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• Nomenclature
• Ionic compounds
• ?? Binary compounds formed from main group
  metals name
• of metal base name of anion ide
• Example RbBr rubidium bromide
• ?? Binary compounds formed from transition
  metals name of
• metal (charge of metal cation) base
  name of anion ide
• Example FeCl3 iron(III) chloride

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• Ionic compounds (continued)
• ?? Compounds formed from polyatomic ions (page
  85 of text)
• ??? Polyatomic cation name of polyatomic
  cation base
• name of anion ide
• Example (NH4)2S ammonium sulfide
• ??? Polyatomic anion name of metal
  (charge of metal
• cation if a transition metal) name
  of polyatomic anion
• Example Mn3(PO4)2
  manganese(II) phosphate
• ??? Polyatomic cation polyatomic anion
• Example (NH4)2CrO4 ammonium
  chromate

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Lesson 6 Sections 3.6 3.8

• Nomenclature (continued)
• Molecular compounds numeric prefix name of
  first element
• numeric prefix base name of second element
  ide
• ?? Numeric prefixes mono 1 hexa 6
• di 2 hepta 7
• tri 3 octa 8
• tetra 4 nona 9
• penta 5 deca 10

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• Molecular compounds (continued)
• ?? Mono is not used if there is only one atom
  of the first
• element
• ?? Examples NO2 nitrogen dioxide
• N2O dinitrogen monoxide
• P4S10 tetraphosphorus decasulfide
• Hydrates the formula contains water molecules
• Example CuSO4?5H2O copper(II) sulfate
  pentahydrate

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• Nomenclature (continued)
• Acids compounds that in water will yield a
  proton (H)
• ?? Binary acids hydro base name of nonmetal
  ic acid
• Example HCl(aq) hydrochloric acid
• ?? Oxyacids (acids in which the anion is
  polyatomic)
• ??? oxyanions ending in ate base name of
  oxyanion ic
• acid
• Example HNO3(aq) nitric acid
• is the nitrate ion

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?? Oxyacids (continued) ??? oxyanions ending
in ite base name of oxyanion ous
acid Example HNO2(aq) nitrous acid

is the nitrite ion
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• Molar mass
• M is the molar (or formula) mass
• ni is the number of atoms per formula unit
• Ai is the atomic mass
• Example C2H4 M 2(12.01 g/mol) 4(1.008 g/mol)
• 28.05 g/mol

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• Composition of a compound
• Px is the mass percent of element x
• nx is the number of atoms of x per
• formula unit
• Ax is the atomic mass of x
• Mcompound is the molar mass of the
• compound

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• Composition of a compound (continued)
• Example calculate the mass percent of N in
  NH4NO3.
• 2(14.01 g/mol) 4(1.008 g/mol) 3(16.00
  g/mol)
• 80.05 g/mol
• 35.00

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Lesson 7 Sections 3.9 3.11

• Determination of the empirical formula from
  chemical analysis
• ?? Examples
• ??? A sample contains the following
  elemental composition
• Se 1.443 g
• Br 5.841 g
• What is the empirical formula?
• ??? Ibuprofen has the following mass
  percent composition
• C 75.69
• H 8.80
• O 15.51
• What is the empirical formula?

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• ?? Combustion analysis used for combustible
  compounds
• ??? Carbon (C) forms carbon dioxide (CO2)
• ??? Hydrogen (H) forms water (H2O)
• ??? Example combustion analysis of a
  13.42-g sample
• containing C, O, and H yields 39.61 g
  of CO2 and
• 9.01 g of H2O. What is the empirical
  formula?
• Rules for balancing chemical equations
• ?? Apply the principle of mass balance the
  number of atoms
• for each element must be the same on either
  side of the
• equation.

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• Rules for balancing chemical equations
  (continued)
• ?? Balance the elements in the more
  complicated formulas
• first.
• ?? Balance molecular elements (e.g. O2, N2)
  last.
• ?? Include state symbols (s, l, g, aq) for all
  species.
• ?? Example C8H18(l) O2(g) ? CO2(g) H2O(l)

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Lesson 8 Sections 4.1 4.3

• Reaction stoichiometry
• The stoichiometric coefficients indicate the
  relative number
• of moles of each substance in the chemical
  reaction.
• ? Reaction stoichiometry conversion chart

Mass
Moles
Moles
Mass
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• Limiting reactant
• The limiting reactant is the one that limits the
  amount of
• product formed.
• The limiting reactant generally is determined by
  trial and
• error.

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• Limiting reactant (continued)
• Conceptual plan
• You are given an amount of A and B for the
  following
• reaction aA bB ? cC

mass A
moles A
moles C
Keep the smaller value of C
moles C
mass C
mass B
moles B
moles C
Whichever reactant yields the smaller value for C
is the limiting reactant.
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• Limiting reactant (continued)
• Yields
• ?? The theoretical yield is determined by the
  limiting reactant.
• ?? The actual yield is calculated by the
  equation
• Y is the percent yield
• AY is the actual yield
• TY is the theoretical yield

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Lesson 9 Sections 4.4 4.6

• Solutions
• Solution solute solvent
• ?? The solute is the component in the least
  concentration.
• ?? The solute can be a dissolved solid, liquid,
  or gas.
• ?? For an aqueous solution, the solvent is
  water.
• Solution concentrations
• ?? Molarity (M) Vsolution is in L

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• Solution concentrations (continued)
• ?? Example calculate the molarity of a
  solution that is made
• by dissolving 45.4 g of NaNO3 in 2.50 L of
  solution.
• Solution dilution equation
• Note For calculating molarities (M1 and M2 )
  using the
• dilution equation, the volumes do not have to
  be in liters.
• For calculating moles using the dilution
  equation, the
• volumes must be in liters.

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• Solution concentrations (continued)
• ?? Solution stoichiometry conversion chart
• Types of solutes
• ?? Electrolytes solutes that ionize in
  solution and conduct
• electricity. Electrolytes include salts,
  acids, and bases.
• ?? Nonelectrolytes solutes that do not ionize
  in solution and
• do not conduct electricity.

Volume of A
Moles of A
Moles of B
Volume of B
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• Types of solutes (continued)
• ?? Interactions of ions with solvent molecules

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• Solutions (continued)
• Solubility of ionic compounds see Table 4.1 on
  page 136 of
• textbook.
• ?? Soluble compounds are not generally
  infinitely soluble.
• ?? Insoluble compounds are in fact slightly
  soluble.

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• Precipitation reactions
• Precipitation reaction a reaction that leads to
  a precipitate
• Example
• PbI2 is not soluble in water.

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• Precipitation reactions (continued)
• Cross pair the ions to determine if a
  precipitate forms
• ?? Examples

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Lesson 10 Sections 4.7 4.9

• Representations of aqueous reactions
• Molecular representation shows the complete
  neutral formulas
• for each compound
• ?? Example

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• Representations of aqueous reactions (continued)
• Complete ionic reaction shows all the
  individual ions
• ?? Example

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• Representations of aqueous reactions (continued)
• Net ionic reaction shows only the ionic species
  that react
• ?? Example

Na(aq) and are called spectator
ions.
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• Acid-base reactions
• Acid a substance that produces H (a proton) in
  solution
• Base a substance that produces OH? in solution
• Examples
• H3O is the hydronium ion.

66

• Acid-base reactions (continued)
• The net ionic reaction for an acid-base reaction
  is
• or
• ?? Example
• acid base water
  salt

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• Gas-evolution reactions
• Reactions that form gases when the reactants are
  mixed
• ?? Examples

68

• Oxidation-reduction (redox) reactions
• Reactions in which the oxidation states of two
  or more elements
• change
• Oxidation states change when electrons are
  gained (reduction)
• or lost (oxidation).
• Rules for assigning oxidation numbers
• ?? The oxidation number of a free element is 0.
• ?? The oxidation number of a monatomic ion is
  its charge.

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• Rules for assigning oxidation numbers
  (continued)
• ?? The sum of the oxidation numbers of all
  atoms in a compound,
• formula unit, or polyatomic ion is the same
  as the overall
• charge.
• ?? Metals always have positive oxidation
  numbers, unless they
• are free elements.
• ??? Group I metals have a 1 oxidation
  number.
• ??? Group II metals have a 2 oxidation
  number.
• ??? Transition metals have oxidation
  numbers that vary.

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• Rules for assigning oxidation numbers
  (continued)
• ?? Nonmetals have oxidation numbers according
  to the following
• hierarchal rules
• ??? Fluorine (F) in any compound is always
  ?1.
• ??? Hydrogen (H) in most compounds is 1.
• ??? Oxygen (O) in most compounds is ?2.
• ??? Group 7A elements in most compounds is
  ?1.
• ??? Group 6A elements in most compounds is
  ?2.
• ??? Group 5A elements in most compounds is
  ?3.

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• Oxidation-reduction (redox) reactions (continued)
• Oxidizing agent the species in the redox
  reaction that is reduced.
• Reducing agent the species in the redox
  reaction that is oxidized.
• Examples What are the oxidation numbers and the
  oxidizing and
• reducing agents in the following reactions?
• ??
• ??

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Lesson 11 Sections 5.1 5.3

• Simple gas laws
• Pressure
• ?? Defined as the force per unit area
• ?? Unit conversions
• 1 atm 101,325 Pa 760 torr (mm Hg)
  14.7 psi

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• Simple gas laws (continued)
• Fundamental state variables P, T, V, n (P T
  are intensive
• properties V n are extensive properties)
• Boyles Law
• if T and n are constant.
• Charless Law
• if P and n are constant.

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• Simple gas laws (continued)
• Avogadros Law
• if T and P are constant.

• Charless Law and absolute zero
• ?? A plot of V vs. T extrapolated to V 0 in
  theory gives
• the value for absolute zero temperature.

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V
-273.15 ? C
-300 -250 -200 -150 -100 -50
0 50 100 150
T/? C
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• Charless Law and absolute zero (continued)
• ?? Kelvin (K) is the absolute (true)
  temperature scale
• Conversion from ?C to K
• Note the Kelvin scale is a shifted
  centigrade scale.

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Lesson 12 Sections 5.4 5.5

• The ideal gas law
• Proportionalities T n constant
• P n constant
• P T constant
• Combining these three proportionalities gives

78

• Application of the ideal gas law
• Density of a gas M is the molar mass
• Molar mass of a gas
• Note These equations apply only for an ideal
  gas.
• Practice problem 5.7 on p. 177

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Lesson 13 Sections 5.6 5.7

• Partial pressures and Daltons Law
• For mixtures of ideal gases
• This equation is Daltons law of partial
  pressures.

80

• Partial pressures and Daltons Law (continued)
• The mole fraction
• The partial pressure in terms of mole fraction
• Practice problems 5.9 on p. 180 and 5.10 on p.
  181

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• Partial pressures and Daltons Law (continued)
• Partial pressures and mole fractions of the
  components of
• dry air
• Nitrogen (N2) 0.78 atm, 0.78
• Oxygen (O2) 0.21 atm, 0.21
• Argon (Ar) 0.009 atm, 0.009
• Carbon dioxide (CO2) 0.0004 atm, 0.0004
• For humid air, the partial pressure of water
  vapor can be up
• to 0.03 atm

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• Reaction stoichiometry for gas reactions
• Conceptual plan
• Practice problem 5.12 on p. 185

P, V, T of B
gas
P, V, T of A
moles of A
moles of B
liquid or solid
mass of B
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Lesson 14 Sections 5.8 5.9

• Kinetic molecular theory
• Theory that explains the behavior of gas
  molecules on a
• molecular level.
• Postulates
• ?? The size of the gas molecules is
  negligible.
• ?? The average kinetic energy of the molecules
  is proportional
• to the temperature.
• ?? Collisions are completely elastic (energy
  is not absorbed by
• the internal modes of the molecule).

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• Kinetic molecular theory (continued)
• Velocity distribution

Fraction of molecules
? ? average
?v?
Molecular speed
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• Kinetic molecular theory (continued)
• Velocity distribution same gas at two different
  temperatures

T1
T1
m1 m2 T2 gt T1 ?v2? gt ?v1?
Fraction of molecules
T2
Molecular speed
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• Kinetic molecular theory (continued)
• Velocity distribution two different gases at
  the same temperature

m1
m1 gt m2 T2 T1 ?v2? gt ?v1?
Fraction of molecules
m2
Molecular speed
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• Kinetic molecular theory (continued)
• Average molar kinetic energy of a gas molecule
• From these equations, it can be shown that
• the root mean square speed

the subscript m is per mole
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• Kinetic molecular theory (continued)
• Average speed
• Root mean square speed
• For calculations of average molecular speeds,
  use
• Practice problem 5.14 on p. 191

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• Kinetic molecular theory (continued)
• Mean free path the average distance traveled by
  a gas
• molecule between collisions.
• Diffusion the spreading out of molecules due to
  concentration
• gradients.
• Effusion the loss of gas molecules as they pass
  through very
• small pinholes.
• Practice problem 5.15 on p. 193

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• Real gases and the van der Waals equation
• Real gases deviate from ideality because
• ?? The size of the gas molecules is not
  negligible.
• ?? The molecules interact with one another.
• van der Waals equation
• b accounts for the finite molecular size a
  accounts for the
• molecular interactions. b tends to make the
  actual pressure
• larger than the ideal pressure a tends to make
  it less than ideal
• (p. 194).

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Lesson 15 Sections 6.1 6.3

• Thermochemistry the study of the relationship
  between
• chemical reactions and energy.
• Energy
• Energy associated with an external reference
  frame
• ?? Kinetic energy energy associated with an
  objects motion
• ?? Potential energy energy associated with an
  objects
• position in an external force field

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• Internal energy
• ?? The energy associated with the thermodynamic
  state of a
• system is called the internal energy.
  Internal energy is
• whats inside the object.
• ?? Internal energy is composed of
• ??? Kinetic energy associated with random
  molecular
• motion (thermal energy)
• ??? Potential energy stored in chemical
  bonds and internal
• modes of molecules
• ?? Internal energy (E) is transferred to and
  from the system
• in the form of heat (q) and work (w).

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• Internal energy (continued)
• ?? The transfer of energy between the system
  and the
• surroundings is in the form of either heat
  or work.

Surroundings
q
System
w
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• Internal energy (continued)
• ?? When energy is transferred as heat, there is
  an increase in the
• thermal energy when energy is transferred
  as work, there is a
• net displacement of an object
• ?? Energy units

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• Internal energy (continued)
• ?? Mathematical expression for the First
• Law of Thermodynamics (energy is
• conserved in any process)
• State function any function of the system
• ?? A change in a state function depends only
  upon the initial
• and final states of the system, not on how
  it got there.
• ?? P, T, V, E are all state functions.
• ?? q and w are not state functions, but are
  ways in which
• energy is transferred between the system
  and the
• surroundings.

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Surroundings
q
System
w
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• Heat (q)
• A transfer of heat changes the level of thermal
  motion of the
• molecules
• A transfer of energy in the form of heat causes
• ?? A change in the phase (i.e. melting,
  boiling)
• ?? A change in the temperature of the system
• C is the heat capacity
• Cm is the molar heat capacity
• Cs is the specific heat

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• A transfer of heat causes (continued)
• ?? The higher the heat capacity, the more heat
  is required to
• raise the temperature of a substance.
• Sign convention
• ?? A process is endothermic if q gt 0.
• ?? A process is exothermic if q lt 0.

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• Work (w)
• Work done on or by a system causes net motion of
  an object.
• A transfer of energy in the form of work causes
• ?? A change in the volume of the system
• ?? A change in the volume of the system
• pressure-volume work

100
Lesson 16 Sections 6.4 6.5

• Constant-volume calorimetry
• If the volume is constant, then
• Energy-balance equation for the bomb calorimeter
  (?V 0)
• ??
• ?? Ccal includes the apparatus and the water
  (p. 217)
• ?? Practice problems 6.4 (p. 218 219)

101

• Enthalpy (H)
• H is a state function.
• At constant pressure
• ??
• ?? ?H and ?E are nearly identical numerically
  unless there
• is significant pressure-volume work that
  occurs from the
• reaction.
• Practice problems 6.5 (p. 221) and 6.6 (p. 222)

102
Lesson 17 Sections 6.6 6.8

• Summary of ?E and ?H
• Constant-pressure calorimetry
• At constant pressure,
• Styrofoam cup method (coffee cup calorimetry, p.
  223)
• ??
• ?? Practice problem 6.7 (p. 224)

103

• Relationships involving ?H
• Multiplication factor A 2B ? C ?H1
• 2A 4B ? 2C
• Reverse factor A 2B ? C ?H1
• C ? A 2B
• Addition factor A 2B ? C ?H1
• C ? 2D ?H2
• A 2B ? 2D
• Hesss law
• Practice problem 6.8 (p. 227)

104
(standard enthalpy) from (the standard
enthalpy of formation)

• Standard state
• ?? For gases the pure gas at 1.00 atm
• ?? For pure liquids and solids the pure
  substance in its most
• stable form at 1.00 atm and (generally)
  25.0 ?C
• ?? For solutions 1.00 M
• ?? Standard enthalpy (?H0) ?H for a process in
  which all
• reactants and products are in their
  standard states.

105

• Standard state (continued)
• ?? Standard enthalpy of formation ( )
  ?H for the
• formation of a compound from elements in
  their standard
• states.
• ?? for a pure substance in its most stable
  state is 0.
• Standard enthalpy change for a reaction (
  )
• Practice problems 6.9 (p. 229), 6.10 (p. 230),
  and 6.11 (p.
• 231)