Review Slides

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Review Slides

• Some pictures, graphs and various concepts you
  may have forgotten

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Advice

• Get some quality sleep and get to 143 at 7 am for
  breakfast
• Look over the practice mc if you have time

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Multiple Choice Advice

• For multiple choice- You have 90 minutres for 75
  questions
• go over the test once and answer the quick
  questions you know
• 2nd pass- go through and answer calculations, and
  more time intensive problems
• DO NOT GUESS unless you can narrow the answers to
  two choices

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• Learn the ions,
• solubility rules, and strong acids and bases once
  and for all!

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• Electrolytic Properties
• Aqueous solutions, solutions in water, have the
  potential to conduct electricity.
• The ability of the solution to conduct depends on
  the number of ions in solution.
• There are three types of solution
• Strong electrolytes,
• Weak electrolytes, and
• Nonelectrolytes.

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Electrolytic Properties
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• Ionic Compounds in Water
• Ions dissociate in water.
• In solution, each ion is surrounded by water
  molecules.
• Transport of ions through solution causes flow of
  current.

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• Molecular Compounds in Water
• Molecular compounds in water (e.g., CH3OH) no
  ions are formed.
• If there are no ions in solution, there is
  nothing to transport electric charge.

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• Titrations

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• Heat Capacity and Specific Heat
• Calorimetry measurement of heat flow.
• Calorimeter apparatus that measures heat flow.
• Heat capacity the amount of energy required to
  raise the temperature of an object (by one
  degree). (J C-1)
• Molar heat capacity heat capacity of 1 mol of a
  substance. (J mol-1 C-1)
• Specific heat specific heat capacity heat
  capacity of 1 g of a substance. (J g-1 C-1)

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Bomb Calorimetry (Constant Volume Calorimetry)

• Reaction carried out under constant volume.
• Use a bomb calorimeter.
• Usually study combustion.

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Example Given Fe2O3 (s) 3 CO (g) ? 2 Fe (s)
3 CO2 (g) ?H -23 kJ 3 Fe2O3 (s) CO (g) ? 2
Fe3O4 (s) CO2 (g) ?H -39 kJ Fe3O4 (s) CO
(g) ? 3 FeO (s) CO2 (g) ?H 18 kJ What is ?H
for the following? FeO (s) CO (g) ? Fe (s)
CO2 (g)
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• Using Enthalpies of Formation of Calculate
  Enthalpies of Reaction
• For a reaction

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• Bohr Model
• Colors from excited gases arise because electrons
  move between energy states in the atom.

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• We can show that
• When ni gt nf, energy is emitted.
• When nf gt ni, energy is absorbed

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Representations of Orbitals
The s-Orbitals
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The p-Orbitals
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• Lattice energy the energy required to completely
  separate an ionic solid into its gaseous ions.
• Lattice energy depends on the charges on the ions
  and the sizes of the ions
• k is a constant (8.99 x 10 9 Jm/C2), Q1 and Q2
  are the charges on the ions, and d is the
  distance between ions.
• Lattice energy increases as
• The charges on the ions increase
• The distance between the ions decreases.

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• Examples Which in each pair would have the
  greatest lattice energy?
• LiF or LiCl
• NaCl or MgCl2
• KBr or KI
• MgCl2 or MgO
• CaO or NaF

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• Examples Which in each pair would have the
  greatest lattice energy?
• LiF or LiCl
• NaCl or MgCl2
• KBr or KI
• MgCl2 or MgO
• CaO or NaF

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• Resonance Structures
• Resonance structures are attempts to represent a
  real structure that is a mix between several
  extreme possibilities.

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• Resonance Structures
• Example in ozone the extreme possibilities have
  one double and one single bond. The resonance
  structure has two identical bonds of intermediate
  character.
• Common examples O3, NO3-, SO42-, NO2, and
  benzene.

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• Mathematically, if ?Hrxn is the enthalpy for a
  reaction, then
• We illustrate the concept with the reaction
  between methane, CH4, and chlorine
• CH4(g) Cl2(g) ? CH3Cl(g) HCl(g) ?Hrxn ?

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• In this reaction one C-H bond and one Cl-Cl bond
  gets broken while one C-Cl bond and one H-Cl bond
  gets formed.
• The overall reaction is exothermic which means
  than the bonds formed are stronger than the bonds
  broken.
• The above result is consistent with Hesss law.

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• When considering the geometry about the central
  atom, we consider all electrons (lone pairs and
  bonding pairs).
• When naming the molecular geometry, we focus only
  on the positions of the atoms.

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• Examples Determine the shape and angles about
  each atom
• 1
• 2

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• Examples Determine the shape and angles about
  each atom
• 1
• 2

Trigonal planar linear
Bent trigonal pyramid Trigonal planar
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Molecular Shape and Molecular Polarity
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• Examples Determine the hybridization on EACH
  atom

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• Examples Determine the hybridization on EACH
  atom

sp2 sp
sp3
sp3
sp2
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• The Pressure-Volume Relationship Boyles Law
• Mathematically
• A plot of V versus P is a hyperbola.
• Similarly, a plot of V versus 1/P must be a
  straight line passing through the origin.

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The Pressure-Volume Relationship Boyles Law
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• The Temperature-Volume Relationship Charless
  Law
• We know that hot air balloons expand when they
  are heated.
• Charless Law the volume of a fixed quantity of
  gas at constant pressure increases as the
  temperature increases (assuming all other factors
  are constant).
• Mathematically

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Collecting Gases over Water
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Example - 150.82 mL of an unknown gas is
collected over water at 27C and 1.032 atm. The
mass of the gas is 0.1644 g. What is the molar
mass of the gas? The vapor pressure of water at
27C is 26.74 torr
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Real Gases Deviations from Ideal Behavior

• The van der Waals Equation
• General form of the van der Waals equation

Corrects for molecular volume
Corrects for molecular attraction
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Surface Tension
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Hydrogen Bonding
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• Vapor Pressure and Boiling Point
• Liquids boil when the external pressure equals
  the vapor pressure.
• Temperature of boiling point increases as
  pressure increases.
• Two ways to get a liquid to boil increase
  temperature or decrease pressure.
• Pressure cookers operate at high pressure. At
  high pressure the boiling point of water is
  higher than at 1 atm. Therefore, there is a
  higher temperature at which the food is cooked,
  reducing the cooking time required.
• Normal boiling point is the boiling point at 760
  mmHg (1 atm).

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The Phase Diagrams of H2O and CO2
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• Covalent-Network Solids
• Forces covalent bonds.
• Atoms held together in large networks.
• Examples diamond, graphite, quartz (SiO2),
  silicon carbide (SiC), and boron nitride (BN).
• In diamond
• each C atom has a coordination number of 4 each
  C atom is tetrahedral there is a
  three-dimensional array of atoms.
• Diamond is hard, and has a high melting point
  (3550 ?C).

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Covalent-Network Solids
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Solute-Solvent Interaction
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• Pressure Effects

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• First Order Reactions

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Second Order Reactions
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Zero Order Rate Law

• Rate kA0 k
• Rate does not change with concentration.
• Integrated A -kt A0
• When A A0 /2 t t1/2
• t1/2 A0 /2k

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Temperature and Rate
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The Equilibrium Constant

• No matter the starting composition of reactants
  and products, the same ratio of concentrations is
  achieved at equilibrium.

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• The Direction of the Chemical Equation and Keq
• An equilibrium can be approached from any
  direction.
• Example
• has

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• In the reverse direction

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• Other Ways to Manipulate Chemical Equations and
  Keq Values
• The reaction
• has
• which is the square of the equilibrium constant
  for

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Strong Base-Strong Acid Titrations
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Buffered Solutions

• Composition and Action of Buffered Solutions
• A buffer consists of a mixture of a weak acid
  (HX) and its conjugate base (X-)
• The Ka expression is

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Gibbs Free Energy

• For a spontaneous reaction the entropy of the
  universe must increase.
• Reactions with large negative ?H values are
  spontaneous.
• How do we balance ?S and ?H to predict whether a
  reaction is spontaneous?
• Gibbs free energy, G, of a state is
• For a process occurring at constant temperature

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?G ?H - T?S
?H ?S ?G
- Always -
- Always
at high T at low T
- - at high T - at low T