CHM 152 Exam 1 (Chapter 11 to 13) – What to Review…

What to Study:

1.    Chapter 11 : States of Matter: Liquids and Solids

1.    Comparison of Gases, Liquids, and Solids

a.     Recall the definitions of gas, liquid, and solid given in Section 1.4.

b.    Compare a gas, a liquid, and a solid using a kinetic molecular theory description.

2.    Phase Transitions

a.     Define change of state (phase transition).

b.    Define melting, freezing, vaporization, sublimation, and condensation.

c.     Define vapor pressure.

d.    Describe the process of reaching a dynamic equilibrium that involves the vaporization of a liquid and condensation of its vapor.

e.     Define boiling point.

f.      Describe the process of boiling.

g.    Define freezing point and melting point.

h.    Define heat (enthalpy) of fusion and heat (enthalpy) of vaporization.

i.      Calculate the heat required for a phase change of a given mass of substance.

3.    Vapor pressure

a.     Describe the general dependence of the vapor pressure (ln P) on the temperature (T).

b.    State the Clausius–Clapeyron equation (the two-point form).

c.     Calculate the vapor pressure at a given temperature.

d.    Calculate the heat of vaporization from vapor pressure.

4.    Phase Diagrams

a.     Define phase diagram.

b.    Describe the melting-point curve and the vapor-pressure curves (for the liquid and the solid) in a phase diagram.

c.     Define triple point.

d.    Define critical temperature and critical pressure.

e.     Relate the conditions for the liquefaction of a gas to its critical temperature.

5.    Properties of Liquids; Surface Tension and Viscosity

a.     Define surface tension.

b.    Describe the phenomenon of capillary rise.

c.     Define viscosity.

6.    Intermolecular Forces; Explaining Liquid Properties

a.     Define intermolecular forces.

b.    Define dipole–dipole force.

c.     Describe the alignment of polar molecules in a substance.

d.    Define London (dispersion) forces.

e.     Relate the properties of liquids to the intermolecular forces involved.

f.      Define hydrogen bonding.

g.    Identify the intermolecular forces in a substance.

h.    Determine relative vapor pressures based on intermolecular attractions.

i.      Determine trends in the boiling point based on intermolecular attractions.

7.    Classification of Solids by Type of Attraction of Units

a.     Define molecular solid, metallic solid, ionic solid, and covalent network solid.

b.    Identify types of solids.

c.     Relate the melting point of a solid to its structure.

8.    Crystalline Solids; Crystal Lattices and Unit Cells

a.     Define crystalline solid and amorphous solid.

b.    Define crystal lattice and unit cell of a crystal lattice.

c.     Define simple cubic unit cell, body-centered cubic unit cell, and face-centered cubic unit cell.

d.    Determine the number of atoms in a unit cell.

e.     Define hexagonal close-packed structure and cubic close-packed structure.

f.      Note the common structures (face-centered cubic and body-centered cubic) of metallic solids.

9.    Calculations Involving Unit-Cell Dimensions

a.     Calculate atomic mass from unit-cell dimension and density.

b.    Calculate unit-cell dimension from unit-cell type and density.

2.    Chapter 12: Solutions

1.    Types of Solutions

a.     Define solute and solvent.

b.    Define miscible fluid.

c.     Provide examples of  gaseous solutions, liquid solutions, and solid solutions

2.    Solubility and the Solution Process

a.     List the conditions that must be present to have a saturated solution, to have an unsaturated solution, and to have a supersaturated solution.

b.    Describe the factors that make one substance soluble in another.

c.     Determine when a molecular solution will form when substances are mixed.

d.    Learn which conditions must be met to create an ionic solution.

3.    Effects of Temperature and Pressure on Solubility

a.     State the general trends for the solubility of gases and solids with temperature.

b.    Explain how the solubility of a gas changes with temperature.

c.     Apply Henry’s law.

4.    Ways of Expressing Concentration

a.     Define colligative property.

b.    Define molarity.

c.     Define mass percentage of solute.

d.    Calculate mass percentage of solute.

e.     Define molality.

f.      Calculate the molality of a solute.

g.    Define mole fraction.

h.    Calculate the mole fractions of components.

i.      Convert molality to mole fraction.

j.      Convert mole fraction to molality.

k.    Convert molality to molarity.

l.      Convert molarity to molality.

5.    Vapor Pressure of a Solution

a.     Explain vapor-pressure lowering of a solvent.

b.    State Raoult’s law.

c.     Calculate vapor-pressure lowering.

d.    Describe an ideal solution.

6.    Boiling-Point Elevation and Freezing-Point Depression

a.     Define boiling-point elevation and freezing-point depression.

b.    Calculate boiling-point elevation and freezing-point depression.

c.     Calculate the molecular mass of a solute from molality.

d.    Calculate the molecular weight from freezing-point depression.

7.    Osmosis

a.     Describe a system where osmosis will take place.

b.    Calculate osmotic pressure.

c.     Calculate the molecular weight from osmotic pressure.

8.    Colligative Properties of Ionic Solutions

a.     Determine the colligative properties of ionic solutions.

3.    Chapter 13: Chemical Kinetics

9.    Definition of a Reaction Rate

a.     Define reaction rate.

b.    Explain instantaneous rate and average rate of a reaction.

c.     Explain how the different ways of expressing reaction rates are related.

d.    Calculate average reaction rate.

10. Experimental Determination of Rate

a.     Describe how reaction rates may be experimentally determined.

11. Dependence of Rate on Concentration

a.     Define and provide examples of a rate law, rate constant, and reaction order.

b.    Determine the order of reaction from the rate law.

c.     Determine the rate law from initial rates.

12. Change of Concentration with Time

a.     Learn the integrated rate laws for first-order, second-order, and zero-order reactions.

b.    Use an integrated rate law.

c.     Define half-life of a reaction.

d.    Learn the half-life equations for first-order, second-order, and zero-order reactions.

e.     Relate the half-life of a reaction to the rate constant.

f.      Plot kinetic data to determine the order of a reaction.

13. Temperature and Rate; Collision and Transition-State Theories

a.     State the postulates of collision theory.

b.    Explain activation energy (E_a).

c.     Describe how temperature, activation energy, and molecular orientation influence reaction rates.

d.    State the transition-state theory.

e.     Define activated complex.

f.      Describe and interpret potential-energy curves for endothermic and exothermic reactions.

14. Arrhenius Equation

a.     Use the Arrhenius equation.

15. Reaction Mechanisms

a.     Define elementary reaction, reaction mechanism, and reaction intermediate.

b.    Determine the rate law from initial rates.

c.     Write the overall chemical equation from a mechanism.

d.    Define molecularity.

e.     Give examples of unimolecular, bimolecular, and termolecular reactions.

f.      Determine the molecularity of an elementary reaction.

g.    Write the rate equation for an elementary reaction.

16.  The Rate Law and the Mechanism

a.     Explain the rate-determining step of a mechanism.

b.    Determine the rate law from a mechanism with an initial slow step.

c.     Determine the rate law from a mechanism with an initial fast, equilibrium step.

17.  Catalysis

a.     Describe how a catalyst influences the rate of a reaction.

b.    Indicate how a catalyst changes the potential-energy curve of a reaction.

c.     Define homogeneous catalysis and heterogeneous catalysis.

d.    Explain enzyme catalysis.