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General and Inorganic Chemistry II

Course NumberCHM 112
Lab Hours45
Lecture Hours60
Course DescriptionPrerequisites: READING LEVEL 2 and WRITING LEVEL 2 and MATH LEVEL 6, and CHM 111 with a minimum grade of C (2.0). Studies chemical equilibria in water, including ionization solubility, complexion, acid-base phenomena, and oxidation reducing equilibria. Discusses the principles of electrochemistry, chemical thermodynamics, chemical kinetics and special topics such as the descriptive study of metals and non-metals. Meets the needs of chemistry majors. Credit may earned in CHM 112 or CHM 112H but not in both. (60-45)

Outcomes and Objectives

Demonstrate an understand of the behavior of liquids and solids.
  1. Describe the different types of solids and their properties.
  2. Describe the concept of crystal lattices and be aware that different types of crystal lattices exist.

Apply information about solution preparation and concentration.
  1. Explain the terms solution, solvent, solute, colligative property, saturated, unsaturated, and supersaturated as they apply to the formation of solutions
  2. Perform and apply calculations using molarity, molality, mole fraction, mole percent, weight percent, and parts per million
  3. Explain the difference between a miscible and immiscible solution
  4. Explain the relationship between lattice energy and enthalpy of hydration to enthalpy of solution
  5. Perform calculations for the solubility of a gas in a solvent using Henry's Law
  6. Apply Le Chatelier's Principle to the change in solute solubility with temperature changes
  7. Use Raoult's law to calculate the mole fraction of a solute or solvent or the effect of a solute on solvent vapor pressure
  8. Define the concept of freezing point depression and boiling point elevation and perform calculations of change in freezing or boiling point, the molal freezing and boiling point elevation constants, and molecular mass using this information
  9. Explain the effect of ionic compounds on colligative properties and use the van't Hoff factor in the appropriate calculations
  10. Define osmotic pressure and use it in calculations of pressure, concentration, and temperature

Perform calculations about chemical kinetics.
  1. Explain the concept of reaction rate
  2. Use experimental information to determine the average and instantaneous reaction rate
  3. Write a rate equation for a reaction when given the appropriate experimental information
  4. Describe and use the relationships between reactant concentration and time for zero-order, first-order, and second-order reactions
  5. Explain the concept of half-life and perform calculations for first and second order reactions
  6. Explain the collision theory of reaction rates
  7. Explain the relationship between activation energy to the rate and thermodynamics of a reaction
  8. Describe the effect of molecular orientation, temperature, and concentration on reaction rate
  9. Apply the Arrhenius equation to calculations
  10. Describe the elementary steps of a mechanism, give their molecularity, and define the rate-determining step
  11. Define homogeneous and heterogeneous catalysts and describe their effects on the activation energy and mechanism of a reaction

Perform calculations involving equilibrium.
  1. Describe the nature and characteristics of the state of equilibrium
  2. Write an equilibrium constant expression for any chemical reaction, realizing that the concentrations of solids and solvents are not included
  3. Explain the difference and relationship between Kc and Kp and calculate each one if the other is known.
  4. Determine the K value for an equation when coefficients are changed or the reaction is reversed
  5. Describe the relationship between the magnitude of K and the position of the equilibrium
  6. Calculate the equilibrium constant when the equilibrium concentrations of the reactants and products are given
  7. Calculate the concentration of a reactant or product at equilibrium when the equilibrium constant is provided
  8. Apply Le Chatelier's principle to predict the effect of a disturbance on a chemical equilibrium
  9. Explain the effect on a reaction mechanism and the kinetics of a reaction if one step in the mechanism involves an equilibrium

Apply concepts concerning acids and bases.
  1. Define and apply the Lewis, Bronsted, and Arrhenius acid-base theories
  2. Recognize and write balanced equations for the ionization of mono and polyprotic acids in water
  3. Define and apply the concept of an amphoteric substance
  4. Describe the autoionization of water and its role in acid-base chemistry
  5. Recognize the Bronsted acids and bases in a reaction and identify the conjugate partner of each
  6. Identify and list common strong and weak acids and bases
  7. Calculate the pH of a solution using the hydroxide ion or hydronium ion concentration
  8. Calculate the pH of a solution from experimental information about ka values and concentrations
  9. Describe the acid-base properties of salts and calculate the pH of a salt solution
  10. Calculate the pH of a polyprotic acid

Predict reactions between acids and bases.
  1. Calculate the pH at the equivalence point in the reaction of a strong acid with a strong or weak base or the reaction between a strong base with a weak acid
  2. Predict the effect of the addition of a common ion on the pH of a weak acid or base solution
  3. Describe the function of a buffer solution and write instructions for preparation of a buffer solution with a designated pH
  4. Use the Henderson-Hasselbach equation to calculate the pH of a buffer with a given composition or predict the change in pH if the composition of the buffer changes
  5. Calculate the pH during an acid-base titration
  6. Explain how an indicator functions in an acid-base titration
  7. Describe the difference between a titration curve for a strong acid/strong base and a titration curve where one of the titrants is weak

Apply principles of solubility and precipitation equilibria.
  1. Write the equilibrium reaction and equilibrium constant expression for any slightly soluble salt
  2. Calculate Ksp, solubility, and the formation of a precipitate from the appropriate experimental information
  3. Calculate the solubility of a salt in the presence of a common ion
  4. Devise a method for separating ions in solution from one another using Ksp values
  5. Calculate the equilibrium constant for the net reaction for a situation in which two or more equilibrium processes are occurring in solution
  6. Understand that hydrolysis increases the solubility of a salt when the anion is the conjugate base of a weak acid
  7. Understand that the solubility of slightly soluble salts may be affected by pH or the formation of a complex ion

Apply principles of thermodynamics.
  1. Describe the difference between the information provided by kinetics and thermodynamics
  2. Define entropy as a measure of disorder in a system, predict the sign and calculate the entropy change for a reaction or change of state
  3. Predict whether a reaction is spontaneous (product-favored) using entropy and enthalpy changes
  4. Use the relationship between Gibbs Free Energy and entropy and enthalpy and perform the appropriate calculations
  5. Describe and apply the relationship between the free energy change and the equilibrium constant for a reaction
  6. Explain how a reactant-favored reaction can become product-favored by coupling with another strongly product-favored reaction

Apply principles of electron transfer reactions.
  1. Define and use the terms electrochemical cell, electrolysis, electrode, electrolyte, salt bridge, anode, and cathode
  2. Balance equations for oxidation-reduction reactions in acidic or basic solutions using the half-reaction approach
  3. Predict what will occur in an electrochemical cell if both half reactions are given as reductions
  4. Calculate the standard electrode potential for a cell reaction and predict whether the reaction will be spontaneous as written
  5. Understand and apply the Nernst equation to calculate the cell potential under nonstandard conditions
  6. Calculate the equilibrium constant for a reaction from the standard electrode potential

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