1-1. Assess a current, technological situation; listing underlying assumptions and consequences.
1-2. Evaluate the consequences of present courses of action.
1-3. Generate alternative actions to the present courses of action.
1-4. Select and explore specific problems and alternatives.
1-5. Evaluate specific solutions to problems and alternatives.
2-1. Differentiate between safe and unsafe procedures, applications, and methods of disposal of chemicals.
2-2. Chose the appropriate safety equipment for specific laboratory situations.
2-3. Decide which safety and emergency procedures to follow in case of particular accidents including fires and hazardous material spills.
2-4. Demonstrate proper methods for carrying and moving chemicals and equipment.
3-1. Identify the base units of the SI system and describe the standards for each.
3-2. Describe the concept of a derived quantity and its units, and identify the dimension (combination of base units) for any derived quantity,initially including area, volume and density.
3-3. Using dimension analysis, determine whether an equation is dimensionally valid, and establish the dimensions of a quantity.
3-4. Explain and give examples of the system of subdivision used in the SI system, including the use of prefixes to represent powers of ten.
3-5.Use conversion factors to convert quantities from one metric unit to another, and also between metric and English units.
4-1. Report measurements to the limit of the measuring instrument.
4-2. Report the degree of uncertainty of a measurement, and carry out mathematical operations with measurements containing stated uncertainties.
4-3. Determine the significant digits in a recorded measurement, and carryout mathematical operations using these measurements with answers rounded off to the correct number of significant digits.
5-1. Describe the general properties of matter.
5-2. Classify matter according to whether it is an element, a compound, or a mixture.
5-3. Define, calculate, and experimentally determine density for a variety of substances.
5-4. Distinguish between physical and chemical properties of matter.
5-5. Determine chemical and physical properties of substances by carrying out physical and chemical changes.
5-6. Use physical methods to separate the components of a mixture.
5-7. Write symbols for 55 common elements.
6-1. Describe the postulates of the modern atomic theory.
6-2. Relate the Laws of Conservation of Mass, Definite Composition, and Multiple Proportions to atomic theory.
6-3. Locate and describe the main components of the atom as used in chemistry.
6-4. Define isotope, and relate atomic number, mass number, and number of atomic particles to each other, and interpret andd write isotope symbols.
6-5. Calculate atomic mass from isotope abundances.
6-6. Design a relative mass scale similar to that of the atomic mass scale
7-1. Identify basic differences between atoms, molecules, and ions and classify compounds as being ionic or molecular.
7-2. Write names of ionic and binary covalent compounds from their formulas using older system of prefixes and suffixes and the newer IUPAC system.
7-3. Use ion-charge method to write formulas for ionic compounds.
7-4. Write formulas for binary covalent compounds.
8-1. Describe the events leading to the modern day arrangement of the periodic table.
8-2. Describe periodic trends of the general characteristics of metals, nonmetals, and metalloids.
8-3. Experimentally determine an activity series of metals.
9-1. Write and balance chemical equations when given reactants and products.
9-2. Classify those equations that come under the heading of synthesis, decomposition, replacement, and ionic reactions.
9-3. Predict the products of chemical reactions when given the reactants.
9-4. Define oxidation and reduction, and identify any species undergoing oxidation or reduction, and identify the oxidizing and reducing agents.
9-5. Use solubility rules to predict the formation of insoluble products, and the activity series to predict the occurrence of replacement reactions.
9-6. Relate complete and incomplete combustion to oxidation.
9-7. Carry out examples of each kind of reaction, and write balanced equations for each.
10-1. Relate Avogadro's number to the atomic mass scale.
10-2. Convert numbers of atoms and molecules to masses by using the mole, and vice versa.
10-3. State the masses of atoms or molecules in terms of molar masses.
10-4. Calculate, and prepare solutions of known molarity.
11-1. Distinguish between empirical and molecular formulas.
11-2. Experimentally determine the empirical formula of an ionic compound.
11-3. Calculate percentage composition of a compound from its formula, and from experimental data.
11-4. Calculate empirical and molecular formulas from experimental data.
12-1. Calculate mass relationships based on balanced chemical equations.
12-2. Determine the limiting reactant, and the theoretical yield for chemical reactions.
12-3. Experimentally determine the mole ratio for a chemical reaction, and use it to determine the equation or the reaction.
13-1. Relate the Law of Conservation of Energy to Chemical Processes.
13-2. Describe the transfer of energy between reaction systems and their surroundings.
13-3. Distinguish between exothermic and endothermic reactions, and relate them to the enthalpy of a system.
13-4. State the three Laws of Thermochemistry, and apply them to calculations of enthalpy changes.
13-5. Calculate H for a reaction using specific heats and heats of formation.
13-6. Experimentally measure heat flow using a calorimeter, and use the measurements to write a thermochemical equation for the reaction.
13-7. Calculate "q" and E for a system according to the First Law of Thermodynamics.
14-1. Define and calculate H and S for a reaction.
14-2. Use the Gibbs-Helmholtz equation to calculate the free energy change for a reaction.
14-3. Given, or having calculated H and S, determine the temperature at which a reaction is at equilibrium at one atmosphere.
14-4. Describe how the signs of H, S, and G relate to the spontaneity of a reaction.
14-5. Determine the maximum efficiency of a heat engine using the Second Law of Thermodynamics.
15-1. State and interpret the postulates of the Quantum Theory.
15-2. Relate energy differences, wavelength, and frequencies of EMR.
15-3. Describe the atomic spectrum of hydrogen in terms of the Bohr model, and calculate energy transitions for Lyman and Balmer series.
15-4. Describe the wave nature of electrons according to deBroglie, Planck, and Schrodinger.
15-5. Identify the four quantum numbers and relate each in terms of energy differences and mathematical interpretation.
15-6. Write electron configurations for elements.
15-7. Use Hund's rule to draw orbital diagrams for electrons in an atom.
15-8. Experimentally determine the wavelengths and frequencies line spectrum of selected elements.
16-1. Write nuclear equations showing a, b and g-emissions.
16-2. Relate nuclear decay to first order kinetics.
16-3. Relate E = mc2 to nuclear thermochemistry.
16-4. Compare nuclear fission to nuclear fusion.
16-5. Consider use of nuclear power with reactor accidents and waste disposal.
17-1. Describe the formation of cations and anions, and relate it to electronegativity and position on the periodic table.
17-2. Relate H of ionic compounds to their lattice energies.
17-3. Write Lewis structures to show the covalent bonding in molecules and polyatomic
17-4. Determine the polarity of covalent bonds from electronegativities.
17-5. Compare bond lengths of covalent bonds.
17-6. Use bond energies to calculate &DeltaH for the formation of molecular compounds.
17-7. Experimentally determine the number of ionizable hydrogens in a compound.
18-1. Use VSEPR model to predict the geometric shape of simple molecules and polyatomic ions.
18-2. Construct models of molecules and polyatomic ions to illustrate their predicted geometric shapes.
18-3. Predict the polarity of molecules by using the VESPR model for molecules containing polar covalent bonds.
17-4. Describe covalent bonding in terms of atomic orbitals: sp, sp2, sp3 hybrid orbitals, sigma and pi bonds, and expanded octets.
18-5. Use the molecular orbital theory to explain the bonding in paramagnetic molecules.
18-6. Experimentally relate solubilities of solutes in solvents to their polarities.
19-1. Write the expression for Kc from the balanced equation for a reaction involving gases.
19-2. Calculate Kc from equilibrium concentrations of all species, or from original concentrations of all species and the equilibrium concentration of one species.
19-3. Predict the direction a chemical system will move to reach equilibrium when the value of Kc is known.
19-4. Predict the equilibrium concentration of one species when given those of all other species when the value of Kc is known.
19-5. Predict the equilibrium concentrations of all species when given their original concentrations and when the value of Kc is known.
19-6. Using LeChatelier's Principle, predict the effect of a change in the number of moles, volume, or temperature upon the position of an equilibrium.
19-7. Experimentally determine Kc for an equilibrium system.
19-8. Relate the standard free energy change for a reaction to the equilibrium constant.
20-1. Relate the acidic and basic properties of aqueous solutions to the dissociation of water.
20-2. Carry out calculations involving pH and pOH.
20-3. Compare strong and weak acids.
20-4. Compare strong and weak bases.
20-5. Predict acidity or basicity of salt solutions (cations and anions).
20-6. Write equations for reactions for reactions between strong acids-strong bases, strong acids-weak bases, and weak acids-strong bases.
20-7. Carry out acid-base titrations and write equations for the reactions.
20-8. Compare Arrhenius, Bronsted-Lowry, and Lewis theories of acids.
21-1. Write the equilibrium expression for dissociation of weak acids and calculate Ka.
21-2. Calculate [H+] in solutions of weak acids when given Ka.
21-3. Calculate [H+] in buffered solutions.
21-4. Write the equilibrium expression for the dissociation of weak bases and calculate Kb.
21-5. Calculate [OH-] in solutions of weak bases when given Ka.
21-6. Relate Ka and Kb.
22-2. Distinguish between electrolytes and nonelectrolytes.
22-3. Carry out calculations involving solution concentrations in mole fractions, molality, and/or molarity.
22-4. Describe the factors that affect the solubility of a solute in a particular solvent.
22-5. Determine the concentration of an unknown solution by using the Spec 20 and Beer's Law
22-6. Describe the colligative properties of solutions.
22-7. Experimentally determine the molar mass of an unknown solute by freezing point depression and boiling point elevation.
23-1. Balance redox reactions by half-reaction method.
23-2. Experimentally carry out a redox titration.
23-3. Relate corrosion to oxidation-reduction and how it may be prevented.
24-1. Express electricical values in terms of SI units.
24-2. Use a cell diagram to represent an electrochemical cell.
24-3. Construct an electrochemical cell and measure its output.
24-4. Use the Nernst Equation in electrochemical calculations.
24-5. Compare electrochemical cells with fuel cells.
25-1. Determine the order of a reaction when given the initial rate as a function of concentration of a reaction.
25-2. Calculate, for a first order reaction, the concentration of a reactant after a given time when given the original concentration and the rate constant.
25-3. Calculate, for a first order reaction, the time required for the concentration to drop by a given amount when given the rate constant.
25-4. When given either the half-life or the rate constant for a first order reaction, claculate the other quantity.
25-5. Experimentally determine the order of a reaction.
25-6. Relate ozone depletion to CFCs and chlorine photochemistry.
25-7. Compare homogeneous with heterogeneous catalysts, and their affects on reaction rates.
26-1. Define pressure and relate to kinetic theory.
26-2. Describe the effect of temperature on pressure and volume of gases.
26-3. Apply mole-volume relationship of gases to gas-phase reactions.
26-4. Describe the relationship between pressure and volume of gases (Boyle's Law).
26-5. Combine Boyle's, Charles, and Avogadro's laws of gases into the ideal gas law.
26-6. Describe difussion of gases and relate to Graham's Law.
26-7. Describe mixtures of gases in terms of Dalton's Law of Partial Pressure.
26-8. Relate density of gases to molar volume and molar mass.
26-9. Describe the operation of mercury barometers.
26-10. Relate motion of molecules to the Boltzman distribution and temperature.
26-11. Compare the behavior of real gases to the ideal and relate to the van der Waals equation.