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Chemistry 400 - General Chemistry Part 1

Course Topics

The topics for this course are typically allocated as follows (approximate number of lecture and lab hours spent on each topic are shown in the left columns):

Lec
Lab
Topic
1
0
Introduction, syllabus.
2
0
The scientific method; precision vs. accuracy; significant figures and scientific notation; unit prefixes and unit analysis; density.
4
0
Matter, element, compound, and mixture definitions; law of conservation of mass, law of constant composition, law of multiple proportions; Dalton’s atomic theory, Thomson’s atomic model, Rutherford’s nuclear atomic model; atomic notation, isotopes, atomic mass concept and calculations; nomenclature; introduction to the periodic table.
5
0
Moles, formula weight/molecular weight/molar mass, Avogadro number; elemental mass % from formula/experimental data; empirical/molecular formulas and combustion analysis of CHO compounds; writing, reading, and balancing chemical equations (formulas and symbols); stoichiometry calculations involving moles, mass, and solution concentration; limiting reactant stoichiometry; percentage yield.
4
0
Ionic compounds dissolving in water; types of electrolyte, ionization and dissociation concept; reaction types; molecular, full ionic, and net ionic equations; half reactions for reduction-oxidation (REDOX) reactions; balancing REDOX reaction with half-reactions.
4.5
0
Kinetic-molecular theory of gases; diffusion and effusion; gas pressure concept, measurement, and unit conversions; gas laws and calculations, including stoichiometry, density and Dalton’s law.
4
0
Internal energy, heat and work, first law of thermodynamics; enthalpy concept and diagrams (exothermic and endothermic); writing equations including enthalpy change and writing reactions for standard enthalpy changes; thermochemical stoichiometry; calorimetry (including heat capacity and specific heat capacity); Hess’s law; enthalpy of reaction calculated from standard heats of formation.
3.5
0
Basic knowledge of electromagnetic (EM) spectrum and energy, frequency, and wavelength calculations; Bohr model of the atom, emission spectrum concept, and calculations; history of modern atomic theory; quantum numbers.
3
0
Orbital diagrams, electronic configurations (including abbreviated), periodic table blocks, Pauli exclusion principle, Hund’s rule, and the Aufbau principle; periodic trends for atomic/ionic size, successive ionization energies, and electron affinity;
3
0
Ionic vs covalent materials (bonding and properties), including electronegativity and bond polarity; Lattice energy and the Born-Haber cycle.
3
0
Lewis structures for molecules and polyatomic ions, formal charges and resonance structures; valence shell electron pair repulsion (VSEPR) theory (AXE notation, electron group arrangement, molecular shape, molecule polarity); valence bond theory.
3
0
Orbital diagrams, hybridization, sigma and pi-bonds.
3
0
Phase changes and phase diagrams; heating curves and calculations; vapor pressure concept and temperature dependence; intermolecular forces and how they affect physical properties of pure substances and solutions.
5
0
Concentration units (molarity, molality, mass %, ppm, volume %, m/v %, mole fraction) and dilution calculations; thermodynamics of the solution process; types of solution; concepts and calculations of colligative properties (vapor pressure lowering, freezing point depression, boiling point elevation, and osmotic pressure).
4
0
Arrhenius and Brønsted-Lowry acids; pH scale concept and calculations; acid/base indicators; pH titration curves and calculations; autoionization of water.
2
0
Final exam
0
6
Determining mass, length, volume, density, assessing precision and accuracy, and volumetric and gravimetric techniques.
0
6
Determining a visible absorption spectrum, spectrophotometric determination of concentration, graphing, linear regression analysis, and dilutions.
0
6
Nomenclature.
0
6
Atomic spectroscopy: determining upper quantum levels, wavelength, frequency, and energy, graphing, and non-linear regression analysis.
0
9
VSEPR, hybridization, Lewis structures, and polarity.
0
9
Reaction types, conservation of mass, observations of chemical reactions, balancing, decantation, measuring, limiting reactants (qualitative), percent yield, and relative percent error.
0
6
Gravimetric decomposition analysis of a mixture of unknown percent composition.
0
6
Gravimetric and stoichiometric analysis of limiting reactants.
0
9
Electrolytes: strong, weak, and non electrolytes, molecular, ionic, net ionic equations, and spectator ions, observing and predicting solubility, measuring conductivity of individual solutions and solutions before and after mixing, predicting chemical reactions based on solubility and electrolyte rules.
0
9
Analysis of a mixture of unknown composition by gravimetric and volumetric methods (gas law), and assessing the quality of the two techniques.
0
9
Calorimetry: Determining the specific heat capacity of a metal, enthalpy of a reaction using Hess's Law, heat of fusion for solid water, and heat of solution.
0
12
Acid/base chemistry: titration, standardization, determining molar mass, strong acid/base titration curve.
0
9
Determining molar mass by freezing point depression.
0
6
Locker check in/out, safety, laboratory reports and procedures, syllabus.

Student Learning Outcomes

Upon completion of this course, the student will be able to:

  • successfully complete laboratory experiments (involving evaluation of experimental data and confirmation of physical constants) in a safe and timely manner, after receiving written and/or verbal instructions.
  • demonstrate the proper collection and recording of scientific measurements in tables with the correct units and number of significant figures (i.e. measuring mass, volume, temperature, length, and pressure), and the recording and evaluation of observations (physical and chemical changes and properties).
  • analyze and then solve chemical calculation problems that involve solids, solutions, or gases, in a clear and logical fashion; for example, stoichiometry, acid-base, and colligative property problems.
  • analyze and then solve chemical calculation problems that involve heat energy transfers in calorimeters or chemical reactions; for example, determining the heat of fusion, heat of solution, heat of reaction, and heat capacity.
  • construct balanced chemical equations from written descriptions of chemical reactions.
  • synthesize data into computer generated graphical outputs, and make predictions by interpolation, using linear and non-linear regression analyses.
  • evaluate errors related to experimental procedures, and assess their effects on experimental results.
  • predict the products of inorganic chemical reactions using solubility rules and the activity series, by assessing electrolyte strength, and by employing the fundamental rules of acid/base chemistry.
  • apply chemical naming rules to inorganic molecular and ionic compounds, acids, and simple straight-chain hydrocarbons.
  • predict changes in solution properties based on colligative property calculations; for example, changes in boiling point, freezing point, vapor pressure, and osmotic pressure.
  • explain how and why substances dissolve in other substances and perform calculations to evaluate solution concentration in various units (Molar, molal, % mass, mole fraction).
  • use the ideal gas law and the empirical or combined gas laws to predict temperature, pressure, volume, mass, or molar quantity of a gas.
  • explain and predict observable properties of gases (pressure, temperature, volume) from an understanding of the behavior of the individual particles in a gas.
  • produce Lewis structures of simple molecules and polyatomic ions and predict their shape and relative polarity.
  • label the parts of a phase diagram, use it to predict the temperature and pressure at which phase changes will occur, and construct heating curves that include phase changes.
  • analyze the structure of an atom and explain the origin of atomic emission spectra.