# Advanced Chemistry

## Course Overview

This Advanced Chemistry course is designed to be the equivalent of the general chemistry course taken during the first year of college. Students successfully completing this course will be endowed with an exceptional understanding of the fundamentals of chemistry and achieve proficiency in solving chemical problems. This course will contribute to the development of each student’s ability to think critically and to express their ideas, in both oral and written fashion, with clarity and logic.

Laboratory investigation is a central pillar of the Advanced Chemistry course. Labs will include an emphasis on experimental procedures. Each week during the first semester students will participate in a Friday lab elective and should expect to spend one to two hours each week on lab work outside of regular class time.

## Course Content

### Unit 1: Lab Safety

The main idea of this unit is to have students work safely when handling chemicals. Students will learn about various types of safety equipment present in a chemistry lab. They will also learn how to safely handle chemicals and equipment in a chemistry lab.

### Unit 2: First Year Chemistry Review

Students will get a brief overview on concepts they learned in their previous chemistry course. This refresher unit quickly covers:

- classifying matter according to a scheme,
- understanding the difference between measured numbers and exact numbers,
- solving problems using various units of measurement,
- relating atomic theory with atomic structure based on indirect evidence,
- describing atomic structure and the properties of atoms, molecules and matter,
- describing key terms,
- using the periodic table,
- comparing empirical formulas from molecular formulas,
- calculating empirical and molecular formulas from experimental data,
- discussing differences between ionic and molecular compounds,
- naming inorganic compounds,
- writing balanced chemical equations to describe a chemical reaction for synthesis, decomposition, single replacement, metathesis, redox, combustion, and acid-base reactions,
- calculating the molar mass of a substance,
- using the molar mass and Avogadro’s number to interconvert among mass, moles and number of particles of a substance, work problems involving mole concepts, molarity, percent composition, empirical formulas, and molecular formulas,
- and solving stoichiometric problems.

### Unit 3: Aqueous Reactions & Solution Stoichiometry

Students study chemical reactions more deeply. Students will describe the nature of aqueous solutions through water as a solvent and strong and weak electrolytes as solutes, identify common strong and weak acids, determine the solubility of ionic compounds from general solubility rules, write molecular, ionic, and net ionic equations, and identify metathesis reactions that go to completion (formation of a gas, precipitate or molecular product). Students will also predict the products for reactions that are redox, neutralization, and precipitation reactions, and perform stoichiometric calculations on acid-base volumetric (titrations), precipitation, and redox reactions.

### Unit 4: Thermochemistry

Students will investigate enthalpy, calorimetry, and Hess’s law. Students will describe the energy flow between a system and its surroundings, explain the significance of the first law of thermodynamics and use the law to calculate ∆E, q, and w, define and distinguish among heat, temperature, work, energy, kinetic and potential energy, calculate the enthalpy change associated with phase changes, determine the enthalpy change or stoichiometric quantities for thermochemical equations, describe a state function, use standard enthalpies of formation to calculate ∆H for a reaction, solve calorimetry problems using q = mc∆T, use Hess’s Law to calculate the enthalpy change for a reaction, use standard enthalpies of formation to calculate ∆H for a reaction, interconvert among calories and Joules.

### Unit 5: Electronic Structure of Atoms

Students will investigate Bohr theory, the wave mechanical behavior of an atom, and quantum mechanics. Students will determine various aspects of the electromagnetic spectrum including relative frequencies, wavelengths and energies, quantitatively and qualitatively relate frequency, wavelength and speed of a wave, describe Planck’s concept of quantized energy and calculate the energy of a photon using the relationship λ = hν, relate Bohr’s model of the atom to the quantum theory, calculate the energy difference resulting from the change in energy levels of an electron, state the meaning and possible values of the quantum numbers and assign the quantum numbers to a given sublevel or orbital, and use the quantum numbers, Aufbau Principle, and Hund’s Rule to assign an electron configuration for a given element or ion.

### Unit 6: Chemical Bonding

Students will investigate periodic trends by interpreting trends within the periodic table in terms of: atomic radii, ionization energy, electron affinity, and ionic radii, distinguishing between metals and nonmetals and semimetals, describe how effective nuclear charge varies with position on the periodic table, and compare the relative energies of atomic energy levels and of sublevels. Students will also use periodic trends and electronegativity to predict bond types, compare and contrast different types of bonding, compare bond strength with ionic sizes of elements on the periodic table, relate the enthalpy dissociation of ionic bonding to bond strength, draw Lewis structures for various atoms, ions, and molecules, draw resonance structures for various molecules, use formal charges to determine the most likely resonance structure, compare oxidation numbers and formal charges for atoms in a molecule, relate electronegativity values to bond polarity, and compare and contrast bond distance and bond energy for single and multiple bonds. While investigating molecular shapes, students will use the VSEPR model to predict molecular geometry, determine molecular polarity using dipole moments of individual bonds, compare VSEPR structures to the hybridization of orbitals, compare and contrast sigma and pi bonds, predict the number of sigma and pi bonds in a structure, compare and contrast valence bond theory with molecular orbital theory, and contrast molecular orbitals (delocalized) and orbitals derived from the valence-bond theory.

### Unit 7: Gases

Students will examine the relationship between pressure, volume and temperature of ideal gases, apply Charles’, Boyle’s, Gay-Lusaac’s, Dalton’s, and the ideal gas laws quantitatively and qualitatively, analyze the kinetic molecular theory, use Graham’s Law to relate the molar masses of gases to their rates or times of effusion, describe how real gases deviate from ideal behavior, show how Van Der Waals’ equation allows for real conditions, use the ideal gas law equation to calculate the density or molar mass of a gas and solve stoichiometric calculations at standard and non-standard conditions, and use the molar volume at STP conditions in calculations.

### Unit 8: Intermolecular Forces

In this unit, students will describe the intermolecular forces such as dipole-dipole, hydrogen bonding, and London dispersion forces, describe the effects that IM forces have on the properties of liquids and solids such as melting point, boiling point, vapor pressure, viscosity, state of matter, phase changes and solubility, characterize the processes of evaporation, condensation, sublimation, fusion at the particle level, distinguish among ionic, molecular, network covalent and metallic solids with regard to particle structure, physical properties, and inter- and intramolecular forces, apply the concepts of unit cells and crystal lattices for solids to calculations involving atomic radii, volume, density or identity, explain the relationship of boiling point to vapor pressure, using phase diagrams, and be able to calculate energy of various phase changes for water and carbon dioxide.

### Unit 9: Solution Chemistry

Students will define and describe solution formation, energy changes, salvation and hydration as they relate to solutions, describe the unique characteristics of water due to its extensive hydrogen bonding, compare and contrast saturated, unsaturated and supersaturated solutions; and be able to interpret graphs and charts of solubility, make calculations involving molarity, molality, mass percent and mole fractions as a means of expressing concentration, analyze the effects of colligative characteristics on the properties of solutions such as electrolytes vs. non-electrolytes, solve problems involving freezing point depression, boiling point elevation, vapor pressure lowering, and increase in osmotic pressure, use Raoult’s Law to relate vapor pressure lowering to solute mole fraction, explain different properties of colloidal systems such as size of particles, Tyndall’s effect, and Brownian’s motion.

### Unit 10: Kinetics

Students will describe the collision theory and the requirements for an effective collision, list factors that affect the rate of a reaction, use experimental data to determine the rate law and rate order of a reaction and to predict a reaction mechanism, interpret graphs of endothermic and exothermic reactions identifying the activation energy, enthalpies, and the reaction course with and without a catalyst, determine a zero, first or second order reaction from graphical analysis of concentration vs. time plots, explain the role of a catalyst in a reaction and distinguish between homogeneous and heterogeneous catalysts, predict how temperature and concentration affect the rate of a reaction over time, use data to calculate the half life of a reaction, and generally describe the meaning and use of the Arrhenius equation and be able to solve problems involving activation energy and the Arrhenius equation.

### Unit 11: Equilibrium & Acid/Base Theory

Students will discuss the concept of equilibrium, write the equilibrium expression for a given equilibrium system in terms of concentrations or pressures, calculate values for any of the equilibrium constants, given Kc or Kp, calculate the equilibrium concentrations for the species in the system, predict the changes in equilibrium that will occur when various stresses are placed on the system (Le Chatelier’s Principle): concentration change, temperature change, pressure change, and addition of a catalyst, calculate pH, pOH, pK, Ka, Kb, ionization constant, percent ionization, Ksp, use the reaction quotient, Q, to determine the initial direction of a reaction needed to establish equilibrium, write the Kw expression for water, explain the common ion effect, identify strong and weak acids and bases and write dissociation equations for each, predict the direction of equilibrium from knowledge of the strength of the acid-base conjugate pair in water, solve problems involving concentrations of substances necessary to produce a precipitate, and concentrations of ions involved in simultaneous equilibrium, graphically determine pKa for a weak acid from a titration curve, given the composition of a buffer system, determine its pH before and after the addition of known amounts of strong acid or base, determine the proportions in which a weak acid and its conjugate base should be mixed to give a buffer of a specified pH, and use the Henderson-Hasselbach equation in equilibrium (buffered) reactions.

### Unit 12: Thermodynamics

Students will discuss the laws of thermodynamics, define entropy, second law of thermodynamics, PV work, enthalpy and free energy, use Hess’s Law to solve problems of energy, entropy and free energy, relate the signs of ∆H and ∆S to determine the direction of a reaction/determine the spontaneity of a reaction, predict the sign of entropy change for a given reaction, apply the relationship between ΔS, surroundings, ΔH, and Temperature, describe how the signs of ΔH, ΔS, and ΔG relate to the spontaneity of a reaction, calculate the free energy change for a given reaction, and calculate ΔS for reactions or phase changes from Absolute Entropy values.

### Unit 13: Electrochemistry

Students will assign oxidation states to various elements in compounds and molecules, recognize reactions that undergo redox reactions by comparing their oxidation states, recognize oxidation/reducing agents in various reactions, balance redox reactions using the half reaction method, balance redox reactions in acidic or basic solutions, draw and label parts of an e-cell showing electron flow, differentiate between galvanic and electrolytic cells, use the Nernst equation to calculate the EMF at non-standard conditions, apply Faraday’s Law to electrolytic cells in calculating amount of products formed, time required, or current required, use the table of standard reduction potentials to determine cell voltages, explain the electrochemical nature of lead storage batteries, corrosion (anode and cathode protection), and fuel cells, define terms such as redox, anode, cathode, oxidizing agent, reducing agent, emf, electrode, Faraday’s Law, voltaic cells, and galvanic cells.

### Unit 14: Nuclear Chemistry

Students will research alpha, beta, and gamma radiation by describing each type of radiation in aspect of mass, charge, penetrating power, and symbol. Students will also write balanced nuclear equations for radioactive decay and nuclear transformations, given a table of nuclear masses, calculate Δ mass for a nuclear reaction and relate it to the energy change, ΔE (binding energy), generally describe the functioning, reactions, and positive and negative aspects of fission and fusion reactors, and generally discuss the rates of radioactive decay.