Topics for Final Exam
CHEM 1411 - General Chemistry I
Houston Community College System
Textbook: Chemistry, Ninth Edition, by Raymond Chang
About the Exam
1. Prepare for 35 multiple choice questions (two points each) and six (five points each) essay/show
work problems. Essay/show work problems will receive partial credit.
2. Bring a #2 lead pencil, a Scantron answer sheet, and a scientific calculator.
3. The final exam will be comprehensive covering chapters 1- 11.
4. Review problems worked out in class as well as those assigned from each chapter.
5. Values of constants and relatively complex equations will be provided. Values from
appendices will be provided. A periodic table, a table of solubility and a table of
electronegativities will be provided.
6. Remember that the final exam is system-wide and is mandatory. The final exam grade will
not be dropped.
CHAPTER 1 - Chemistry:The Study of Change
 Definition and examples of: matter, states of matter, mixtures (homogeneous and
heterogeneous), pure substances, elements, and compounds
 Law of constant composition/definite proportion
 Names and symbols of common elements
 Chemical and physical properties
 Chemical and physical changes
 Units of measurements (metric, SI, and British systems)
 Dimensional analysis and conversion factors in problem solving
 Density = Mass / Volume
 Temperature relationships and related problems: 0C = 5/9(0F - 32), K = 0C + 273.15
 Rules for determination of significant figures in a measured quantity
CHAPTER 2- Atoms, Molecules, and Ions
 Law of conservation of mass
 Atomic mass unit
 Modern view of atomic structure (electrons, protons and neutrons)
 Nuclear symbol for an atom
 Atomic number, mass number, and isotopes
 Molecules and ions
 Cations, anions, monoatomic and polyatomic ions
 Periodic table (periods, groups, metals, nonmetals, transition metals, and metalloids)
 Molecular, empirical and structural formula
 Chemical bonding: ionic bonds (metal and nonmetal), covalent bonds (among nonmetals)
 Naming ionic compounds, ionic charges: Group 1A: +1, 2A: +2, 3A: +3, 5A: -3, 6A: -2, 7A: -1
 Naming binary covalent compounds (CO2, N2O4, etc.)
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CHAPTER 3 – Mass Relationships in Chemical Reactions
 Balancing chemical equations
 Determination of average atomic mass (atomic weight) with isotopic distribution
 Calculation of formula weight (molar mass)
 The mole and related problems (Avogadro’s number = 6.022 x 1023)
 Calculation of percent composition
 Calculation of empirical formula and molecular formula from the percent composition
 Calculations based on the stoichiometry of chemical equations including limiting reactant,
theoretical yield and percent yield
CHAPTER 4 – Reactions in Aqueous Solutions
 Electrolytes (classification and example)
 Acids, bases, and salts (definition, classifications and examples)
 Double displacement and metathesis reactions (definitions and examples)
 Solubility tables and net ionic equations
 Oxidation and reduction ("redox") reactions
 Single displacement reactions and the activity series
 Oxidation numbers
 Solution concentration: Molarity = moles of solute per liter of solution; M = moles/L
 Dilution (V1M1 = V2M2)
 Titrations and related problems, mLA MA nH+ = mLB MB nOH–
CHAPTER 5 - Gases
 Kinetic-Molecular Theory
 Atmospheric pressure, units of pressure (atm, mm Hg or torr, Pa)
 Gas laws: Boyle’s law, Charles’s law, Avogadro’s law, Ideal-Gas equation, Gas densities and
molar mass, Dalton’s law of partial pressure, Graham’s law of effusion and diffusion
 Definition of STP conditions
 At STP, one mole of any ideal gas occupies 22.414 L
 Deviations from ideal behavior
CHAPTER 6 - Thermochemistry
 Definition of energy, kinetic energy and potential energy
 Units, joules and calories, 1 cal = 4.184 J
 First law of Thermodynamics
 Internal energy (E), E = q + w
 State functions
 Enthalpy (H), H = q at constant pressure, exothermic processes (negative H value) and
endothermic processes (positive H value)
 Hess’s law and related problems
 Specific heat and heat capacity
 H = q = specific heat x grams of substance x T
 Heat = heat capacity x T
 Calorimetry, problems dealing with coffee-cup and bomb calorimeter
 Calculating enthalpies of reactions using enthalpies of formation (Hfo)
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CHAPTER 7 – Quantum Theory and the Electronic Structure of Atoms
 Electromagnetic radiation
 c = wavelength x frequency (c = λν)
 Relationship between the energy of radiation and its frequency or wavelength
 Planck’s quantum theory, E = hν
 Photoelectric effect and the photon
 Bohr’s model of the hydrogen atom
 Principal quantum number n, En = –RH (1/ n2), n = 1, 2, 3, . . .
 Electronic transitions, E = RH (1/ni2 – 1/nf2 )
 Wave functions and electron density
 Quantum mechanical description of orbitals, their shapes and energies
 Shells (energy levels) and subshells (sublevels)
 Four quantum numbers (n, l, ml, ms)
 Pauli exclusion principle, Aufbau principle, and Hund’s rule
CHAPTER 8 – Periodic Relationships Among the Elements
 Atomic sizes, ionization energies and electron affinities
 Effective nuclear charge, Zeff
 The periodic table, groups and periods
 Physical and chemical properties of metals, nonmetals, and metalloids
 Overview of chemistry of the representative elements (Groups 1A–8A)
CHAPTER 9 – Chemical Bonding I: Basic Concepts
 Ionic bonding, lattice energy, ionic sizes (cations, anions, and isoelectronic ions)
 Covalent bonding, Lewis dot structures, the octet rule, formal charges, and resonance
 Exception to octet rule, less than an octet (e.g., BF3), "expanded octet" (e.g., PF5, SF6)
 Bond dissociation energy and related problems
 Electronegativity
 Polar and nonpolar covalent bonds
CHAPTER 10 – Chemical Bonding II: Molecular Geometry and Hybridization of Atomic Orbitals
 Molecular shape, bond angle and bond length
 The VSEPR model
 Electron domain (or “electron pair") geometry
 Molecular geometry
 Polarity in molecules; dipole moment
 Valence-bond theory, sigma and pi bonds
 Hybrid orbitals (sp, sp2, sp3, sp3d, sp3d2)
 Molecular orbital (MO) description of covalent bonding
 Molecular orbital energy level diagrams
 Bond order, diamagnetic and paramagnetic molecules
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CHAPTER 11- Intermolecular Forces and Liquids and Solids
 Characteristic properties of the three states of matter
 Intermolecular forces: ion-dipole forces, dipole-dipole forces, London dispersion forces, and
hydrogen bonding
 Phase changes and dynamic equilibrium
 Properties of liquids: vapor pressure, volatility, boiling point, surface tension, and viscosity
 Phase diagrams
 Heat of fusion and heat of vaporization
 The Clausius-Clapeyron equation
 Heating curves and related calculations
 Structures of solids: crystalline vs. amorphous solids
 Unit cells (primitive cubic, body-centered cubic, face-centered cubic)
 Calculating the density of a solid from the type of unit cell and the unit cell dimensions
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