General Chemistry Practice Exam Study Guide

Atomic Structure and Chemical Bonding

Orbital Overlap and Sigma Bonds

Atomic orbitals combine to form molecular bonds, with sigma (σ) bonds resulting from head-on overlap. The type of orbitals involved determines the bond's characteristics.

• Sigma Bond Formation: Occurs via direct overlap of orbitals along the internuclear axis.

• Example: In CH3CH3, the C–C sigma bond is formed by the overlap of two sp3 hybrid orbitals.

• Hybridization: Atoms can hybridize their s and p orbitals to form sp, sp2, or sp3 hybrids, affecting molecular geometry.

Equation:

• $\text{sp}^3 + \text{sp}^3 \rightarrow \sigma \text{ bond}$

Lone Pairs and Orbital Types

Lone pairs are non-bonding pairs of electrons localized on an atom, often occupying specific hybrid orbitals.

• Example: In PCl3, the lone pair on phosphorus occupies an sp3 orbital.

Electronic Structure and Quantum States

Ground State vs. Excited State

The ground state of an atom or molecule is its lowest energy configuration, with electrons in the lowest available orbitals.

• Excited State: Electrons occupy higher energy orbitals than in the ground state.

• Graphical Representation: The ground state is shown as the lowest point on an energy diagram.

Phase Changes and Equilibrium

Dynamic Equilibrium in Vaporization

Dynamic equilibrium occurs when the rate of vaporization equals the rate of condensation in a closed system.

• Key Point: At equilibrium, the rates of forward and reverse processes are equal, and the concentrations of vapor and liquid remain constant.

• False Statement Example: "The rates of evaporation and condensation are equal at equilibrium." (This is actually true; a false statement would be that only evaporation occurs.)

Intermolecular Forces and Surface Tension

Ranking Surface Tension

Surface tension depends on the strength of intermolecular forces. Molecules with stronger hydrogen bonding or dipole-dipole interactions have higher surface tension.

• Example: Water (H2O) has higher surface tension than ethanol (C2H5OH) due to stronger hydrogen bonding.

Thermochemistry

Endothermic and Exothermic Processes

Endothermic processes absorb heat from the surroundings, while exothermic processes release heat.

• Endothermic Example: Vaporization (liquid to gas) is endothermic.

• Exothermic Example: Condensation (gas to liquid) is exothermic.

Calculating Heat Required

To calculate the heat required for a phase change, use the enthalpy of vaporization or fusion.

• Equation:

$q = m \times \Delta H$

• Where q is heat, m is mass, and ΔH is enthalpy change per gram or mole.

Solutions and Concentration Calculations

Molarity and Dilution

Molarity (M) is the number of moles of solute per liter of solution. Dilution calculations use the formula:

• $M_1V_1 = M_2V_2$

• Example: To make 0.0500 M HCl from 250 mL of 10.0 M HCl, use the equation above to solve for the final volume.

Preparing Solutions

To prepare a solution of a given molarity, calculate the required mass of solute using:

• $\text{mass} = \text{moles} \times \text{molar mass}$

Stoichiometry and Chemical Quantities

Mole Calculations

The mole is a fundamental unit in chemistry, representing $6.022 \times 10^{23}$ entities (Avogadro's number).

• Number of Molecules: $\text{Number of molecules} = \text{moles} \times 6.022 \times 10^{23}$

• Empirical Formula: The simplest whole-number ratio of atoms in a compound.

Percent Composition and Empirical Formula

Percent composition is the mass percentage of each element in a compound. The empirical formula is determined from percent composition data.

• Example: A compound with 40% C, 6.7% H, and 53.3% O has the empirical formula CH2O.

Molecular Structure and Bonding

Lewis Structures and Bond Types

Lewis structures show the arrangement of atoms and electrons in a molecule. Sigma (σ) and pi (π) bonds are types of covalent bonds.

• Sigma Bonds: Formed by head-on overlap of orbitals.

• Pi Bonds: Formed by side-to-side overlap of p orbitals.

• Example: CH3CHO has 5 sigma bonds and 1 pi bond.

Hybridization and Molecular Geometry

Hybridization describes the mixing of atomic orbitals to form new hybrid orbitals, which determine molecular geometry.

• sp3 Hybridization: Tetrahedral geometry (e.g., methane).

• sp2 Hybridization: Trigonal planar geometry (e.g., ethene).

• sp Hybridization: Linear geometry (e.g., acetylene).

Molecular Orbital Theory

Bond Order Calculation

Bond order indicates the number of chemical bonds between a pair of atoms.

• Equation:

$\text{Bond Order} = \frac{\text{Number of bonding electrons} - \text{Number of antibonding electrons}}{2}$

• Example: For N2, bond order is 3.

Periodic Table and Chemical Symbols

Periodic Table Usage

The periodic table organizes elements by atomic number and properties. It is essential for identifying element symbols, atomic masses, and electron configurations.

• Groups: Vertical columns with similar chemical properties.

• Periods: Horizontal rows indicating energy levels.

Reference Tables

Abbreviations, Constants, and Formulas

Reference tables provide essential abbreviations, constants, and formulas for solving chemistry problems.

Abbreviation	Meaning
mol	mole
M	molarity
g	gram
L	liter
atm	atmosphere
K	Kelvin

Constant	Value
Avogadro's Number	$6.022 \times 10^{23}$
Gas Constant (R)	$0.0821 \text{ L atm mol}^{-1} \text{K}^{-1}$

Formula	Purpose
$PV = nRT$	Ideal Gas Law
$M_1V_1 = M_2V_2$	Dilution
$q = m \times \Delta H$	Heat Calculation

Sample Calculations

Stoichiometry Example

• How many grams of Na2SO4 are needed to make 1.50 mol of 0.300 M solution?

• Calculate moles: $1.50 \text{ mol} \times 142.04 \text{ g/mol} = 213.06 \text{ g}$

Empirical Formula Example

• Given percent composition, determine the empirical formula:

• Convert percentages to grams, then to moles, and find the simplest ratio.

Additional info:

• Some questions reference molecular orbital diagrams and hybridization, which are covered in chapters on chemical bonding and molecular structure.

• Reference tables and periodic table are standard tools for General Chemistry exams.