General Chemistry Practice Exam Study Guide: Thermochemistry, Bonding, Gases, and Molecular Structure

Thermochemistry and Chemical Energy

Internal Energy and Enthalpy

The study of thermochemistry involves understanding how energy is transferred in chemical reactions and physical changes. Internal energy (E) is the total energy contained within a system, including both kinetic and potential energy at the molecular level.

• Internal energy of a system: The sum of the potential and kinetic energies of all the components.

• Enthalpy (H): A thermodynamic quantity defined as $H = E + PV$, where P is pressure and V is volume.

• Exothermic process: Releases heat to the surroundings; $riangle H$ is negative.

• Endothermic process: Absorbs heat from the surroundings; $riangle H$ is positive.

• Standard enthalpy change ($riangle H^ ext{0}$): The enthalpy change measured under standard conditions (1 atm, 25°C).

Example: The combustion of titanium with oxygen to produce titanium dioxide is an exothermic reaction.

Calculating Enthalpy Changes

• Hess's Law: The total enthalpy change for a reaction is the sum of the enthalpy changes for each step.

• Standard enthalpy of formation ($riangle H_f^ ext{0}$): The enthalpy change when one mole of a compound is formed from its elements in their standard states.

• Bond enthalpy: The energy required to break one mole of a particular type of bond in a gaseous molecule.

Equation:

$riangle H_ ext{reaction} = riangle H_ ext{products} - riangle H_ ext{reactants}$

Example: Calculating $riangle H$ for the reaction $2H_2(g) + O_2(g) \rightarrow 2H_2O(g)$ using bond enthalpies.

Heat Capacity and Calorimetry

• Specific heat capacity (c): The amount of heat required to raise the temperature of 1 gram of a substance by 1°C.

• Calorimetry: The measurement of heat flow in a chemical reaction, often using a coffee-cup or bomb calorimeter.

Equation:

$q = mc riangle T$

Where q = heat absorbed or released, m = mass, c = specific heat, $riangle T$ = change in temperature.

Chemical Bonding and Molecular Structure

Types of Chemical Bonds

Chemical bonds are the forces that hold atoms together in compounds. The main types are ionic, covalent, and metallic bonds.

• Ionic bond: Formed by the transfer of electrons from one atom to another, resulting in oppositely charged ions.

• Covalent bond: Formed by the sharing of electrons between atoms.

• Bond order: The number of chemical bonds between a pair of atoms (single, double, triple).

Example: The maximum number of bonds that can be formed between two atoms is three (triple bond).

Lewis Structures and Resonance

• Lewis structure: A diagram showing the arrangement of valence electrons among atoms in a molecule.

• Resonance: Occurs when more than one valid Lewis structure can be drawn for a molecule.

Example: The best Lewis structure for OCl2 and resonance forms for NCO-.

Hybridization and Molecular Geometry

• Hybridization: The mixing of atomic orbitals to form new hybrid orbitals suitable for bonding.

• Common types: sp, sp2, sp3, sp3d, sp3d2

• VSEPR theory: Predicts the shape of molecules based on electron pair repulsion.

• Bond angles: Determined by the geometry (e.g., tetrahedral = 109.5°, trigonal planar = 120°).

Example: The geometry of SF4 is described as seesaw according to VSEPR theory.

Gases: Properties and Behavior

Gas Laws

The behavior of gases is described by several laws relating pressure, volume, temperature, and amount.

• Boyle's Law: $P_1V_1 = P_2V_2$ (at constant temperature)

• Charles's Law: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$ (at constant pressure)

• Ideal Gas Law: $PV = nRT$

• Avogadro's Law: Equal volumes of gases at the same temperature and pressure contain equal numbers of molecules.

Example: Calculating the volume of 1.0 mol of gas at STP (Standard Temperature and Pressure).

Gas Stoichiometry

• Molar volume at STP: 22.4 L/mol for an ideal gas.

• Density of gases: $\text{Density} = \frac{PM}{RT}$, where M is molar mass.

Example: Determining the density of CO2 at STP.

Intermolecular Forces and Phase Changes

Types of Intermolecular Forces

• London dispersion forces: Weak forces present in all molecules, especially nonpolar ones.

• Dipole-dipole interactions: Occur between polar molecules.

• Hydrogen bonding: Strong dipole-dipole interaction involving H bonded to N, O, or F.

Example: H2O exhibits hydrogen bonding, which affects its boiling point.

Phase Changes

• Boiling point: The temperature at which a liquid becomes a gas.

• Melting point: The temperature at which a solid becomes a liquid.

Example: Comparing boiling points of PH3, H2S, and H2O.

Atomic Structure and Periodicity

Electronic Structure of Atoms

• Electron configuration: The arrangement of electrons in an atom's orbitals.

• Paramagnetic: Atoms or ions with unpaired electrons; attracted to magnetic fields.

• Diamagnetic: Atoms or ions with all electrons paired; weakly repelled by magnetic fields.

Example: O2 is paramagnetic due to unpaired electrons in its molecular orbital diagram.

Periodic Trends

• Atomic radius: Decreases across a period, increases down a group.

• Ionization energy: Increases across a period, decreases down a group.

• Electronegativity: Tendency of an atom to attract electrons in a bond.

Example: The trend in decreasing F-X-F angle in compounds such as BeF2, BF3, and CF4.

Tables

Sample Table: Standard Enthalpy of Formation Values

Compound	$\triangle H_f^0$ (kJ/mol)
H2O(l)	-285.8
CO2(g)	-393.5
CH4(g)	-74.8
O2(g)	0

Additional info: Table values inferred from standard enthalpy tables commonly used in general chemistry.

Additional Info

• Some questions involve calculations using calorimetry, gas laws, and enthalpy changes. Practice with these equations is essential for exam success.

• Understanding molecular geometry and hybridization is key for predicting molecular shapes and properties.

• Recognizing intermolecular forces helps explain physical properties such as boiling and melting points.