ACS General Chemistry Practice Test 1A – Study Guide

Acids & Bases

Definitions and Properties

Acids and bases are fundamental concepts in chemistry, described by several theories including Arrhenius, Brønsted-Lowry, and Lewis definitions.

• Arrhenius Acid: Produces H+ ions in aqueous solution.

• Arrhenius Base: Produces OH- ions in aqueous solution.

• Brønsted-Lowry Acid: Proton donor.

• Brønsted-Lowry Base: Proton acceptor.

• Lewis Acid: Electron pair acceptor.

• Lewis Base: Electron pair donor.

Example: NH3 acts as a Lewis base by donating a pair of electrons to H+ to form NH4+.

Atomic Structure

Atoms, Ions, and Isotopes

Atoms consist of protons, neutrons, and electrons. The atomic number (Z) is the number of protons, and the mass number (A) is the sum of protons and neutrons.

• Isotopes: Atoms with the same number of protons but different numbers of neutrons.

• Ions: Atoms or molecules that have gained or lost electrons, resulting in a net charge.

Example: The mass number of an atom with 17 protons and 18 neutrons is 35.

Bonding & Molecular Structure

Covalent and Ionic Bonds

Chemical bonds form when atoms share or transfer electrons to achieve stable electron configurations.

• Ionic Bond: Transfer of electrons from a metal to a nonmetal.

• Covalent Bond: Sharing of electrons between nonmetals.

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

Example: NaCl is an ionic compound, while H2O is covalent.

Chemical Reactions & Equations

Balancing and Types of Reactions

Chemical equations must be balanced to obey the law of conservation of mass. Common reaction types include synthesis, decomposition, single replacement, double replacement, and combustion.

• Balancing Equations: Ensure the same number of each atom on both sides of the equation.

• Combustion Reaction: Hydrocarbon reacts with O2 to produce CO2 and H2O.

Example: $2H_2 + O_2 \rightarrow 2H_2O$

Chemical Quantities & Stoichiometry

Mole Concept and Calculations

The mole is a counting unit in chemistry. Stoichiometry involves calculations based on balanced chemical equations.

• Molar Mass: The mass of one mole of a substance (g/mol).

• Percent Composition: The percent by mass of each element in a compound.

Example: $\text{Percent N in NaNO}_3 = \frac{14.0}{85.0} \times 100\%$

Gases

Gas Laws and Properties

Gases are described by several laws relating pressure, volume, temperature, and amount.

• Ideal Gas Law: $PV = nRT$

• Boyle's Law: $P_1V_1 = P_2V_2$ (at constant T and n)

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

Example: Calculate the pressure of a gas in a manometer using the difference in mercury levels.

Thermochemistry

Heat, Work, and Enthalpy

Thermochemistry studies energy changes in chemical reactions, focusing on heat (q), work (w), and enthalpy (ΔH).

• First Law of Thermodynamics: $\Delta E = q + w$

• Enthalpy Change (ΔH): Heat change at constant pressure.

Example: Exothermic reactions release heat (ΔH < 0), endothermic absorb heat (ΔH > 0).

Chemical Kinetics

Reaction Rates and Catalysts

Kinetics deals with the speed of chemical reactions and factors affecting them.

• Rate Law: $\text{Rate} = k[A]^m[B]^n$

• Catalyst: Substance that increases reaction rate without being consumed.

Example: Increasing temperature generally increases reaction rate.

Chemical Equilibrium

Dynamic Equilibrium and Le Châtelier's Principle

At equilibrium, the rates of forward and reverse reactions are equal. Le Châtelier's Principle predicts how a system at equilibrium responds to disturbances.

• Equilibrium Constant (K): $K = \frac{[\text{products}]}{[\text{reactants}]}$

• Le Châtelier's Principle: A system at equilibrium will shift to counteract a change in concentration, temperature, or pressure.

Example: Increasing the concentration of a reactant shifts equilibrium toward products.

Electrochemistry

Redox Reactions and Electrochemical Cells

Electrochemistry involves oxidation-reduction (redox) reactions and the generation of electrical energy.

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Standard Reduction Potential (E°): Indicates the tendency of a species to be reduced.

Example: In a galvanic cell, electrons flow from the anode (oxidation) to the cathode (reduction).

Periodic Properties

Trends in the Periodic Table

The periodic table organizes elements by increasing atomic number and similar chemical properties.

• Atomic Radius: Increases down a group, decreases across a period.

• Ionization Energy: Energy required to remove an electron; increases across a period.

• Electronegativity: Increases across a period, decreases down a group.

Example: Fluorine is the most electronegative element.

Laboratory Techniques & Safety

Common Procedures and Precautions

Laboratory safety and proper technique are essential for accurate and safe experimentation.

• Always add acid to water, not water to acid.

• Use appropriate glassware for measurements.

• Dispose of chemicals according to safety guidelines.

Example: To prepare a solution, dissolve the solute in less than the final volume, then dilute to the mark.

Molecular Geometry

VSEPR Theory and Molecular Shapes

The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the shapes of molecules based on electron pair repulsion.

• Linear: 180° bond angle (e.g., CO2).

• Tetrahedral: 109.5° bond angle (e.g., CH4).

• Trigonal Planar: 120° bond angle (e.g., BF3).

Example: SO2 has a bent molecular geometry due to lone pairs on the central atom.

Sample Table: Electronegativity Values

Element	Electronegativity
F	4.0
O	3.5
N	3.0
Cl	3.0
H	2.1

Sample Table: Standard Reduction Potentials

Half-Cell	E° (V)
F2 + 2e- → 2F-	2.87
Cu2+ + 2e- → Cu	0.34
Zn2+ + 2e- → Zn	-0.76
Na+ + e- → Na	-2.71

Additional info: This study guide covers the main topics found in a typical General Chemistry ACS practice test, including acids and bases, atomic structure, bonding, chemical reactions, stoichiometry, gases, thermochemistry, kinetics, equilibrium, electrochemistry, periodic properties, laboratory techniques, and molecular geometry. Tables are included for electronegativity and standard reduction potentials as referenced in the test.