General Chemistry Exam Study Guide: Key Concepts and Applications

General Chemistry Exam Study Guide

Introduction

This study guide covers essential topics in General Chemistry, including unit conversions, chemical nomenclature, atomic structure, stoichiometry, molar mass calculations, chemical reactions, and properties of matter. Each section provides definitions, explanations, formulas, and examples to help students prepare for exams.

Unit Conversions and Physical Quantities

Speed and Unit Conversion

Unit conversion is a fundamental skill in chemistry, allowing for the translation of measurements between different systems.

• Speed Conversion: To convert miles per hour (mi/h) to meters per second (m/s), use the following relationships:

$1 \text{ mi} = 1.609 \text{ km} = 1609 \text{ m}$ $1 \text{ hour} = 3600 \text{ s}$ $\text{Speed (m/s)} = \text{Speed (mi/h)} \times \frac{1609 \text{ m}}{3600 \text{ s}}$

• Example: Convert 3060 mi/h to m/s: $3060 \times \frac{1609}{3600} = 1368.775 \text{ m/s}$

Density

Density is a physical property defined as mass per unit volume.

• Formula: $\rho = \frac{m}{V}$

• Units: Commonly expressed in g/cm3 or kg/m3.

• Application: Used to compare substances and solve mass-volume problems.

Chemical Nomenclature

Oxyanions and Acids

Chemical nomenclature involves systematic naming of compounds, including oxyanions and acids.

• Oxyanions: Polyatomic ions containing oxygen. Naming depends on the number of oxygens:

  • -ate: More oxygen (e.g., SO42-: sulfate)

  • -ite: Fewer oxygen (e.g., SO32-: sulfite)

  • per- ... -ate: Most oxygen (e.g., ClO4-: perchlorate)

  • hypo- ... -ite: Least oxygen (e.g., ClO-: hypochlorite)

• Acids: Named based on the anion:

  • -ide: Hydro- ... -ic acid (e.g., HCl: hydrochloric acid)

  • -ate: ... -ic acid (e.g., H2SO4: sulfuric acid)

  • -ite: ... -ous acid (e.g., HNO2: nitrous acid)

• Example: ClO3- is correctly named as chlorate.

Atomic Structure and Isotopes

Neutral Atoms and Ions

Atoms are composed of protons, neutrons, and electrons. Neutral atoms have equal numbers of protons and electrons.

• Representation: $^{A}_{Z}\text{X}$, where A = mass number, Z = atomic number, X = element symbol.

• Ions: Atoms or molecules with a net charge due to loss or gain of electrons.

• Requirement: An ion must have at least one electron.

Isotopes and Atomic Weight

Isotopes are atoms of the same element with different numbers of neutrons. Atomic weight is the weighted average of isotopic masses.

• Calculation: $\text{Atomic weight} = \sum (\text{isotope mass} \times \text{fractional abundance})$

• Example: For iron with 64.5% $^{56}$Fe and 35.5% $^{59}$Fe: $\text{Atomic weight} = (56 \times 0.645) + (59 \times 0.355) = 57.065$ amu

Stoichiometry and Chemical Reactions

Balancing Chemical Equations

Balancing equations ensures the conservation of mass and charge.

• Steps:

  1. Write the unbalanced equation.

  2. Count atoms of each element on both sides.

  3. Add coefficients to balance atoms.

• Example: $\text{CH}_4(g) + \text{Br}_2(l) \rightarrow \text{CBr}_4(s) + \text{HBr}(g)$ Balance Br: 4 Br in CBr4, so coefficient for Br2 is 2.

Limiting Reactant and Excess Calculation

In reactions, the limiting reactant is consumed first, determining the maximum product formed. Excess reactant remains after the reaction.

• Steps:

  1. Convert masses to moles.

  2. Determine mole ratio from balanced equation.

  3. Identify limiting reactant.

  4. Calculate excess reactant remaining.

• Example: 4.00 g Ga reacts with 5.50 g As to form GaAs. Calculate excess.

Mole Concept and Molar Mass

Mole Calculations

The mole is a fundamental unit for counting particles in chemistry. Avogadro's number ($N_A$) defines the number of entities per mole.

• Avogadro's Number: $N_A = 6.022 \times 10^{23}$

• Mole Calculation: $\text{Moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

• Example: 12.0 g of LiCl ($\text{molar mass} = 42.4$ g/mol): $\text{Moles} = \frac{12.0}{42.4} = 0.283$ moles

Molar Mass Determination

Molar mass is the mass of one mole of a substance, calculated by summing atomic masses.

• Example: Potassium nitrite (KNO2): $\text{Molar mass} = 39.10 + 14.01 + (2 \times 16.00) = 85.11$ g/mol

Percent Composition and Empirical/Molecular Formulas

Percent by Mass

Percent composition expresses the mass percentage of each element in a compound.

• Formula: $\%\text{Element} = \frac{\text{mass of element in 1 mol compound}}{\text{molar mass of compound}} \times 100$

• Example: Nitrogen in Ca(NO3)2:

  • Mass of N: $2 \times 14.01 = 28.02$ g

  • Molar mass: $40.08 + 2 \times (14.01 + 3 \times 16.00) = 164.10$ g/mol

  • Percent N: $\frac{28.02}{164.10} \times 100 = 17.08\%$

Empirical and Molecular Formulas

The empirical formula shows the simplest whole-number ratio of atoms; the molecular formula shows the actual number of atoms in a molecule.

• Steps:

  1. Convert percent composition to grams (assume 100 g sample).

  2. Convert grams to moles.

  3. Divide by smallest number of moles to get ratio.

  4. Multiply to get whole numbers if necessary.

  5. Compare empirical formula mass to molar mass to find molecular formula.

• Example: Cyclohexane: 85.6% C, 14.4% H, molar mass 84.2 g/mol. Empirical formula: CH2; Molecular formula: C6H12

Properties of Matter

Comparison of Densities

Density comparisons are useful for predicting mass relationships between substances.

• Formula: $\text{Relative mass} = \frac{\text{density}_1}{\text{density}_2}$

• Example: Titanium (4.506 g/cm3) vs. Lead (11.34 g/cm3): $\frac{11.34}{4.506} = 2.517$ times heavier

Electrochemistry and Fundamental Constants

Charge, Current, and Frequency

Electrochemistry involves the movement of electrons and the relationship between charge, current, and frequency.

• Charge of Electron: $1.602 \times 10^{-19}$ C

• Current: $I = \frac{q}{t}$

• Frequency: $f = \frac{I}{e}$, where $e$ is the elementary charge

• Example: 4.0 A current: $f = \frac{4.0}{1.602 \times 10^{-19}} = 2.5 \times 10^{19}$ Hz

Tables

Provided Formulae

Formula	Description
$\rho = \frac{m}{V}$	Density: mass divided by volume
$N_A = 6.022 \times 10^{23}$	Avogadro's number: entities per mole

Sample Polyatomic Ions and Acids

Ion/Acid	Name
ClO3-	Chlorate
SO42-	Sulfate
NO2-	Nitrite
HNO2	Nitrous acid
HBrO4	Perbromic acid

Sample Isotope Abundance Calculation

Isotope	Abundance (%)	Mass (amu)
$^{56}$Fe	64.5	56
$^{59}$Fe	35.5	59

Atomic weight: $56 \times 0.645 + 59 \times 0.355 = 57.065$ amu

Additional info:

• Some questions reference alternate universes or hypothetical scenarios to test understanding of isotope calculations.

• All calculations should use significant figures as appropriate for exam settings.