Study Guide: CHEM 134 Exam 1 (Chapters 1–4) – Matter, Atoms, Molecules, and Chemical Reactions

Exam 1 Study Guide Overview

This study guide covers foundational topics in General Chemistry, focusing on matter, measurement, atomic structure, molecules, compounds, and chemical reactions. Mastery of these concepts is essential for understanding the principles of chemistry and for success in subsequent topics.

Chapter 1: Matter, Measurement, and Problem Solving

Classification and Properties of Matter

• Matter is anything that has mass and occupies space. It can be classified as a pure substance or a mixture.

• Pure substances have a fixed composition (elements or compounds), while mixtures contain two or more substances physically combined.

• Physical changes do not alter the chemical composition (e.g., melting, boiling), while chemical changes result in new substances (e.g., combustion, rusting).

Measurement and Units

• SI base units include the meter (length), kilogram (mass), and second (time).

• Measurements must include both a number and a unit.

• Significant figures reflect the precision of a measurement. When performing calculations, the number of significant figures in the result should reflect the least precise measurement.

Problem Solving and Dimensional Analysis

• Dimensional analysis (factor-label method) is used to convert between units using conversion factors.

• Example: To convert 5.0 cm to meters: $5.0\ \text{cm} \times \frac{1\ \text{m}}{100\ \text{cm}} = 0.050\ \text{m}$

Energy and Its Forms

• Energy is the capacity to do work or transfer heat. It exists as kinetic energy (motion) and potential energy (position or composition).

• Law of Conservation of Energy: Energy cannot be created or destroyed, only transformed.

Chapter 2: Atoms and Elements

Atomic Theory and Structure

• Atoms are the fundamental units of matter, composed of protons, neutrons, and electrons.

• Key experiments: Thomson (discovery of the electron), Millikan (electron charge), Rutherford (nuclear model of the atom).

• Atomic number ($Z$): Number of protons in the nucleus.

• Mass number ($A$): Total number of protons and neutrons.

• Isotopes: Atoms of the same element with different numbers of neutrons.

Elements and the Periodic Table

• Elements are organized in the periodic table by increasing atomic number.

• Groups (columns) and periods (rows) reflect recurring chemical properties.

• Metals, nonmetals, and metalloids are classified based on their properties.

Atomic Mass and Calculations

• Atomic mass is the weighted average of all naturally occurring isotopes of an element.

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

Key Particles and Ions

• Ions are atoms or molecules with a net charge due to loss or gain of electrons.

• Cations: Positively charged ions (loss of electrons).

• Anions: Negatively charged ions (gain of electrons).

Chapter 3: Molecules and Compounds

Chemical Bonds and Formulas

• Chemical bonds hold atoms together in compounds. Main types: ionic (transfer of electrons) and covalent (sharing of electrons).

• Molecular formula: Shows the exact number of each type of atom in a molecule.

• Empirical formula: Shows the simplest whole-number ratio of atoms.

• Example: Hydrogen peroxide, $\text{H}_2\text{O}_2$ (molecular), $\text{HO}$ (empirical).

Naming Compounds

• Ionic compounds: Name the cation first, then the anion (e.g., NaCl: sodium chloride).

• Covalent compounds: Use prefixes to indicate the number of atoms (e.g., CO2: carbon dioxide).

• Acids: Named based on the anion (e.g., HCl: hydrochloric acid).

Types of Chemical Reactions

• Synthesis (combination), decomposition, single displacement, double displacement, and combustion.

• Example: Combustion of methane: $\text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O}$

Chapter 4: Chemical Reactions and Chemical Quantities

Balancing Chemical Equations

• Law of Conservation of Mass: The number of atoms of each element must be the same on both sides of the equation.

• Steps: Write formulas, count atoms, balance one element at a time, check work.

• Example: $2\text{H}_2 + \text{O}_2 \rightarrow 2\text{H}_2\text{O}$

Stoichiometry

• Relates quantities of reactants and products using balanced equations.

• Mole concept: $1\ \text{mol} = 6.022 \times 10^{23}$ particles (Avogadro's number).

• Use molar mass to convert between grams and moles: $\text{moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

• Limiting reactant: The reactant that is completely consumed first, limiting the amount of product formed.

• Theoretical yield: Maximum amount of product possible.

• Percent yield: $\text{Percent yield} = \frac{\text{actual yield}}{\text{theoretical yield}} \times 100\%$

Types of Chemical Equations

• Be able to write and interpret balanced equations for synthesis, decomposition, combustion, and reactions of alkali metals and halogens.

Summary Table: Key Concepts and Calculations

Concept	Definition/Formula	Example
Atomic Number ($Z$)	Number of protons in nucleus	$Z_{\text{C}} = 6$ for carbon
Mass Number ($A$)	Number of protons + neutrons	$A_{\text{C}} = 12$ for $^{12}\text{C}$
Mole	$1\ \text{mol} = 6.022 \times 10^{23}$ particles	1 mol H2O = $6.022 \times 10^{23}$ molecules
Molar Mass	g/mol	H2O: 18.02 g/mol
Percent Yield	$\frac{\text{actual yield}}{\text{theoretical yield}} \times 100\%$	If actual = 8.0 g, theoretical = 10.0 g, percent yield = 80%

Additional info:

• This guide is based on a study guide for CHEM 134, covering the first four chapters of a standard General Chemistry curriculum.

• Students should be familiar with basic laboratory techniques, significant figures, and the use of scientific notation.