General Chemistry Exam 1 Review: Measurements, Atoms, Compounds, and Reactions

Exam 1 Review: General Chemistry

Overview

This study guide covers foundational topics in General Chemistry, including measurements, atomic theory, chemical compounds, chemical reactions, and quantitative problem solving. Each section summarizes key concepts, definitions, and formulas essential for exam preparation.

Chapter 2: Measurements and Problem Solving

Types of Measurements

• Qualitative Measurements: Descriptions without numerical values (e.g., color, texture).

• Quantitative Measurements: Measurements with numerical values (e.g., mass, volume).

Units and SI System

• Base Units: Standard units for measurement (meter, kilogram, second, mole, liter).

• Volume: Measured in liters (L) or milliliters (mL).

• SI Prefixes: Multipliers such as kilo- (103), centi- (10-2), milli- (10-3).

Uncertainty, Accuracy, and Precision

• Uncertainty: All measurements have some degree of uncertainty.

• Accuracy: How close a measurement is to the true value.

• Precision: How close repeated measurements are to each other.

Significant Figures

• Definition: Digits in a measurement that are known with certainty plus one estimated digit.

• Rules: Nonzero digits are always significant; zeros may or may not be significant depending on their position.

Scientific Notation

• Expresses numbers as a product of a coefficient and a power of ten (e.g., $6.02 \times 10^{23}$).

Dimensional Analysis

• Conversion Factors: Used to convert between units.

• Dimensional Analysis: Systematic approach to problem solving using conversion factors.

• Physical Property: Characteristics that can be measured without changing the substance's identity.

Chapter 4: Atoms & Elements

Atomic Theory

• Atoms: The smallest unit of an element that retains its properties.

• Subatomic Particles:

  • Protons: Positively charged, found in the nucleus.

  • Neutrons: Neutral, found in the nucleus.

  • Electrons: Negatively charged, found outside the nucleus.

Atomic Number and Mass Number

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

• Mass Number (A): Total number of protons and neutrons ($A = Z + N$).

Isotopes

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

Chemical Symbols and Periodic Table

• Element Symbol: One or two-letter abbreviation for an element.

• Periodic Table: Organizes elements by atomic number and properties.

• Groups/Families: Vertical columns with similar chemical properties.

• Periods: Horizontal rows.

Classification of Elements

• Main Group Elements (A): Groups 1A-8A.

• Transition Metals (B): Groups in the center of the table.

• Metals, Nonmetals, Metalloids:

  • Metals: Conduct electricity, malleable, shiny.

  • Nonmetals: Poor conductors, brittle.

  • Metalloids: Properties intermediate between metals and nonmetals.

Common Groups

• Alkali Metals (IA)

• Alkaline Earth Metals (IIA)

• Halogens (VIIA)

• Noble Gases (VIIIA)

Monatomic Ions

• Single atom ions (e.g., Na+, Cl-).

Chapter 6: Ionic & Molecular Compounds and Chemical Naming

Types of Compounds

• Ionic Compounds: Formed from cations and anions.

• Molecular Compounds: Formed from nonmetals sharing electrons.

Formulas and Nomenclature

• Cations: Positively charged ions (Type 1: fixed charge, Type 2: variable charge).

• Anions: Negatively charged ions.

• Naming Ionic Compounds:

  • Type 1: Cation + Anion (ending in -ide).

  • Type 2: Cation (charge in Roman numerals) + Anion (ending in -ide).

• Polyatomic Ions: Ions composed of multiple atoms (e.g., SO42-).

Naming Acids

• Binary Acids: Hydrogen + nonmetal (e.g., HCl).

• Oxyacids: Hydrogen + polyatomic ion (e.g., HNO3, H2SO4).

• Naming Rules:

  • -ate → -ic acid (e.g., nitrate → nitric acid)

  • -ite → -ous acid (e.g., nitrite → nitrous acid)

  • per- → per-ic acid; hypo- → hypo-ous acid

Naming Molecular Compounds

• Use prefixes to indicate number of atoms (mono-, di-, tri-, etc.).

• First element + prefix (if more than one) + second element (prefix + -ide).

• Example: CO2 is carbon dioxide.

Chapter 6: Chemical Composition

Avogadro's Number and the Mole

• Avogadro's Number: $6.02 \times 10^{23}$ units/mol.

• Mole: The amount of substance containing Avogadro's number of entities.

Molar Mass and Formula Weight

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

• Formula Weight: Sum of atomic weights in a chemical formula.

Percent Composition

• Definition: Percent by mass of each element in a compound.

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

Empirical and Molecular Formulas

• Empirical Formula: Simplest whole-number ratio of atoms.

• Molecular Formula: Actual number of atoms in a molecule.

• Calculation: Based on percent composition and molar mass.

Chapter 6: Lewis Structures

Lewis Structures and Covalent Bonding

• Lewis Structure: Diagram showing valence electrons and bonds in a molecule.

• Octet Rule: Atoms tend to have eight electrons in their valence shell.

• Bond Types:

  • Single Bond: One shared pair of electrons.

  • Double Bond: Two shared pairs.

  • Triple Bond: Three shared pairs.

• Lone Pairs: Non-bonding pairs of electrons.

Electronegativity and Bond Polarity

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

• Bond Polarity: Difference in electronegativity leads to polar covalent bonds.

Molecular Shape

• Determined by electron pair geometry (VSEPR theory).

Chapter 7: Chemical Reactions

Types of Chemical Reactions

• Combination (Synthesis): $A + B \rightarrow AB$

• Decomposition: $AB \rightarrow A + B$

• Combustion: $C_xH_yO_z + O_2 \rightarrow CO_2 + H_2O$

• Displacement:

  • Single Displacement: $A + BC \rightarrow AC + B$

  • Double Displacement: $AB + CD \rightarrow AD + CB$

• Redox (Oxidation-Reduction): Electron transfer reactions.

  • LEO: Loss of Electrons is Oxidation

  • GER: Gain of Electrons is Reduction

Balancing Chemical Equations

• Law of Conservation of Matter: Atoms are neither created nor destroyed.

• Coefficients are used to balance equations.

Chapter 7: Reaction Quantities and Stoichiometry

Stoichiometry

• Definition: Quantitative relationship between reactants and products.

• Mole Ratios: Derived from balanced chemical equations.

Stoichiometric Conversions

• Use mole ratios to convert between amounts of reactants and products.

• General Steps:

  1. Convert given quantity to moles.

  2. Use mole ratio to find moles of desired substance.

  3. Convert moles to desired units (grams, molecules, etc.).

Limiting Reactant and Percent Yield

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

• Theoretical Yield: Maximum amount of product possible.

• Percent Yield: $\%\,\text{Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100$

Appendix: Common Polyatomic Ions

Name	Formula	Charge
Nitrate	NO3	-1
Sulfate	SO4	-2
Phosphate	PO4	-3
Ammonium	NH4	+1
Hydroxide	OH	-1
Carbonate	CO3	-2

Additional info: Some details and examples have been expanded for clarity and completeness based on standard General Chemistry curriculum.