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

Exam 1 Review: General Chemistry

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

Chapter 2: Measurements and Problem Solving

Qualitative vs. Quantitative Observations

• Qualitative observations: Describe the quality or character of a substance (e.g., color, smell, texture).

• Quantitative observations: Involve measurements and express the quantity or amount (e.g., mass, volume, temperature).

Parts of a Measurement

• Magnitude: The numerical value of the measurement.

• Unit: The scale or standard of the measurement (e.g., grams, liters).

• Uncertainty: The reliability or precision of the measurement, often indicated by significant figures.

Scientific Notation

• Expresses numbers as a product of a coefficient and a power of 10.

• Example: $4,500 = 4.5 \times 10^3$

Significant Figures

• Definition: Digits in a measurement that carry meaningful information about its precision.

• Addition/Subtraction: Result has the same number of decimal places as the measurement with the fewest decimal places.

• Multiplication/Division: Result has the same number of significant figures as the measurement with the fewest significant figures.

• Rounding: Round to the correct number of significant figures; if the digit to be dropped is 5 or greater, round up.

Units and Conversions

• SI Units: Standard units for scientific measurements (e.g., meter, kilogram, second, mole).

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

• Conversion Factors: Ratios used to convert between units.

• Dimensional Analysis: A method for solving problems using conversion factors to ensure units cancel appropriately.

Physical and Chemical Properties

• Physical property: Can be observed without changing the substance's identity (e.g., melting point, density).

• Chemical property: Describes a substance's ability to undergo chemical changes (e.g., flammability).

Chapter 4: Atoms & Elements

Atomic Theory and Structure

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

• Subatomic particles:

  • Protons: Positively charged, found in the nucleus.

  • Neutrons: Neutral, found in the nucleus.

  • Electrons: Negatively charged, found in orbitals around the nucleus.

• Atomic number (Z): Number of protons in the nucleus; defines the element.

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

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

• Nuclear symbol: Notation showing the element, atomic number, and mass number (e.g., $^{14}_6\text{C}$).

Periodic Table and Periodic Law

• Periodic table: Organizes elements by increasing atomic number and similar properties.

• Groups/Families: Vertical columns; elements in a group have similar chemical properties.

• Periods: Horizontal rows; properties change progressively across a period.

• Main group elements (A groups): Groups 1A–8A.

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

• Metals, nonmetals, metalloids: Classified by physical and chemical properties.

Chapter 6: Ionic & Molecular Compounds and Chemical Equations

Ions and Ionic Compounds

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

• Anions: Negatively charged ions.

• Monatomic ions: Single-atom ions (e.g., Na+, Cl-).

• Polyatomic ions: Ions composed of multiple atoms (e.g., NO3-, SO42-).

• Ionic compounds: Formed from cations and anions; electrically neutral.

Naming Ionic Compounds

• Type 1: Metal (cation) + nonmetal (anion with -ide ending).

• Type 2: Transition metal (cation with charge in Roman numerals) + nonmetal (anion with -ide ending).

• Salts: General term for ionic compounds.

Naming Acids

• Binary acids: Hydrogen + nonmetal (e.g., HCl = hydrochloric acid).

• Oxoacids: Hydrogen + polyatomic ion (e.g., HNO3 = nitric acid).

• Naming rules:

  • -ic acid: Polyatomic ion ends in -ate (e.g., sulfate → sulfuric acid).

  • -ous acid: Polyatomic ion ends in -ite (e.g., sulfite → sulfurous acid).

  • Prefixes: per- (more oxygen), hypo- (less oxygen).

Naming Molecular Compounds

• Use prefixes to indicate the number of each atom (e.g., CO2 = carbon dioxide).

• First element keeps its name; second element ends in -ide.

• Common prefixes: mono-, di-, tri-, tetra-, penta-, hexa-, etc.

Chemical Formulas and Composition

• Empirical formula: Simplest whole-number ratio of atoms in a compound.

• Molecular formula: Actual number of atoms of each element in a molecule.

• Structural formula: Shows how atoms are connected.

Avogadro's Number and Molar Mass

• Avogadro's number: $6.02 \times 10^{23}$ particles per mole.

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

• Calculating molar mass: Sum the atomic masses of all atoms in the formula.

Percent Composition

• Percent by mass of each element in a compound.

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

Lewis Structures and Covalent Bonding

• Lewis structures: Diagrams showing valence electrons and bonding in molecules.

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

• Bond types:

  • Single bond: 1 shared electron pair

  • Double bond: 2 shared electron pairs

  • Triple bond: 3 shared electron pairs

• 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 and Quantities

Chemical Reactions and Equations

• Chemical change: Substances are transformed into new substances.

• Chemical equation: Symbolic representation of a chemical reaction.

• Law of Conservation of Matter: Matter is neither created nor destroyed in a chemical reaction.

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

Types of Chemical Reactions

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

• Decomposition: $AB \rightarrow A + B$

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

• Redox (Oxidation-Reduction): Involves electron transfer; LEO (Lose Electrons = Oxidation), GER (Gain Electrons = Reduction).

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

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

Stoichiometry and Reaction Quantities

• Stoichiometry: Quantitative relationships between reactants and products in a chemical reaction.

• Mole ratios: Derived from coefficients in balanced equations.

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

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

• Theoretical yield: Maximum amount of product that can be formed from given reactants.

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

Example Stoichiometry Problem

• Given: $2H_2 + O_2 \rightarrow 2H_2O$

• If you have 4 moles of $H_2$ and excess $O_2$, how many moles of $H_2O$ can be produced?

• Solution: $4\,\text{mol}\,H_2 \times \frac{2\,\text{mol}\,H_2O}{2\,\text{mol}\,H_2} = 4\,\text{mol}\,H_2O$

Summary Table: Types of Chemical Reactions

Type	General Equation	Description
Combination (Synthesis)	$A + B \rightarrow AB$	Two or more substances combine to form one product.
Decomposition	$AB \rightarrow A + B$	One substance breaks down into two or more products.
Combustion	$C_xH_yO_z + O_2 \rightarrow CO_2 + H_2O$	Hydrocarbon reacts with oxygen to produce carbon dioxide and water.
Single Displacement	$A + BC \rightarrow AC + B$	One element replaces another in a compound.
Double Displacement	$AB + CD \rightarrow AD + CB$	Exchange of ions between two compounds.
Redox	Electron transfer	Oxidation and reduction occur simultaneously.

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