Exam 1 Review: Measurements, Atoms & Elements, Compounds, and Chemical Reactions

Chapter 2: Measurements and Problem Solving

Types of Measurements

Understanding measurements is fundamental in chemistry, as it allows for the quantification and comparison of substances and reactions.

• Qualitative Measurements: Describe qualities or characteristics (e.g., color, texture) without numerical values.

• Quantitative Measurements: Involve numerical values and units (e.g., mass, volume, temperature).

Units and Uncertainty

• SI Units: The International System of Units is used for standardization (e.g., meter, kilogram, second, mole).

• Uncertainty: All measurements have some degree of uncertainty due to limitations in measurement tools.

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

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

Scientific Notation and Significant Figures

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

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

Metric System and Conversions

• Metric Prefixes: Used to express multiples or fractions of units (e.g., kilo-, centi-, milli-).

• Common Units: Mass (grams), volume (liters), length (meters).

• Conversion Factors: Ratios used to convert between units.

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

Density

• Definition: Density is the mass per unit volume of a substance.

$\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

Chapter 4: Atoms & Elements

Atomic Theory and Structure

The atomic theory describes the nature of atoms, the fundamental building blocks of matter.

• Subatomic Particles: Protons (positive), neutrons (neutral), electrons (negative).

• 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.

• Element Symbol: One- or two-letter abbreviation for an element (e.g., H for hydrogen).

Periodic Table and Element Classification

• Groups/Families: Vertical columns; elements in the same group have similar properties.

• Periods: Horizontal rows.

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

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

• Metals, Nonmetals, Metalloids: Classified by physical and chemical properties.

Type	Properties
Metals	Conductive, malleable, ductile, shiny
Nonmetals	Poor conductors, brittle, dull
Metalloids	Properties intermediate between metals and nonmetals

Ions and Atomic Mass

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

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

• Atomic Mass: Weighted average mass of an element's isotopes.

Chapter 6: Ionic & Molecular Compounds and Chemical Nomenclature

Compounds and Formulas

• Compound: Substance formed from two or more elements chemically combined in fixed ratios.

• Chemical Formula: Representation of a compound using element symbols and subscripts.

Ions and Ionic Compounds

• Cations: Type 1 (fixed charge, e.g., Na+), Type 2 (variable charge, e.g., Fe2+, Fe3+).

• Anions: Typically nonmetals, end in "-ide" (e.g., Cl- is chloride).

• Naming Ionic Compounds:

  • Type 1: Cation + Anion (e.g., NaCl = sodium chloride)

  • Type 2: Cation (charge in Roman numerals) + Anion (e.g., FeCl2 = iron(II) chloride)

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

Acids and Nomenclature

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

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

• Naming Oxyacids:

  • "-ate" → "-ic acid" (e.g., H2SO4, sulfate → sulfuric acid)

  • "-ite" → "-ous acid" (e.g., H2SO3, sulfite → sulfurous acid)

  • "per-...-ate" → "per-...-ic acid" (e.g., HClO4, perchloric acid)

  • "hypo-...-ite" → "hypo-...-ous acid" (e.g., HClO, hypochlorous acid)

Molecular Compounds

• Naming: Prefix + 1st atom + prefix + 2nd atom (ending in "-ide").

• Prefixes: mono-, di-, tri-, tetra-, penta-, etc.

• Example: CO2 = carbon dioxide

Chemical Composition and Moles

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

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

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

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

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

Lewis Structures and 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 (1 pair), double (2 pairs), triple (3 pairs) of shared electrons.

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

• Polarity: Unequal sharing of electrons leads to polar bonds and molecules.

• Molecular Shape: Determined by electron pair repulsion (VSEPR theory).

Chapter 7: Chemical Reactions and Quantities

Chemical Reactions and Equations

• Chemical Change: Process in which 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: Ensuring 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: Calculation of reactants and products in chemical reactions using mole ratios.

• Mole Ratios: Derived from coefficients in balanced equations.

• Stoichiometric Conversions: Use of conversion factors to relate moles, mass, and number of particles.

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

Example Stoichiometry Problem

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

• Question: How many grams of water are produced from 4.0 g of hydrogen gas?

• Solution Steps:

  1. Convert grams H2 to moles H2.

  2. Use mole ratio to find moles H2O.

  3. Convert moles H2O to grams H2O.

Additional info: These notes are structured to provide a comprehensive review for a GOB Chemistry Exam 1, covering foundational concepts in measurement, atomic structure, chemical bonding, nomenclature, and reaction stoichiometry.