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

Exam 1 Review

Overview

This study guide covers foundational topics in GOB Chemistry, including measurements, atomic theory, chemical compounds, and chemical reactions. 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, liter).

• Metric Multipliers: Prefixes such as kilo-, centi-, milli- used to express multiples or fractions of units.

Accuracy, Precision, and Uncertainty

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

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

• Uncertainty: The degree of doubt in a measurement, often expressed as ± value.

Significant Figures

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

• Rules: All nonzero digits are significant; zeros between nonzero digits are significant; leading zeros are not significant; trailing zeros are significant only if there is a decimal point.

Scientific Notation

• Purpose: Expresses very large or small numbers in the form $a imes 10^n$.

• Example: $6,200 = 6.2 imes 10^3$

Volume Measurements

• Liters (L): Standard unit for measuring liquid volume.

Problem Solving and Dimensional Analysis

• Conversion Factors: Ratios used to convert between units.

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

• Physical Property: A characteristic that can be measured without changing the substance's identity (e.g., density).

Chapter 4: Atoms & Elements

Atomic Theory

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

• Subatomic Particles:

  • Protons: Positively charged particles in the nucleus.

  • Neutrons: Neutral particles in the nucleus.

  • Electrons: Negatively charged particles orbiting 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 ($A = Z + N$).

Element Symbols and Periodic Table

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

• 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: Shiny, conductive, malleable.

  • Nonmetals: Dull, poor conductors.

  • Metalloids: Properties intermediate between metals and nonmetals.

Common Groups

• Alkali Metals (IA)

• Alkaline Earth Metals (IIA)

• Halogens (VIIA)

• Noble Gases (VIIIA)

Monatomic Ions

• Definition: Ions consisting of a single atom with a positive or negative charge.

Chapter 6: Ionic & Molecular Compounds and Chemical Equations

Types of Compounds

• Ionic Compounds: Formed from cations (positive ions) and anions (negative ions).

• Molecular Compounds: Formed from nonmetals sharing electrons (covalent bonds).

Formulas and Nomenclature

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

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

Cations and Anions

• 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 (with charge in Roman numerals) + Anion (ending in -ide).

Polyatomic Ions

• Definition: Ions composed of two or more atoms covalently bonded, carrying a charge (e.g., $NO_3^-$, $SO_4^{2-}$).

Naming Acids

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

• Oxyacids: Hydrogen + polyatomic ion (e.g., $H_2SO_4$).

• Naming Rules:

  • If the polyatomic ion ends in -ate, the acid ends in -ic.

  • If the polyatomic ion ends in -ite, the acid ends in -ous.

  • Prefixes: Per- (more oxygen), Hypo- (less oxygen).

Naming Molecular Compounds

• Rule: Prefix for number of atoms + first element name + prefix for number of atoms + second element name ending in -ide (e.g., $CO_2$ is carbon dioxide).

Chapter 6: Chemical Composition

Avogadro's Number and the Mole

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

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

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.

• Calculation: Add atomic masses of all atoms in the formula.

Percent Composition

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

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

Empirical and Molecular Formulas

• Empirical Formula: Simplest ratio of elements.

• 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 Compounds

• Lewis Structure: Diagram showing valence electrons as dots around atoms.

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

• Bonding:

  • Single Bond: One shared pair of electrons.

  • Double Bond: Two shared pairs of electrons.

  • Triple Bond: Three shared pairs of electrons.

Electronegativity and Bond Polarity

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

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

Molecular Shape

• VSEPR Theory: Predicts molecular shape based on electron pair repulsion.

Chapter 7: Chemical Reactions

Types of Chemical Reactions

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

• Decomposition: $AB \rightarrow A + B$

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

• Redox (Oxidation/Reduction): Electron transfer reactions (LEO: Lose Electrons = Oxidation, GER: Gain Electrons = Reduction).

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

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

Balancing Chemical Equations

• Law of Conservation of Matter: Matter is neither created nor destroyed; equations must be balanced.

Chapter 7: Chemical Quantities and Stoichiometry

Stoichiometry

• Definition: Quantitative relationship between reactants and products in a chemical reaction.

• Mole Ratios: Derived from coefficients in balanced equations.

Stoichiometric Conversions

• Conversions:

  • Grams A $\rightarrow$ Moles A $\rightarrow$ Moles B $\rightarrow$ Grams B

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

• Theoretical Yield: Maximum amount of product possible from given reactants.

Example Table: Types of Chemical Reactions

Type	General Equation	Description
Combination	$A + B \rightarrow AB$	Two substances form one product
Decomposition	$AB \rightarrow A + B$	One substance breaks into two or more products
Combustion	$C_xH_yO_z + O_2 \rightarrow CO_2 + H_2O$	Hydrocarbon reacts with oxygen
Single Displacement	$A + BC \rightarrow AC + B$	Element replaces another in a compound
Double Displacement	$AB + CD \rightarrow AD + CB$	Exchange of ions between two compounds

Additional info: Some content was expanded for clarity and completeness, including definitions, formulas, and examples.