Introductory Chemistry Final Exam Review: Key Concepts and Calculations

Final Exam Review: Introductory Chemistry

Concepts

Chemical and Physical Properties and Changes

Chemical and physical properties describe the characteristics of substances, while chemical and physical changes refer to the processes substances undergo.

• Chemical Properties: Characteristics that determine how a substance reacts with other substances (e.g., flammability, reactivity).

• Physical Properties: Characteristics that can be observed without changing the substance's identity (e.g., melting point, density, color).

• Chemical Change: A process that alters the chemical composition of a substance (e.g., rusting, burning).

• Physical Change: A process that does not alter the chemical composition (e.g., melting, boiling).

• Example: Ice melting is a physical change; iron rusting is a chemical change.

Naming Compounds

Chemical compounds are named according to systematic rules based on their composition.

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

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

• Acids: If the anion ends in -ide, the acid name begins with hydro- and ends with -ic (e.g., HCl is hydrochloric acid).

• Example: KNO3 is potassium nitrate.

Isotopes and Atomic Notation

Isotopes are atoms of the same element with different numbers of neutrons.

• Atomic Notation: $^A_ZX$ where X is the element symbol, A is the mass number, and Z is the atomic number.

• Example: $^{14}_6C$ is carbon-14.

Periodic Law and Families of Elements

The periodic law states that properties of elements repeat periodically when arranged by atomic number.

• Families: Groups of elements with similar properties (e.g., alkali metals, halogens, noble gases).

• Example: Group 1 elements are alkali metals.

Phases of Matter and Changes of State

Matter exists in different phases: solid, liquid, and gas. Changes of state occur with energy changes.

• Melting: Solid to liquid

• Evaporation/Boiling: Liquid to gas

• Condensation: Gas to liquid

• Freezing: Liquid to solid

• Sublimation: Solid to gas

• Deposition: Gas to solid

Trends in Electronegativity

Electronegativity is the ability of an atom to attract electrons in a bond.

• Trend: Increases across a period (left to right), decreases down a group.

• Example: Fluorine is the most electronegative element.

Balancing Reactions

Chemical equations must be balanced to obey the law of conservation of mass.

• Steps: Count atoms of each element on both sides, adjust coefficients to balance.

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

Single and Double Replacement Reactions

Replacement reactions involve the exchange of elements between compounds.

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

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

• Example: $Zn + 2HCl \rightarrow ZnCl_2 + H_2$ (single replacement)

Net Ionic Equations and Spectator Ions

Net ionic equations show only the species that participate in the reaction.

• Spectator Ions: Ions that do not change during the reaction.

• Example: For $AgNO_3 + NaCl \rightarrow AgCl + NaNO_3$, the net ionic equation is $Ag^+ + Cl^- \rightarrow AgCl$.

Polar Bonds and Polar Molecules

Polarity depends on differences in electronegativity and molecular geometry.

• Polar Bond: Unequal sharing of electrons between atoms.

• Polar Molecule: Molecule with an uneven distribution of charge (e.g., H2O).

Nuclear Reactions

Nuclear reactions involve changes in the nucleus, such as radioactive decay.

• Alpha Decay: $^A_ZX \rightarrow ^{A-4}_{Z-2}Y + ^4_2He$

• Beta Decay: $^A_ZX \rightarrow ^A_{Z+1}Y + ^0_{-1}e$

Properties of Acids and Bases, pH

Acids and bases have characteristic properties and are measured by pH.

• Acids: Taste sour, turn litmus red, donate H+ ions.

• Bases: Taste bitter, turn litmus blue, accept H+ ions.

• pH: $pH = -\log[H^+]$

• Example: pH 7 is neutral; below 7 is acidic, above 7 is basic.

Valence Electrons

Valence electrons are the outermost electrons involved in bonding.

• Group Number: For main group elements, the group number equals the number of valence electrons.

• Example: Oxygen (Group 16) has 6 valence electrons.

Electron Configuration for Atoms and Ions

Electron configuration shows the arrangement of electrons in an atom or ion.

• Example: Sodium: $1s^2 2s^2 2p^6 3s^1$

• For ions: Remove/add electrons from/to the outermost shell.

Dot Structures and Molecular Shapes

Lewis dot structures represent valence electrons; molecular shapes are predicted by VSEPR theory.

• Example: H2O has a bent shape.

• Common Shapes: Linear, bent, trigonal planar, tetrahedral.

Hydrogen Bonding

Hydrogen bonding is a strong intermolecular force between molecules with H bonded to N, O, or F.

• Example: Water molecules form hydrogen bonds.

Solubility: Like Dissolves Like, Effect of Temperature and Pressure

Solubility depends on the nature of solute and solvent, temperature, and pressure.

• Like Dissolves Like: Polar solvents dissolve polar solutes; nonpolar dissolves nonpolar.

• Temperature: Solubility of solids increases with temperature; gases decrease.

• Pressure: Solubility of gases increases with pressure (Henry's Law).

Ideal Gases

The ideal gas law relates pressure, volume, temperature, and moles of a gas.

• Equation: $PV = nRT$

• STP: Standard temperature (0°C) and pressure (1 atm).

Naming Hydrocarbons and Identifying Organic Functional Groups

Hydrocarbons are named based on the number of carbons and type of bonds; functional groups determine reactivity.

• Alkanes: Single bonds (e.g., methane, ethane).

• Alkenes: Double bonds (e.g., ethene).

• Alkynes: Triple bonds (e.g., ethyne).

• Functional Groups: Alcohol (-OH), carboxylic acid (-COOH), amine (-NH2), etc.

Isomers

Isomers are compounds with the same molecular formula but different structures.

• Structural Isomers: Differ in connectivity.

• Stereoisomers: Differ in spatial arrangement.

Calculations

Metric Conversions

Metric conversions use conversion factors to change units.

• Example: 1 km = 1000 m; 1 mL = 0.001 L

Significant Figures

Significant figures reflect the precision of a measurement.

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

• In Calculations: For multiplication/division, use the least number of significant figures; for addition/subtraction, use the least number of decimal places.

Density Calculations

Density is mass per unit volume.

• Formula: $Density = \frac{Mass}{Volume}$

• Use as Conversion Factor: To convert between mass and volume.

Mole Calculations and Molar Mass

The mole is a counting unit; molar mass is the mass of one mole of a substance.

• Formula: $Moles = \frac{Mass}{Molar\ Mass}$

• Example: 18 g H2O × (1 mol / 18 g) = 1 mol H2O

Nuclear Half-Life

Half-life is the time required for half of a radioactive sample to decay.

• Formula: $N = N_0 \left( \frac{1}{2} \right)^{\frac{t}{t_{1/2}}}$

Mass Percent

Mass percent is the mass of a component divided by the total mass, multiplied by 100%.

• Formula: $\%\ mass = \frac{mass\ of\ component}{total\ mass} \times 100\%$

Molarity

Molarity is the concentration of a solution in moles per liter.

• Formula: $M = \frac{moles\ of\ solute}{liters\ of\ solution}$

Dilution Calculations

Dilution involves adding solvent to decrease concentration.

• Formula: $M_1V_1 = M_2V_2$

Ideal Gas Law and Combined Gas Law

Relate pressure, volume, temperature, and moles of gas.

• Ideal Gas Law: $PV = nRT$

• Combined Gas Law: $\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$

Percent Composition

Percent composition is the percent by mass of each element in a compound.

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

Empirical and Molecular Formulas

The empirical formula is the simplest whole-number ratio; the molecular formula is the actual number of atoms.

• Steps: Find moles of each element, divide by smallest, multiply to get whole numbers.

• Example: Empirical formula CH2O; molecular formula C6H12O6.

Stoichiometry

Stoichiometry involves quantitative relationships in chemical reactions.

• Mole to Mole: Use coefficients from balanced equation.

• Mass to Mass: Convert mass to moles, use mole ratio, convert back to mass.

• Mass to Volume (Gas at STP): 1 mol gas = 22.4 L at STP.

• Limiting Reactant: The reactant that is used up first, limiting the amount of product.

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

Summary Table: Key Equations and Concepts

Topic	Key Equation or Concept
Density	$Density = \frac{Mass}{Volume}$
Molarity	$M = \frac{moles\ of\ solute}{liters\ of\ solution}$
Ideal Gas Law	$PV = nRT$
Percent Yield	$\%\ yield = \frac{actual\ yield}{theoretical\ yield} \times 100\%$
Empirical Formula	Find moles, divide by smallest, multiply to whole numbers
Half-Life	$N = N_0 \left( \frac{1}{2} \right)^{\frac{t}{t_{1/2}}}$
pH	$pH = -\log[H^+]$