Introductory Chemistry Final Exam Review: Concepts and Calculations

Chemical and Physical Properties and Changes

Definitions and Examples

• Chemical Properties: Characteristics that describe a substance's ability to undergo chemical changes (e.g., flammability, reactivity).

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

• Chemical Change: A process where a substance transforms into a new substance (e.g., rusting of iron).

• Physical Change: A process that alters the form or appearance but not the composition (e.g., melting ice).

• Example: Burning wood is a chemical change; melting ice is a physical change.

Naming Compounds

Rules for Ionic and Molecular Compounds

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

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

• Acids: If the anion ends in -ide, use "hydro-" and "-ic" (e.g., HCl: hydrochloric acid).

• Example: FeCl3 is iron(III) chloride.

Isotopes and Atomic Notation

Understanding Isotopes

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

• Atomic Notation: $^{A}_{Z}X$ where A = mass number, Z = atomic number, X = element symbol.

• Example: $^{14}_{6}C$ is an isotope of carbon.

Periodic Law and Families of Elements

Periodic Trends and Group Names

• Periodic Law: Properties of elements repeat periodically when arranged by increasing atomic number.

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

• Example: Group 1: Alkali metals; Group 17: Halogens.

Phases of Matter and Changes of State

Solid, Liquid, Gas, and Phase Changes

• Phases: Solid (fixed shape), liquid (fixed volume), gas (variable shape and volume).

• Changes of State: Melting, freezing, vaporization, condensation, sublimation.

• Example: Water boiling is vaporization.

Trends in Electronegativity

Periodic Table Trends

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

• Trend: Increases across a period, decreases down a group.

• Example: Fluorine is the most electronegative element.

Balancing Reactions

Steps and Examples

• Balancing: Adjust coefficients to ensure equal numbers of each atom on both sides.

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

Single and Double Replacement Reactions

Types and Examples

• Single Replacement: One element replaces another in a compound (e.g., $Zn + CuSO_4 \rightarrow ZnSO_4 + Cu$).

• Double Replacement: Exchange of ions between two compounds (e.g., $AgNO_3 + NaCl \rightarrow AgCl + NaNO_3$).

Net Ionic Equations and Spectator Ions

Writing Net Ionic Equations

• Net Ionic Equation: Shows only the species that change during the reaction.

• Spectator Ions: Ions that do not participate in the reaction.

• Example: $Ag^+ + Cl^- \rightarrow AgCl$ (net ionic equation for AgNO3 + NaCl).

Polarity: Polar Bonds and Molecules

Identifying Polarity

• Polar Bond: Unequal sharing of electrons due to difference in electronegativity.

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

• Example: CO2 has polar bonds but is a nonpolar molecule.

Nuclear Reactions

Types and Notation

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

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

• Example: $^{14}_{6}C \rightarrow ^{14}_{7}N + ^{0}_{-1}e$

Acids, Bases, and pH

Properties and Calculations

• Acids: Produce H+ ions in solution; taste sour.

• Bases: Produce OH- ions; taste bitter.

• pH: $pH = -\log[H^+]$

• Example: pH of 1 is strongly acidic; pH of 13 is strongly basic.

Valence Electrons

Definition and Importance

• Valence Electrons: Electrons in the outermost shell; determine chemical reactivity.

• Example: Sodium has 1 valence electron.

Electron Configuration for Atoms and Ions

Writing Configurations

• Electron Configuration: Arrangement of electrons in shells and subshells.

• Example: Na: $1s^2 2s^2 2p^6 3s^1$

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

Lewis Dot Structures and Molecular Shapes

Drawing and Naming Shapes

• Lewis Dot Structure: Shows valence electrons as dots around element symbol.

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

• Example: H2O is bent.

Hydrogen Bonding

Definition and Effects

• Hydrogen Bond: Strong intermolecular force between H and N, O, or F.

• Example: Water molecules form hydrogen bonds.

Solubility: Like Dissolves Like, Temperature and Pressure Effects

Solubility Principles

• 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.

• Example: Sugar dissolves in water; CO2 dissolves better under pressure.

Ideal Gases

Properties and Equations

• Ideal Gas Law: $PV = nRT$

• Variables: P = pressure, V = volume, n = moles, R = gas constant, T = temperature.

• Example: Calculate volume of 1 mole at STP.

Hydrocarbons and Organic Functional Groups

Naming and Identifying

• Hydrocarbons: Compounds of C and H; alkanes, alkenes, alkynes.

• Functional Groups: Specific groups of atoms (e.g., alcohol, carboxylic acid, amine).

• Example: Ethanol has an alcohol group (-OH).

Isomers

Definition and Recognition

• Isomers: Compounds with same formula but different structures.

• Example: Butane and isobutane.

Metric Conversions and Significant Figures

Measurement and Calculation

• Metric Conversions: Use conversion factors (e.g., 1 km = 1000 m).

• Significant Figures: Digits that reflect precision; rules for counting and rounding.

• Example: 0.00450 has 3 significant figures.

Density Calculations

Formula and Applications

• Density: $D = \frac{m}{V}$

• Use: As a conversion factor between mass and volume.

• Example: If mass = 10 g, volume = 2 mL, density = 5 g/mL.

Mole Calculations and Molar Mass

Conversions and Formulas

• Mole: Amount containing $6.022 \times 10^{23}$ particles.

• Molar Mass: Mass of 1 mole of a substance (g/mol).

• Conversions: $\text{moles} = \frac{\text{mass}}{\text{molar mass}}$

• Example: 18 g H2O = 1 mole.

Nuclear Half-Life

Calculating Remaining Amount

• Half-Life: Time for half of a sample to decay.

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

• Example: After 2 half-lives, 25% remains.

Mass Percent and Percent Composition

Calculations

• Mass Percent: $\% = \frac{\text{mass of component}}{\text{total mass}} \times 100$

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

• Example: In H2O, % H = 11%, % O = 89%.

Molarity and Dilution Calculations

Concentration and Dilution

• Molarity: $M = \frac{\text{moles}}{\text{liters}}$

• Dilution: $M_1V_1 = M_2V_2$

• Example: Diluting 1.0 M to 0.5 M by doubling volume.

Gas Laws: Ideal and Combined

Equations and Applications

• Ideal Gas Law: $PV = nRT$

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

• Example: Calculate new volume when pressure and temperature change.

Empirical and Molecular Formulas

Determining Formulas

• Empirical Formula: Simplest ratio of elements.

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

• Example: Empirical: CH2; Molecular: C2H4.

Stoichiometry

Types and Calculations

• Mole-to-Mole: Use balanced equation to convert between moles.

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

• Mass-to-Volume (Gas at STP): $1 \text{mol} = 22.4 \text{L}$ at STP.

• Limiting Reactant: Reactant that runs out first; determines maximum product.

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

• Example: Calculate product from given reactant masses.

Summary Table: Families of Elements

Classification of Element Groups

Family Name	Group Number	Key Properties
Alkali Metals	1	Highly reactive, 1 valence electron
Alkaline Earth Metals	2	Reactive, 2 valence electrons
Halogens	17	Very reactive, 7 valence electrons
Noble Gases	18	Inert, full valence shell

Summary Table: Organic Functional Groups

Common Functional Groups

Functional Group	Structure	Example
Alcohol	-OH	Ethanol
Carboxylic Acid	-COOH	Acetic acid
Amine	-NH2	Methylamine
Alkene	C=C	Ethene
Alkyne	C≡C	Ethyne

Summary Table: Gas Laws

Comparison of Gas Laws

Law	Equation	Variables
Ideal Gas Law	$PV = nRT$	P, V, n, T
Combined Gas Law	$\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$	P, V, T
Boyle's Law	$P_1V_1 = P_2V_2$	P, V
Charles' Law	$\frac{V_1}{T_1} = \frac{V_2}{T_2}$	V, T

Additional info: Academic context and examples were added to ensure completeness and clarity for exam preparation.