Chemistry 1151: Comprehensive Study Guide

Chemical Reactions & Stoichiometry

Balancing Chemical Equations

Chemical equations must be balanced to obey the law of conservation of mass. This means the number of atoms of each element must be the same on both sides of the equation.

• Balancing Steps: Write the correct formulas, count atoms, adjust coefficients, and check your work.

• Example: Balance: C3H8 + O2 → CO2 + H2O

Balanced Equation:

$\mathrm{C_3H_8 + 5O_2 \rightarrow 3CO_2 + 4H_2O}$

Mole Ratios and Stoichiometry

Mole ratios from balanced equations allow calculation of product or reactant amounts.

• Example: If 2.0 mol O2 reacts, how many moles of CO2 form?

• From the balanced equation, 5 mol O2 produce 3 mol CO2:

$\mathrm{2.0\ mol\ O_2 \times \frac{3\ mol\ CO_2}{5\ mol\ O_2} = 1.2\ mol\ CO_2}$

Types of Chemical Reactions

• Synthesis: Two or more substances combine.

• Decomposition: One substance breaks down.

• Single Replacement: One element replaces another.

• Double Replacement: Ions exchange between compounds.

• Combustion: Substance reacts with O2, producing energy.

• Example: Zn + CuSO4 → ZnSO4 + Cu is a single replacement reaction.

Limiting Reactant

The limiting reactant is the substance that is completely used up first, limiting the amount of product formed.

Solutions & Concentration

Molarity (M)

Molarity is the number of moles of solute per liter of solution.

$\mathrm{M = \frac{mol\ solute}{L\ solution}}$

• Example: 1.0 mol NaCl in 2.0 L solution: $\mathrm{M = \frac{1.0}{2.0} = 0.50\ M}$

Dilution Formula

Used to calculate the volume or concentration after dilution.

$\mathrm{M_1V_1 = M_2V_2}$

• Example: To make 1.0 L of 1.0 M from 3.0 M stock: $\mathrm{V_1 = \frac{M_2V_2}{M_1} = \frac{1.0 \times 1.0}{3.0} = 0.33\ L}$

Percent Solutions

Percent solutions express concentration as a percentage (mass/volume, mass/mass, or volume/volume).

Ion Concentrations

• Example: [Cl-] in 0.50 M CaCl2: Each CaCl2 yields 2 Cl- ions, so [Cl-] = 1.0 M.

Atomic Structure & Periodic Trends

Subatomic Particles

• Protons: Positive charge, found in nucleus.

• Neutrons: Neutral, found in nucleus.

• Electrons: Negative charge, found in electron cloud.

Mass Number

Mass number = protons + neutrons.

• Example: Mass 40, atomic number 20: Neutrons = 40 - 20 = 20.

Valence Electrons

Valence electrons are electrons in the outermost shell, important for bonding.

• Example: Phosphorus (P) has 5 valence electrons.

Periodic Groups

• Alkali Metals: Group 1 (e.g., Na)

• Alkaline Earth Metals: Group 2 (e.g., Mg)

• Example: Na is an alkali metal.

Bonding, Naming, and Structure

Ionic vs Covalent Bonds

• Ionic: Transfer of electrons, metal + nonmetal (e.g., CaCl2).

• Covalent: Sharing of electrons, nonmetal + nonmetal (e.g., CO2).

Naming Compounds

• Ionic: Name cation, then anion (e.g., FeCl3 is iron(III) chloride).

• Covalent: Use prefixes (e.g., CO2 is carbon dioxide).

Polarity

• Polar: Unequal sharing of electrons (e.g., H2O).

• Nonpolar: Equal sharing (e.g., O2).

• Example: H2O is polar.

Intermolecular Forces & Properties

Types of Intermolecular Forces

• Hydrogen Bonding: Strong dipole-dipole force between H and N, O, or F (e.g., NH3).

• Dipole-Dipole: Between polar molecules.

• London Dispersion: Weak, in all molecules.

• Example: NH3 has hydrogen bonding; CH4 and CO2 do not.

Boiling Point Changes

• Adding salt increases boiling point (boiling point elevation).

• Bubbles in boiling water contain water vapor.

Gases & Gas Laws

Gas Laws

• Boyle's Law: $P_1V_1 = P_2V_2$ (at constant T, n)

• Charles's Law: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$ (at constant P, n)

• Ideal Gas Law: $PV = nRT$

• Example: If pressure doubles, volume halves (Boyle's Law).

• Convert 25°C to K: $25 + 273 = 298\ K$

• Total pressure = 2.0 atm, O2 = 1.2 atm, other gas = 0.8 atm.

Energy & Thermochemistry

Heat Calculations

Heat required to change temperature:

$q = mc\Delta T$

• Example: Heat to raise 100 g water by 10°C (c = 4.18 J/g°C): $q = 100 \times 4.18 \times 10 = 4180\ J$

Exothermic vs Endothermic

• Exothermic: Releases heat; products lower in energy.

• Endothermic: Absorbs heat; products higher in energy.

• Example: Melting ice is endothermic.

Acids, Bases, and pH

Acids and Bases

• Acid: Donates H+ (proton).

• Base: Accepts H+.

pH and Strength

$\mathrm{pH = -\log[H^+]}$

• [H+] = 1 × 10-3 is acidic (pH = 3).

• Stronger acid: higher Ka (e.g., Ka = 10-2 is stronger than 10-6).

Conjugate Acid-Base Pairs

• Conjugate base of H2CO3 is HCO3-.

Redox (Oxidation-Reduction)

Identifying Oxidation and Reduction

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Example: In Zn + Cu2+ → Zn2+ + Cu, Zn is oxidized.

Miscellaneous Concepts

Density

$\mathrm{Density = \frac{mass}{volume}}$

• Example: 20 g in 10 mL: $\mathrm{Density = \frac{20}{10} = 2.0\ g/mL}$

Significant Figures

• 0.00450 has 3 significant figures; 4500 has 2 (unless specified with a decimal).

Mixtures

• Homogeneous: Uniform composition (e.g., saltwater).

• Heterogeneous: Non-uniform composition.