CHEM 110 Final Exam Study Guide: Key Concepts and Topics

Chemical Bonding

Sharing vs Transfer of Electrons

Chemical bonds form through the interaction of electrons between atoms. The nature of this interaction determines the type of bond.

• Sharing of electrons: Occurs in covalent bonds, where atoms share pairs of electrons.

• Transfer of electrons: Occurs in ionic bonds, where one atom donates electrons to another, resulting in ions.

• Example: Sodium chloride (NaCl) forms by transfer; water (H2O) forms by sharing.

Ionic vs Covalent Compounds

Ionic and covalent compounds differ in their bonding and properties.

• Ionic compounds: Formed from metals and nonmetals; consist of positive and negative ions.

• Covalent compounds: Formed from nonmetals; consist of molecules with shared electrons.

• Example: NaCl (ionic), CO2 (covalent).

Core vs Valence Electrons

Electrons in atoms are classified based on their energy levels.

• Core electrons: Inner electrons not involved in bonding.

• Valence electrons: Outermost electrons involved in chemical reactions and bonding.

• Example: In sodium (Na), 1 valence electron, 10 core electrons.

Octet Rule

The octet rule states that atoms tend to gain, lose, or share electrons to achieve eight valence electrons.

• Key Point: Stable electron configuration similar to noble gases.

• Example: Oxygen forms two bonds to complete its octet.

Single vs Double vs Triple Bonds

Covalent bonds can involve one, two, or three pairs of shared electrons.

• Single bond: One pair of electrons shared (e.g., H–H).

• Double bond: Two pairs shared (e.g., O=O).

• Triple bond: Three pairs shared (e.g., N≡N).

Dipole Moment

A dipole moment arises when there is a separation of charge in a molecule due to differences in electronegativity.

• Key Point: Indicates molecular polarity.

• Example: Water (H2O) has a strong dipole moment.

Electron and Molecular Geometries

The arrangement of atoms and electron pairs determines molecular shape.

• Electron geometry: Based on all electron groups.

• Molecular geometry: Based on only atoms (ignoring lone pairs).

• Example: Methane (CH4) is tetrahedral.

Gases

All Gas Laws

Gas laws describe the relationships between pressure, volume, temperature, and amount of gas.

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

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

• Avogadro's Law: $\frac{V_1}{n_1} = \frac{V_2}{n_2}$ (at constant P and T)

• Ideal Gas Law: $PV = nRT$

Intermolecular Forces and Properties

All Intermolecular Attractive Forces

Intermolecular forces are attractions between molecules, affecting physical properties.

• Types: London dispersion, dipole-dipole, hydrogen bonding, ion-dipole.

• Example: Water exhibits hydrogen bonding.

Polarizability

Polarizability is the ability of an electron cloud to be distorted, influencing intermolecular forces.

• Key Point: Larger atoms/molecules are more polarizable.

Polar vs Nonpolar

Molecules can be polar (unequal sharing of electrons) or nonpolar (equal sharing).

• Polar: Water (H2O)

• Nonpolar: Methane (CH4)

Hydrophilic vs Hydrophobic

Hydrophilic substances interact well with water; hydrophobic substances do not.

• Hydrophilic: Polar or ionic compounds.

• Hydrophobic: Nonpolar compounds.

Solutions and Solubility

Homogenous Solutions

A homogenous solution is a mixture with uniform composition throughout.

• Example: Salt water.

Formula and Experimental Mass Analysis

Mass analysis involves determining the composition of compounds.

• Formula mass: Sum of atomic masses in a formula.

• Experimental mass: Determined by measurement.

Insoluble vs Soluble

Solubility describes whether a substance dissolves in a solvent.

• Soluble: Dissolves easily (e.g., NaCl in water).

• Insoluble: Does not dissolve (e.g., AgCl in water).

Viscosity

Viscosity is a measure of a fluid's resistance to flow.

• High viscosity: Honey

• Low viscosity: Water

Polyatomic Ions and Acids

Polyatomic Ions

Polyatomic ions are charged species composed of multiple atoms.

• Example: Nitrate ($NO_3^-$), sulfate ($SO_4^{2-}$)

Oxy Acids

Oxy acids contain hydrogen, oxygen, and another element.

• Example: Sulfuric acid ($H_2SO_4$)

Redox and Acid-Base Reactions

Oxidation/Reduction

Redox reactions involve the transfer of electrons.

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Example: $2Na + Cl_2 \rightarrow 2NaCl$

Neutralization Reactions

Neutralization occurs when an acid reacts with a base to form water and a salt.

• Example: $HCl + NaOH \rightarrow NaCl + H_2O$

Phase Changes and Properties

Evaporation vs Condensation

Evaporation is the process of liquid turning into gas; condensation is gas turning into liquid.

• Evaporation: Endothermic process.

• Condensation: Exothermic process.

Chemical Formulas and Reactions

Empirical vs Molecular Formula

The empirical formula shows the simplest ratio of elements; the molecular formula shows the actual number of atoms.

• Example: Glucose: Empirical (CH2O), Molecular (C6H12O6)

Decomposition

Decomposition reactions break down a compound into simpler substances.

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

Combustion

Combustion reactions involve a substance reacting with oxygen to produce energy.

• Example: $CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O$

Atomic Structure

Orbitals

Orbitals are regions in an atom where electrons are likely to be found.

• Types: s, p, d, f

Electron Configuration

Electron configuration describes the arrangement of electrons in an atom.

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

Orbital Diagram

Orbital diagrams use boxes and arrows to show electron placement and spin.

• Key Point: Follows the Pauli exclusion principle and Hund's rule.

Acid-Base Chemistry

Acid-Base Reactions

Acid-base reactions involve the transfer of protons (H+).

• Example: $NH_3 + H_2O \rightarrow NH_4^+ + OH^-$

Conjugate Pairs

Conjugate pairs consist of an acid and its corresponding base after donating or accepting a proton.

• Example: $H_2O$ (acid) and $OH^-$ (conjugate base)

Thermochemistry

Thermochemistry

Thermochemistry studies energy changes in chemical reactions.

• Key Point: Endothermic (absorbs energy), exothermic (releases energy).

• Equation: $q = mc\Delta T$

Lewis Structures

Lewis Dot Structures

Lewis dot structures represent valence electrons as dots around element symbols.

• Key Point: Used to predict bonding and molecular structure.

• Example: Water: H:O:H

Bond Type	Electron Sharing	Example
Single	1 pair	H–H
Double	2 pairs	O=O
Triple	3 pairs	N≡N

Gas Law	Equation	Variables
Boyle's Law	$P_1V_1 = P_2V_2$	Pressure, Volume
Charles's Law	$\frac{V_1}{T_1} = \frac{V_2}{T_2}$	Volume, Temperature
Avogadro's Law	$\frac{V_1}{n_1} = \frac{V_2}{n_2}$	Volume, Moles
Ideal Gas Law	$PV = nRT$	Pressure, Volume, Moles, Temperature

Intermolecular Force	Strength	Example
London Dispersion	Weak	CH4
Dipole-Dipole	Moderate	HCl
Hydrogen Bonding	Strong	H2O
Ion-Dipole	Very Strong	Na+ in water

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