Comprehensive Study Guide: Moles, Compounds, Stoichiometry, Gases, and Acids/Bases

Unit 5: The Mole

Definition and Significance of the Mole

The mole is the SI unit for the amount of substance, defined as containing exactly 6.022 x 1023 particles (Avogadro's number). This unit allows chemists to count atoms, molecules, or ions by weighing them.

• Key Point: 1 mole of any substance contains 6.022 x 1023 particles (atoms, molecules, ions, etc.).

• Example: How many molecules are in 2.5 moles of H2O? $2.5 \text{ moles} \times 6.022 \times 10^{23} = 1.51 \times 10^{24}$ molecules$

• Key Point: In 1 mole of H2O, there are 2 moles of H atoms and 1 mole of O atoms.

Relative Masses and Molar Mass

The molar mass is the mass of one mole of a substance, expressed in grams per mole (g/mol). Different substances have different molar masses due to differences in atomic or molecular composition.

• Key Point: 1 mole of Ne and 1 mole of He contain the same number of atoms, but their masses differ due to different atomic masses.

• Example: Mass of 1 mole of Ne (20.18 g) vs. 1 mole of He (4.00 g).

Conversions: Moles, Grams, and Particles

Conversions between moles, grams, and number of particles use Avogadro's number and molar mass as conversion factors.

• Key Point: $\text{Number of particles} = \text{moles} \times 6.022 \times 10^{23}$

• Key Point: $\text{Mass (g)} = \text{moles} \times \text{molar mass (g/mol)}$

• Key Point: $\text{Moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

• Example: Calculate the number of molecules in 2.31 g of CH4: $\text{Moles CH}_4 = \frac{2.31}{16.05} = 0.144$ $\text{Molecules} = 0.144 \times 6.022 \times 10^{23} = 8.67 \times 10^{22}$ molecules$

Unit 6: Binary Compounds, Bonds

Naming Ions and Compounds

Chemical nomenclature follows international rules for naming anions and cations.

• Anions: Monoatomic (e.g., Cl- is chloride), Polyatomic (e.g., NO3- is nitrate).

• Cations: All monoatomic except NH4+ (ammonium) and H3O+ (hydronium).

• Example: The N3- ion is called nitride.

Types of Chemical Bonds

Chemical bonds form due to interactions between electrons of atoms.

• Ionic Bonds: Formed by transfer of electrons from metals to nonmetals.

• Covalent Bonds: Formed by sharing of electrons between nonmetals.

• Metallic Bonds: Involve a 'sea of electrons' shared among metal atoms.

• Example: Ionic bonds are formed between a metal and a nonmetal.

Physical Properties of Compounds

• Ionic Compounds: High melting points, brittle, conduct electricity when molten.

• Covalent Compounds: Low melting points, poor conductors.

• Metals: Malleable, ductile, good conductors.

Chemical Formulas and Lewis Structures

Chemical formulas represent the composition of compounds. Lewis dot structures show valence electrons.

• Ionic Formula: Shows the simplest ratio of ions (formula unit).

• Molecular Formula: Shows the actual number of atoms in a molecule.

• Lewis Dot Notation: Represents valence electrons around an element symbol.

Molecules: Monatomic, Diatomic, Polyatomic

• Monatomic: Single atom (e.g., He).

• Diatomic: Two atoms (e.g., N2, O2, F2, Cl2, Br2, I2, H2).

• Polyatomic: More than two atoms (e.g., O3).

Bond Polarity and the Octet Rule

• Polar Bonds: Electrons are shared unequally due to differences in electronegativity.

• Nonpolar Bonds: Electrons are shared equally.

• Octet Rule: Atoms tend to have 8 electrons in their valence shell (exceptions: H, He, B, etc.).

Molecular Geometry

Molecular shape is determined by the number of electron domains (bonding and nonbonding pairs) around the central atom.

Geometry	Bond Angle	Example
Linear	180°	CO2
Trigonal Planar	120°	H2CO
Tetrahedral	109.5°	CH4
Bent	109.5°	H2O
Trigonal Pyramidal	109.5°	NH3

Bonding in Metals

Metals are held together by a 'sea of electrons' that allows for conductivity and malleability.

Unit 7: Polyatomic Ions, Ternary Compounds, Stoichiometry

Naming Ionic and Molecular Compounds

• Ionic Compounds: Name the cation (with Roman numeral if variable charge), then the anion.

• Molecular Compounds: Use Greek prefixes (mono-, di-, tri-, etc.) to indicate the number of each atom.

• Greek Prefixes Table:

Number	Prefix
1	mono-
2	di-
3	tri-
4	tetra-
5	penta-
6	hexa-
7	hepta-
8	octa-
9	nona-
10	deca-

Stoichiometry and Chemical Equations

• Reactants: Substances consumed in a reaction.

• Products: Substances formed in a reaction.

• Coefficients: Indicate the number of molecules or moles; Subscripts: Indicate the number of atoms in a molecule.

• Balancing Equations: Ensures the same number of each atom on both sides.

• Stoichiometry: The calculation of reactants and products in chemical reactions using mole ratios.

Limiting and Excess Reactants

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

• Excess Reactant: The reactant that remains after the reaction is complete.

Unit 8: Gas Laws

Kinetic Molecular Theory and Gas Laws

The behavior of gases is described by the kinetic molecular theory and summarized by several gas laws.

• Kinetic Molecular Theory: Gas particles move rapidly, are far apart, and do not attract or repel each other.

• Combined Gas Law: $\frac{P_1 V_1}{T_1} = \frac{P_2 V_2}{T_2}$

• Ideal Gas Law: $PV = nRT$

• Standard Temperature and Pressure (STP): 0°C (273 K) and 1 atm; 1 mole of gas occupies 22.4 L at STP.

Partial Pressures

• Dalton's Law: The total pressure of a gas mixture is the sum of the partial pressures of each component.

• Example: $P_{\text{total}} = P_{O_2} + P_{Ne} + P_{CO_2}$

Mini Unit 9: Naming Acids and Bases, Neutralization Reactions

Naming Acids

• Binary Acids: "hydro-" + anion root + "ic acid" (e.g., HCl is hydrochloric acid).

• Ternary (Oxo) Acids: Based on polyatomic anion: - "-ate" ion → "-ic acid" (e.g., HNO3 is nitric acid) - "-ite" ion → "-ous acid" (e.g., HNO2 is nitrous acid)

Naming Bases

• Bases: Most common bases are metallic cations with hydroxide (OH-), e.g., NaOH is sodium hydroxide.

Additional info: This guide covers core topics from general chemistry including the mole concept, chemical bonding, nomenclature, stoichiometry, gas laws, and acid/base naming conventions. Practice problems and tables are included for reference and review.