GOB Chemistry Study Guide: Ionic Compounds, Chemical Bonding, Chemical Equations, and Gas Laws

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Ions and Ionic Compounds

Formation and Naming of Ions

Ions are atoms or molecules that have gained or lost electrons, resulting in a net electrical charge. The number of protons determines the element, while the number of electrons determines the charge.

Cation: An ion with a positive charge, formed by losing electrons.
Anion: An ion with a negative charge, formed by gaining electrons.
Example: Nitride ion (N3−) has gained 3 electrons; Calcium ion (Ca2+) has lost 2 electrons.

Key Equations:

Charge = Number of protons − Number of electrons

Writing Ionic Formulas

Ionic compounds are formed from the electrostatic attraction between cations and anions. The formula is written so that the total positive and negative charges balance.

Example: Lithium ion (Li+) and Nitride ion (N3−) combine to form Li3N.

Naming Ionic Compounds

Ionic compounds are named by stating the cation first, followed by the anion. For transition metals, the charge is indicated in Roman numerals.

Example: Mg3N2 is named Magnesium Nitride.
Example: Cu2O is named Copper(I) Oxide.

Chemical Bonding and Lewis Structures

Lewis Structures

Lewis structures represent the arrangement of electrons in molecules, showing bonds and lone pairs.

Example: Cl2O: Each chlorine atom is bonded to the oxygen atom, with lone pairs shown on each atom.

Electronegativity and Bond Classification

Bonds are classified based on the difference in electronegativity between atoms:

Nonpolar Covalent: Electrons are shared equally (e.g., Br–Br).
Polar Covalent: Electrons are shared unequally, creating partial charges (δ+ and δ−).
Ionic: Electrons are transferred, resulting in full charges.
Example: Na–O is ionic; P–Cl is polar covalent.

VSEPR Theory and Molecular Shapes

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the shapes of molecules based on electron pair repulsion.

Example: SiCl4 has a tetrahedral shape.

Chemical Equations and Stoichiometry

Balancing Chemical Equations

Balanced equations have equal numbers of each atom on both sides. This is essential for stoichiometric calculations.

Example: Na3PO4(aq) + MgCl2(aq) → Mg3(PO4)2(s) + NaCl(aq)

Counting Atoms in Equations

Each compound's formula indicates the number of atoms present. For example, Fe2S3 contains 2 iron atoms and 3 sulfur atoms per formula unit.

Stoichiometry and Mole Calculations

Stoichiometry involves using balanced equations to calculate the amounts of reactants and products.

Example: In 2Fe(s) + 3S(s) → Fe2S3(s), 1.42 mol Fe requires (1.42 mol Fe) × (3 mol S / 2 mol Fe) = 2.13 mol S.

Mole-to-Mass Conversions

To convert moles to grams, use the molar mass:

Example: 0.750 mol Ag × 107.87 g/mol = 80.90 g Ag

Gas Laws and Calculations

Pressure Units and Conversions

Pressure can be measured in mmHg, atm, or other units. Conversion is often necessary:

Example: 4820 mmHg × (1 atm / 760 mmHg) = 6.34 atm

Gas Volume and Pressure Relationships

Boyle's Law describes the inverse relationship between pressure and volume for a fixed amount of gas at constant temperature:

Example: If pressure increases, volume decreases proportionally.

Gas Law Applications

Gas laws are used to solve problems involving changes in pressure, volume, and temperature.

Example: Calculating the final volume of oxygen in a tank when pressure changes.

Partial Pressure and Mixtures

Dalton's Law of Partial Pressures states that the total pressure of a mixture of gases is the sum of the partial pressures of each gas:

Example: If total pressure is 7.00 atm and oxygen is 1140 mmHg (1.50 atm), helium is 7.00 atm − 1.50 atm = 5.50 atm.

Summary Table: Types of Chemical Bonds

Bond Type	Electronegativity Difference	Electron Sharing	Example
Nonpolar Covalent	0	Equal	Br–Br
Polar Covalent	0.4–1.7	Unequal	P–Cl
Ionic	>1.7	Transferred	Na–O

Additional info:

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