CHEM 1104 Exam 2 Study Guide: Ionic & Molecular Compounds, Chemical Reactions, and Gases

Chapter 6: Ionic and Molecular Compounds

Octet Rule

The octet rule states that atoms tend to gain, lose, or share electrons to achieve a stable configuration of eight valence electrons, similar to noble gases.

• Key Point: Atoms with a full octet are generally more stable.

• Example: Sodium (Na) loses one electron to achieve the electron configuration of neon.

Formation of Ions

Atoms form ions by gaining or losing electrons:

• Positive ions (cations): Formed when atoms lose electrons (usually metals).

• Negative ions (anions): Formed when atoms gain electrons (usually nonmetals).

• Example: Na → Na+ + e-; Cl + e- → Cl-

Groups and Their Ions

• Group 1A, 2A, 3A metals: Form +1, +2, +3 ions respectively.

• Group 5A, 6A, 7A nonmetals: Form -3, -2, -1 ions respectively.

• Example: Mg (Group 2A) forms Mg2+; O (Group 6A) forms O2-

Ionic Bonds and Ionic Compounds

An ionic bond is the electrostatic attraction between oppositely charged ions. Ionic compounds are formed from metals and nonmetals.

• Example: Na+ and Cl- combine to form NaCl.

Formulas and Charge Balance

Formulas for ionic compounds must be electrically neutral. The total positive charge must balance the total negative charge.

• Example: Al3+ and O2- combine to form Al2O3.

Naming Ionic Compounds

• Binary ionic compounds: Name the metal first, then the nonmetal with an "-ide" ending (e.g., NaCl = sodium chloride).

• Metals with variable charges: Use Roman numerals to indicate the charge (e.g., FeCl2 = iron(II) chloride).

Polyatomic Ions

Polyatomic ions are ions composed of more than one atom. Memorize the following:

Name	Formula	Charge
Ammonium	NH4+	+1
Hydroxide	OH-	-1
Cyanide	CN-	-1
Nitrate	NO3-	-1
Bicarbonate	HCO3-	-1
Carbonate	CO32-	-2
Sulfate	SO42-	-2
Phosphate	PO43-	-3

Naming Ionic Compounds with Polyatomic Ions

• Name the cation first, then the polyatomic ion (e.g., NaNO3 = sodium nitrate).

Covalent Bonds and Molecular Compounds

Covalent bonds involve the sharing of electrons between nonmetals. Molecular compounds are composed of molecules formed by covalent bonds.

• Diatomic molecules: H2, N2, O2, F2, Cl2, Br2, I2

• Double bonds: Two pairs of electrons shared (e.g., O2); Triple bonds: Three pairs shared (e.g., N2).

Naming Covalent Compounds

• Use prefixes to indicate the number of each atom: di- (2), tri- (3), tetra- (4), penta- (5), hexa- (6).

• Example: CO2 = carbon dioxide; PCl3 = phosphorus trichloride.

Identifying Ionic vs. Covalent Compounds

• Ionic: Metal + nonmetal or contains polyatomic ions.

• Covalent: Only nonmetals.

Electronegativity and Bond Polarity

Electronegativity is the ability of an atom to attract shared electrons. The difference in electronegativity determines bond type:

• Nonpolar covalent: Electronegativity difference < 0.5

• Polar covalent: Difference 0.5–1.9

• Ionic: Difference > 1.9

Molecular Geometry (VSEPR Theory)

The shape of molecules is determined by the number of bonding and lone pairs around the central atom:

Geometry	Example
Linear	BeF2, CO2
Trigonal planar	BH3, BCl3
Trigonal pyramidal	PCl3, NH3
Tetrahedral	CH4, CCl4
Bent	SO2, H2O

Intermolecular Forces

• Dispersion forces: Weakest, present in all molecules.

• Dipole-dipole attractions: Between polar molecules.

• Hydrogen bonding: Strongest, occurs when H is bonded to N, O, or F.

• Example: H2O molecules exhibit hydrogen bonding.

Physical Properties and Intermolecular Forces

• Stronger intermolecular forces lead to higher boiling and melting points.

• Physical state (solid, liquid, gas) depends on the strength of these forces.

Chapter 7: Chemical Quantities and Reactions

Mole Concept

One mole contains $6.02 \times 10^{23}$ particles (Avogadro's number).

• Example: 1 mole of H2O contains $6.02 \times 10^{23}$ molecules.

Converting Moles and Particles

• To particles: $\text{moles} \times 6.02 \times 10^{23} = \text{number of particles}$

• To moles: $\text{particles} \div 6.02 \times 10^{23} = \text{moles}$

Molar Mass

Molar mass is the mass of one mole of a substance (g/mol).

• Calculation: Add the atomic masses of all atoms in the formula.

• Example: Molar mass of H2O = 2(1.01) + 16.00 = 18.02 g/mol

Conversions Using Molar Mass

• Grams to moles: $\text{grams} \div \text{molar mass} = \text{moles}$

• Moles to grams: $\text{moles} \times \text{molar mass} = \text{grams}$

Balancing Chemical Equations

Equations must have the same number of each atom on both sides.

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

Types of Chemical Reactions

• Combination: $A + B \rightarrow AB$

• Decomposition: $AB \rightarrow A + B$

• Single replacement: $A + BC \rightarrow AC + B$

• Double replacement: $AB + CD \rightarrow AD + CB$

• Combustion: Hydrocarbon + O2 → CO2 + H2O

Oxidation-Reduction (Redox) Reactions

• Oxidation: Loss of electrons

• Reduction: Gain of electrons

• Example: $2Na + Cl_2 \rightarrow 2NaCl$ (Na is oxidized, Cl is reduced)

Mole-to-Mole Conversions

• Use coefficients from balanced equations to convert between moles of reactants and products.

• Example: $2H_2 + O_2 \rightarrow 2H_2O$ (2 moles H2 produce 2 moles H2O)

Conditions for a Reaction to Occur

• Reactant particles must collide.

• Collisions must have sufficient energy (activation energy).

• Particles must be oriented properly.

Exothermic and Endothermic Reactions

• Exothermic: Release heat ($\Delta H$ is negative)

• Endothermic: Absorb heat ($\Delta H$ is positive)

Catalysts

• A catalyst speeds up a reaction by lowering the activation energy, without being consumed.

Chapter 8: Gases

Kinetic Molecular Theory

Describes the behavior of gases:

• Gases consist of small particles in constant, random motion.

• Collisions are elastic; energy is conserved.

• Volume of gas particles is negligible compared to container volume.

• No attractive forces between particles.

Gas Properties

• Pressure (P): Force per unit area, measured in atm, mmHg, or kPa.

• Volume (V): Space occupied, usually in liters (L).

• Temperature (T): Must be in Kelvin (K).

• Amount (n): Number of moles.

Direct and Inverse Relationships

• Direct: Both variables increase or decrease together (e.g., V ∝ T at constant P).

• Inverse: One variable increases as the other decreases (e.g., P ∝ 1/V at constant T).

Gas Laws

• 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)

• Gay-Lussac's Law: $\frac{P_1}{T_1} = \frac{P_2}{T_2}$ (at constant V and n)

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

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

Standard Temperature and Pressure (STP)

• STP: 0°C (273 K) and 1 atm pressure.

Molar Volume

• At STP, 1 mole of any gas occupies 22.4 L.

Dalton's Law of Partial Pressures

• The total pressure of a gas mixture is the sum of the partial pressures of each component.

• Equation: $P_{\text{total}} = P_1 + P_2 + P_3 + \ldots$

Additional info: Where the original notes were brief, standard textbook explanations and examples have been added for clarity and completeness.