Chapters 6–8: Ions, Molecular Compounds, Chemical Reactions, and Gases – Study Notes

Chapter 6: Ions and Molecular Compounds

6.1 Ions: Transfer of Electrons

Ions are formed when atoms gain or lose electrons to achieve a stable electron configuration. This process is fundamental to the formation of ionic compounds.

• Metals lose electrons to form cations (positively charged ions).

• Nonmetals gain electrons to form anions (negatively charged ions).

• Ionic bonds are the electrostatic attractions between cations and anions.

• Covalent bonds involve the sharing of electrons between nonmetals.

Octet Rule: Atoms tend to gain, lose, or share electrons to achieve eight valence electrons (a noble gas configuration).

Example: Sodium (Na) loses one electron to form Na+; chlorine (Cl) gains one electron to form Cl-.

6.2 Ionic Compounds

Ionic compounds are composed of cations and anions held together by ionic bonds. They are electrically neutral overall.

• Properties: High melting points, crystalline solids, conduct electricity when dissolved in water.

• Formulas: The ratio of ions in the formula ensures overall neutrality (e.g., NaCl, MgCl2).

Example: Magnesium chloride: Mg2+ and Cl- combine in a 1:2 ratio to form MgCl2.

6.3 Naming and Writing Ionic Compounds

The name of an ionic compound consists of the name of the cation followed by the name of the anion. For transition metals, Roman numerals indicate the charge.

• Monatomic ions: Use the element name for cations and the root + “-ide” for anions (e.g., sodium chloride).

• Polyatomic ions: Use the ion name as is (e.g., ammonium sulfate).

Example: FeCl3 is iron(III) chloride.

6.4 Polyatomic Ions

Polyatomic ions are ions composed of two or more atoms covalently bonded, carrying a net charge.

• Examples: NO3- (nitrate), SO42- (sulfate), NH4+ (ammonium).

• Polyatomic ions often end in “-ate” or “-ite.”

Formula	Name	Common Use/Meaning
Zn(C2H3O2)2	zinc acetate	cold remedy
Mg3(PO4)2	magnesium phosphate	antacid
NH4Cl	ammonium chloride	expectorant (coughs)
NaHCO3	sodium hydrogen carbonate	baking soda
NaNO2	sodium nitrite	used in processed meat

6.5 Molecular Compounds: Sharing Electrons

Molecular compounds consist of nonmetals sharing electrons to achieve stability. Covalent bonds are formed when atoms share one or more pairs of electrons.

• Formulas: Use prefixes to indicate the number of each atom (e.g., CO2 is carbon dioxide).

• Examples: H2O (water), NH3 (ammonia).

6.7 Electronegativity and Bond Polarity

Electronegativity is the ability of an atom to attract electrons in a bond. The difference in electronegativity determines bond polarity.

• Nonpolar covalent bond: Electrons are shared equally (e.g., H2).

• Polar covalent bond: Electrons are shared unequally (e.g., HCl).

• Ionic bond: Large difference in electronegativity (e.g., NaCl).

Example: In H2O, the O-H bonds are polar, making water a polar molecule.

6.8 Molecular Shape and Polarity

The shape of molecules is determined by the number of electron groups around the central atom (VSEPR theory). Molecular shape affects polarity.

• Linear: 2 electron groups (e.g., CO2).

• Trigonal planar: 3 electron groups (e.g., BF3).

• Tetrahedral: 4 electron groups (e.g., CH4).

Example: NH3 is trigonal pyramidal due to a lone pair on nitrogen.

6.9 Intermolecular Forces in Compounds

Intermolecular forces are attractions between molecules that affect physical properties.

• Dipole-dipole interactions: Between polar molecules.

• Hydrogen bonds: Strong dipole interactions involving H bonded to N, O, or F.

• Dispersion forces: Weak attractions present in all molecules, stronger in larger molecules.

Example: Water has hydrogen bonding, making it have a high boiling point.

Chapter 7: Chemical Reactions and Quantities

7.1 Equations for Chemical Reactions

Chemical equations represent the reactants and products in a chemical reaction. They must be balanced to obey the law of conservation of mass.

• Reactants: Substances present before the reaction.

• Products: Substances formed by the reaction.

• Balancing equations: Adjust coefficients to have the same number of each atom on both sides.

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

7.2 Types of Chemical Reactions

• Combination: Two or more substances form one product. $A + B \rightarrow AB$

• Decomposition: One substance breaks into two or more products. $AB \rightarrow A + B$

• Single Replacement: One element replaces another. $A + BC \rightarrow AC + B$

• Double Replacement: Exchange of ions between two compounds. $AB + CD \rightarrow AD + CB$

• Combustion: A substance reacts with O2 to produce energy, CO2, and H2O.

7.3 Oxidation-Reduction Reactions

Redox reactions involve the transfer of electrons. Oxidation is the loss of electrons; reduction is the gain of electrons.

• Oxidizing agent: Causes oxidation, is reduced.

• Reducing agent: Causes reduction, is oxidized.

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

7.4 The Mole

The mole is a counting unit for atoms, molecules, and ions. One mole contains $6.02 \times 10^{23}$ particles (Avogadro's number).

• Molar mass: The mass of one mole of a substance (g/mol).

• Conversions: Use molar mass to convert between grams and moles.

Example: $\text{Moles} = \frac{\text{grams}}{\text{molar mass}}$

7.5 Calculations with Moles and Molar Mass

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

• Use conversion factors from the coefficients in the balanced equation.

• Calculate theoretical yield, limiting reactant, and percent yield.

Example: If 2 mol H2 react with 1 mol O2, 2 mol H2O are produced.

Chapter 8: Gases

8.1 Properties of Gases

Gases have unique properties: they expand to fill their container, are compressible, and have low density.

• Atmosphere: The layer of gases surrounding Earth.

• Kinetic Molecular Theory: Gas particles move rapidly and randomly, with negligible volume and no attractive forces.

8.2 Pressure and Volume (Boyle's Law)

Boyle's Law describes the inverse relationship between pressure and volume at constant temperature.

$P_1V_1 = P_2V_2$

Example: If the volume of a gas decreases, its pressure increases.

8.3 Temperature and Volume (Charles's Law)

Charles's Law states that the volume of a gas is directly proportional to its temperature (in Kelvin) at constant pressure.

$\frac{V_1}{T_1} = \frac{V_2}{T_2}$

8.4 Temperature and Pressure (Gay-Lussac's Law)

Gay-Lussac's Law states that the pressure of a gas is directly proportional to its temperature (in Kelvin) at constant volume.

$\frac{P_1}{T_1} = \frac{P_2}{T_2}$

8.5 The Combined Gas Law

The combined gas law relates pressure, volume, and temperature for a fixed amount of gas.

$\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$

8.6 Volume and Moles (Avogadro's Law)

Avogadro's Law states that the volume of a gas is directly proportional to the number of moles at constant temperature and pressure.

$\frac{V_1}{n_1} = \frac{V_2}{n_2}$

8.7 The Ideal Gas Law

The ideal gas law combines all the gas laws into one equation:

$PV = nRT$

• P = pressure (atm)

• V = volume (L)

• n = moles

• R = ideal gas constant (0.0821 L·atm/mol·K)

• T = temperature (K)

8.8 Partial Pressures (Dalton's Law)

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

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

Additional info:

• These notes include diagrams, tables, and worked examples to illustrate key concepts.

• Practice problems and summary tables are provided for reinforcement.