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 compound is such that the total positive charge equals the total negative charge.

Example: NaCl (sodium chloride) consists of Na+ and Cl- ions in a 1:1 ratio.

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.

• Polyatomic ions: Use the common name (e.g., sulfate, nitrate).

Example: FeCl2 is iron(II) chloride; 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).

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 are formed when two or more nonmetals share electrons to achieve stability.

• Covalent bonds: Shared pairs of electrons between atoms.

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

Example: H2O (water) is a molecular compound with covalent bonds between hydrogen and oxygen.

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: Electrons are transferred (large electronegativity difference).

Example: In H2O, the O-H bonds are polar covalent.

6.8 Molecular Shape and Polarity

The shape of a molecule is determined by the number of electron groups around the central atom (VSEPR theory).

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

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

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

Polarity of molecules: Determined by both bond polarity and molecular shape.

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-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 substances involved in a chemical reaction, showing reactants and products.

• Law of Conservation of Mass: Atoms are neither created nor destroyed in a chemical reaction.

• Balancing equations: The number of atoms of each element must be the same on both sides of the equation.

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 splits into two or more products. $AB \rightarrow A + B$

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

• Double Replacement: Two compounds exchange ions. $AB + CD \rightarrow AD + CB$

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

7.3 Oxidation-Reduction (Redox) Reactions

Redox reactions involve the transfer of electrons.

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

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. Avogadro's number is $6.02 \times 10^{23}$ particles per mole.

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

Example: 1 mole of H2O has a mass of 18.0 g.

7.5 Calculations with Moles and Molar Mass

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

• Use conversion factors: grams ↔ moles ↔ particles.

Example: To find grams of product from grams of reactant, convert grams to moles, use the mole ratio, then convert back to grams.

7.10 Energy in Chemical Reactions

Chemical reactions involve changes in energy.

• Exothermic: Release energy (products have lower energy than reactants).

• Endothermic: Absorb energy (products have higher energy than reactants).

Example: Combustion reactions are exothermic.

Chapter 8: Gases

8.1 Properties of Gases

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

• 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.

$P_1V_1 = P_2V_2$

8.3 Temperature and Volume (Charles's Law)

Charles's Law states that volume is directly proportional to 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 pressure is directly proportional to temperature 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.

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

8.6 Volume and Moles (Avogadro's Law)

Avogadro's Law states that volume 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 = 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.

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

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