Unit 4 Chemistry

Unit 4 Practice Test Notes

1. Types of Bonds

  • Main Types of Bonds:

    • Ionic Bonds

    • Polar Covalent Bonds

    • Metallic Bonds: Characterized by a sea of delocalized electrons surrounding positively charged metal ions, allowing for conductivity and malleability.

  • Elements that Form Each Type:

    • Ionic Bonds: Typically form between metals and nonmetals.

    • Polar Covalent Bonds: Typically form between two nonmetals with different electronegativities.

    • Metallic: Typically form between nonmetals with similar electronegativities.

  • Electronegativity Ranges for Bonds:

    • Ionic Bond: Electronegativity difference > 1.7

    • Polar Covalent Bond: Electronegativity difference between 0.4 and 1.7

    • Non-polar covalent bond : Electronegativity difference < 0.4

  • Methane (CH4):

    • Methane is a nonpolar covalent molecule.

    • Explanation: The electronegativity difference between carbon (C) and hydrogen (H) is approximately 0.4, indicating a nonpolar bond.

2. Intermolecular Forces (IMFs)

  • Properties in Strong IMF Molecules:

    • Melting Point: High

    • Boiling Point: High

    • Surface Tension: High

  • Relation of IMFs to Particle Attraction:

    • Stronger IMFs result in greater attraction between particles, leading to higher melting/boiling points and surface tension.

  • Types of Intermolecular Forces:

    1. London Dispersion Forces:

      • Weakest, present in all molecules (nonpolar)

    2. Dipole-Dipole Interactions:

      • Occur in polar molecules, stronger than London forces

    3. Hydrogen Bonding:

      • Strongest, occurs with H bonded to F, O, or N

3. Naming Compounds

  • Ionic Compounds:

    • Naming Rules:

      • Cation name + Anion name (with -ide for simple anions)

      • Example: NaCl (Sodium Chloride)

  • Covalent Compounds:

    • Naming Rules:

      • Use prefixes to indicate the number of each atom

      • Example: CO2 (Carbon Dioxide)

  • Acids:

    • Naming Rules:

      • If anion ends in -ide, use prefix "hydro-" and suffix "-ic". If anion ends in -ate, suffix is "-ic". If anion ends in -ite, suffix is "-ous".

      • Example: HCl (Hydrochloric Acid), H2SO4 (Sulfuric Acid)

4. Types of Reactions

  • General Equations for Types of Reactions:

    1. Combustion:

      • Hydrocarbon + O2 → CO2 + H2O

    2. Single Replacement:

      • A + BC → B + AC

    3. Double Replacement:

      • AB + CD → AD + CB

    4. Decomposition:

      • AB → A + B

    5. Synthesis:

      • A + B → AB

5. Endothermic vs Exothermic Reactions

  • Definitions:

    • Endothermic:

      • Absorbs heat, resulting in a temperature decrease in the surrounding environment.

    • Exothermic:

      • Releases heat, resulting in a temperature increase in the surrounding environment.

  • Energy Diagrams:

    • Endothermic Reaction:

      • Reactants start lower, end higher indicating heat absorption.

    • Exothermic Reaction:

      • Reactants start higher, end lower indicating heat release.

6. Balancing Equations

  • Equations to Balance:

    1. Equation: ( \text{Ca(OH)}_2 + \text{Al}_2(\text{SO}_4)_3 \rightarrow \text{CaSO}_4 + \text{Al(OH)}_3 )

      • Balance:

    2. Equation: ( \text{C}_2\text{H}_4 + \text{O}_2 \rightarrow \text{CO}_2 + \text{H}_2\text{O} )

      • Balance:

7. Stoichiometry (Bonus Questions)

  • Moles of Sodium Phosphate Needed:

    • From reaction (Na_3PO_4 + 3 KOH \rightarrow 3 NaOH + K_3PO_4):

      • To produce 6.8 moles of NaOH, need 2.267 moles of Na3PO4.

  • Moles in 140 Grams of Sodium Phosphate:

    • Molar mass calculation needed; typically: ( \text{Molar Mass of Na3PO4} \approx 163.94 g/mol ).

    • Result: ( \approx 0.854 \text{ moles} )

  • Mass of Sodium Hydroxide Produced:

    • Calculate based on moles from Na3PO4 to NaOH conversion.

    • Total: Mass from stoichiometry based on production from 140 grams of Na3PO4.