Unit 6 quick test

Here’s a detailed study guide covering the Learning Targets (LTs) you mentioned. Use this as a structured reference to ace your understanding of electronegativity, polarity, intermolecular forces (IMFs), and their applications.

1. Differences in Electronegativity and Polarity

• Electronegativity: A measure of an atom’s ability to attract electrons in a chemical bond.

• Atoms with higher electronegativity (e.g., F, O, N) pull electrons closer to themselves.

• Polarity:

• When two atoms in a bond have different electronegativities, the shared electrons are unequally distributed, creating a polar bond.

• Polar molecules have a partial positive charge (δ⁺) and a partial negative charge (δ⁻) on opposite ends. Example: H₂O.

• Nonpolar bonds occur when the electronegativity difference is very small or zero (e.g., Cl₂, CH₄).

2. How Polarity Affects Physical Properties

• Physical Properties Influenced by Polarity:

• Melting and Boiling Points: Polar substances have higher melting/boiling points because their molecules experience strong intermolecular forces.

• Solubility:

• “Like dissolves like”: Polar substances dissolve in polar solvents (e.g., water), and nonpolar substances dissolve in nonpolar solvents (e.g., oil).

• Surface Tension: Higher in polar substances (e.g., water) due to strong cohesive forces.

• Capillary Action: Polar molecules climb surfaces because of adhesion to polar materials.

3. Types of Intermolecular Forces (IMFs) and Their Relative Strengths

• IMFs (In order of strength):

1. Hydrogen Bonding (Strongest IMF in covalent substances):

• Occurs when H is bonded to F, O, or N (e.g., H₂O, NH₃). Extremely strong dipole-dipole interaction.

2. Dipole-Dipole Forces:

• Found between polar molecules. Example: HCl.

3. London Dispersion Forces (LDF) (Weakest):

• Present in all molecules (polar and nonpolar).

• Dominant in nonpolar molecules due to temporary dipoles. Example: CH₄.

4. Identifying IMFs Based on Bonding

• Ionic Bonding:

• Strong electrostatic forces between ions. Example: NaCl.

• Covalent Bonding:

• Nonpolar covalent bonds: Only LDFs (e.g., Cl₂).

• Polar covalent bonds: LDFs + Dipole-Dipole or Hydrogen Bonding (e.g., H₂O).

• Steps to Identify IMFs:

1. Check if the molecule is polar or nonpolar.

2. Look for bonds involving H with F, O, or N for hydrogen bonding.

3. Determine if it’s ionic, which will have much stronger interactions than IMFs.

5. Relating Bulk Properties to IMFs

• Strong IMFs = Higher melting/boiling points.

• Substances with hydrogen bonds (e.g., water) often have anomalously high boiling points compared to other molecules of similar size.

• Viscosity: Higher in substances with strong IMFs.

• Volatility: Low in substances with strong IMFs; they evaporate slowly.

6. Describing Hydrogen Bonds

• Definition: A strong attraction between a hydrogen atom bonded to F, O, or N and another F, O, or N atom with a lone pair.

• Key Characteristics:

• Strongest IMF among covalent compounds.

• Responsible for water’s unique properties (e.g., high boiling point, surface tension).

7. Practicing Applications of IMFs

• Example 1: Rank in boiling points: CH₄, NH₃, H₂O.

• CH₄ (LDF) < NH₃ (Hydrogen Bonding) < H₂O (Stronger Hydrogen Bonding).

• Example 2: Predict solubility:

• NaCl dissolves in water (polar), not in oil (nonpolar).

• CH₄ dissolves in oil, not water.

• Practice Problems:

1. Identify the strongest IMF in the following: HCl, H₂O, CH₄.

2. Explain why ethanol (C₂H₅OH) is soluble in water but hexane (C₆H₁₄) is not.

3. Compare the boiling points of H₂S and H₂O and justify the difference.

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