AP Chemistry: Ionic vs Covalent Bonding and Lattice Energy

AP Chemistry – Ionic vs Covalent Bonding and Lattice Energy

AP Chemistry Study Guide: Ionic vs. Covalent Bonding & Lattice Energy

What This Is

Ionic and covalent bonding explain how atoms stick together to form compounds—ionic bonds involve electron transfer (metal + nonmetal), while covalent bonds involve electron sharing (nonmetal + nonmetal). Lattice energy measures the strength of ionic bonds and determines properties like melting point and solubility. This topic is high-yield on the AP exam because it connects to intermolecular forces, thermodynamics, and real-world applications (e.g., why salt dissolves in water but diamond doesn’t). Historical example: In 1916, Gilbert Lewis proposed the "cubical atom" model to explain covalent bonding, laying the foundation for modern valence bond theory.

Key Terms & Concepts

• Ionic Bond: A bond formed by the complete transfer of electrons from a metal (low ionization energy) to a nonmetal (high electron affinity), creating oppositely charged ions (e.g., Na?Cl?).
• Covalent Bond: A bond formed by the sharing of electron pairs between nonmetals (e.g., H?O, CO?). Can be polar (unequal sharing, e.g., HCl) or nonpolar (equal sharing, e.g., O?).
• Electronegativity (EN): A measure of an atom’s ability to attract shared electrons. Trend: Increases up and right on the periodic table (F is most electronegative).
• Bond Polarity: Determined by the difference in EN between two atoms:
• 0–0.4: Nonpolar covalent (e.g., Cl?)
• 0.5–1.7: Polar covalent (e.g., H?O)
• >1.7: Ionic (e.g., NaCl)
• Lattice Energy (?H_lattice): The energy released when gaseous ions form a solid ionic lattice (always exothermic, -?H). Higher lattice energy = stronger ionic bond = higher melting point.
• Formula: ?H_lattice? (|q?q?|) / r
  • q?, q? = ion charges (e.g., +1, -2)
  • r = distance between ion centers (sum of ionic radii)
• Coulomb’s Law: Describes the force between charged particles. Key idea: Lattice energy increases with higher ion charges and smaller ion sizes.
• Born-Haber Cycle: A thermodynamic cycle used to calculate lattice energy indirectly (AP exam may ask you to interpret this, not calculate).
• Resonance: When a molecule can be represented by multiple valid Lewis structures (e.g., O?, CO?²?). The actual structure is a hybrid of these.
• Formal Charge: Used to determine the most stable Lewis structure.
• Formula: Formal charge = (valence e?) – (nonbonding e?) – ½(bonding e?)
• Goal: Minimize formal charges (0 is ideal; negative charges should be on more electronegative atoms).
• Octet Rule: Atoms tend to gain, lose, or share electrons to achieve 8 valence electrons (exceptions: H, He, B, Be, and expanded octets for Period 3+).
• Dipole Moment (?): A measure of bond polarity (vector quantity). Unit: Debye (D). Example: H?O has a dipole moment because its bent shape creates a net charge separation.

Step-by-Step / Process Flow

1. Predicting Bond Type (Ionic vs. Covalent)

1. Identify the elements: Metal + nonmetal-ionic. Nonmetal + nonmetal-covalent.
2. Check electronegativity difference (?EN):
3. ?EN > 1.7-ionic
4. ?EN 0.5–1.7-polar covalent
5. ?EN < 0.5-nonpolar covalent
6. Example: MgO (?EN = 3.5 – 1.2 = 2.3-ionic). CO? (?EN = 3.5 – 2.5 = 1.0-polar covalent).

2. Drawing Lewis Structures for Covalent Compounds

1. Count total valence electrons (add/subtract for ions).
2. Draw a skeleton structure (least electronegative atom in the center; H and halogens are usually terminal).
3. Connect atoms with single bonds (2 e? per bond).
4. Distribute remaining electrons to satisfy the octet rule (start with terminal atoms).
5. Check for octets: If an atom lacks an octet, form double/triple bonds.
6. Calculate formal charges and adjust if needed (e.g., move lone pairs to form multiple bonds).
7. Example: CO?-O=C=O (total valence e? = 16; double bonds satisfy octets).

3. Comparing Lattice Energies

1. Compare ion charges first: Higher charges (e.g., Mg²?O²? vs. Na?Cl?)-stronger lattice energy.
2. If charges are equal, compare ion sizes: Smaller ions (e.g., LiF vs. NaCl)-stronger lattice energy (shorter distance between ions).
3. Example: Which has higher lattice energy, NaF or MgO?
4. MgO (charges: +2/-2) > NaF (charges: +1/-1)-MgO wins.

4. Using the Born-Haber Cycle (Conceptual)

1. Understand the steps: Sublimation (solid-gas), ionization energy, bond dissociation, electron affinity, lattice energy.
2. AP trick: You won’t calculate it, but you may be asked to identify which step corresponds to lattice energy (always the last step, forming the solid lattice).
3. Example: For NaCl, lattice energy is the energy released when Na?(g) + Cl?(g)-NaCl(s).

Common Mistakes

• Mistake: Assuming all metal-nonmetal bonds are ionic.

• Correction: Check ?EN! AlCl? (?EN = 1.5) is polar covalent, not ionic.

• Mistake: Forgetting that lattice energy is exothermic (negative ?H).

• Correction: Lattice energy is released when ions form a solid, so it’s always negative (e.g., -787 kJ/mol for NaCl).

• Mistake: Thinking larger ions always mean weaker lattice energy.

• Correction: Charge matters more! MgO (small ions, +2/-2) has higher lattice energy than NaCl (larger ions, +1/-1).

• Mistake: Drawing Lewis structures with incorrect formal charges.

• Correction: The most stable structure has formal charges closest to zero (e.g., CO? is O=C=O, not C?O-O).

• Mistake: Confusing bond polarity with molecular polarity.

• Correction: A molecule can have polar bonds but be nonpolar overall if dipoles cancel (e.g., CO? is linear-nonpolar; H?O is bent-polar).

AP Exam Insights

1. FRQ Favorite: You’ll often be asked to:
2. Compare lattice energies (e.g., "Which has a higher lattice energy, CaO or KCl? Explain.").
3. Draw Lewis structures and justify them with formal charges (e.g., "Draw two possible structures for N?O and identify the more stable one.").

4. Explain bond polarity (e.g., "Why is H?O polar but CO? is not?").

5. Multiple-Choice Traps:

6. Tricky distinction: "Which has the highest melting point?"-Look for highest lattice energy (charge > size).

7. Misleading options: A question might list bond energy (covalent) instead of lattice energy (ionic) as the answer choice.

8. Lab Connection: Lattice energy explains solubility trends (e.g., why Mg(OH)? is less soluble than NaOH—higher lattice energy = harder to break apart).

9. Thermodynamics Link: Lattice energy is part of Hess’s Law problems (e.g., calculating ?H° for a reaction using bond energies and lattice energy).

Quick Check Questions

1. Multiple Choice

Which of the following compounds has the highest lattice energy? (A) NaCl (B) MgO (C) KBr (D) CaS

Answer: (B) MgO Explanation: MgO has +2/-2 charges, while the others have +1/-1 or +2/-2 with larger ions (CaS). Higher charge = stronger lattice energy.

2. Short FRQ

a) Draw the Lewis structure for the nitrate ion (NO). Include resonance structures. b) Explain why the actual structure is a hybrid of these resonance forms.

Answer: a)

   O
   ||
O—N—O?
   |
   O?

(Show all 3 resonance structures with double bonds rotating between N and each O.) b) The actual structure is a hybrid because the electrons are delocalized across all N-O bonds, giving each bond a bond order of 1.33 (not single or double).

3. Multiple Choice

Which of the following molecules is nonpolar? (A) NH? (B) CH?Cl? (C) CO? (D) H?O

Answer: (C) CO? Explanation: CO? is linear (O=C=O), so its dipoles cancel. NH? and H?O are bent (polar), and CH?Cl? is tetrahedral with asymmetrical dipoles.

Last-Minute Cram Sheet

1. Ionic bond: Metal + nonmetal, electron transfer, high ?EN (>1.7).
2. Covalent bond: Nonmetal + nonmetal, electron sharing, ?EN < 1.7.
3. Lattice energy? (|q?q?|) / r-charge matters more than size!
4. Higher lattice energy = higher melting point, lower solubility.
5. Polar covalent: Unequal sharing (e.g., HCl); nonpolar covalent: equal sharing (e.g., O?).
6. Formal charge = valence e? – (nonbonding + ½ bonding e?)-minimize formal charges!
7. Resonance: Multiple valid Lewis structures-actual structure is a hybrid.
8. Octet rule exceptions: H (2 e?), Be (4 e?), B (6 e?), Period 3+ (expanded octets).
9. Dipole moment: Polar bonds + asymmetrical shape-polar molecule (e.g., H?O).
10. Lattice energy is exothermic (negative ?H)! Don’t mix it up with ionization energy (endothermic).