Chemical Bonding, Polarity, Reactions and Intermolecular Forces
Endothermic Reactions
An endothermic reaction absorbs energy.
- Heat (kJ) is written on the reactant side.
- Bonds are broken; no new bonds are formed.
- On Table I, the heat of reaction is positive (+).
- The surroundings feel colder because energy is absorbed.
Exothermic Reactions
An exothermic reaction releases energy.
- Heat (kJ) is written on the product side.
- Bonds are formed.
- On Table I, the heat of reaction is negative (−).
- The surroundings feel warmer because energy is released.
Electronegativity
- Electronegativity is the ability of an atom to attract electrons.
- The greater the electronegativity, the stronger the attraction for electrons in a bond.
- Differences in electronegativity determine bond type and polarity.
Formation of Ions
Metals
- Have 1, 2, or 3 valence electrons.
- Lose electrons.
- Become positive ions (cations).
- Ion size becomes smaller than the original atom.
Nonmetals
- Have 5, 6, or 7 valence electrons.
- Gain electrons.
- Become negative ions (anions).
- Ion size becomes larger than the original atom.
Electron Configuration
- Ions always achieve the electron configuration of a noble gas (Group 18).
Octet Rule
- Atoms gain, lose, or share electrons to obtain 8 valence electrons.
- This gives them the stability of a noble gas.
- Exceptions:
- Hydrogen, lithium, and beryllium are stable with 2 valence electrons (helium configuration).
Ionic Bonds
- Form between a metal and a nonmetal.
- The metal loses electrons.
- The nonmetal gains electrons.
- Electrons are transferred, not shared.
- Oppositely charged ions attract strongly, forming an ionic bond.
Covalent Bonds
- Form between two nonmetals.
- Electrons are shared.
- Both atoms want electrons, so they share as a compromise.
Types of Covalent Bonds
- Single bond: 1 pair of shared electrons (Cl–Cl).
- Double bond: 2 pairs of shared electrons (O=O).
- Triple bond: 3 pairs of shared electrons (N≡N).
Covalent Bond Polarity
Nonpolar Covalent Bonds
- Equal sharing of electrons.
- Equal electronegativity values.
- Equal attraction for electrons.
- Electronegativity difference is 0 or very small (≈ 0–0.4).
Polar Covalent Bonds
- Unequal sharing of electrons.
- Unequal electronegativity between atoms.
- One atom pulls electrons closer, creating partial charges.
- Characterized by a larger electronegativity difference.
Electronegativity Difference Rules
- Greater difference → more polar, more ionic, less covalent.
- Smaller difference → less polar, less ionic, more covalent.
Example:
- HCl is more polar than HBr because H and Cl have a larger electronegativity difference than H and Br.
Determining Bond Type
- Metal + nonmetal = ionic.
- Nonmetal + nonmetal = covalent.
- Diatomic molecules are always nonpolar.
- Compounds with polyatomic ions contain both ionic and covalent bonds.
Molecules
- Molecules are compounds made of covalent bonds.
- Molecules can be polar or nonpolar.
Nonpolar Molecules
- Symmetrical shape.
- Equal charge distribution.
- Examples: CH₄, CCl₄, CO₂, all diatomic molecules.
Polar Molecules
- Asymmetrical shape.
- Unequal charge distribution.
- Examples: H₂O, H₂S, NH₃, CO, HCl.
Key Reminder
- Bond polarity → check electronegativity difference.
- Molecular polarity → check symmetry.
Intermolecular Forces (IMF)
- Forces between molecules, not within molecules.
- Do not form compounds.
- Hold molecules together in the liquid and solid phases.
- Higher melting and boiling points = stronger IMF.
- When comparing substances, the one with the higher melting point (m.p.) or boiling point (b.p.) has stronger intermolecular forces.
Hydrogen Bonding
- A special, strong intermolecular force.
- Stronger than many other IMFs.
- Causes higher melting and boiling points.
- Must contain: hydrogen bonded to oxygen, nitrogen, or fluorine.
- Examples: H₂O, NH₃, HF.
Metallic Bonds
- Found in pure metals only.
- Do not form compounds.
- Electrons are free-moving (sea of mobile electrons).
- Explains why metals conduct electricity in solid, liquid, and molten states.