Covalent Bonds, Electronegativity, and Lewis Structures: General Chemistry Study Notes

Ionization Energy, Electron Affinity, and Lattice Energy

Ionization Energy

Ionization energy is the energy required to remove an electron from a gaseous atom or ion. It is a fundamental property that influences chemical reactivity and bonding.

• Definition: The minimum energy needed to remove the outermost electron from a neutral atom in the gas phase.

• Trends:

  • Increases across a period (left to right) due to increasing nuclear charge.

  • Decreases down a group due to increasing atomic radius and electron shielding.

• Example: The ionization energies for carbon are:

  Ionization	Ionization Energy (kJ/mol)
  1	1086.5
  2	2352.6
  3	4620.5
  4	6222.7

Electron Affinity

Electron affinity is the energy change that occurs when an electron is added to a gaseous atom. It reflects an atom's tendency to gain electrons.

• Definition: The energy released (or required) when an atom in the gas phase accepts an electron.

• Trends:

  • Halogens have large, negative electron affinities (they readily gain electrons).

  • Noble gases have positive electron affinities (they do not readily gain electrons).

  • Alkaline earth metals have electron affinities close to zero.

Lattice Energy

Lattice energy is the energy released when ions come together to form an ionic solid. It is a key factor in the stability of ionic compounds.

• Definition: The energy required to separate one mole of an ionic solid into its gaseous ions.

• Relationship: Lattice energies and ionic bond strengths are directly proportional.

• Estimation: Lattice energies can be estimated using the Born-Haber cycle.

Covalent Bonds, Electronegativity, and Lewis Structures

Why Are Covalent Bonds Necessary?

Not all bonds are ionic; covalent bonds form when sharing electrons is more energetically favorable than transferring them. This is often the case when the energy required to form ions is too high to be compensated by lattice energy.

• Example: In CCl4, the energy required to form C4+ is over 14,000 kJ/mol, which is not available under typical reaction conditions.

• Conclusion: Covalent bonds form in preference to ionic bonds when electron sharing is energetically favored.

Localized Bonds

Covalent bonding is localized, meaning the properties of a covalent bond depend almost entirely on the two atoms involved.

• Bond Dissociation Energy: The energy required to break a bond; the negative of this is the energy released when a bond forms.

• Bond Length: The internuclear distance at which the energy is minimized.

• Table: Bond Dissociation Energies and Lengths

  Molecule	Bond	Bond Dissociation Energy (kJ/mol)	Bond Length (pm)
  C2H6	C–H	376	109.1
  C3H8	C–H	356	109.6
  C4H10	C–H	352	111.7

Realistic Bonds: Ionic Character and Bonding Continuum

Bonds exist on a continuum between complete electron transfer (ionic) and equal sharing (covalent). The degree of electron transfer is quantified as percent ionic character.

• Electrostatic Potential Maps: Visual representations of electron distribution in molecules.

• Examples:

  • Cl2: 0% ionic

  • HCl: 17% ionic

  • NaCl: 80% ionic

Bond Polarity

Bond polarity refers to the degree to which bonding electrons are transferred or shared unequally between atoms.

• Percent Ionic Character: Quantifies the extent of electron transfer.

• Symbols: δ+ and δ− indicate partial positive and negative charges, respectively.

• Table: Bond Polarity Comparison

  Compound	Percent Ionic Character	Bond Type	Representation of Bond Polarity
  Cl2	0	Nonpolar covalent	Cl–Cl
  HCl	17	Polar covalent	δ+H–Clδ−
  NaCl	80	Ionic	Na+Cl−

Electronegativity

Electronegativity is the ability of an atom to attract shared electrons in a chemical bond. Differences in electronegativity between atoms lead to bond polarity.

• Definition: A measure of the tendency of an atom to attract electrons in a bond.

• Trends:

  • Increases across a period; decreases down a group.

  • Fluorine (F), oxygen (O), and nitrogen (N) have the highest electronegativities.

  • Hydrogen's electronegativity is close to that of carbon.

• Bond Polarity and Electronegativity: The greater the difference in electronegativity, the more polar the bond.

Lewis Structures

Purpose and Importance

Lewis structures are diagrams that represent the sharing of electrons between atoms in covalent or polar covalent bonds. They are essential for understanding molecular structure and properties.

• Key Information: Shows which atoms are bonded and the nature of those bonds (single, double, triple).

• Application: Used to predict molecular geometry, reactivity, and physical properties.

Drawing Lewis Structures of Atoms

Lewis structures for atoms show the valence electrons as dots around the chemical symbol.

• Steps:

  1. Write the chemical symbol (represents nucleus and core electrons).

  2. Place dots around the symbol to represent valence electrons (up to 8).

  3. Arrange the first four dots singly; pair remaining dots.

• Example: The Lewis structure for oxygen (O) has six dots around the symbol.

Covalent Bonding Examples

• Cl2: Two chlorine atoms share a pair of electrons to satisfy the octet rule.

• H2O: Oxygen shares electrons with two hydrogens, forming two covalent bonds and two lone pairs.

• NH3: Nitrogen shares electrons with three hydrogens, forming ammonia.

Predicting Formulas of Covalent Compounds

Lewis structures and the octet rule can be used to predict the formulas of covalent compounds.

• Example: Silicon and fluorine combine to form SiF4, where silicon shares electrons with four fluorine atoms.

Multiple Covalent Bonds

Some atoms share more than one pair of electrons, resulting in double or triple bonds.

• Example: O2 forms a double bond by sharing two pairs of electrons.

• Triple bonds: Occur in molecules like N2.

Guidelines for Drawing Lewis Structures

• Arrange atoms: The atom lowest and leftmost in the periodic table is usually central.

• Hydrogen and halogens are typically terminal atoms.

• Determine total number of valence electrons (add for negative ions, subtract for positive ions).

• Connect terminal atoms to the central atom with single bonds (subtract two electrons per bond).

• Distribute remaining electrons in pairs to satisfy the octet rule (terminal atoms first, then central atom).

• If the central atom lacks an octet, convert lone pairs from terminal atoms into multiple bonds.

Lewis Structure Examples

• CCl4: Carbon is central, bonded to four chlorines, each with three lone pairs.

• NH4+: Nitrogen is central, bonded to four hydrogens; brackets and superscript indicate charge.

• CO2: Carbon is central, double bonded to two oxygens.

Practice Drawing Lewis Structures

• Examples for practice: PCl3, H2CO, O3, C2H4.

• Follow the guidelines above for each molecule.

Summary of Key Concepts

• Covalent bonding is localized; bond properties depend on the two atoms involved.

• Bonding exists on a continuum from complete electron transfer (ionic) to equal sharing (covalent).

• Bond polarity is determined by the difference in electronegativity between atoms.

• Lewis structures are essential for visualizing electron sharing and predicting molecular formulas.