General Chemistry I: Exam 4 Review Study Guide – Bonding, Structure, and Periodic Trends

Section 1: Bond Length and Bond Strength

Bond Length and Strength Overview

Chemical bonds vary in length and strength depending on the atoms involved and the type of bond (single, double, triple). Understanding these properties is essential for predicting molecular behavior and reactivity.

• Bond Length: The distance between the nuclei of two bonded atoms. Generally, triple bonds are shortest, double bonds are intermediate, and single bonds are longest.

• Bond Strength: The energy required to break a bond. Triple bonds are strongest, followed by double, then single bonds.

• Example: In nitrogen molecules, $\mathrm{N \equiv N}$ (triple bond) is shorter and stronger than $\mathrm{N = N}$ (double bond) or $\mathrm{N - N}$ (single bond).

Bond Order and Bond Energy

• Bond Order: Number of shared electron pairs between two atoms. Higher bond order means shorter and stronger bonds.

• Bond Energy: The energy needed to break one mole of bonds in gaseous molecules.

• Formula: $\text{Bond Energy} = \sum \text{Energy of bonds broken} - \sum \text{Energy of bonds formed}$

Section 2: Electronegativity

Electronegativity and Bond Polarity

Electronegativity is the tendency of an atom to attract electrons in a chemical bond. It determines bond polarity and affects molecular properties.

• Definition: Electronegativity is a measure of an atom's ability to attract shared electrons.

• Trends: Electronegativity increases across a period and decreases down a group.

• Bond Polarity: A bond between atoms of different electronegativities is polar; the more electronegative atom attracts electrons more strongly.

• Example: In HCl, Cl is more electronegative than H, so the bond is polar.

Ionic Character and Electronegativity Differences

• Ionic Character: The greater the difference in electronegativity between two atoms, the more ionic the bond.

• Periodic Trends: Bonds between metals and nonmetals tend to be more ionic.

Section 3: Atomic Radius and Ionization Energy

Periodic Trends in Atomic Radius

Atomic radius is the distance from the nucleus to the outermost electron shell. It varies predictably across the periodic table.

• Trend: Atomic radius decreases across a period (left to right) and increases down a group (top to bottom).

• Example: Cs > Na > Li (Cs has the largest radius).

Ionization Energy

• Definition: Ionization energy is the energy required to remove an electron from a gaseous atom.

• Trend: Increases across a period, decreases down a group.

• Successive Ionization Energies: Removing each additional electron requires more energy.

• Example: $\mathrm{Ca^+ (g) \rightarrow Ca^{2+} (g) + e^-}$ is the second ionization of calcium.

Section 4: Lattice Energy

Lattice Energy Concepts

Lattice energy is the energy released when ions form a crystalline solid. It is a measure of the strength of the ionic bonds in a solid.

• Definition: Lattice energy is the energy required to separate one mole of an ionic solid into gaseous ions.

• Factors Affecting Lattice Energy: Charge and size of ions; higher charge and smaller size increase lattice energy.

• Formula: $U = k \frac{Q_1 Q_2}{r}$ (where $Q_1$ and $Q_2$ are ion charges, $r$ is the distance between ions)

Section 5: Bond Energy Calculations

Calculating Enthalpy Change ($\Delta H_{rxn}$) Using Bond Energies

Bond energies can be used to estimate the enthalpy change of a reaction.

• Formula: $\Delta H_{rxn} = \sum \text{Bond energies of bonds broken} - \sum \text{Bond energies of bonds formed}$

• Example: For $\mathrm{C_2H_4 + H_2 \rightarrow C_2H_6}$, calculate using given bond energies.

Section 6: Lewis Structures

Drawing Lewis Structures and the Octet Rule

Lewis structures represent the arrangement of electrons in molecules. The octet rule states that atoms tend to have eight electrons in their valence shell.

• Steps:

  1. Count total valence electrons.

  2. Arrange atoms and connect with single bonds.

  3. Distribute remaining electrons to complete octets.

• Exceptions: Some atoms (e.g., B, Be, S, P) can have less or more than eight electrons.

• Expanded Octet: Elements in period 3 or higher can have more than eight valence electrons.

Section 7: VSEPR Geometry

Valence Shell Electron Pair Repulsion (VSEPR) Theory

VSEPR theory predicts the shapes of molecules based on electron pair repulsion around a central atom.

• Common Geometries:

  • Linear: 180°

  • Trigonal planar: 120°

  • Tetrahedral: 109.5°

  • Trigonal bipyramidal: 90°, 120°

  • Octahedral: 90°

• Example: $\mathrm{BF_3}$ is trigonal planar; $\mathrm{H_2O}$ is bent.

Section 8: Polarity

Molecular Polarity and Dipole Moments

Molecular polarity depends on bond polarity and molecular geometry. Polar molecules have net dipole moments.

• Determining Polarity: Assess bond dipoles and overall molecular shape.

• Example: $\mathrm{H_2O}$ is polar; $\mathrm{CO_2}$ is nonpolar.

Section 9: Formal Charge

Calculating Formal Charge

Formal charge helps determine the most stable Lewis structure.

• Formula: $\text{Formal Charge} = \text{Valence electrons} - \text{Nonbonding electrons} - \frac{1}{2} \times \text{Bonding electrons}$

• Application: Structures with formal charges closest to zero are preferred.

Section 10: Sigma and Pi Bonds

Bond Types in Molecules

Sigma ($\sigma$) and pi ($\pi$) bonds are types of covalent bonds formed by orbital overlap.

• Sigma Bond: End-to-end overlap of orbitals; all single bonds are sigma bonds.

• Pi Bond: Side-to-side overlap of p orbitals; present in double and triple bonds.

• Example: Ethylene ($\mathrm{C_2H_4}$) has one sigma and one pi bond between the carbons.

Section 11: Hybridization

Hybridization and Molecular Geometry

Hybridization describes the mixing of atomic orbitals to form new hybrid orbitals for bonding.

• Types:

  • sp: Linear geometry, 180°

  • sp2: Trigonal planar, 120°

  • sp3: Tetrahedral, 109.5°

• Example: Carbon in methane ($\mathrm{CH_4}$) is sp3 hybridized.

Cis-Trans Isomerism

• Cis-Trans Isomers: Occur in molecules with restricted rotation, such as double bonds.

• Example: $\mathrm{CHCl=CHCl}$ can have cis and trans forms.

Additional info: These study notes expand upon the original exam review questions, providing definitions, explanations, and formulas for key concepts in chemical bonding, molecular structure, and periodic trends. All topics are relevant to General Chemistry I, covering chapters on chemical bonding, molecular geometry, periodic properties, and related calculations.