General Chemistry Exam 4 Review: Bonding, Structure, and Periodic Properties

Section 1: Bond Length and Bond Strength

Bond Length

Bond length refers to the distance between the nuclei of two bonded atoms. It is influenced by the type of bond (single, double, triple), atomic size, and bond order.

• Single bonds are generally longer than double or triple bonds.

• Bond order increases as the number of shared electron pairs increases, resulting in shorter bonds.

• Example: $\text{N} \equiv \text{N}$ (triple bond) < $\text{N} = \text{N}$ (double bond) < $\text{N} - \text{N}$ (single bond) in terms of bond length.

Bond Strength

Bond strength is the energy required to break a bond between two atoms. Stronger bonds are shorter and have higher bond order.

• Triple bonds are the strongest, followed by double, then single bonds.

• Example: $\text{Bond strength:} \ \text{Triple} > \text{Double} > \text{Single}$

Ordering Bonds

• To order bonds by length or strength, consider bond order and atomic radii.

• Example: $\text{NO} < \text{NO}_2 < \text{NO}_3$ (increasing bond length)

Section 2: Electronegativity

Definition and Trends

Electronegativity is the ability of an atom to attract electrons in a chemical bond. It is a key factor in determining bond polarity and molecular properties.

• Electronegativity increases across a period (left to right) and decreases down a group (top to bottom).

• Example: $\text{F} > \text{O} > \text{N} > \text{C}$

Bond Character and Polarity

• Bonds between atoms with large electronegativity differences are more ionic.

• Polar covalent bonds have unequal sharing of electrons, resulting in partial charges.

• Nonpolar covalent bonds have equal sharing of electrons.

Periodic Trends

• Order bonds by ionic character or elements by electronegativity using periodic trends.

• Example: $\text{Si-F} > \text{Si-Cl} > \text{Si-Br} > \text{Si-I}$ (increasing ionic character)

Section 3: Atomic Radius and Ionization Energy

Atomic Radius

The atomic radius is the distance from the nucleus to the outermost electron shell. It varies with position on the periodic table.

• Atomic radius increases down a group and decreases across a period.

• Example: $\text{K} > \text{Na} > \text{Li}$

Ionization Energy

Ionization energy is the energy required to remove an electron from a gaseous atom or ion.

• First ionization energy decreases down a group and increases across a period.

• Successive ionization energies increase for the same atom.

• Equation: $\text{M}(g) \rightarrow \text{M}^+(g) + e^-$

• Second ionization energy: $\text{M}^+(g) \rightarrow \text{M}^{2+}(g) + e^-$

Section 4: Lattice Energy

Definition and Trends

Lattice energy is the energy released when gaseous ions form an ionic solid. It is a measure of the strength of the ionic bonds in a crystal lattice.

• Lattice energy increases with higher ionic charges and smaller ionic radii.

• Equation: $\Delta H_{\text{lattice}} = \text{M}^+(g) + \text{X}^-(g) \rightarrow \text{MX}(s)$

• Example: $\text{Mg}^{2+} + \text{Se}^{2-} \rightarrow \text{MgSe}(s)$

Section 5: Bond Energy Calculations

Bond Energies and Enthalpy Change

Bond energy is the energy required to break a specific bond in one mole of gaseous molecules. The enthalpy change of a reaction can be estimated using bond energies.

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

• Use provided bond energies to estimate $\Delta H_{\text{rxn}}$ for reactions.

Section 6: Lewis Structures

Drawing Lewis Structures

Lewis structures represent the arrangement of valence electrons in molecules. They help predict molecular shape, bonding, and reactivity.

• Follow the octet rule: atoms tend to have eight electrons in their valence shell.

• Some atoms (e.g., P, S, Cl) can have expanded octets.

• Count valence electrons to construct the structure.

• Example: XeCl2 has 22 valence electrons (Xe: 8, Cl: 7 each).

Bonding in Polyatomic Ions

• CO32- has resonance structures with delocalized electrons.

• Bonding can involve single and double bonds, or resonance hybrids.

Section 7: VSEPR Geometry

Electron and Molecular Geometry

The VSEPR (Valence Shell Electron Pair Repulsion) model predicts molecular shapes based on electron pair repulsion.

• Electron geometry considers all electron groups (bonding and lone pairs).

• Molecular geometry considers only the arrangement of atoms.

• Example: H2O has tetrahedral electron geometry, bent molecular geometry.

• Bond angles for tetrahedral: $109.5^\circ$

Section 8: Polarity

Molecular Polarity

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

• Symmetrical molecules may be nonpolar even if bonds are polar.

• Example: CO2 is linear and nonpolar; H2O is bent and polar.

Section 9: Formal Charge

Calculating Formal Charge

Formal charge helps determine the most stable Lewis structure.

• Equation: $\text{Formal charge} = \text{Valence electrons} - \text{Nonbonding electrons} - \frac{1}{2}(\text{Bonding electrons})$

• Structures with formal charges closest to zero are most stable.

Section 10: Sigma and Pi Bonds

Bond Types

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

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

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

• Example: Ethylene (C2H4): 5 sigma, 1 pi bond.

Section 11: Hybridization

Hybrid Orbitals

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

• sp: Linear geometry, 180° bond angle.

• sp2: Trigonal planar, 120° bond angle.

• sp3: Tetrahedral, 109.5° bond angle.

• Electron groups (bonds and lone pairs) determine hybridization.

Cis-Trans Isomerism

• Occurs in molecules with double bonds and different substituents.

• Example: CHCl=CHCl has cis and trans isomers.

Bond Type	Bond Order	Bond Length	Bond Strength
Single	1	Longest	Weakest
Double	2	Intermediate	Intermediate
Triple	3	Shortest	Strongest

Hybridization	Electron Groups	Geometry	Bond Angle
sp	2	Linear	180°
sp2	3	Trigonal planar	120°
sp3	4	Tetrahedral	109.5°

Additional info: These study notes expand on the practice questions by providing definitions, examples, and key equations relevant to the topics covered in General Chemistry chapters on chemical bonding, molecular structure, and periodic properties.