Study Guide: Atomic Properties, Chemical Bonding, and Molecular Structure

Chapter 3: Periodic Properties of the Elements

Ionization Energy

Ionization energy is a fundamental property of atoms that describes the energy required to remove an electron from an atom or ion in the gas phase.

• Definition: The first ionization energy is the energy needed to remove the first electron from a neutral atom.

• Sequential Ionization Energies: Successive ionization energies refer to the energies required to remove additional electrons after the first. Each subsequent ionization energy is higher due to increased effective nuclear charge.

• Trends: Ionization energies generally increase across a period and decrease down a group.

• Identifying Cation Charges: Large jumps in successive ionization energies help identify the most stable cation charge for an element.

• Example: For magnesium, the first and second ionization energies are relatively low, but the third is much higher, indicating Mg2+ is the stable ion.

Chapter 4: Molecules and Compounds

Types of Chemical Bonds

Chemical bonds are the forces that hold atoms together in compounds. The main types are ionic and covalent bonds.

• Ionic Bonds: Formed between metals and nonmetals; involve transfer of electrons.

• Covalent Bonds: Formed between nonmetals; involve sharing of electrons.

• Comparison: Ionic bonds result in charged ions, while covalent bonds result in neutral molecules.

• Example: NaCl (ionic), H2O (covalent).

Chemical Representations

Various representations are used to describe molecules and compounds.

• Chemical Formula: Shows the types and numbers of atoms (e.g., H2O).

• Structural Formula: Shows how atoms are connected (e.g., O-H-H).

• Ball & Stick / Space-Filling Models: Visualize 3D structure.

• Line Structures: Used for organic molecules; lines represent bonds, vertices represent carbon atoms.

Energy of Bond Formation

Bond formation lowers the energy of a system.

• Coulomb's Law: Describes the energy of interaction between charged particles.

• Equation: $E = \frac{kQ_1Q_2}{r}$

• Ionic Bond: Energy is lowered when oppositely charged ions form a pair.

• Covalent Bond: Energy is lowered by sharing electrons, stabilizing the atoms.

• Bond Formation: Always exothermic; energy is released.

Lewis Symbols and the Octet Rule

Lewis symbols represent valence electrons and help predict bonding.

• Lewis Symbols: Dots around element symbols represent valence electrons.

• Octet Rule: Atoms tend to gain, lose, or share electrons to achieve eight valence electrons.

• Application: Helps determine formulas of ionic compounds and number of bonds in covalent compounds.

Nomenclature and Ionic Charges

Naming compounds requires knowledge of ionic charges and conventions.

• Common Ionic Charges: Group 1 (+1), Group 2 (+2), Group 13 (+3), Group 16 (–2), Group 17 (–1).

• Polyatomic Ions: Ammonium (NH4+), Hydroxide (OH–), Carbonate (CO32–), Hydrogen Carbonate (HCO3–), Nitrate (NO3–), Sulfate (SO42–), Phosphate (PO43–).

• Roman Numerals: Used to indicate charge of transition metal cations (e.g., Fe(III)).

• Writing Formulas: Combine ions to achieve charge neutrality.

• Naming: Use systematic rules for ionic and covalent compounds.

Bonding Definitions

• Lone-Pair Electrons: Non-bonding pairs of electrons on an atom.

• Single, Double, Triple Bonds: One, two, or three pairs of shared electrons.

• Valence: Number of bonds an element typically forms (C: 4, N: 3, O: 2, F: 1).

Molar Mass and Percent Composition

Molar mass is used to relate mass and moles, and to determine percent composition.

• Molar Mass: Sum of atomic masses in a compound (g/mol).

• Conversion: $\text{moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

• Percent Composition: $\%\text{element} = \frac{\text{mass of element in 1 mol}}{\text{molar mass}} \times 100$

• Empirical Formula: Simplest ratio of elements, determined from mass percent.

• Molecular Formula: Actual formula, determined from empirical formula and molar mass.

Chapter 5: Chemical Bonding I

Electronegativity and Bond Polarity

Electronegativity is the tendency of an atom to attract electrons in a bond.

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

• Polar Covalent Bonds: Occur when atoms have different electronegativities.

• Polarity Designation: The more electronegative atom is marked with δ–, the less with δ+.

• Example: In H–Cl, Cl is δ–, H is δ+.

Electron Distribution and Electrostatic Potential Maps

Electrostatic potential maps visualize electron density and polarity in molecules.

• Color Coding: Red indicates regions of high electron density (negative), blue indicates low (positive).

Lewis Structures and Resonance

Lewis structures show bonding and lone pairs; resonance describes delocalized electrons.

• Drawing Lewis Structures: Follow the octet rule, but some molecules have exceptions.

• Resonance: Multiple valid Lewis structures; electrons are delocalized.

• Example: Ozone (O3) absorbs UV-B due to resonance, unlike O2.

Formal Charge

Formal charge helps identify the most stable Lewis structure.

• Definition: $\text{Formal charge} = \text{valence electrons} - (\text{nonbonding electrons} + \frac{1}{2}\text{bonding electrons})$

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

VSEPR Theory and Molecular Shape

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes.

• Repelling: Electron pairs (bonding and lone pairs) repel each other.

• Shape Measurement: Only nuclei are considered, not lone pairs.

• Bond Angles: Lone pairs occupy more space, reducing bond angles.

• Identifying Shapes: Use Lewis structure and VSEPR to determine geometry (e.g., linear, bent, trigonal planar, tetrahedral).

Molecular Polarity

Polarity depends on shape and bond polarities.

• Determination: Use Lewis structure, VSEPR, and bond polarities.

• Example: CO2 is nonpolar (linear), H2O is polar (bent).

Bond Energy and Bond Length

Bond energy is the energy required to break a bond; bond length is the distance between nuclei.

• Single, Double, Triple Bonds: Triple bonds are shortest and strongest; single bonds are longest and weakest.

• Estimating Reaction Energy: Use tabulated bond energies to calculate energy change.

• Equation: $\Delta H = \sum \text{(bond energies of bonds broken)} - \sum \text{(bond energies of bonds formed)}$

Energy Carrier

An energy carrier is a molecule or ion that transports energy within a system (from Lecture 20).

Chapter 6: Chemical Bonding II

Wave Interference and Bond Formation

Constructive and destructive interference of electron waves is fundamental to chemical bonding.

• Constructive Interference: Builds up electron density between nuclei, stabilizing the bond.

• Destructive Interference: Creates nodes, destabilizing the bond.

Valence Bond Theory and Hybrid Orbitals

Valence bond theory explains how atomic orbitals overlap to form bonds; hybridization addresses shape challenges.

• Hybrid Orbitals: Formed by mixing s and p orbitals.

• Types: sp (linear), sp2 (trigonal planar), sp3 (tetrahedral).

• Determining Hybridization: Use Lewis structure to identify hybrid orbitals, especially for carbon.

• Unused p-Orbitals: Remain for pi bonding.

Sigma and Pi Bonds

Sigma (σ) and pi (π) bonds differ in their formation and properties.

• Sigma Bonds: Formed by head-on overlap; all single bonds are sigma bonds.

• Pi Bonds: Formed by side-on overlap of p orbitals; present in double and triple bonds.

• Identification: Sigma bonds are cylindrical; pi bonds have electron density above and below the bond axis.

Cis- and Trans- Isomers

Double bonds can lead to geometric isomers.

• Cis-Isomer: Substituents on the same side of the double bond.

• Trans-Isomer: Substituents on opposite sides.

• Role in Vision: Cis-trans isomerization is key in the chemistry of vision.

Molecular Orbital Theory

Molecular orbital (MO) theory describes how atomic orbitals combine to form molecular orbitals.

• Overlap: Two atomic orbitals form two molecular orbitals: one bonding, one antibonding.

• Bonding Orbital: Builds electron density between nuclei.

• Antibonding Orbital: Has a node between nuclei.

• Energy Diagram: Shows relative energies of atomic and molecular orbitals.

• Stability: H2 is stable (bonding orbital filled), He2 is not (bonding and antibonding orbitals filled).

Constants and Periodic Table

Fundamental Constants

• Avogadro's Number: $N_a = 6.022 \times 10^{23}$

• Speed of Light: $c = 2.998 \times 10^{8} \text{ m/s}$

• Planck's Constant: $h = 6.63 \times 10^{-34} \text{ J·s}$

Periodic Table

The periodic table organizes elements by atomic number and groups with similar properties.

• Groups: Vertical columns; elements in a group have similar chemical properties.

• Periods: Horizontal rows; properties change progressively across a period.

• Key Groups: Alkali metals (Group 1), alkaline earth metals (Group 2), halogens (Group 17), noble gases (Group 18).

Common Polyatomic Ions (Table 4.4)

These ions are frequently encountered in general chemistry.

Ion Name	Formula	Charge
Ammonium	NH4	+1
Hydroxide	OH	-1
Carbonate	CO3	-2
Hydrogen Carbonate	HCO3	-1
Nitrate	NO3	-1
Sulfate	SO4	-2
Phosphate	PO4	-3

Additional info: Line structures, energy carrier, and regions of attraction/repulsion graphs are included based on lecture content not found in the textbook.