Study Guide: Atomic Properties, Chemical Bonding, and Molecular Structure (Chapters 3–6)

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 in the gas phase.

• Definition: Ionization energy is the energy needed to remove one mole of electrons from one mole of gaseous atoms or ions.

• Sequential Ionization Energies: The first ionization energy removes the first electron; the second removes the next, and so on. Each successive ionization energy is higher due to increased effective nuclear charge.

• Trends: Ionization energy generally increases across a period and decreases down a group.

• Identifying Cation Charges: Large jumps in successive ionization energies indicate the removal of electrons from a new shell, helping to identify the typical charge of cations.

• Example: The first ionization energy of sodium is much lower than the second, indicating sodium forms a +1 cation.

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; electrons are transferred from one atom to another.

• Covalent Bonds: Formed between nonmetals; electrons are shared between atoms.

• Comparison: Ionic compounds are usually crystalline solids with high melting points; covalent compounds can be gases, liquids, or solids with lower melting points.

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

Chemical Representations

Different representations help visualize and interpret chemical 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: 3D representations of molecules.

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

Energy of Ionic Bonds (Coulomb's Law)

Ionic bonds lower the energy of the system by electrostatic attraction between oppositely charged ions.

• Coulomb's Law: The energy of interaction between two ions is given by:

$E = \frac{k \cdot q_1 \cdot q_2}{r}$

• Where k is a constant, q1 and q2 are charges, and r is the distance between ions.

• Ion Pairs: The formation of ion pairs stabilizes the compound.

Lewis Symbols and the Octet Rule

Lewis symbols represent valence electrons and help predict bonding in compounds.

• 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 the formula of ionic compounds and the number of bonds in covalent compounds.

• Example: Na (1 valence electron) loses one to Cl (7 valence electrons) to form NaCl.

Bond Formation and Energy

Bond formation is an exothermic process because it releases energy as atoms achieve more stable configurations.

• Covalent Bond: Sharing electrons lowers the energy of the system.

• Exothermic Process: Energy is released when bonds form.

Ionic Charges and Nomenclature

Understanding ionic charges is essential for writing formulas and naming compounds.

• Group Charges: Elements in groups 1, 2, 13, 16, and 17 have predictable charges.

• Polyatomic Ions: Common ions include ammonium (NH4+), hydroxide (OH–), carbonate (CO32–), hydrogen carbonate (HCO3–), nitrate (NO3–), sulfate (SO42–), phosphate (PO43–).

• Nomenclature: Use Roman numerals to indicate charge for transition metals (e.g., Fe(III) chloride).

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

Bonding Definitions

• Lone-Pair Electrons: Electrons not involved in bonding.

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

• 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 convert between grams 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 number of atoms, determined from empirical formula and molar mass.

Chapter 5: Chemical Bonding I

Electronegativity and Bond Polarity

Electronegativity is the ability of an atom to attract electrons in a bond, leading to bond polarity.

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

• Polar Covalent Bonds: Unequal sharing of electrons due to differences in electronegativity.

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

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

Electron Distribution and Electrostatic Potential Maps

Electrostatic potential maps visually represent 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 (e.g., NO, BF3) do not fully obey it.

• 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: Formal charge = (valence electrons) – (nonbonding electrons) – (1/2 bonding electrons).

$\text{Formal charge} = \text{valence electrons} - \text{nonbonding electrons} - \frac{1}{2} \text{bonding electrons}$

• 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 based on electron pair repulsion.

• 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.

• Example: Water (H2O) is bent due to two lone pairs on oxygen.

Polarity of Molecules

Polarity depends on both bond polarities and molecular shape.

• Lewis Structure + VSEPR: Used to predict overall molecular polarity.

• 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:

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

Energy Carrier

An energy carrier is a molecule or ion that transports energy within a system (e.g., ATP in biology).

• Example: ATP carries energy in biological systems.

Additional info: Energy carriers are not typically covered in general chemistry textbooks but are important in biochemistry and physical chemistry.

Chapter 6: Chemical Bonding II

Wave Interference and Bonding

Constructive and destructive interference of waves explains how atomic orbitals combine.

• Constructive Interference: Waves add, increasing electron density between nuclei (bonding).

• Destructive Interference: Waves subtract, creating nodes (antibonding).

Valence Bond Theory and Hybrid Orbitals

Valence bond theory describes how atomic orbitals overlap to form bonds; hybridization explains molecular geometry.

• Hybrid Orbitals: s and p orbitals mix to form sp, sp2, and sp3 hybrids.

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

• Determining Hybridization: Based on the number of electron domains around an atom.

• Example: Carbon in methane (CH4) is sp3 hybridized.

Sigma and Pi Bonds

Sigma (σ) and pi (π) bonds are types of covalent bonds formed by orbital overlap.

• Sigma Bond: Head-on overlap; all single bonds are sigma bonds.

• Pi Bond: Side-on overlap; present in double and triple bonds.

• Identification: Sigma bonds are along the internuclear axis; pi bonds are above and below.

Cis- and Trans- Isomers

Double bonds restrict rotation, leading to cis (same side) and trans (opposite side) isomers.

• Definition: Cis isomers have substituents on the same side; trans isomers on opposite sides.

• Role in Vision: Isomerization of retinal (cis-trans) is key in vision chemistry.

Molecular Orbital Theory

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

• Bonding and Antibonding Orbitals: Bonding orbitals increase electron density between nuclei; antibonding orbitals have a node.

• Energy Diagram: When two s-orbitals interact, two MOs form: one bonding (lower energy), one antibonding (higher energy).

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

Molecule	Bonding MO Electrons	Antibonding MO Electrons	Stable?
H2	2	0	Yes
He2	2	2	No

Constants and Periodic Table

Fundamental Constants

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

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

• Speed of Light: $c = 2.998 \times 10^{8} \text{ m} \cdot \text{s}^{-1}$

Additional info: The periodic table is included for reference during the test.