Exam 2 Study Guide: Chemical Bonding, Compounds, and Chemical Reactions

Chemical Bonding

Types of Chemical Bonds

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

• Ionic Bonds: Formed when electrons are transferred from one atom to another, typically between metals and nonmetals. The resulting ions are held together by electrostatic attraction.

• Covalent Bonds: Formed when two atoms share one or more pairs of electrons, usually between nonmetals.

• Predicting Bond Type: Use the periodic table: metals + nonmetals → ionic; nonmetals + nonmetals → covalent.

• Similarities and Differences: Both types involve the attraction between atoms, but ionic bonds involve electron transfer and covalent bonds involve electron sharing.

Example: NaCl (sodium chloride) is ionic; H2O (water) is covalent.

Ionic Compounds

Structure, Properties, and Nomenclature

Ionic compounds consist of positive and negative ions arranged in a crystal lattice. Their properties and naming conventions are distinct from molecular compounds.

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

• Predicting Ion Charges: Use group number on the periodic table (e.g., Group 1 forms +1 ions, Group 17 forms -1 ions).

• Writing Formulas: Combine ions in ratios that result in a neutral compound (e.g., Mg2+ and Cl- form MgCl2).

• Forces in Ionic Compounds: Held together by strong electrostatic (Coulombic) forces.

• Lattice Energy: The energy required to separate one mole of an ionic solid into gaseous ions. Higher lattice energy means stronger ionic bonds.

• Naming: Name the cation first, then the anion. For transition metals, indicate charge with Roman numerals (e.g., FeCl2: iron(II) chloride).

• Polyatomic Ions: Ions composed of more than one atom (e.g., SO42-, NO3-).

Example: Na2SO4 is sodium sulfate; Ca(NO3)2 is calcium nitrate.

Molecular Compounds

Structure, Properties, and Nomenclature

Molecular compounds are formed by covalent bonding between nonmetals. Their formulas and names follow specific rules.

• Formulas: Can be structural, molecular, or empirical.

• Diatomic Molecules: Elements that exist as two-atom molecules (e.g., H2, O2, N2).

• Naming: Use prefixes to indicate the number of each atom (e.g., CO2: carbon dioxide).

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

• Conversions: Use Avogadro's number ($6.022 \times 10^{23}$) to convert between moles and number of molecules or atoms.

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

• Empirical vs. Molecular Formula: Empirical is the simplest ratio; molecular is the actual number of atoms.

Example: C6H12O6 (glucose) has an empirical formula of CH2O.

Lewis Structures and Bonding Theories

Lewis Dot Structures

Lewis structures represent valence electrons as dots and show how atoms are bonded in a molecule or ion.

• Electronegativity: The ability of an atom to attract electrons in a bond. Increases across a period, decreases down a group.

• Polar vs. Nonpolar Bonds: Determined by the difference in electronegativity between atoms.

• Drawing Lewis Structures: Show all valence electrons, including lone pairs and bonding pairs.

• Resonance: Some molecules have multiple valid Lewis structures; the actual structure is a hybrid.

• Octet Rule Exceptions: Some atoms can have fewer or more than eight electrons (e.g., BF3, SF6).

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

• Bond Order: $\text{Bond order} = \frac{\text{number of bonding pairs}}{\text{number of bonds}}$

Example: O3 (ozone) has resonance structures with delocalized electrons.

VSEPR Theory and Molecular Geometry

Predicting Shapes and Polarity

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the 3D shape of molecules based on electron group repulsion.

• Electron Groups: Regions of electron density (bonds or lone pairs) around a central atom.

• Common Geometries: Linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.

• Bond Angles: Determined by geometry; lone pairs reduce bond angles.

• Axial vs. Equatorial: In trigonal bipyramidal, axial positions are 180° apart; equatorial are 120° apart.

• Polarity: Determined by shape and bond polarities; asymmetrical shapes with polar bonds are usually polar molecules.

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

Valence Bond Theory and Hybridization

Describing Bonding in Molecules

Valence Bond (VB) theory explains bonding as the overlap of atomic orbitals. Hybridization describes the mixing of atomic orbitals to form new, equivalent orbitals.

• Orbital Overlap: Bonds form when atomic orbitals overlap, creating a region of increased electron density.

• Hybridization: Mixing of s, p, and sometimes d orbitals to form hybrid orbitals (e.g., sp3, sp2, sp).

• σ (Sigma) and π (Pi) Bonds: σ bonds are formed by head-on overlap; π bonds by side-on overlap.

• Geometry and Hybridization: The number of electron groups determines hybridization (e.g., 4 groups = sp3).

Example: In methane (CH4), carbon is sp3 hybridized.

Molecular Orbital Theory

Bonding in Diatomic Molecules

Molecular Orbital (MO) theory describes bonding by combining atomic orbitals to form molecular orbitals that are delocalized over the entire molecule.

• Molecular Orbitals: Can be bonding (lower energy) or antibonding (higher energy).

• σ, σ*, π, π* Orbitals: σ and π are bonding; σ* and π* are antibonding.

• Bond Order: $\text{Bond order} = \frac{\text{number of bonding electrons} - \text{number of antibonding electrons}}{2}$

• Magnetism: Molecules with unpaired electrons are paramagnetic; all electrons paired means diamagnetic.

• Comparison to VB Theory: MO theory explains properties like magnetism and why some molecules (e.g., He2) do not exist.

Example: O2 is paramagnetic due to two unpaired electrons in π* orbitals.

Chemical Reactions and Stoichiometry

Balancing and Interpreting Chemical Equations

Chemical equations represent the reactants and products in a chemical reaction. Balancing equations ensures the law of conservation of mass is obeyed.

• Physical vs. Chemical Changes: Physical changes do not alter chemical identity; chemical changes do.

• Components of Equations: Reactants, products, phases (s, l, g, aq), and coefficients.

• Balancing Methods: Use inspection and inventory to ensure equal numbers of each atom on both sides.

• Stoichiometry: The calculation of reactants and products in chemical reactions using mole ratios.

• Conversions: Use molar mass and Avogadro's number for mass-mole-particle conversions.

Example: $2H_2 + O_2 \rightarrow 2H_2O$

Limiting Reactant, Theoretical Yield, and Percent Yield

In reactions, the limiting reactant determines the maximum amount of product formed. Theoretical yield is the calculated maximum, while actual yield is what is obtained experimentally.

• Limiting Reactant: The reactant that is completely consumed first, limiting the amount of product.

• Theoretical Yield: The maximum amount of product possible, calculated from stoichiometry.

• Percent Yield: $\text{Percent yield} = \frac{\text{actual yield}}{\text{theoretical yield}} \times 100\%$

• Excess Reactant: The reactant left over after the reaction is complete.

Example: If 10 g of A reacts with 10 g of B, and A is limiting, use A to calculate the maximum product.

Common Polyatomic Ions

Names and Formulas

Polyatomic ions are ions composed of more than one atom. Memorizing their names, formulas, and charges is essential for writing and naming compounds.

Name	Formula	Charge
Ammonium	NH4+	+1
Nitrate	NO3-	-1
Sulfate	SO42-	-2
Phosphate	PO43-	-3
Hydroxide	OH-	-1
Carbonate	CO32-	-2
Acetate	C2H3O2-	-1
Chlorate	ClO3-	-1

Additional info: This is a partial list; refer to your textbook for a complete table.

Unit Prefixes and Conversions

Metric Prefixes

Metric prefixes are used to express units in different scales. Memorize common prefixes for conversions.

Prefix	Symbol	Factor
Kilo	k	103
Centi	c	10-2
Milli	m	10-3
Micro	μ	10-6
Nano	n	10-9

Example: 1 mL = 0.001 L

Summary Table: Key Concepts for Exam 2

Topic	Key Points
Chemical Bonding	Ionic vs. covalent, predicting bond type, properties
Ionic Compounds	Octet rule, charges, formulas, lattice energy, naming
Molecular Compounds	Formulas, naming, molar mass, percent composition
Lewis Structures	Drawing, resonance, formal charge, bond order
VSEPR Theory	Geometry, bond angles, polarity
Valence Bond Theory	Orbital overlap, hybridization, σ and π bonds
Molecular Orbital Theory	MO diagrams, bond order, magnetism
Chemical Reactions	Balancing, stoichiometry, limiting reactant, yield
Polyatomic Ions	Names, formulas, charges
Unit Prefixes	Conversions between metric units

Additional info: For exam preparation, focus on understanding concepts, practicing problems, and memorizing key tables and values as indicated.