Chemical Bonding and Molecular Structure

Lewis Structures and Covalent Bonding

Lewis structures are diagrams that represent the arrangement of valence electrons among atoms in a molecule. They are essential for visualizing covalent bonding and predicting molecular properties.

• Lewis Structure: A representation showing how valence electrons are distributed among atoms in a molecule.

• Covalent Bond: A chemical bond formed when two atoms share one or more pairs of electrons.

• Octet Rule: Atoms tend to form bonds to achieve eight electrons in their valence shell (except for hydrogen and helium, which follow the duet rule).

• Duet Rule: Hydrogen and helium are stable with two valence electrons.

Example: The Lewis structure for water (H2O) shows two single bonds between oxygen and hydrogen, with two lone pairs on oxygen.

Exceptions to the Octet Rule

Some elements do not follow the octet rule due to their electron configurations or the presence of expanded valence shells.

• Incomplete Octet: Elements like boron (B) and beryllium (Be) may have fewer than eight electrons.

• Expanded Octet: Elements in period 3 or higher (e.g., phosphorus, sulfur) can have more than eight valence electrons.

• Odd-Electron Molecules: Molecules with an odd number of electrons (e.g., NO) cannot have all atoms with complete octets.

Resonance Structures

Resonance structures are different Lewis structures for the same molecule that show the possible arrangements of electrons.

• Resonance: Occurs when more than one valid Lewis structure can be drawn for a molecule.

• Electron Delocalization: Resonance structures indicate that electrons are delocalized across multiple atoms.

Example: The nitrate ion (NO3-) has three resonance structures, each with a different N–O double bond.

Naming Covalent Compounds

Covalent compounds are named using prefixes to indicate the number of each type of atom present.

• Prefixes: mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-

• Naming Rule: The first element keeps its name; the second element ends with “-ide.” Prefixes are used for both elements (except mono- is usually omitted for the first element).

Example: CO2 is named carbon dioxide; N2O4 is dinitrogen tetroxide.

Electronegativity and Bond Polarity

Electronegativity is the ability of an atom to attract shared electrons in a bond. Differences in electronegativity determine bond polarity.

• Polar Covalent Bond: A bond where electrons are shared unequally due to differences in electronegativity.

• Nonpolar Covalent Bond: A bond where electrons are shared equally.

• Bond Polarity: Determined by the difference in electronegativity between bonded atoms.

Example: In HCl, chlorine is more electronegative than hydrogen, making the bond polar.

Predicting Molecular Geometry and Bond Angles

The shape of a molecule is determined by the number of electron pairs (bonding and lone pairs) around the central atom, using the Valence Shell Electron Pair Repulsion (VSEPR) theory.

• Electron Geometry: The arrangement of all electron pairs (bonding and lone pairs) around the central atom.

• Molecular Geometry: The arrangement of only the atoms (ignoring lone pairs).

• Bond Angle: The angle between adjacent bonds in a molecule.

Example: Methane (CH4) has a tetrahedral geometry with bond angles of 109.5°.

Recognizing Chemical Formulas

Chemical formulas can represent atomic elements, molecular elements, molecular compounds, and ionic compounds.

• Atomic Element: Exists as single atoms (e.g., Na, Fe).

• Molecular Element: Exists as molecules with two or more atoms of the same element (e.g., O2, N2).

• Molecular Compound: Composed of molecules formed by covalent bonds (e.g., H2O, CO2).

• Ionic Compound: Composed of cations and anions held together by ionic bonds (e.g., NaCl, CaF2).

Chemical Reactions and Stoichiometry

Chemical Changes and Equations

Chemical reactions involve the rearrangement of atoms, breaking old bonds and forming new ones. Chemical equations represent these changes symbolically.

• Reactants: Substances present before the reaction.

• Products: Substances formed as a result of the reaction.

• Chemical Equation: Uses formulas, coefficients, and symbols to represent a reaction.

• Physical States: Indicated by (s), (l), (g), (aq).

• Arrow (→): Separates reactants from products.

Example:

Balancing Chemical Equations

Balancing ensures the conservation of atoms in a chemical reaction.

• Coefficients: Numbers placed before formulas to indicate the number of units.

• Subscripts: Indicate the number of atoms in a molecule; never change subscripts to balance equations.

Steps to Balance:

1. Write the unbalanced equation.

2. Count atoms of each element on both sides.

3. Add coefficients to balance atoms.

4. Check your work.

Precipitation Reactions

Precipitation reactions occur when two aqueous solutions combine to form an insoluble solid (precipitate).

• Precipitate: An insoluble solid formed in a chemical reaction.

• Predicting Products: Use solubility rules to determine if a precipitate will form.

Example:

Types of Equations for Precipitation Reactions

• Molecular Equation: Shows all reactants and products as compounds.

• Total Ionic Equation: Shows all strong electrolytes as ions.

• Net Ionic Equation: Shows only the ions and molecules directly involved in the reaction.

• Spectator Ions: Ions that do not participate in the actual chemical change.

Example:

• Molecular:

• Total Ionic:

• Net Ionic:

The Mole Concept and Avogadro’s Number

The mole is a counting unit for atoms, molecules, or ions. Avogadro’s number defines the number of particles in one mole.

• Mole (mol): particles.

• Avogadro’s Number: particles/mol.

Formula Mass, Molar Mass, and Conversions

Formula mass is the sum of atomic masses in a compound’s formula. Molar mass is the mass of one mole of a substance.

• Formula Mass: Sum of atomic masses (in amu) for all atoms in a formula.

• Molar Mass: Mass (in grams) of one mole of a substance.

Example: Molar mass of H2O = 2(1.01) + 16.00 = 18.02 g/mol

Conversions Involving Moles, Mass, and Number of Particles

• Mass to Moles:

• Moles to Mass:

• Moles to Number of Particles:

Stoichiometry and Chemical Equations

Stoichiometry uses balanced equations to relate amounts of reactants and products.

• Mole Ratio: Ratio of coefficients in a balanced equation.

• Mass Relationships: Use mole ratios and molar masses to convert between masses of reactants and products.

Percent Yield and Limiting Reactant

• Theoretical Yield: Maximum amount of product possible from given reactants.

• Actual Yield: Amount of product actually obtained.

• Percent Yield:

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

Oxidation-Reduction (Redox) Reactions

Redox reactions involve the transfer of electrons between substances.

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Oxidizing Agent: Substance that gains electrons (is reduced).

• Reducing Agent: Substance that loses electrons (is oxidized).

Example: (Na is oxidized, Cl is reduced)

Additional info: This study guide covers key concepts from chemical bonding, molecular geometry, and chemical reactions, including stoichiometry, precipitation reactions, and redox processes, as outlined in the provided notes.