Unit 3: Chemical Bonding

Electron Configurations and Valence Electrons

Understanding electron configurations is essential for predicting chemical behavior and bonding. Valence electrons are the outermost electrons and determine an element's reactivity.

• Electron Configuration: The arrangement of electrons in an atom's orbitals, typically written using the Aufbau principle, Hund's rule, and the Pauli exclusion principle.

• Valence Electrons: Electrons in the outermost shell of an atom, important for chemical bonding.

• Example: The electron configuration of sodium (Na, atomic number 11) is 1s2 2s2 2p6 3s1. It has one valence electron.

Lewis Dot Structures

Lewis dot structures visually represent the valence electrons of atoms and show how atoms bond in molecules.

• Lewis Dot Symbol: A symbol showing the element and its valence electrons as dots.

• Octet Rule: Atoms tend to gain, lose, or share electrons to achieve eight valence electrons (a full outer shell).

• Steps to Draw Lewis Structures:

  1. Count total valence electrons.

  2. Arrange atoms and connect with single bonds.

  3. Distribute remaining electrons to satisfy the octet rule.

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

Ions and Ionic Compounds

Ions are charged particles formed when atoms gain or lose electrons. Ionic compounds are formed from the electrostatic attraction between cations and anions.

• Cation: A positively charged ion (e.g., Na+).

• Anion: A negatively charged ion (e.g., Cl-).

• Formula Unit: The simplest ratio of ions in an ionic compound.

• Naming: Name the cation first, then the anion (e.g., NaCl is sodium chloride).

Covalent Compounds and Lewis Structures

Covalent compounds are formed when atoms share electrons. Lewis structures help visualize shared pairs and lone pairs.

• Covalent Bond: A bond formed by sharing electrons between atoms.

• Typical Number of Covalent Bonds: Determined by the number of electrons needed to complete the octet.

• Example: Carbon typically forms four covalent bonds (as in CH4).

Polyatomic Ions

Polyatomic ions are ions composed of two or more atoms covalently bonded, carrying a net charge.

• Common Polyatomic Ions: Examples include sulfate (SO42-), nitrate (NO3-), and ammonium (NH4+).

• Naming: Use standard names (e.g., sodium sulfate for Na2SO4).

Writing and Naming Compounds

Correctly writing formulas and naming compounds is fundamental in chemistry.

• Ionic Compounds: Combine cations and anions in ratios that yield a neutral compound.

• Covalent Compounds: Use prefixes to indicate the number of atoms (e.g., CO2 is carbon dioxide).

• Polyatomic Compounds: Use parentheses if more than one polyatomic ion is needed (e.g., Ca(NO3)2).

Formula and Molecular Weights

The formula weight (for ionic compounds) or molecular weight (for covalent compounds) is the sum of the atomic masses of all atoms in the formula.

• Formula Weight: For ionic compounds, sum the atomic masses of all ions in the empirical formula.

• Molecular Weight: For covalent compounds, sum the atomic masses of all atoms in the molecule.

• Example: The molecular weight of H2O is g/mol.

Unit 4: The Mole and Chemical Reactions

The Mole Concept

The mole is a fundamental unit in chemistry for counting particles (atoms, molecules, ions).

• Definition: 1 mole = particles (Avogadro's number).

• Conversions: Moles to mass, mass to moles, moles to number of particles.

• Example: 2 moles of H2O contains molecules.

Percent Composition and Empirical Formulas

Percent composition shows the mass percentage of each element in a compound. The empirical formula is the simplest whole-number ratio of atoms.

• Percent Composition Formula:

• Empirical Formula: Determined from percent composition by converting masses to moles and finding the simplest ratio.

Balancing Chemical Equations

Balancing equations ensures the law of conservation of mass is obeyed.

• Steps:

  1. Write the unbalanced equation.

  2. Balance atoms one element at a time using coefficients.

  3. Check that all atoms are balanced.

• Example:

Stoichiometry

Stoichiometry involves calculations based on balanced chemical equations to relate quantities of reactants and products.

• Mole Ratios: Derived from coefficients in the balanced equation.

• Example: From , 2 moles of H2 produce 2 moles of H2O.

Types of Chemical Reactions

Chemical reactions can be classified into several types based on the rearrangement of atoms.

• Combination (Synthesis): Two or more substances combine to form one product.

• Decomposition: A single compound breaks down into two or more products.

• Single Replacement: An element replaces another in a compound.

• Double Replacement: Exchange of ions between two compounds.

• Combustion: A substance reacts with oxygen, releasing energy.

Redox Reactions

Redox (reduction-oxidation) reactions involve the transfer of electrons between species.

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Oxidizing Agent: Causes oxidation, is reduced.

• Reducing Agent: Causes reduction, is oxidized.

• Example: In , Na is oxidized, Cl2 is reduced.

Unit 5: Thermodynamics and Intermolecular Forces

VSEPR Theory and Molecular Geometry

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the shapes of molecules based on electron pair repulsion.

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

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

• Example: Methane (CH4) is tetrahedral.

Polarity of Molecules

Molecular polarity depends on the difference in electronegativity and the geometry of the molecule.

• Polar Molecule: Has a net dipole moment due to unequal sharing of electrons.

• Nonpolar Molecule: No net dipole moment; electron distribution is symmetrical.

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

Intermolecular Forces

Intermolecular forces are attractions between molecules, affecting physical properties like boiling and melting points.

• Types:

  • London Dispersion Forces: Weak, present in all molecules.

  • Dipole-Dipole Interactions: Between polar molecules.

  • Hydrogen Bonding: Strong dipole-dipole interaction involving H bonded to N, O, or F.

• Intramolecular vs. Intermolecular: Intramolecular forces are within molecules (covalent, ionic bonds); intermolecular are between molecules.

• Example: Water has hydrogen bonding, leading to high boiling point.