Principles of Chemistry I: Ionic and Covalent Bonding

Ionic Bonding

Formation of cations and anions through electron transfer.

Cation

Positively charged ion formed by losing electrons.

Anion

Negatively charged ion formed by gaining electrons.

Ionic Compound

Compound formed from ionic bonds between ions.

Example of NaCl

Sodium gives one electron to chlorine.

Example of CaCl2

Calcium gives two electrons to two chlorines.

Properties of Ionic Compounds

High melting points, solid, conductive when molten.

Electronegativity

Atom's tendency to attract electrons in bonds.

Covalent Bonding

Atoms share electrons to form molecular compounds.

Pure Covalent Bond

Equal sharing of electrons between identical atoms.

Polar Covalent Bond

Unequal sharing of electrons between different atoms.

Electron Affinity (EA)

Energy released when an atom gains an electron.

Water (H₂O)

Essential molecule with two hydrogen and one oxygen.

Methane (CH₄)

Natural gas component with one carbon and four hydrogens.

Glucose (C₆H₁₂O₆)

Simple sugar vital for cellular energy.

Sulfur Dioxide (SO₂)

Gas with one sulfur and two oxygen atoms.

Lewis Symbols

Representation of valence electrons around an atom.

Lewis Structures

Diagrams showing bonding and lone pairs in molecules.

Lone Pairs

Non-bonding pairs of electrons in Lewis structures.

Bonding Pairs

Shared electron pairs represented by dashes.

Balancing Charges

Ensuring overall ionic compound neutrality.

Monatomic Ion

Ion formed from a single atom.

Polyatomic Ion

Ion composed of multiple atoms.

Conductivity of Ionic Compounds

Conductive when molten, nonconductive in solid state.

Electronegativity Difference

Increased difference leads to more ionic character.

Practice Problems

Exercises to predict bond types and electronegativity.

Octet Rule

Atoms form bonds to achieve eight valence electrons.

Valence Electrons

Electrons in the outermost shell of an atom.

Double Bond

Two pairs of electrons shared between atoms.

Triple Bond

Three pairs of electrons shared between atoms.

Lewis Structure

Diagram showing electron arrangement in molecules.

Skeleton Structure

Initial arrangement of atoms in a molecule.

Electron Deficient Molecules

Central atom has fewer electrons than needed.

Hypervalent Molecules

Atoms with more than eight valence electrons.

Odd-Electron Molecules

Molecules with an unpaired electron.

Free Radicals

Molecules containing an odd number of electrons.

Formal Charge

Hypothetical charge from electron redistribution.

Formal Charge Formula

Valence Electrons - Lone Pairs - 1/2 Bonding Electrons.

Resonance

Multiple Lewis structures represent a molecule's electrons.

Resonance Hybrid

Actual electron distribution averaged from resonance forms.

Bond Strength

Energy required to break a bond.

Bond Energy

Energy needed to break one mole of bonds.

Enthalpy (H)

Measurement of energy in thermodynamic systems.

Endothermic Process

Chemical process that absorbs heat.

Bond Length

Distance between two bonded nuclei.

Bond Dissociation Energy

Standard enthalpy change for breaking a bond.

Strength and Number of Bonds

More bonds increase bond strength.

Bond Length and Strength Relationship

Stronger bonds have shorter lengths.

Example of Bond Energy

CH4 bond energy is 415 kJ/mol.

Molecular Structure

Arrangement of atoms in a molecule.

Charge Distribution Preference

Prefer structures with minimal formal charges.

Least Electronegative Element

Typically placed at the center of structures.

Enthalpy Change (ΔH)

Energy change during a chemical reaction.

Exothermic Reaction

ΔH negative; heat produced, stronger product bonds.

Endothermic Reaction

ΔH positive; heat absorbed, weaker product bonds.

Lattice Energy (ΔHlattice)

Energy to separate one mole of ionic solid.

Born-Haber Cycle

Series of steps for ionic solid formation.

Ionization Energy (IE)

Energy needed to remove an electron from an atom.

Enthalpy of Sublimation (ΔHs)

Energy required to convert solid to gas.

Bond Dissociation Energy (D)

Energy required to break a bond in a molecule.

VSEPR Theory

Predicts molecular structure based on electron pair repulsion.

Bond Angle

Angle between two bonds at a common atom.

Bond Distance

Distance between nuclei of two bonded atoms.

Ångstrom (Å)

Unit of length; 1 Å = 10⁻¹⁰ m.

Picometer (pm)

Unit of length; 1 pm = 10⁻¹² m.

Electron Density

Regions of high electron concentration around atoms.

Lone Pair

Pair of valence electrons not involved in bonding.

Molecular Polarity

Distribution of electrical charge across a molecule.

Dipole Moment (µ)

Vector representing charge separation in a molecule.

Nonpolar Molecule

Molecule without a net dipole moment.

Bond Moment

Vector quantity representing bond dipole.

Partial Charge (δ+ or δ-)

Charge distribution in polar covalent bonds.

Electric Field Alignment

Polar molecules align in an electric field.

Like Dissolves Like

Polar dissolves polar; nonpolar dissolves nonpolar.

Trigonal Planar Geometry

Three regions of electron density around a central atom.

Tetrahedral Geometry

Four regions of electron density around a central atom.

Trigonal Bipyramidal Geometry

Five regions of electron density around a central atom.

Octahedral Geometry

Six regions of electron density around a central atom.