Chemical Bonds and Ionic Bonds

Chemical Bonds

• Noble gases are in Group 18 of the periodic table, with complete valence electron shells. This makes them stable and unreactive. Examples include Helium (He), Argon (Ar), and Xenon (Xe).

• Valency is the combining power of an element, determined by the number of electrons in the outermost shell. It shows how many electrons an atom will gain, lose, or share when bonding.

• Lewis Structures are diagrams that illustrate bonding between atoms and lone electron pairs. The octet rule (most atoms strive to have eight electrons in their valence shell) is important in these structures. Dots represent electrons; lines/pairs of dots, bonds.

Ionic Bonds

• Ionic bonds form when one atom donates electrons to another, creating ions (charged atoms). The resulting electrostatic attraction between oppositely charged ions holds the compound together.

• Formation of an ionic compound: Metal + Nonmetal → Ionic Compound

Examples of Ionic Compounds

• Sodium Chloride (NaCl) : Sodium loses one electron, Chlorine gains one.

• Magnesium Oxide (MgO): Magnesium loses two electrons, Oxygen gains two.

• Calcium Fluoride (CaF2): Calcium (Ca²⁺) and Fluorine (F⁻).

Covalent Bonds

• Covalent bonds form when nonmetal atoms share electrons to achieve a full outer shell. Bonds can be single, double, or triple, depending on the number of shared electron pairs.

• Single Bond: One pair of electrons shared

• Double Bond: Two pairs of electrons shared

• Triple Bond: Three pairs of electrons shared.

Examples of Covalent Molecules

• Water (H₂O): Oxygen shares electrons with two hydrogens (two single bonds).

• Methane (CH₄): Carbon shares electrons with four hydrogens (four single bonds).

• Carbon Dioxide (CO₂): Carbon forms double bonds with each of two oxygen atoms.

• Nitrogen (N₂): Two nitrogen atoms form a triple bond.

• Oxygen (O₂): Two oxygen atoms form a double bond.

Polarity

• Electronegativity is a measure of an atom's ability to attract electrons in a chemical bond. The greater the difference in electronegativity between atoms, the more polar the bond.

• Polarity of a molecule depends on individual bond polarities and the overall molecular shape. Symmetrical shapes can cancel out bond dipoles, resulting in a nonpolar molecule.

Molecular Geometry

• Molecular geometry is the three-dimensional arrangement of atoms within a molecule. This is determined by repulsions between electron pairs. Different shapes are possible.

Intermolecular Forces

• Intermolecular forces are weaker attractions between molecules.

• Intramolecular forces are strong bonds that hold atoms together within a molecule (e.g., ionic, covalent).

• Examples include dipole-dipole interactions, hydrogen bonds, and London dispersion forces.

Hydrogen Bonding

• Hydrogen bonding is a special type of dipole-dipole interaction. It occurs when hydrogen is bonded to a highly electronegative atom (like oxygen, nitrogen, or fluorine) and is attracted to a lone pair on another electronegative atom.

Ion-Dipole Forces

• Ion-dipole forces occur between an ion and a polar molecule. These are important in solutions where salts dissolve in polar solvents.

Ion-Induced Dipole Forces and Dipole-Induced Dipole Forces

• Ion-induced dipole forces occur when an ion distorts the electron cloud in a neighboring nonpolar molecule, inducing a temporary dipole. Dipole-induced forces occur when a polar molecule causes a dipole in a nonpolar molecule.

London Dispersion Forces

• London dispersion forces are weak, temporary attractive forces that arise from temporary dipoles in atoms or molecules. They exist in all molecules and are the only intermolecular force in nonpolar molecules.

Acids and Bases

• In acid-base reactions, the acid donates a proton (H⁺) and becomes its conjugate base; the base accepts a proton and becomes its conjugate acid. The resulting acid-base pairs are conjugate pairs. pH is the negative logarithm of hydrogen ion concentration.