Ionic Bonds: Formation and Properties

• An ionic bond forms via electrostatic attraction between oppositely charged ions.
• Ions form through electron transfer from one atom to another.

Formation of Ions

• Cations are positively charged ions created when an atom loses electrons.
• Metals tend to form cations to achieve a stable electron configuration.
• Anions are negatively charged ions created when an atom gains electrons.
• Nonmetals typically form anions to achieve stable electron configurations.
• The number of electrons lost or gained dictates the ion's charge; losing two results in a +2 cation.

The Ionic Bond

• Electrostatic attraction between oppositely charged ions defines the ionic bond.
• This strong attraction extends in all directions, forming a crystal lattice structure.
• Ionic compounds are electrically neutral overall because the total positive charge of cations equals the total negative charge of anions.

Properties of Ionic Compounds

• High Melting and Boiling Points: Strong electrostatic forces between ions necessitate significant energy to overcome, leading to high melting and boiling points.
• Solubility: Many ionic compounds dissolve in polar solvents like water, where water molecules disrupt the crystal lattice by surrounding and separating ions.
• Electrical Conductivity: Solid-state ionic compounds are poor conductors due to fixed ion positions, but they become good conductors when dissolved or melted because ions are free to move and carry charge.
• Brittleness: Ionic compounds are brittle because mechanical stress can force like-charged ions into proximity, causing repulsion and crystal fracture.

Examples of Ionic Compounds

• Sodium Chloride (NaCl): Formed from sodium (Na) reacting with chlorine (Cl); Na loses an electron to become Na+, and Cl gains one to become Cl-.
• Magnesium Oxide (MgO): Formed from magnesium (Mg) reacting with oxygen (O); Mg loses two electrons to become Mg2+, and O gains two to become O2-.

Covalent Bonds

• A covalent bond forms by sharing one or more electron pairs between atoms.
• They usually form between two nonmetal atoms.

Formation of Covalent Bonds

• Atoms share electrons to achieve a stable electron configuration, like that of a noble gas.
• Shared electrons attract to both atoms' nuclei, effectively bonding them.
• Single covalent bonds involve sharing one electron pair; double bonds, two pairs; and triple bonds, three pairs.
• Bond Order: It defines the number of shared electron pairs between two atoms; single bonds have a bond order of 1, double bonds of 2, and triple bonds of 3.

Properties of Covalent Compounds

• Lower Melting and Boiling Points: Covalent compounds generally have lower melting and boiling points than ionic compounds because intermolecular forces are weaker than electrostatic forces.
• Solubility: Solubility varies; polar covalent compounds dissolve in polar solvents, while nonpolar covalent compounds dissolve in nonpolar solvents.
• Electrical Conductivity: Covalent compounds typically conduct electricity poorly because they lack freely moving ions or electrons.
• Softness or Flexibility: Covalent compounds exist as gases, liquids, or solids, with solids often softer and more flexible than ionic compounds.

Examples of Covalent Compounds

• Water (H2O): Each hydrogen atom shares an electron with the oxygen atom, forming two single covalent bonds.
• Methane (CH4): The carbon atom shares electrons with four hydrogen atoms, forming four single covalent bonds.
• Carbon Dioxide (CO2): The carbon atom shares two electron pairs with each oxygen atom, forming two double covalent bonds.

Bond Polarity and Electronegativity

• Electronegativity measures an atom's ability to attract shared electrons in a chemical compound.
• Electronegativity differences between atoms in a covalent bond create bond polarity.
• Nonpolar Covalent Bond: Electrons are shared equally when the electronegativity difference is very small (typically less than 0.4).
• Polar Covalent Bond: Unequal electron sharing occurs, giving the more electronegative atom a partial negative charge (δ-) and the less electronegative atom a partial positive charge (δ+).
• Dipole Moment: It measures bond or molecule polarity, equaling charge magnitude times the distance between charges.

Molecular Polarity

• Molecular polarity depends on both individual bond polarity and molecular geometry.
• A molecule with polar bonds can be nonpolar overall if symmetrical geometry causes bond dipoles to cancel.
• Carbon dioxide (CO2) has two polar bonds, but its linear shape cancels the dipoles, rendering the molecule nonpolar.
• Water (H2O) has polar bonds and a bent shape, so dipoles do not cancel, making it polar.

Comparison of Ionic and Covalent Bonds

• Nature of Interaction: Ionic bonds involve electron transfer and electrostatic attraction, while covalent bonds involve electron sharing.
• Elements Involved: Ionic bonds typically form between metals and nonmetals, whereas covalent bonds form between two nonmetals.
• Physical Properties: Ionic compounds usually have high melting and boiling points, are brittle, and conduct electricity when dissolved or melted; covalent compounds typically have lower melting and boiling points and are generally poor conductors.
• Polarity: Ionic bonds are highly polar, while covalent bonds range from nonpolar to very polar, depending on electronegativity differences.

Explore ionic bonds, formed by electrostatic attraction between oppositely charged ions. Understand how cations (positive ions) and anions (negative ions) are created through electron transfer. Learn about the properties and formation of crystal lattices in ionic compounds.