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Ionic and Covalent Bonds: Structure, Naming, and Properties

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Ionic and Covalent Bonds

Introduction to Ions and Ionic Bonds

Ions are atoms or groups of atoms that have gained or lost electrons, resulting in a net electrical charge. Ionic bonds form between metals and non-metals through the transfer of electrons, creating cations (positively charged ions) and anions (negatively charged ions).

  • Cations: Positively charged ions formed by the loss of valence electrons. Typically formed by metals in Groups 1A, 2A, and 3A, which have low ionization energies and readily lose electrons.

  • Anions: Negatively charged ions formed by the gain of valence electrons. Typically formed by non-metals in Groups 5A, 6A, and 7A, which have high electron affinities and tend to gain electrons to fulfill the octet rule.

  • Ionization Energy: The energy required to remove an electron from an atom. Lower for metals, higher for non-metals.

  • Example: Sodium (Na) loses one electron to form Na+; Fluorine (F) gains one electron to form F-.

Trends in Ionic Charge and Size

The charge of ions is determined by their group number, and the size of ions changes upon gaining or losing electrons.

  • Group Trends:

    • Group 1A: +1 charge (e.g., Na+)

    • Group 2A: +2 charge (e.g., Mg2+)

    • Group 3A: +3 charge (e.g., Al3+)

    • Group 5A: -3 charge (e.g., N3-)

    • Group 6A: -2 charge (e.g., O2-)

    • Group 7A: -1 charge (e.g., F-)

  • Size Changes:

    • Cations are smaller than their neutral atoms due to loss of electrons and reduced electron-electron repulsion.

    • Anions are larger than their neutral atoms due to gain of electrons and increased electron-electron repulsion.

  • Example: Neutral Be vs. Be2+ (size decreases); Neutral F vs. F- (size increases).

Naming Ions and Ionic Compounds

The naming of ions and ionic compounds follows specific conventions based on the type of ion and the elements involved.

  • Cations: Named by adding -ium to the element name (e.g., Barium: Ba2+).

  • Anions: Named by adding -ide to the root of the element name (e.g., Fluoride: F-).

  • Ionic Compounds: Name the metal (cation) first, then the non-metal (anion) with -ide ending (e.g., Magnesium chloride: MgCl2).

  • Criss-Cross Method: Used to determine the formula of ionic compounds by balancing charges.

    • Example: Mg2+ and Cl- combine to form MgCl2.

Transition Metals and Variable Charges

Transition metals can form more than one type of cation with different charges. Their charge is indicated using Roman numerals in parentheses.

  • Examples:

    • Cu+ (Copper(I)), Cu2+ (Copper(II))

    • Fe2+ (Iron(II)), Fe3+ (Iron(III))

  • Naming: Iron(III) oxide: Fe2O3; Iron(II) sulfide: FeS

  • Finding Charge: Use the formula: (number of metal atoms) × (charge of metal) + (number of non-metal atoms) × (charge of non-metal) = 0

    • Example: SnO2: (1 × x) + (2 × -2) = 0 ⇒ x = +4 (Tin(IV) oxide)

Polyatomic Ions

Polyatomic ions are ions composed of two or more atoms covalently bonded, carrying a net charge. Their names and formulas must be memorized.

Name

Formula

Notes

Nitrate

NO3-

"-ate" ending, more oxygen

Chlorate

ClO3-

"-ate" ending, more oxygen

Carbonate

CO32-

Sulfate

SO42-

Sulfite

SO32-

"-ite" ending, less oxygen

Phosphate

PO43-

Hydroxide

OH-

Ammonium

NH4+

Cation

  • Adding Hydrogen: Adds +1 to the charge (e.g., CO32- → HCO3-).

  • Criss-Cross with Polyatomics: Use the same method as with simple ions.

    • Example: Sodium sulfite: Na+ and SO32- combine to form Na2SO3.

Molecular (Covalent) Compounds

Introduction to Covalent Bonds

Molecular compounds are formed when two or more non-metals share electrons, creating covalent bonds. These compounds are named using prefixes to indicate the number of each atom present.

  • Prefixes:

    • 1: mono-

    • 2: di-

    • 3: tri-

    • 4: tetra-

    • 5: penta-

    • 6: hexa-

    • 7: hepta-

    • 8: octa-

    • 9: nona-

    • 10: deca-

  • Naming: Prefix + name of first non-metal (no "mono-" for the first element), prefix + name of second non-metal with -ide ending.

  • Examples:

    • NCl3: Nitrogen trichloride

    • H2O: Dihydrogen monoxide

Lewis Dot Structures

Lewis structures represent the arrangement of valence electrons in molecules, showing how atoms share electrons to achieve stable configurations (octet rule).

  • Steps to Draw Lewis Structures:

    1. Identify the central atom (usually the least electronegative or the one with the highest bonding capacity).

    2. Count the total number of valence electrons.

    3. Distribute electrons to give each atom its typical valence.

    4. Form covalent bonds by sharing electron pairs.

    5. Check for octet completion; use double or triple bonds if necessary.

    6. Assign formal charges if needed.

    7. Consider expanded octets for elements in period 3 or below (e.g., P, S, Cl).

  • Example: H2O: Oxygen is the central atom, forms two single bonds with hydrogen, and has two lone pairs.

Electronegativity and Bond Polarity

Electronegativity is the ability of an atom to attract shared electrons in a covalent bond. The difference in electronegativity between atoms determines bond polarity.

  • Electronegativity Scale: Values increase across a period and decrease down a group. Fluorine is the most electronegative element.

  • Bond Polarity:

    • Non-polar covalent: (electrons shared equally)

    • Polar covalent: (electrons shared unequally)

    • Ionic: (electrons transferred)

  • Example: H–F bond is highly polar due to large electronegativity difference.

Shapes of Molecules (VSEPR Theory)

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

  • Steps to Determine Shape:

    1. Count the number of electron groups around the central atom.

    2. Determine the electronic geometry (e.g., linear, trigonal planar, tetrahedral).

    3. Determine the molecular geometry by considering the number of bonding pairs and lone pairs.

    4. Draw the 3D structure to visualize the shape.

  • Common Geometries:

    • 2 groups: Linear

    • 3 groups: Trigonal planar

    • 4 groups: Tetrahedral

    • 3 bonding, 1 lone pair: Trigonal pyramidal

  • Example: NH3 (ammonia) has a trigonal pyramidal shape due to one lone pair on nitrogen.

Polarity of Molecules

The overall polarity of a molecule depends on both the polarity of its bonds and its molecular geometry. If polar bonds are arranged symmetrically, their dipoles may cancel, resulting in a non-polar molecule.

  • Non-polar Molecules: Symmetrical arrangement of polar bonds (e.g., CO2).

  • Polar Molecules: Asymmetrical arrangement, resulting in a net dipole moment (e.g., H2O).

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