Ionic and Covalent Bonds: Ions, Naming, and Molecular Structure

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) and anions (negatively charged).

• 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.

• 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 ionization energies and seek to fulfill the octet rule.

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

• Octet Rule: Atoms tend to gain, lose, or share electrons to achieve a full set of eight valence electrons.

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

Trends in Ionic Size

The size of ions changes compared to their neutral atoms due to the gain or loss of electrons.

• Cations are smaller than their parent atoms because they lose electrons, reducing electron-electron repulsion.

• Anions are larger than their parent atoms because they gain electrons, increasing electron-electron repulsion.

Example: Be (beryllium) forms Be2+, which is smaller than neutral Be; F forms F-, which is larger than neutral F.

Naming Ions and Ionic Compounds

The naming of ions and ionic compounds follows specific conventions:

• Cations: Named by adding -ium to the element name (e.g., sodium, barium).

• Anions: Named by replacing the ending with -ide (e.g., fluoride, chloride).

• Ionic Compounds: Name the cation first, then the anion (e.g., magnesium chloride for MgCl2).

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

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

Transition Metals and Variable Charges

Transition metals can have more than one possible charge (multivalent). Their charge is indicated using Roman numerals in parentheses.

• Example: Fe2+ is iron(II); Fe3+ is iron(III).

• Formula Example: Iron(III) oxide is Fe2O3; Iron(II) sulfide is FeS.

• Finding the Charge: Use the total charge of the compound to deduce the charge on the metal ion.

Example Calculation: For SnO2: Let x = charge on Sn, O = 2- each, so x + 2(-2) = 0 ⇒ x = +4. Thus, tin(IV) oxide.

Polyatomic Ions

Polyatomic ions are ions composed of more than one atom. They often contain oxygen and have characteristic endings:

• -ate: More oxygen atoms (e.g., nitrate NO3-, sulfate SO42-).

• -ite: One fewer oxygen atom (e.g., nitrite NO2-, sulfite SO32-).

• per- ... -ate: One more oxygen than -ate (e.g., perchlorate ClO4-).

• hypo- ... -ite: One fewer oxygen than -ite (e.g., hypochlorite ClO-).

• Adding Hydrogen: Adds a hydrogen atom and increases the charge by +1 (e.g., CO32- to HCO3-).

• Common Polyatomic Ions: hydroxide (OH-), ammonium (NH4+), carbonate (CO32-), phosphate (PO43-).

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 through 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, prefix + name of second non-metal (ending in -ide). The prefix 'mono-' is omitted for the first element.

Examples:

• NCl3: Nitrogen trichloride

• H2O: Dihydrogen monoxide

Lewis Dot Structures

Lewis structures represent the arrangement of valence electrons in molecules. They help visualize bonding and lone pairs, following the octet rule.

• Steps to Draw Lewis Structures:

  1. Identify the central atom (usually the least electronegative).

  2. Count total valence electrons.

  3. Distribute electrons to satisfy the octet rule for each atom.

  4. Form covalent bonds (shared pairs) between atoms.

  5. Check for double or triple bonds if needed to satisfy octets.

  6. Assign formal charges if necessary.

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

Example: For PCl3:

• P is the central atom (5 valence electrons).

• Each Cl has 7 valence electrons; 3 × 7 = 21.

• Total valence electrons = 5 + 21 = 26.

• Form three P–Cl bonds and assign remaining electrons as lone pairs.

Electronegativity and Bond Polarity

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

• Non-polar Covalent Bonds: Electronegativity difference ($\Delta EN$) = 0–0.4; electrons shared equally.

• Polar Covalent Bonds: $\Delta EN$ = 0.5–1.8; electrons shared unequally, creating partial charges.

• Ionic Bonds: $\Delta EN$ > 1.8; electrons transferred completely.

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 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 based on 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. Asymmetrical arrangements lead to polar molecules.

• Non-polar Molecules: Symmetrical shape, dipoles cancel (e.g., CO2).

• Polar Molecules: Asymmetrical shape, dipoles do not cancel (e.g., H2O).

Example: H2O is polar due to its bent shape and polar O–H bonds.

Summary Table: Common Polyatomic Ions

Name	Formula	Charge
Nitrate	NO3-	-1
Chlorate	ClO3-	-1
Carbonate	CO32-	-2
Sulfate	SO42-	-2
Sulfite	SO32-	-2
Phosphate	PO43-	-3
Hydroxide	OH-	-1
Ammonium	NH4+	+1

Additional info: Some steps and explanations were expanded for clarity and completeness, including the VSEPR model and the process for drawing Lewis structures.