chapter 7 part 1

CHAPTER 7 – CHEMICAL BONDING AND MOLECULAR GEOMETRY

7.1 Ionic Bonding

Ionic bonding is the electrostatic attraction between oppositely charged ions in an ionic compound. This type of bond typically forms between metals and nonmetals, where electrons are transferred from one atom to another, resulting in the formation of cations and anions.

• Formation of Ionic Bonds: Atoms achieve noble gas configurations by losing or gaining electrons. For example, sodium (Na) loses one electron to become Na+ (isoelectronic with Ne), while chlorine (Cl) gains one electron to become Cl- (isoelectronic with Ar).

• Example: $2\ \mathrm{Mg}(s) + \mathrm{O}_2(g) \rightarrow 2\ \mathrm{MgO}(s)$ Each Mg atom loses two electrons to become Mg2+, and each O atom gains two electrons to become O2-.

• Other Examples:

  • K2S: Each K loses 1 electron (becomes K+), S gains 2 electrons (becomes S2-).

  • AlCl3: Al loses 3 electrons (becomes Al3+), each Cl gains 1 electron (becomes Cl-).

  • NH3: Each H loses 1 electron, N gains 3 electrons (becomes isoelectronic with Ne).

Key Properties of Ionic Compounds:

• Usually solids at room temperature

• High melting points

• Many are soluble in water and are strong electrolytes

• Molten ionic compounds conduct electricity due to mobile ions

7.2 Covalent Bonding

Covalent bonding involves the sharing of electron pairs between atoms. This type of bond typically forms between nonmetals.

• Pure Covalent Bonds: Occur only between identical atoms (e.g., H2, F2, Cl2, Br2, I2, N2, O2), where electrons are shared equally.

• Polar Covalent Bonds: Occur between different atoms (e.g., HF, HBr, IF3, CCl4, CHCl3), where electrons are shared unequally due to differences in electronegativity.

Electronegativity and Bond Polarity:

• Electronegativity is the ability of an atom to attract electrons in a chemical bond toward itself. Fluorine (F) is the most electronegative element.

• Electronegativity increases from left to right across a period and from bottom to top in a group (with exceptions among transition metals).

• Bond Polarity: The difference in electronegativity ($\Delta$EN) determines bond type:

  • $\Delta$EN > 1.6: Ionic bond

  • $0.5 \leq \Delta$EN \leq 1.6$: Polar covalent bond

  • $\Delta$EN < 0.5$: Pure covalent bond

• Percent Ionic Character: Pure ionic bonds have 100% ionic character (not observed in practice); pure covalent bonds have 0% ionic character.

• Difference between Electron Affinity (EA) and Electronegativity: EA is the ability of an isolated atom to attract an additional electron, while electronegativity refers to the ability to attract electrons in a bond. EA is measurable; electronegativity is estimated.

Examples:

• NaCl: Na (0.9), Cl (3.0), $\Delta$EN = 2.1 (ionic)

• CO: C (2.5), O (3.5), $\Delta$EN = 1.0 (polar covalent)

• Cl2: $\Delta$EN = 0.0 (pure covalent)

• NH3: N (3.0), H (2.1), $\Delta$EN = 0.9 (polar covalent)

• Hydrogen peroxide (H2O2): O (3.5), H (2.1), $\Delta$EN = 1.4 (polar covalent)

7.3 Lewis Dot Symbols and Structures

Lewis dot symbols represent valence electrons as dots around the chemical symbol. Lewis structures show covalent bonding and lone pairs in molecules.

• Valence Electrons: Only the outermost electrons are involved in bonding. The number of valence electrons equals the group number for main group elements.

• Bonding and Lone Pairs: Shared electron pairs (bonds) are shown as lines or pairs of dots; non-bonding electrons (lone pairs) are shown as pairs of dots.

• Bond Length: The distance between nuclei of two covalently bonded atoms. Multiple bonds (double, triple) are shorter and stronger than single bonds.

Bond Lengths (pm):

Bond	Length (pm)	Bond	Length (pm)
C–H	107	C–N	143
C–O	143	C=N	138
C=O	121	C≡N	116
C–C	154	N–O	136
C=C	133	N=O	122
C≡C	120	O–H	96

Physical Properties:

• Covalent compounds: Usually gases, liquids, or low-melting solids; low melting points; poor conductors of electricity.

• Ionic compounds: Usually solids; high melting points; conduct electricity when molten or dissolved in water.

Steps for Drawing Lewis Structures:

1. Write the skeletal structure, placing the least electronegative atom in the center (except H and F, which are always terminal).

2. Count the total number of valence electrons (add for anions, subtract for cations).

3. Draw single bonds (2 electrons each) between the central atom and surrounding atoms. Complete octets for surrounding atoms, then place remaining electrons on the central atom as lone pairs.

4. If the central atom lacks an octet, form double or triple bonds as needed using lone pairs from surrounding atoms.

Examples to Practice: CO2, NH3, H2CO, NH4+, CO, PCl3, P2O5, O2, N2

Exceptions to the Octet Rule

• Incomplete Octet: Some elements (e.g., Be, B) form stable compounds with fewer than 8 electrons (e.g., BeH2, BH3, BCl3, BF3).

• Odd-Electron Molecules: Molecules with an odd number of electrons (e.g., NO, NO2) are called free radicals and are highly reactive. They often dimerize to pair up unpaired electrons (e.g., $\mathrm{NO}_2 + \mathrm{NO}_2 \rightarrow \mathrm{N}_2\mathrm{O}_4$).

• Expanded Octet: Elements in period 3 and beyond can have more than 8 electrons due to available d orbitals (e.g., SF6, PCl5, SO42-, XeF2, XeF4).

• Coordinate Covalent Bond: A bond where both electrons come from the same atom (e.g., in F3B–NH3, the N donates both electrons to B).

7.4 Formal Charges and Resonance

Formal charge helps determine the most plausible Lewis structure for a molecule or ion. It is calculated as the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in the Lewis structure.

Formal Charge Formula:

• For each atom: $\mathrm{Formal\ Charge} = \mathrm{Valence\ electrons} - (\mathrm{Nonbonding\ electrons} + \frac{1}{2} \times \mathrm{Bonding\ electrons})$

Procedure:

1. Count the number of nonbonding (lone pair) electrons on the atom.

2. Count the number of bonding electrons (divide by 2).

3. Add these two values and subtract from the number of valence electrons for that atom.

Examples:

Structure	O	C	N	Sum	Preferred?
[O–C≡N]-	-1	0	0	-1	Yes (best, negative on O)
[O=C=N]-	0	0	-1	-1	Yes
[O≡C–N]-	+1	0	-2	-1	No (less preferred)

Structure A is preferred because the negative charge is on the more electronegative atom (O).

Structure	N (left)	N (center)	O	Sum	Preferred?
N=N=O	-1	+1	0	0	Yes
N≡N–O	0	+1	-1	0	Yes (best, negative on O)
N–N≡O	-2	+1	+1	0	No

Structure B is preferred because the negative charge is on the more electronegative atom (O).

Resonance: Some molecules (e.g., NO2) have more than one valid Lewis structure. The actual structure is a resonance hybrid of all possible forms.

Summary Table: Types of Chemical Bonds

Bond Type	Electron Sharing/Transfer	Typical Elements	Example
Ionic	Transfer	Metal + Nonmetal	NaCl
Pure Covalent	Equal Sharing	Same Nonmetals	Cl2
Polar Covalent	Unequal Sharing	Different Nonmetals	HF
Coordinate Covalent	Both electrons from one atom	Nonmetals	F3B–NH3

Additional info: For more complex molecules, resonance and formal charge analysis are essential for predicting reactivity and stability. Expanded octets are possible for elements in period 3 and beyond due to available d orbitals.