Week 8 Lecture 2 - Topic 5 (bonding)

Chemical Bonding

Electronegativity and Bond Types

Bonds classified by electronegativity differences:
- Covalent bond: Difference of 0 to 0.3.
- Polar covalent bond: Difference of 0.5 to 1.6.
- Ionic bond: Difference greater than 2.

Valency Electrons

Valency electrons are in the outermost energy level.
Example: Nitrogen (N) with atomic number 7.
- Electronic configuration: 1s² 2s² 2p³
- Valence electrons: 5 (2 in 2s and 3 in 2p).

Common Atoms and Bond Formation

Hydrogen (H) always forms one bond.
Oxygen (O) and Sulfur (S) form two covalent bonds.
- Can be one double bond or two single bonds.
Carbon (C) forms four bonds.
- Examples: Two double bonds (CO₂), four single bonds (CH₄).
Nitrogen (N) forms three bonds.
- One triple bond (N₂), three single bonds (NH₃), or one double and one single bond.
Carbon (C) is versatile and forms different types of four bonds.
- Two double bonds, four single bonds (Methane - CH₄), one triple and one single bond, or one double and two single bonds.

Drawing Lewis Structures: Rules

Count Total Valence Electrons:
- For anions, add electrons equal to the charge.
- For cations, subtract electrons equal to the charge.
- Example: H₂O (2 Hydrogen, 1 Oxygen), none are ionized.
- For ionic compounds like Sodium Chloride (NaCl), ions are formed when Na loses an electron to Cl, but the overall molecule has a neutral charage.
Skeleton Structure:
- Place the least electronegative element in the center.
- Typically, the first element in the chemical formula.
- Join terminal atoms with single bonds.
- Example: Carbon Dioxide (CO₂), Carbon is less electronegative and written first.
- Sodium (Na) is written first in Sodium Chloride (NaCl) because it is a cation and has lower electronegativity.
Outer Atoms:
- Place three pairs of electrons on each outer atom (except Hydrogen, which only needs one pair to be stable).
- Central atom is the least electronegative, surrounded by outer atoms.
Extra Electrons:
- Assign any extra electrons to the central atom.
Formal Charge Calculation:
- Formal charge = (Valence electrons of atom) - (Non-bonding electrons + 1/2 bonding electrons).
Minimize Formal Charges:
- Move electrons to minimize formal charge, considering the octet rule.

Formal Charge Details

Formal charge helps understand bonding in a molecule.
Atoms prefer minimal formal charge in covalent molecules.
Formal Charge = (Total valence electrons) - (Lone pair electrons + 1/2 Shared electrons).

Example: Oxygen Molecule (O₂)

Total Valence Electrons:
- Each oxygen has 6 valence electrons.
- Total: 6 x 2 = 12.
- Electronic configuration: 1s² 2s² 2p⁴. Six electrons in the second energy level.
Skeleton Structure:
- Since both atoms are oxygen, either can be the central atom.
Place Electrons on Outer Atom:
- Draw three pairs of electrons on the terminal oxygen atom.
Extra Electrons on Central Atom:
- Add extra electrons to the central atom to satisfy its octet if possible.
Formal Charge Calculation:
- Terminal Oxygen: 6 - (6 + 1/2 * 2) = -1.
- Central Oxygen: 6 - (4 + 1/2 * 2) = +1.
Minimize Formal Charges:
- Move electrons to form a double bond, reducing formal charges to zero.
- Each oxygen atom now has an octet.

Example: Water (H₂O)

Total Valence Electrons:
- Hydrogen: 1 valence electron each (1s¹).
- Oxygen: 6 valence electrons (1s² 2s² 2p⁴).
- Total: (2 x 1) + 6 = 8.
Skeleton Structure:
- Oxygen is central (least electronegative).
- Connect two hydrogen atoms to the oxygen with single bonds.
Place Electrons on Outer Atoms:
- Hydrogen is complete with one bond; no extra electrons needed.
Extra Electrons on Central Atom:
- Draw two lone pairs of electrons on oxygen.
Formal Charge Calculation:
- Hydrogen: 1 - (0 + 1/2 * 2) = 0.
- Oxygen: 6 - (4 + 1/2 * 4) = 0.
  - Formal Charge Calculation for Oxygen:
    - Total valence electrons: 6
    - Unshared electrons: 4 (two lone pairs)
    - Shared electrons: 4 (two single bonds, 2 electrons each)
    - Formal Charge = 6 - 4 - $frac{1}{2}$ * 4 = 0
Minimize Formal Charges:
- All formal charges are already zero, so the structure is stable.

Hydrogen and Stability

Hydrogen is stable with just two electrons (duet rule), unlike other atoms needing eight (octet rule).
Final Lewis structure shows central oxygen, two lone pairs, and bonds to two hydrogen atoms.

Example: Carbonate Ion (CO₃²⁻)

Total Valence Electrons:
- Carbon: 4 valence electrons (1s² 2s² 2p²).
- Oxygen: 6 valence electrons each.
- Total: 4 + (3 x 6) + 2 (from 2- charge) = 24.
Skeleton Structure:
- Carbon is the central atom.
- Connect three oxygen atoms with single bonds to carbon.
Place Electrons on Outer Atoms:
- Add three pairs of electrons (6 dots) to each terminal oxygen atom.
Extra Electrons on Central Atom:
- Add any remaining electrons to the central carbon atom.
Formal Charge Calculation:
- Carbon: 4 - (0 + 1/2 * 6) = +1.
- Oxygen: 6 - (6 + 1/2 * 2) = -1.
  - The formal charge for each oxygen is calculated as follows:
    - Valence electrons: 6
    - Unshared electrons: 6 (three lone pairs, 6 dots)
    - Shared electrons: 2 (single bond shared with carbon)
    - $Formal Charge = 6 - 6 - \frac{1}{2} * 2 = -1$
Minimize Formal Charges and Resonance:
- Rearrange electrons to form a double bond between one oxygen and the central carbon.
- This creates resonance structures with equal probabilities.
- Because there are three oxygen atoms, three resonance structures are possible.

Resonance and Bond Order

Resonance: Equal probabilities of double bond formation among multiple oxygen atoms.
Bond Order: Total number of bonds / Number of bond groups = 4/3 = 1.33.

Nitrate Ion Example

Nitric acid (HNO₃) ionizes to form nitrate (NO₃⁻).
Nitrate also exhibits resonance with three possible structures.
Each structure features nitrogen double-bonded to one oxygen, and single-bonded to two others.
The actual structure is a mixture of these forms.
Resonance stabilizes these structures through electron delocalization.
When discussing the carbonate example with her student, they brought up the following question:
- The student asked why oxygen has 6 valence electrons, but then also appears to have a "seventh" when contributing to the oxygen carbon bond.
- The instructor clarifies, noting that normally oxygen needs two more electrons to reach the octet state so the electrons are shared through a covalent bond.

Valence Shell Electron Pair Repulsion (VSEPR) Theory

Molecular shape determined by repulsion between electron pairs.
Forces in molecules: Attraction and repulsion.
- Attraction: Between electrons and nuclei.
- Repulsion: Between electrons, and between nuclei.
To minimize repulsion, electron pairs around the central atom are positioned as far apart as possible.
Goal: Stable molecular shape.
Ronald Gillespie contributed to VSEPR theory.
Regions of high electron density oriented as far as possible to minimize repulsion.
*Example: Methane (CH₄) and Carbon Dioxide (CO₂).

VSEPR Theory

Number of Regions of High Electron Density dictate structure and depend on number of regions.
- Linear orientation
  - Example: Carbon dioxide with a 180 degree angle.
- Trigonal Planar
  - Bond angle of 120.
  - Example: Boron Trifluoride (BF₃).
- Tetrahedral
  - Example: Methane, or Carbon Tetrafluoride CF₄.
- Trigonal Bipyramidal
  - Bond angle of 120 and 90 degrees.
  - Example: PFFI.
- Octahedral
  - Bond angle of 90.
  - Example: Sulfur Hexafluoride (SF₆).

Steps for Determining Molecular Shape

Draw the Lewis Structure.
Count bonding pairs and lone pairs of electrons.
Use the electron pair number to determine the geometry.
Modify geometry as needed, considering the repulsion order: Lone Pair-Lone Pair > Lone Pair-Bonding Pair > Bonding Pair-Bonding Pair.

BP: Bonding Pairs, shared electrons.
LP: Lone Pairs, unshared electrons.

Geometry: Two Sets of Electron Pairs

Linear Geometry/Shape.
Bond angle is 180 degrees.
Example: Carbon Dioxide (CO₂).

Geometry: Three Sets of Electron Pairs

Trigonal Planar Geometry.
Bond angle is 120 degrees.
Example: Boron Trifluoride (BF₃).
Recognize the difference between geometry and shape.
- Geometry is the arrangement of sets of electrons.
- Shape refers to the position of the atoms.
- Example: Nitrogen Dioxide (NO₂).
  - Atoms have a bent shape.
  - Bond angle is less than what is theoretically expected (115 instead of 120).
  - There are three bonding pairs like BF₃, but also a non-bonding lone pair.

Geometry: Four Sets of Electrons

Tetrahedral Geometry.
- Bonding pairs = 4
- Lone pairs = 0
Example:
- CH4, bonds formed between electrons in the hybrid orbital of Carbon and the electrons of Hydrogen.
Ammonia (NH₃): Trigonal Pyramidal with 1 lone pair.
Water (H₂O): Bent Shape due to 2 lone pairs, with bond angle 104.5 (should be 109 in tetrahedral geometry).
- Bent shaped is used to avoid repulsion.

More Complex Geometries

Five Sets of Electrons

Trigonal Bipyramidal; many probabilities according to number of bonding pairs and lone pairs.
- PCl₅ is an example.
  - Lone pairs = 0
  - Trigonal Bipyramidal for both geometry and shape.
SF₄
- Geometry is Trigonal Bipyramidal, but shape is seesaw.
ClF₃
- Two long pairs, and three bonding pairs.
- Has a T-shape.
ArI₃
- Three long pairs and two bonding pairs.
- Has a linear shape.

Six Sets as Geometry

Octahedral
- Most famous example is SF six with no long layers with an octahedral geometry.
- CLF F5 is also square pyramidal.

Properties of Covalent Bonds

Dipole Moment

Polar covalent bonds (δ+ and δ-).
Dipole moment depends on:
- Bond polarity.
- Molecular shape.
  - The greater the difference in electronegativity, the larger the dipole moment.

Bond Lengths

Nuclear separation distance where the molecule is most stable.
Single bonds are longer than double bonds, which are longer than triple bonds.
- Carbon-Carbon Bonds Trend (in descending order)
  - Length of Carbon-Carbon bond
  - Length of Carbon Double Bond Carbon.
  - Length of Carbon Triple Bond Carbon.

Bond Energy

The energy needed to break the bond between two atoms.
Amount of energy supplied to break a chemical bond.
Bond energies increase as more electrons are shared, and as electronegativity increases.
Bond energies decrease as bond length increases.
Inverse relationship between bond length and bond energy: Single bonds have lower bond energy compared to double bonds.

Orbital Overlap

Bonds formed by orbital overlap (e.g., H₂, 1s-1s overlap).
Electrons assigned to specific orbitals.
Valency orbitals primarily describe bonding.

Hybridization

Bonds formed between orbitals to obtain new hybridized orbitals.

Methane Example

CH₄: Carbon is central, bonds to four Hydrogen.
C: 1s² 2s² 2p², four valence electrons.
Hybridized orbitals: sp³.
Four sp³ hybridized orbitals (one s and three p).

Sigma bonds and hybridization

Sigma bonds formed through orbital overlap.
- Valence orbitals are 2s and 2p.
In each Carbon, new hybrid orbitals are mixed together (blended) to simplify before bonding to each Hydrogen atom.

SB Two

*Mix one 1s and 2p to form three SB2

SB2

Mix it by one s and two p to form 3sB2.

SB

Formed between one s and one p, and the bond angle is 180.
Mixing 1s with 1p, it will form 2sb ready to form bonds with other atoms.

Types of Hybridization (Summary)

Hybridization to form stable sharing of electrons between orbitals of different atoms.
- Two sigma bonds (SP): Linear geometry (180°).
  - Example: ethyne (C₂H₂).
- Three sigma bonds (SP²): Trigonal planar geometry (120°).
  - Example: ethylene (C₂H₄).
- Four sigma bonds (SP³): Tetrahedral geometry (109.5°).
  - Example: methane (CH₄).