Lewis Structures and Bonding

Chapter 1: Introduction

Lewis Structures: Used to represent valence electrons around an element's symbol.
- Dots are placed around d the element symbol to represent valence electrons.
Valence Electrons of First 20 Elements
- The number of valence electrons corresponds to the column number (for main group elements).
  - First column elements have one valence electron.
  - Second column elements have two valence electrons.
  - And so on.
Electron Configuration and Orbitals
- Electrons are located in s and p orbitals.
- There is one s orbital and three p orbitals (px, py, pz).
- The first two electrons fill the s orbital, and the remaining electrons fill the p orbitals in order.
Transition Metals
- The outer electrons of transition metals are generally two or one; the number isn't as predictable as in main group elements.
Valence Electrons and Position on the Periodic Table
- Elements in the outer columns have valence electrons equal to their group number (1, 2, 3, 4, 5, 6, etc.).
  - Carbon has four valence electrons.
  - Nitrogen has five valence electrons.
  - Oxygen has six valence electrons.
  - Fluorine has seven valence electrons.
Maximum Electron Capacity per Side
- Each side of the atom can have a maximum of two dots (electrons) due to orbital capacity.
Neon's Electron Configuration
- Neon has eight valence electrons, with two electrons on each side of the atom when drawing Lewis structures.

Chapter 2: Sharing Electrons

Lewis Structures for Bonding: Representation differs when elements bond.
Types of Bonds
- Single Bond: Represented by a single line; consists of two electrons.
- Double Bond: Represented by two lines; consists of four electrons.
- Triple Bond: Represented by three lines; consists of six electrons.
- $(\text{No Quadruple Bond})$
Covalent Compounds: Involve the sharing of electrons.
Diatomic Molecules: Sharing of electrons allows both atoms to achieve a stable electron configuration (octet or duet).
Hydrogen Molecule (H2)
- Each hydrogen atom has one electron and needs one more to achieve a stable configuration of two electrons.
- Two hydrogen atoms share their electrons, forming a single bond (H-H).
- The shared electrons are counted towards both hydrogen atoms, satisfying their need for two electrons each.
Fluorine Molecule (F2)
- Each fluorine atom has seven electrons and needs one more to achieve a stable configuration of eight electrons.
- Two fluorine atoms share one electron each, forming a single bond (F-F).
- Each fluorine atom now has eight electrons (six non-bonding and two shared).

Chapter 3: Many Valence Electrons

Predicting Bond Types: Recognizing when to use single, double, or triple bonds will come with practice, guided by specific rules.
Hydrogen and Halogens: Hydrogen (H) and the halogens (F, Cl, Br, I) typically form only one bond because they need only one additional electron to complete their valence shells.
Terms: Shared Electrons and Lone Pairs
- Shared Electrons: Electrons involved in a bond.
- Lone Pairs: Electrons not involved in bonding, existing as pairs on an atom.
Hydrogen (H)
- Forms only one bond.
- Cannot have lone pairs.
Ammonia (NH3)
- Nitrogen (N) bonds with three hydrogen (H) atoms.
- Each H atom shares one electron with the N atom.
- N shares one electron with each H atom.
- The resulting molecule has N bonded to three H atoms and one lone pair on the N atom.
Rules for Drawing Lewis Structures: These rules help determine how to draw molecules accurately.
Rule 1: Count the Total Number of Valence Electrons
- Example: Phosphorus Trichloride (PCl3)
  - Phosphorus (P) has five valence electrons.
  - Each Chlorine (Cl) atom has seven valence electrons.
  - Total valence electrons: $5 + (3 \times 7) = 26$

Chapter 4: Subtract The Electrons

Rules for Drawing Lewis Structures (Continued)
Rule 2: Pick the Central Atom and Form Single Bonds
- The central atom is usually the least electronegative element.
- Electronegativity: The ability of an atom to attract electrons.
- Metals are less electronegative than nonmetals.
  - Fluorine is the most electronegative element.
  - Electronegativity decreases as you move down and to the left on the periodic table.
- In PCl3, phosphorus (P) is the central atom because it is less electronegative than chlorine (Cl).
- Place single bonds from the central atom to the other atoms.
- Subtract the electrons used in the single bonds from the total valence electrons.
  - Each single bond uses two electrons.
  - For PCl3: 3 single bonds use $3 \times 2 = 6$ electrons.
  - Remaining electrons: $26 - 6 = 20$
Rule 3: Complete Octets on Outer Atoms
- Add electrons to the outer atoms to give each an octet (8 electrons).
- Subtract the electrons used to complete the octets from the remaining valence electrons.
  - Each chlorine atom in PCl3 needs 6 more electrons to complete its octet.
  - $3 \times 6 = 18$ electrons are used.
  - Remaining electrons: $20 - 18 = 2$

Chapter 5: Electrons On Central

Rule 4: Place Extra Electrons on the Central Atom
- Put any remaining electrons on the central atom as lone pairs.
- After placing the extra electrons, there should be zero electrons left.
- In PCl3, the remaining two electrons are placed on the central phosphorus atom.
- Subtract these two electrons: $2 - 2 = 0$ . Zero electrons left.
Check for Octets: Verify that all atoms have an octet (or duet for hydrogen).
- If the central atom does not have an octet, form double or triple bonds.
- In PCl3, check if the phosphorus atom has an octet.
  - Phosphorus has two electrons from the lone pair and six electrons from the three single bonds, totaling eight electrons.
  - All atoms have octets; the structure is correct.
Alternative Method for Determining Structure: By understanding the valence electrons of elements, you can predict the structure.
- Elements like phosphorus need three additional electrons to complete their octets, while chlorine needs one.
- Therefore, each chlorine atom will form a single bond with the phosphorus atom.
Lone Pairs and Bonding Pairs: For PCl3, the central phosphorus atom has one lone pair and three bonding pairs.
Water (H2O)
- Hydrogen cannot be a central atom because it can only hold two electrons and form one bond.
- Total valence electrons: Oxygen has six valence electrons, and each hydrogen has one, totaling $6 + (2 \times 1) = 8$ electrons.
- Place oxygen as the central atom and form single bonds to each hydrogen atom.
- Subtract the electrons used in the single bonds from the total valence electrons.
  - $8 - (2 \times 2) = 4$ electrons remaining.

Chapter 6: No More Electrons

Water (H2O) Lewis Structure Conclusion: The remaining four electrons are placed on the central oxygen atom as two lone pairs.
Satisfying Octets: Check if all atoms are "happy" (have the required number of electrons).
- Hydrogen is satisfied with two electrons.
- Oxygen has eight electrons (two from the bonds with hydrogen and six from the two lone pairs) and is satisfied.
Iodine (I2)
- Total valence electrons: Each iodine atom has seven, totaling $7 \times 2 = 14$ electrons.
- Form a single bond between the two iodine atoms: I-I. This uses two electrons, leaving 12.
- Place six electrons (three lone pairs) on one iodine atom to complete its octet. This uses six electrons, leaving six.
- Place the remaining six electrons (three lone pairs) on the other iodine atom.
- Each iodine atom now has eight electrons (one single bond and three lone pairs) and is stable.

Chapter 8: VSEPR Theory

Valence-Shell Electron Pair Repulsion

First # → Total Domains — Bonded and Lone Pairs
Second # → Bonded Domains
Third # → Lone Pair Domains

Molecular Geometry is based on the shape of the entire molecule

Electron-Domain Geometry is based on the electrons in the molecule

If the first number of the VSEPR is 2, the geometry is linear.

If the first number of the VSEPR is 3, the geometry is trigonal planar.

If the first number of the VSEPR is 4, the geometry is tetrahedral.

If the first number of the VSEPR is 5, the geometry is trigonal bipyramidal.

If the first number of the VSEPR is 6, the geometry is octahedral.