CHM135H1 Study Guide - Comprehensive Final Guide: Valence Bond Theory, Orbital Hybridisation, Tetrahedron

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CHM139H1
FINAL EXAM
STUDY GUIDE
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Ch. 1 Introduction to Organic Chemistry
Structure of an Atom:
- Nucleus: positively charged (protons & neutrons)
- Electron cloud: negatively charged electrons
- Atomic number: # protons
- Mass number: # protons & neutrons
- Isotopes: # protons & different # neutrons
- Atomic mass: average weighted mass (of different isotopes)
Carbon Atom:
- Tetravalent- 4 bonds
o 4 valence electrons available for bonding
o Most stable configuration
- Chains & rings
Quantum mechanics: describes electron energies & locations with wave functions
- Psy2: where electrons are most likely to be found
- Electron cloud has no specific boundary- only most probable areas
Atomic orbitals:
- 4 different orbitals- s, p, d, f
- Orbitals grouped in shells of increasing size & energy
o S- spherical
o P- dumbbell
o D- cloverleaf
Atomic structure: Electron configurations
1. Aufbau’s Principle: lowest-energy orbitals fill first
2. Pauli Exclusion Principle: only 2 electrons in each shell & have opposite
3. Hund’s Rule: fill empty shells w/ 1 electron first
- Possible to promote an electron from lower energy shell to higher
Ex. Carbon
Indicating covalent bonds:
- Covalent bonds: bonding through sharing electrons between two atoms
o Dashed bond- into the paper
o Wedged bond- out of the paper
o Line bond- in the plane of the page
- Lewis structures: atoms dots
- Kekule structures: line b/w atoms indicating electron sharing
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**Octet rule- atoms are most stable when they have 8 electrons around their atoms**
- Often dictates how an atom will react. If a reaction breaks the octet ruleunlikely to happen
Valency Rules (& Examples):
1. Atoms with # of valence electrons- generally form same # of bonds
2. Atoms w/ 4+ valence electrons- form as many bonds as needed to fill s & p orbitals
a. Satisfy octet rule
Ex. Carbon- 4 valence electrons tends for form 4 bonds
Ex. Nitrogen- 5 valence electrons form 3 bonds w/ 2 lone pair electrons
Ex. Oxygen- 6 valence electrons forms 2 bonds, w/ 4 lone pair electrons
- Only 4 bonds around carbon, NEVER 5 (b/c doesn’t have the orbitals)
Bond Formation Theories
I. Valence Bond Theory
II. Molecular Orbitals Theory
Valence Bond Theory:
- Covalent bonds form: singly occupied electron orbitals OVERLAP with another singly occupied orbital
- Bonds formation: releases energy- the molecule is in a more stable state
- Bond length: the maximum distance b/w 2 nuclei for maximum stability
o Too close- atoms repel = more energy needs to be inputted to keep atoms close
o Too far- attraction is too weak
- Hybrid orbitals:
o Will only hybridize the # of orbitals needed for bonding w/ atoms
o Bond w/ 4 different atoms- SP3 orbitals (needs all 4 hybrids for different atoms)
o Bond w/ 3 different atoms- SP2 orbitals (only need 3 hybrids for different atoms)
- S-S overlap (sigma bonds)
o Shared electrons reside between nuclei
- P-P overlap (pi bond)
o Shared electrons are away from the 2 nuclei
o Significantly weaker than sigma bonds
- Sp3 hybrids: 109.5 (tetrahedral angle) the angle for maximum stability
- Sp2 hybrids: 120
o Often formed for double bonds
o Remaining P orbital- perpendicular to the plane
- Double bonds:
o Shorter & stronger than single bonds
o Need sp2 hybridization to occur
- Sp hybrids: S & P orbitals hybridize
o 2 p orbitals remain the same
o Form triple bonds
1 sigma bond
2 pi bonds
Hybridized Atoms other than CARBON:
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