Chapter 1: Structure & Bonding

organic chemistry

the study of carbon compounds

chemical reaction

a process in which substances transform into different substances through bond breaking and forming

what we need to know to predict if/when a reaction will occur

location of electrons and their energy levels, density, where they want to move to

structure of an atom

positively charged nucleus (protons + neutrons), orbiting electrons in defined energy levels

Heisenberg’s Uncertainty Principle

you cannot know the exact position and momentum (energy, speed) of a particle at the same time

Schrodinger’s Wave Equation

a mathematical model that describes the wave properties of electrons and predicts the probability of finding them in specific locations around the nucleus

quantum mechanics

describes electron energies and locations by a wave equation

organic chemists use

solutions to the wave equation (orbitals) to describe where electrons in an atom or compound are

Aufbau (“Build Up”) Principle

orbitals fill in order of increasing energy from lowest energy to highest energy

Pauli Exclusion Principle

no more than two electrons may be present in an orbital; if two electrons are present, their spins must be paired (Q# ½ , - ½ )

Hund’s Rule

when orbitals of equal energy are available one electrons is added to each orbital before a second electron is added to any one of them; the spins of the electrons in degenerate orbitals should be aligned

orbitals

where electrons occupy around the nucleus

s orbitals shape

spherical, nucleus in the center

p orbital shape

dumbbell, nucleus at middle

covalent bonds

electron pair is shared between atoms

valence bond theory

electron sharing occurs by overlap of two atomic orbitals; electrons are attracted to nuclei of both atoms

molecular orbital (MO) theory

bonds result from combination of atomic orbitals to give molecular orbitals, which belong to the entire molecule; can predict when bonds will/won’t form

quantum numbers

n, l, m_l, m_s

n (principal energy level)

electron energy level or shell number (1,2,3, 4, 5, 6, 7,…)

l (subshells)

labels orbital type (s, p, d, f); n-1

m_l

labels orbital sub-type (s-1 orbital, p-3 orbitals, d-5 orbitals, f-7 orbitals); each orbitals can hold 2 electrons

degenerate orbitals

electron orbitals that have the same energy level

node

lobes of a p orbital separated by region of zero electron density

bond length

distance between nuclei that leads to maximum stability; too close = repulsion bc both are positively charged, too far apart = bonding is weak

bond overlap

symmetrical and asymmetrical

types of bonds

sigma and pi

symmetrical overlap

along nuclear axis

asymmetrical overlap

above/below nuclear axis (only pi bonding)

valence shell

the outermost occupied electron shell of an atom

valence electrons

in the valence shell of an atom; are used to form chemical bonds and in chemical reactions

Lewis dot structure

the symbol of an element represents the nucleus and all inner shell electrons; dots represent electrons in the valence shell of the atom

nonbonding electrons (lone-pair)

valence electrons not used in bonding

problem with valence bond theory

struggles to explain magnetic properties (like dioxygen's paramagnetism), predict molecular geometries for complex molecules, accurately calculate bond energies, and fully describe delocalized electrons in species with resonance

Linus Pauling

proposed hybrid orbitals

hybrid orbitals

interact to form bonds by overlapping with orbitals from other atoms; formed by combinations of atomic orbitals (using MO theory) by a process called hybridization (combing wave functions)

sp³ hybridization

one 2s atomic orbitals and three 2p atomic orbitals forms four equivalent sp³ hybrid orbitals; four single bonds, tetrahedral, 109.5°

sp² hybridization

one 2s atomic orbitals and two 2p atomic orbitals forms three equivalent sp² hybrid orbitals; double bonds, trigonal planar, 120°

sp hybridization

one 2s atomic orbitals and one 2p atomic orbital forms two equivalent sp hybrid orbitals; triple bonds, linear, 180°

molecular orbital (MO)

where electrons are most likely to be found (specific energy and general shape) in a molecule

additive combination (bonding)

MO is lower in energy (constructive)

subtractive combination (antibonding)

MO is higher energy (destructive)

C valence requirements

valency of 4

valency

the combining capacity of an atom or element in forming chemical bonds, indicating the number of bonds an atom typically forms

N valence requirements

valency of 3

O valence requirements

valency of 2

H valence requirements

valency of 1

halogens (F, Cl, Br) valence requirements

valency of 1

ground state of carbon

electrons are placed in accordance with the quantum chemistry principles (Aufbau, Hund’s Rule, Pauli Exclusion, etc.) that dictate the lowest energy form of carbon

excited state of carbon

if electrons are placed in a different manner (ex: one electron in the 2s and three electrons in the 2p) there would be a higher energy level (excited); when the electrons are rearranged back into the ground state, energy is released