Study Guide Lectures

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CH 220 1st Edition
Exam # 1 Study Guide Lectures: 1 - 6
Lecture 1

This class will focus on H, C, N, O, F, Cl, Br, I (and sometimes P & S)
Atomic Structure - Positively charged nucleus that contains protons and neutrons
Atomic Number - Often written as (Z), is the number of protons in the atom’s nucleus. Atoms of a given element always have the same atomic number
Atomic Mass - Often written as (A), is the number of protons plus neutrons. Isotopes are atoms of the same element that have different numbers of neutrons
Orbitals describe where electrons are most likely to be with reference to the nucleus, it is a ‘probable’ area and not a definite one
Orbitals further away from the nucleus have the most energy
Orbitals are grouped into shells, which can be made up of different types of orbitals, with each orbital being able to be occupied by two electrons
Carbon always has four bonds, and adopts a tetrahedron shape so that the bonds have sufficient spacing – the angles between the bonds are approximately 109 degrees
Atoms form bonds because the compound that results is more stable than the separate atoms
Lecture 2

In neutral atoms the number of protons is equal to the number of electrons
Valence electrons are involved in chemical bonding & reactions, as well as achieving a noble gas configuration
Organic chemistry focuses on nonmetal-nonmetal bonding, known as covalent bonding. Covalently bonded compounds share electrons to achieve their octet
In a covalent bond, the nuclei are attracted to the electron density between atoms, which forms a directional bond
Lewis structures – bars represent bonded electrons, dots represent lone pairs
VSEPR model – Valence Shell Electron Pair Repulsion. The directional nature of bonds dictates the shape of molecules and influences the reactivity of the molecule. Atoms that obey the octet rule can only have 2, 3, or 4 electron regions. The electron regions around atoms repel each other and they get as far away from each other as they can while still being tethered (attracted).
One s orbital hybridizes with one p orbital > 2 electron regions > bond angle is 180
One s orbital hybridizes with two p orbitals > 3 electron regions (sp2) > bond angle is 120
One s orbital hybridizes with all three p orbitals > 4 electron regions (sp3) > bond angle is ~109
Lecture 3
Functional groups - collection of atoms at a site that have a characteristic behavior in all molecules where it occurs
i. Alkenes have a carbon carbon double bond ii. Alkynes have a carbon carbon triple bond iii. Arenes have an alternating single and double bonds arranged in a ring, sometimes known as an aromatic ring (it doesn’t mean it smells good, it just means that it smells) iv. Alkyl halide has a bond to F, Cl, Br or I (halogens)
v. Alcohol has a bond to OH vi. Ether has an oxygen bridge between 2 carbons vii. Some functional groups have a carbon-oxygen double bonds (these have a family name of Carbonyls) – see diagrams in the inside cover of textbook
1. Aldehyde
2. Ketone
3. Carboxylic acid
4. Ester
5. Amide
6. Acid chloride
Alkanes - compounds with carbon-carbon single bonds and carbon hydrogen bonds, no functional groups (aka hydrocarbons). Formula for alkane with no rings must be CnH2n+2
Alkanes with carbons connected to no more than 2 other carbons are straight chain or