Test 1 Review
Lewis Structures and Bonding
∙ Know the number of valence electrons (group number) i.e., oxygen has 6 valence electrons because it is in group 6
∙ Carbon is usually ALWAYS the central atom in a Lewis Structure
∙ Heteroatoms are any atom that is not hydrogen or carbon in organic chemistry. ∙ Draw lone pairs!
∙ A carbon with 3 bonds will always have a positive charge.
∙ A carbon with 3 bonds and 1 lone pair will always have a negative charge.
Formal charge= #valence electrons- (1/2 bonding electrons)- nonbonding electrons ∙ Don’t include the charge when calculating formal charge.
∙ S orbitals are symmetric.
∙ P orbitals have 2 lobes with a node (looks like 2 balloons attached to each other). ∙ Hybridized orbitals are the average density of electrons in lobes focused on one side about the nucleus.
∙ Sigma* indicates the area of the orbital that is empty. This anti-bonding combination experiences reactivity.
∙ Sigma bonds are the strongest.
∙ Lone pairs have an effect on the shape/arrangement of atoms (their geometry). They decrease the angle length between bonds.
∙ As the hybridization changes, geometry changes.
Atoms + Lone pairs
4 + 0 Lone pairs
3 + 1 lone pair
2 + 2 Lone pairs
3 + 0 Lone pairs
2 + 1 lone pair
2 + 0 Lone pairs
If you want to learn more check out Define the krebs cycle.
We also discuss several other topics like What is being observed? where is it being observed? how is it being observed?
∙ The polarity of bonds determine reactivity.
∙ Induction is the tendency of a higher electronegative atom to pull electrons toward itself ∙ Both lone pairs and the types of atoms in a molecule change the direction of a dipole - Larger dipole moment from more pull on electrons towards the electronegative atom (ie, Fluorine, Oxygen, etc)
- A shorter bonds leads to a stronger dipole because the pull is closer.
∙ Neutral molecules are attracted through dipole-dipole interactions, hydrogen bonds, and London dispersion forces.
∙ Protic compounds are compounds containing hydrogen atoms that are capable of forming hydrogen bonds.
∙ Hydrogen bonds are strong (relatively), increase the boiling point of a compound, and occur when hydrogen is bound to a higher electronegative atom (O, N, F). If you want to learn more check out How do you make decisions?
∙ London dispersion forces are weak attractive forces between slightly covalent bonds; occur when there is a brief dipole in the molecule.
- Bigger molecules with London dispersion forces have higher boiling points - If branching of the molecule is increased, the surface area and boiling points decrease
∙ Isomers are molecules with the same chemical formula, but they are different compounds. - Constitutional isomers have the same molecular formula, but different connectivity of atoms
- Stereoisomers have the same molecule formula, but are arranged differently in space. ∙ Bond-line formulas are a short-hand way to write/draw structures.
∙ Always draw groups on the end of the structures, and include lone pairs on atoms *halogens are usually on the end of a molecule because they can only form one bond* ∙ General Guidelines for Drawing Bond-Line Structures
1. Carbons are at intersections or ends
2. Hydrogens can be implicit (not drawn) Don't forget about the age old question of One sometimes hears that so many species resolve their contests via mostly harmless threat signals to reduce the number of injuries and thereby protect the breeding adults needed to produce the next generation of offspring. what’s the problem with this hy
3. Use wedge/dash notation to indicate 3-demensionality
4. Only 4 bonds on a neutral carbon atom
5. All heteroatoms must be shown
6. All lone pairs are shown
7. Show all hydrogens on heteroatoms
8. Bonds always drawn to bonding atoms
9. Geometry MUST be correct
∙ Alkanes are saturated hydrocarbons (single bonds) C-H
∙ Alkenes (olefins) are unsaturated hydrocarbons (double bonds) C=H
∙ Alkynes are carbon triple bonds We also discuss several other topics like Who wrote manifesto?
∙ Arene (benzene) is a 6-membered ring with 3 double bonds in the ring; “aryl” group ∙ Alkyl halide: R-X, X is a halogen
∙ Alcohol: R-OH
∙ Ether: R-O-R
∙ Thiol: R-S-H
∙ Sulfide: R-S-R
∙ Amine: R-N-R2
∙ Amide: R NH2 O
∙ Ketone:R R If you want to learn more check out What are the three body senses?
; H can be R groups
∙ Aldehyde:RH O
∙ Ester:R OR
∙ Carboxylic Acid: R OH
∙ Anhydride: R O O O
O O S
∙ Substituents are the atoms/compounds hanging off the parent molecule
Bonded to 1 R group
Bonded to 2 R groups
Bonded to 3 R groups
Bonded to 4 R groups
Atom attached to an alkene (C=C)
Atom next to a double bond (one bond away)
One bond away off a benzene ring
Substituent attached to an arene
∙ Resonance allows us to see multiple dimensions of the molecule
∙ Rules for Resonance
1. Fulfill octet rule for all atoms
2. Formal charge matches electronegativity of the charged atom
3. Minimize overall charge
4. Minimize charge separation
*NEVER BREAK SIGMA BONDS*
∙ Move electrons toward positive charge and away from negative charge.
∙ The major contributor is the most stable resonance structure
∙ ALWAYS draw hydrogens on resonance structure
∙ For nitrogen to participate in resonance, it must have a delocalized electron pair, meaning the lone pair on the N atom MUST be allylic to a pi bond (double bond). Localized electrons will NOT participate in resonance.
∙ Always sp2 hybridized; carbon has positive charge
∙ Hyperconjugation occurs when more R groups surround a carbon atom than hydrogen atoms; can share/contribute to the charge, increasing the stability.
∙ Most stable carbocations: tertiary> Secondary>>Primary
∙ Benzylic is slightly more stable than allylic
∙ Conjugated is when there are double bonds “next” to each other, separated by on sigma bond Carbanions
∙ Inversely related to carbocation stability
∙ Most stable: primary>secondary>>tertiary
∙ Destabilized by more substituents surrounding the carbanion
∙ Electron withdrawing atom will make the molecule more stable; induction pulls the electron towards the more electronegative atom to spread the charge across the atom
Acids and Bases
∙ Bronsted-Lowry Definition: proton donators (acids) and acceptors (bases)
∙ Lewis Definition: electron donators (base) and electron acceptors (acid)
∙ If an acid is much stronger than its conjugate base, equilibrium favors the weaker side (conjugate side).
∙ The length of equilibrium arrows depict the magnitude of favoring sides in equilibrium. ∙ Protonation is receiving a proton, while deprotonation is donating a proton. ∙ Always draw out bonds in acid/base reactions!
∙ Acids differ in their ability to donate H+ atoms, or their degree of ionization ∙ Strong acids have a low pKa, while weak acids have a high pKa
∙ Stronger acids are more willing to give up a proton because they are more stable as an ion. ∙ Predicting Acid Strength
1. Which atom bears the charge?
- Electronegativity: across a row; higher electronegativity indicates a stronger acid (more stable conjugate)
- Atomic Radius: down a column; larger in sizes indicates a stronger acid (more stable because able to spread out the charge more effectively)
- A conjugate base that has resonance is more stable because it can spread out the charge 3. Induction
- More electron withdrawing atoms provide more stability, and, therefore, a stronger acid - Increasing the dipole moment leads to a stronger acid
- Sp>sp2>sp3(most stable to least stable)
∙ An adduct is the product of a Lewis acid-base reaction
∙ Used to identify functional groups and compounds based on signals received by an IR reading.