Organic Chemistry Exam 1 (Ch.1,2,3)

Orbitals

Orbitals

mathematical predictions of where electrons might be

What orbitals will first-row orbitals have? (H, He)

S orbitals only

What orbitals will second-row orbitals have? (B,C,N,O,F)

1s,2s,2p orbitals

What is the octet rule?

Max of 8 electrons

ionic bonding

electrostatic interaction of 2 ions (one positive and one negative)

Cation

positive ion

Anion

negative ion

The most common ionic bonding is between

metal and nonmetal

covalent bonding

Not all bonds are the same, and NOT all are equal

formal charge equation

valence electrons minus [# of lone pairs + (1/2 # of shared electron bonds)]

Molecular shape

angles will change with lone pairs and bondage

Bond length

average length between 2 bonded atoms

What does bond length do when atomic radius increases?

Bond length decreases with atomic radius

Atomic radius decreases as you move

left to right across periodic table

The atomic radius increases as you move

down the periodic table

Bond angle

# of groups (lone pairs or bonded atoms) will determine geometry

2 groups is what geometry

linear, 180

3 groups is what geometry

trigonal planar, 120

4 groups is what geometry

tetrahedral, 109.5

Heteroatom

any atom other then C or H

Hybridization

rearrangement of atomic orbitals into molecular orbitals

2 groups and 2 hybrid orbitals means a hybridization of

sp

3 groups and 3 hybrid orbitals means a hybridization of

sp^2

4 groups and 4 hybrid orbitals means a hybridization of

sp^3

TRUE or FALSE. Lone pairs count as a group when determining hybridzation.

True

As bond length increases, bond strength…

decreases

Shorter bond is a

stronger bond

TRUE or FALSE. Triple bonds are the strongest.

True

More s-character means a

stronger and shorter bond

TRUE or FALSE. sp is the shortest and strongest bond.

true

Polarity

unequal sharing of electron density

High electron density likely to have

electron between 2 atoms

An electronegative atom connected to an electropositive atom leads to

a polar molecule

Nonpolar molecule

does NOT contain electronegative atoms

Polar molecules have a net

dipole

Resonance structure

2 or more Lewis Structures with some placement of atoms, but different arrangement of electrons (bonds and lone pairs)

Curved arrow notation

move electrons in pairs in resonance structures

sigma bond

one bond

pi bond

double bond

one sigma bond and 2 pi bonds

triple bond

conjugation

occurs where p-orbitals (pi bonds) overlap on 3 or more adjacent carbons

Special resonance allows for

delocalization (spreading out) of electron density

aromaticity

special type of conjugation, whatever type of charge you start with is the charge you must end with

The better resonance form has

lower energy

Better resonance forms have more bonds and less charges which leads to more

stablity

Hydrocarbons

made up of only hydrogens and carbons

Aliphatic hydrocarbons

alkanes, sigma bond

alkenes, 1 sigma, and 1 pi bond

alkynes, 1 sigma and 2 pi bonds

R is a substituent that means

any hydrocarbon

1 degree

primary carbon, bound to 1 other carbon

2 degree

secondary carbon, bound to 2 carbons

3 degree

tertiary carbon, bound to 3 carbons

4 degree

quaternary carbon, bound to 4 carbons

Alkyl halides

R-X

X is F, Cl, Br, I

Alcohols (hydroxyl groups or -OH)

R-OH

Ex: methanol

Ethers (alkoxy groups)

R1-O-R2 (bent structure)

Ex: Dimethyl ether

H3C-O-CH3 (bent structure)

Amines (amino group) 1, 2, or 3 amines

Ex: methyl amine

H3C-NH2

Thiols (marcato groups)

R-SH

Ex: H3C-SH

Sulfide (alkylthio group)

R1-S-R2

Ex: Dimethyl sulfide

H3C-S-CH3

Aldehyde (end in “al”)

Ex: ethanal

Must have hydrogen

Ketone (carbonyl groups)

Ex: acetone

Ketone structure

can also be R1 and R2

carboxylic acid (carboxy groups)

Ex: acetic acid

carboxylic acid structure

Esters

Ex: ethyl acetate

Ester structure

amide

Ex: methyl amide

amide structure

The R2 and R3 can also be H’s

acid chloride

Ex: acetyl chloride

acid chloride structure

resonance structures are better when they have

fewer charges and more bonds

full octets

negative charges on electronegative atoms

Bronsted Lowry acid

donates H+ (proton) or H3O+

Bronsted Lowry base

accepts H+ (lone pair electron or pi bond)

Example of Bronsted Lowry acid

H-Cl

H-Br

H2SO4

H3O+

H2O

Acetic Acid

Example of Bronsted Lowry Bases

H2O (4 lone pairs on Oxygen)

H3N (2 lone pairs on Nitrogen)

Hydroxide, negative charge

Methoxide, negative charge

Acetone

Reaction of Bronsted Lowry acids and bases

H-A + B ⇆ A^- + H-B^+

acid base C.B C.A

Another example of Bronsted Lowry acids and bases

H-Br + Methanol ⇆ Br^- + (Methanol + extra H)

Acid Base C.B C.A

Lewis acid

accepts an electron pair (lone pair or pi bond)

Lewis base

donates an electron pair

TRUE or FALSE. Transition metals can accept a lone pair of electrons.

true. B, Al, Ga, etc.

Lewis acid and base reaction example

H-A + water ⇆ A +H3O

Acid Base C.B C.A

What is the K equilibrium formula

products divided by starting material

Example of K equilibrium formula

[H30+][A-] divided by [H-A][H2O], water is liquid therefore not included so equation is

[H3O+][A-] divided by [HA]

Ka is the

acidity constant

A bigger Ka means

equilibrium lies towards the right (product-favored)

A higher Ka means

a stronger HA (acid)

pKa =

-log(Ka)

smaller pKa means a

stronger acid

Larger pKa means a

weaker acid

Strong acid wants to

donate a proton

A strong acid makes a

weak conjugate base

A strong base wants to

accept a H+ (Become B-H)

A strong base makes a

weak conjugate acid

TRUE or FALSE. Equilibrium favors the side with the weaker acids/bases because they don’t react well.

true

Equilibrium favors the

higher pKa value (weaker acid)

Acidity increases from

left to right across periodic table

The strongest conjugate base is the least

stable anion

A larger ionic radius allows us to better distribute the

negative charge over the surface area

distribute charge = more stable anion

increase acidity

The most negative pKa value has the

biggest ionic radius

H-F < H-Cl < H-Br < H-I

Inductive effect

presence of a more electronegative atom has an electron-withdrawing effect

electron density being towards pulled the more electronegative atom

How do you determine acid strength?

element effects

inductive effects

resonance effects

hybridization effects