CHEM 201 Exam 1

Ch 1-3

Organic molecules

most organic molecules are formed by atoms of 1 or more of these elements: C, H, N, O, S, P, F, Cl, Br, I, Si

contain 1 or more carbon atoms BUT not all carbon compounds are organic for example CO₂ is inorganic

Octet Rule

8 valence electrons (filled shell of electrons = stability)

expanded octets (8 or more valence electrons) happen with elements in 3rd row and higher

H is the exception with it having 2 valence electrons

Covalent bonds

pairs of electrons are shared between atoms

Polarity

Non-polar vs polar

non-polar example is C and H

polar example is H₂O

Lewis structures

includes all covalent bonds, lone pairs of electrons and charges present in a molecule of ion

H — H, hydrogen only has 1 chemical bond

H⁺ = no electrons and is just called a proton

H^- = 2 electrons, called hydride

in order to get the total number of valence electrons of atoms in the molecule/ion add up number of valence electrons for each (add electrons if negative charge or subtract if positive charge)

if more electrons need to be added, first look at outer atoms when and and to check if the octet rule is satisfied then look at central atom to see if octet rule is satisfied

Formal Charge = # of valence electrons - non bonding electrons - 1/2(bonding electrons) or valence electrons - nonbonding electrons - bonds = formal charge

Formal Charge

formal charge = # of valence electrons - non bonding electrons - 1/2(bonding electrons) or # of valence electrons - nonbonding electrons - bonds = formal charge

Hydrogen: 1 bond + 0 lone pairs = neutral, 0 bonds + 0 lone pairs = +1, 0 bonds + 1 lone pair = -1

Carbon: 4 bonds + 0 lone pairs = neutral, 3 bonds + 0 lone pairs = +1, 3 bonds + 1 lone pair = -1

Oxygen: 2 bonds + 2 lone pairs = neutral, 3 bonds + 1 lone pair = +1, 1 bond and 3 lone pair = -1

Nitrogen: 3 bonds + 1 lone pair = neutral, 4 bonds and 0 lone pairs = +1, 2 bonds + 2 lone pairs = -1

Halogens are neutral with 1 chemical bond and 3 lone pairs

Lewis Dot Formulas

shows all different elements

shows all connectivity

Condensed Formula (C_xH_w)

Ex: C₃H₈

no information on connectivity

Condensed structural formulas

Ex: (CH₃)₂CHCH₂OH

• (CH₃)₂ are both attached to the same central atom which is the C next to it in the formula

shows you the order, gives a better idea on how things are connected

no individual bonds are shown

if there are 2 or more identical groups a parenthesis and subscript are used

when multiple bonds are part of a molecule they’re represented as in lewis structures

Other ways to write formulas

using R which shows 1 or more C’s/something else

R — CHO, R — CO₂H

CHO = aldehyde, 2 bonds between C + O and 1 bond between C + H

CO₂H = carboxylic acid, on O has 2 bonds between C + O, other oxygen has 1 bond between C + O and 1 bond between O + H

Line angle formulas

there are carbon atoms where the line starts and ends

there is carbon atom at the intersection between lines

hydrogen is the ONLY atom implied with carbon

if a chain ends on CH₂ then there’s a double bonds

hydrogen needs to be shown when attached to atoms other than C

Molecular geometry

electron domain: anything attached to central atom (bonds and lone pairs)

E = lone pair of electrons

molecular geometry does not look at lone pair of electrons

solid = towards you and below x axis, dashed = away from you

AX₂: 2 e^- domains, linear e^- domain geometry, linear molecular geometry, ex: CO₂

AX₃: 3 e^- domains, trigonal planar e^- domain geometry, trigonal planar molecular geometry, ex: COH₂

AX₂E: 3 e^- domains, trigonal planar e^- domain geometry, bent molecular geometry, ex: SO₂

AX₄: 4 e^- domains, tetrahedral e^- domain geometry, tetrahedral molecular geometry, ex: CH₄

AX₃E: 4 e^- domains, tetrahedral e^- domain geometry, trigonal pyramidal molecular geometry, ex: NH₃

AX₂E₂: 4 e^- domains, tetrahedral e^- domain geometry, bent molecular geometry, ex: H₂O

Molecular Orbital (MO) Theory

with A — A bonds (2 electrons shared between them)

• when electrons move synchronously as a wave then it’s a favorable interaction for example + + + = big + bonding MO or - + - = big - bonding MO, it’s a larger orbital (more stable)

• when those 2 electrons move asynchronously then it is unfavorable and results in an antibonding MO, there is also a node in between which is a space where finding electron possibility is 0 or nearly 0

*for waves, above x axis = + and below = -

Atomic Orbitals chart

• 1s—1s sigma bond

• 1s—2p sigma bond

• 2p—2p sigma bond

• 2p—2p pi bond (very strong orbital)

sigma is stronger than pi

Hybrid Orbitals

*atoms internally rearrange to get ready to form MO

p orbitals lose energy and gains in order to match each other

sp³: s+p+p+p → 4 sp³ orbitals, all sigma bonds (hybrid forms sigma bond)

• 25% s character

• 75% p character

• geometry tetrahedral: 109.5 deg

sp^2: s+p+p → 3 sp² orbitals (3sp² and 1 2p orbital)

• 33.3% s character

• 66.7% p character

• trigonal planar: 120 deg

*when p orbital next to each other and parallel then pi bond forms

sp: s+p → 2 sp orbitals (2 sp and 2 2p orbitals)

• 50% s character

• 50% p character

• linear: 180 deg

Covalent bonds

can be formed when EN is same or similar

Electronegativity

non-polar = 0 - 0.4

polar = > 0.4 and < 2.0

ionic = >2.0

the longer the bond the more polar it is

Dipole moment measurement

if μ = 0 then nonpolar (ex: different directions of vectors but same magnitude)

if μ > 0 = polar

μ = is measured in Debyes (D)

amount of charge at either end of dipole + distance between charges = measurement

Molecular polarity

determined by bond dipole moments and molecular geometry

dipole moment of a molecule = vector sum of the individual bond dipole moments (magnitude and direction)

analyzed based off of the central atom

vectors cancel out (when sp³ with all same atoms attached then it’ll count, ex: CH₄ and CF₄)

any hydrocarbon is nonpolar (C and H only)

lone pairs also contribute to dipole moment

Resonance

when a molecule has 2 or more valid lewis-dot structures

electrons aren’t static, resonance structures can tell you where electrons are more concentrated

when resonance structures have the same energy and stability then they are equivalent

structures are enclosed in brackets

separated by a single double-headed arrow

electrons are the only particles that can participate

multiple bonds and lone pairs are the ones that are most commonly shifted

more stable structures are major resonance contributors and the less stable ones are minor contributors

overall charge remains constant

good contributors have all octets satisfied, as many bonds as possible and as little charge separation as possible

negative charges are more stable on more EN atoms

resonance hybrid = combination of all resonance contributors

• partial bond is shown with dotted line

• more EN atom should have negative charge

• less EN atom should have positive charge

• ideally want to see full octets

• can only move double bonds if the neighbor has a positive charge or double bond or lone pair (from more negative side to more positive)

• single electrons aka radicals also participate

*sigma bonds don’t change, they’re constant

*don’t move electrons to an sp³ atom only sp² and sp

*a ring involved with pi bonds can always participate in resonance

*the ket to moving bonds and lone pairs is to have a p orbital available

*for hybridization, look at the one with more pi bonds to determine that hybridization of atoms in resonance structures (even if there are sp³ but also sp² then must put down as sp² for that atom)

hydrogen donor = ex: OH, NH (where hydrogen is together with O, N or F)

hydrogen acceptor: just O, N, F

Intermolecular Forces

attractive force between molecules: dipole-dipole, hydrogen bonds, and london dispersion forces

intermolecular forces are weaker than chemical bonds and have a direct effect on physical properties like melting point, boiling point and solubility

solubility increases with strong molecular forces

LDF < d-d < H-bd

Dipole-dipole: happens when a molecule is polar and has permanent dipole

• Ex: H — Cl, CHCL₃

Hydrogen bonding interactions: very strong dipole-dipole

• Ex: H — O (like water), H — N, and H — F

• doesn’t have to be the same molecule it can be H — O — H and H — O — CH₃ and because of hydrogen bonding the O atoms of the other molecule will want to bond to the H of the other molecule

• H — O — H and CH₃CH₂CH₂CH₂CH₃ (pentane) is not soluble, remember hydrocarbons aren’t polar

• polar dissolves polar and nonpolar dissolves nonpolar

• partial positive charge on H strong affinity for nonbonding electrons

London dispersion forces: nonpolar molecules, weaker than H-bd and d-d

• Ex: hydrocarbons, CCL₄, Br₂

• induced dipoles: when another nonpolar moleculae is close to one that has dipole for a fraction of a second then it also will have a dipole-dipole momet

• dipoles are present for a fraction of a second

• surface area: linear structures have increase surface area contact with each other as opposed to branched

  • greater surface area = stronger LDF and therefore higher BP

  • depends on close surface contact of 2 molecules

• temporary dipole moments induced in a molecule by other nearby molecules (the electrons are not always evenly distributed)

*if a nonpolar region grows, that part doesn’t interact with water and makes the molecule/compound less soluble

Arrhenius acids and bases

arrhenius acid: H₃O⁺ is formed in aqueous solution

arrhenius base: OH^- is formed in aqueous solution

Bronsted-Lowry acids and bases

bronsted-lowry acid: able to donate H⁺

bronsted-lowry base: able to accept H⁺

NH₃ + HCl → NH₄⁺ + Cl^-

• from left to right: base, acid, conjugate acid, conjugate base (left side of reaction is acid and base, right side is for conjugates)

• the stronger an acid the weaker its conjugate base; the weaker the acid, the stronger its conjugate base

Strong acids: HA + H₂O → A^- + H₃O⁺

Acidity: HA + H₂O → ← A^- + H₃O⁺

• acid-dissociation constant: K_a = [A^-][H₃O⁺]/[HA]

  • lower pK_a and higher K_a = stronger HA (acid)

  • pK_a = -log(K_a)

  • if A^- is stable then HA is a good acid, increases/higher acidity

effects on acidity: electronegativity, size (polarizability), hybridization, resonance and inductive effect

Electronegativity acidity effect

when comparing atoms in a row/across a period the more electronegative element bears a negative charge more easily and will have more stable conjugate base

higher EN = more stable

lower EN = less stable

Size acidity effect

when looking at elements down a column/group

the negative charge of an anion is more stable if it is spread over a larger region of space

the bigger the anion = the more stable because it’s more polarizable

again look at conjugate base

Hybridization acidity effect

more s character, more electronegative atom

the greater the s character the better because it stabilizes the negative charge better since it’s closer to the nucleus

sp³ < sp² < sp

Resonance acidity effect

the negative charge of a conjugate base may be delocalized over 2 or more atoms by resonance, which is often the dominant effect helping to stabilize an anion

when the conjugate base has resonance then it’s more stable

Inductive effect (on acidity)

electron withdrawing atoms and groups can also stabilize a conjugate base through the sigma bonds of the molecule

the effect that substitutes have on the stability of the conjugate base

Electron withdrawing group (EWG): NO₂, SO₃H, CN, carbonyl, and halogens

• pulls electron density towards their center so it makes the conjugate base more stable and a stronger acid since not all of the negative charge is distributed over 1 atom

• look at # of EWG, distance between the atom and EWG (less sigma bonds = closer which = stronger)

Electron donating group (EDG): alkyl groups, lowers acidity

Lewis acids and bases

lewis acids: species able to accept pairs of electrons

• metals (for the most part) because they have available orbitals

• also Boron (B)

lewis bases: species able to donate pairs of electrons

• atoms with lone pairs

lewis base attacks the lewis acid

arrow goes from base to acid B - - → A⁺ or Nu - - → E⁺ (nucleophile to electrophile)

• nucleophile = something with lone pairs (neutral or negative) wanting to donate electrons

atom is electron rich when it has a negative charge

in a molecule a lewis base attacks where somewhere will want to receive an electron so a carbon would want to receive an electron when the EN of the attached atom like Br causes the carbon to be delta positive and then Br would leave

Functional groups

a group of atoms within a molecule that has a characteristic chemical behavior (atoms and how they’re connected and their function/reactivity)

Functional groups to memorize: alkene, alkyne, arene, allene, alcohol, ether, carbonyl, amine (1°), 2° amine, 3° amine, phenol, aldehyde, ketone, carboxylic acid, ester, amide, acid chloride, acid anhydride, nitrile, and nitro

Alkanes

hydrocarbons (C and H only)

only sigma bonds are present

sp³ carbons (no double or triple bonds)

non polar

C_nH_2n+2

straight-chain alkanes

branched alkanes (have substituents which is the branched part)

• branched = when there’s carbons outside the main chain/longest straight chain

Types of isomers

isomers: different compounds with the same condensed formula

• Constitutional or structural: different connectivity

• Stereoisomers

Types of sp³ carbons based on the number of substituents

1°: carbon with only 1 carbon attached (and 3 hydrogens)

2°: carbon with 2 carbons attached (and 2 hydrogens)

3°: carbon with 3 carbons attached (and 1 hydrogen)

4°: carbon with 4 carbons attached (no hydrogens)

Nomenclature for straight chain alkanes

1 C = methane CH₄

2 C = ethane C₂H₆

3 C = propane C₃H₈

4 C = butane C₄H₁₀

5 C = pentane C₅H₁₂

6 C = hexane C₆H₁₄

7 C = heptane C₇H₁₆

8 C = octane C₈H₁₈

9 C = nonane C₉H₂₀

10 C = decane C₁₀H₂₂

11 C = undecane C₁₁H₂₄

12 C = dodecane C₁₂H₂₆

13 C = tridecane C₁₃H₂₈

20 C = icosane C₂₀H₄₂

30 C = triacontane C₃₀H₆₂

Common names:

• butane: n-butane if just straight chain, isobutane if y-shaped

• pentane: n-pentane if just straight chain, isopentane if there’s a y-shape, neopentane when it kind of looks like an x

• y shape = iso

IUPAC rules

1. find longest carbon chain (main chain)

2. number carbons from the end closest to the first substituent (if 2 substituents are the same distance from both ends then start alphabetically or if there are more than 2 substituents then start with whatever’s closest to the next one)

3. name substituents (number—substituent) (if they’re the same substituent then separate numbers with comma)

4. add the root name (the name of the hydrocarbon chain)

*repeat of substituent prefixes: 2x = di- 3x = tri- 4x = tetra-, 5x = penta-, 6x = hexa-, 7x = hepta-

Alkyl groups

alkyl groups are formed when a hydrogen atom is removed from an alkane that are parts of larger compounds

when a CH₃ is attached to the main chain it had to lose a hydrogen so it isn’t methane but it’s methyl

when it’s a CH₂CH₃ then it isn’t a ethane it’s ethyl

3 carbons alkane group = propane but for alkyl group:

• straight chain = propyl

• y-shaped = isopropyl

4 carbons alkane group = butane but for alkyl group:

• straight chain = butyl

• y-shaped = isobutyl

• carbon attached to main chain is 2° = sec-butyl or s-butyl (don’t count the carbon on the main chain)

• carbon attached to the main chain is 3° = tert-butyl or t-butyl