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General Chemistry Biology

General Chemistry 1: Atomic Structure

Rutherford Model: 1911. Electrons surround a nucleus.

Bohr Model: 1913. Described orbits in more detail. Farther orbits = Energy Photon emitted when n¯, absorbed when n

Heisenberg Uncertainty: It is impossible to know the momentum and position simultaneously.

Hund’s Rule: e-^ only double up in orbitals if all orbitals first have 1 e-^.

Pauli Exclusion Principle: (^) Paired e- (^) must be + "

,^ −^

" #.

Scientist Contributions

AHED Mnemonic A bsorb light H igher potential E xcited D istant from nucleus

A

X

A = Mass number = protons + neutrons

Z Z = Atomic number = # of protons

Note: Atomic Weight = weighted average

Constants Light Energy

𝐸 = ( l^ ) 𝐸 = h 𝑓

𝑓 = frequency h = Planck^8 s constant c = speed of light

Diamagnetic: ¯

All electrons are paired REPELLED by an external magnetic field

Paramagnetic: 1 or more unpaired electrons PULLED into an external magnetic field Follow Hund’s rule to build the atom’s electron configuration. If 1 or more orbitals have just a single electron, the atom is paramagnetic. If there are no unpaired electrons, then the atom is diamagnetic.

Examples: He = 1s 2 = diamagnetic and will repel magnetic fields. C = 1s 2 2s 2 2p 2 = paramagnetic and will be attracted to magnetic fields.

Diamagnetic vs. Paramagnetic

Quantum Number Name^ What it Labels^

Possible Values Notes

n Principal^ e^

• (^) energy level or shell number

1, 2, 3, … Except for d- and f-orbitals, the shell # matches the row of the periodic table.

l Azimuthal^ 3D shape of orbital^ 0, 1, 2, …, n-1^ 0 =1 =^ sp^ orbitalorbital

2 = d orbital 3 = f orbital 4 = g orbital

ml Magnetic^ Orbital sub-type^ Integers– l ® + l

ms Spin^ Electron spin^ +

"

,^ −^

"

Maximum e-^ in terms of n = 2 n^2 Maximum e-^ in subshell = 4 l + 2

Free Radical: An atom or molecule with an unpaired electron.

Quantum Numbers

Avogadro’s Number: (^) 6. 022 × 10 #F^ = 1 mol

Planck’s ( h ): (^) 6. 626 × 10 HFI^ J•s

Speed of Light ( c ) (^) 3. 0 × 10 K m s

3D shapes of s, p, d, and f orbitals

Atomic Orbitals on the Periodic Table

The Aufbau Principle

General Chemistry 2: The Periodic Table

Alkali Metals

Alkaline Earth Metals

Transition Metals

Post Transition Metals

Metalloids

Non-metals

Halogens

Noble Gases

Z (^) eff

Unchanged

Pull between nucleus & valence e-

IE

Lose e-

1 st^ Ionization energies

EA

Gain e-

DHrxn < 0 when gaining e-

but EA is reported as positive value

8A

Noble Gases have no affinity for e -. It would take energy to force an e -^ on them

EN

Force the atom exerts

on an e-^ in a bond

Of the Noble Gases, only Kr and Xe have an EN

Common Electronegativities

H C N O F Exact (^) 2.20 2.55 3.04 3.44 3.

Atomic

Size

Only trend this direction

Cations < Neutral < Anions

0

Kr

Xe

Rare Earth Metal Rows

General Chemistry 4: Compounds and Stoichiometry

Combination: Two or more reactants forming one product 2H2 (g) + O2 (g) ® 2H 2 O (^) (g)

Decomposition: Single reactant breaks down 2HgO (^) (s) ® 2Hg (^) (l) + O2 (g)

Combustion: Involves a fuel, usually a hydrocarbon, and O2 (g) Commonly forms CO 2 and H 2 O CH4 (g) + 2O2 (g) ® CO2 (g) + H 2 O (^) (g)

Single-Displacement: An atom/ion in a compound is replaced by another atom/ion Cu (^) (s) + AgNO3 (aq) ® Ag (^) (s) + CuNO3 (aq)

Double-Displacement: (metathesis)

Elements from two compounds swap places CaCl2 (aq) + 2AgNO3 (aq) ® Ca(NO 3 )2 (aq) + 2AgCl (^) (s)

Neutralization: A type of double-replacement reaction Acid + base ® salt + H 2 O HCl (aq) + NaOH (aq) ® NaCl (aq) + H 2 O (l)

Equivalents & Normality Equivalent Mass:

Mass of an acid that yields 1 mole of H +^ or mass of a base that reacts with 1 mole of H +.

GEW = !"#$%^ !$''

!"# ()^ "% *+

Equivalents = !$''^ ",^ -"!."/ 234

Normality =^35 6

For acids, the # of equivalents (n) is the # of H +^ available from a formula unit.

Molarity = 0"%!$# !"# ()^ "% *+

Compound Formulas

Types of Reactions

Naming Ions

For elements (usually metals) that can form more than one positive ion, the charge is indicated by a Roman numeral in parentheses following the name of the element

Fe2+^ Iron(II) Fe3+^ Iron(III) Cu+^ Copper(I) Cu2+^ Copper(II)

Older method: –ous and –ic to the atoms with lesser and greater charge, respectively

Fe2+^ Ferrous Fe3+^ Ferric Cu+^ Cuprous Cu2+^ Cupric

Monatomic anions drop the ending of the name and add –ide

H-^ Hydride F-^ Fluoride O2-^ Oxide S2-^ Sulfide N3-^ Nitride P3-^ Phosphide

Oxyanions = polyatomic anions that contain oxygen. MORE Oxygen = –ate LESS Oxygen = –ite

NO 3 -^ Nitrate NO 2 -^ Nitrite SO 4 2-^ Sulfate SO 3 2-^ Sulfite

In extended series of oxyanions, prefixes are also used. MORE Oxygen = Hyper- (per-) LESS Oxygen = Hypo-

ClO-^ Hypochlorite ClO 2 -^ Chlorite ClO 3 -^ Chlorate ClO 4 -^ Perchlorate

Polyatomic anions that gain H+^ to for anions of lower charge add the word Hydrogen or dihydrogen to the front.

HCO 3 -^ Hydrogen carbonate or bicarbonate HSO 4 -^ Hydrogen sulfate or bisulfate H 2 PO 4 -^ Dihydrogen phosphate

Empirical: Simplest whole-number ratio of atoms.

Molecular: Multiple of empirical formula to show exact # of atoms of each element.

Acid Names

-ic: Have one MORE oxygen than -ous.

-ous: Has one FEWER oxygen than -ic.

General Chemistry 5: Chemical Kinetics

Equations

Arrhenius: (^) 𝑘 = 𝐴 × 𝑒& )'*(

Definition of Rate: For^ a A +^ b B^ ®^ c C +^ d D

Rate = − D /[D- 0 ] = − D 2 [D^10 ] = D 4 [D^30 ] = D 6 [^5 D 0 ]

Rate Law: rate^ =^ 𝑘^ [A]=^ [B]? Radioactive Decay: [A] 0 = [A]@ × 𝑒A

Reaction Mechanisms Overall Reaction: A 2 + 2B ® 2AB

Step 1: A 2 + B ® A 2 B slow Step 2: A 2 B + B ® 2AB fast

A 2 B is an intermediate Slow step is the rate determining step

Types of Reactions

Combination: Two or more reactants forming one product. 2H2 (g) + O2 (g) ® 2H 2 O (^) (g)

Decomposition: Single reactant breaks down. 2HgO (^) (s) ® 2Hg (^) (l) + O2 (g)

Combustion: Involves a fuel, usually a hydrocarbon, and O2 (g). Commonly forms CO 2 and H 2 O. CH4 (g) + 2O2 (g) ® CO2 (g) + H 2 O (^) (g)

Single-Displacement: An atom or ion in a compound is replaced by another atom or ion. Cu (s) + AgNO3 (aq) ® Ag (s) + CuNO3 (aq) Double-Displacement: (metathesis)

Elements from two compounds swap places. CaCl2 (aq) + 2AgNO3 (aq) ® Ca(NO 3 )2 (aq) + 2AgCl (^) (s)

Neutralization: A type of double-replacement reaction. Acid + base ® salt + H 2 O HCl (^) (aq) + NaOH (^) (aq) ® NaCl (^) (aq) + H 2 O (^) (l) Hydrolysis: Using water to break the bonds in a molecule.

Arrhenius Equation

Arrhenius: (^) 𝑘 = 𝐴 × 𝑒& )'*(

k = rate constant A = frequency factor E a = activation energy R = gas constant = 8. 314

G HIJ K T = temp in K Trends: (^) A Þ k

T Þ k (Exponent gets closer to 0. Exponent becomes less negative)

Gibbs Free Energy

∆G = EO − EO PQR

−∆G = Exergonic

+∆G = Endergonic

Zeroth Order Reaction First Order Reaction Second Order Reaction

[A] (^) ln [A] 1 [A]

m Order Rate Law Integrated Rate Law Half Life Units of Rate Constant

0 zeroth order^ 𝑅 = 𝑘 [A] = [A]@ − 𝑘 𝑡 𝑡^ _

[A]@

1 first order^ 𝑅 = 𝑘 [A] [A] = [A]@ × 𝑒&A^0 𝑡^ _

ln ( 2 ) 𝑘

2 second order^ 𝑅 = 𝑘 [A]_^1 [A]

[A]@^ +^ 𝑘𝑡^ 𝑡^ _^ =^

𝑘 [A]@

Reaction Order and Michaelis-Menten Curve: At low substrate concentrations, the reaction is approximately FIRST-ORDER. At very high substrate concentration, the reaction approximates ZERO-ORDER since the reaction ceases to depend on substrate concentration.

General Chemistry 7: Thermochemistry

Systems and Processes

Isolated System: Exchange neither matter nor energy with the environment. Closed System: Can exchange energy but not matter with the environment. Open system: Can exchange BOTH energy and matter with the environment. Isothermal Process: Constant temperature. Adiabatic Process: Exchange no heat with the environment. Isobaric Process: Constant pressure. Isovolumetric: (isochoric)

Constant volume.

States and State Functions

State Functions: Describe the physical properties of an equilibrium state. Are pathway independent. Pressure, density, temp, volume, enthalpy, internal energy, Gibbs free energy, and entropy. Standard Conditions: 298 K, 1 atm, 1 M Note that in gas law calculations, Standard Temperature and Pressure (STP) is 0°C, 1 atm. Fusion: (^) Solid ® liquid

Freezing: Liquid ® solid Vaporization: Liquid ® gas Sublimation: Solid ® gas Deposition: (^) Gas ® solid

Triple Point: Point in phase diagram where all 3 phases exist.

Supercritical Fluid: Density of gas = density of liquid, no distinction between those two phases.

Temperature ( T ) and Heat ( q )

Temperature ( T ): Scaled measure of average kinetic energy of a substance. Celsius vs Fahrenheit: ℉ = ( % & ℃ ) + 32

0 °C = 32°F Freezing Point H 2 O 25 °C = 75°F Room Temp 37 °C = 98.6°F Body Temp Heat ( q ): The transfer of energy that results from differences of temperature. Hot transfers to cold.

Enthalpy ( H )

Enthalpy ( H ): A measure of the potential energy of a system found in intermolecular attractions and chemical bonds. Phase Changes: (^) Solid ® Liquid ® Gas: ENDOTHERMIC since gases have more heat energy than liquids and liquids have more heat energy than solids.

Gas ® Liquid ® Solid: EXOTHERMIC since these reactions release heat. Hess’s Law: Enthalpy changes are additive.

D𝐻-./^ °^ from heat of formations ∆𝑯𝐫𝐱𝐧^ °^ = ∆𝑯𝐩𝐫𝐨𝐝𝐮𝐜𝐭𝐬^ °^ − ∆𝑯𝐫𝐞𝐚𝐜𝐭𝐚𝐧𝐭𝐬^ °

D𝐻-./^ °^ from bond dissociation energies ∆𝑯𝐫𝐱𝐧^ °^ = ∆𝑯𝐫𝐞𝐚𝐜𝐭𝐚𝐧𝐭𝐬^ °^ − ∆𝑯𝐩𝐫𝐨𝐝𝐮𝐜𝐭𝐬^ °

Entropy ( S )

Entropy ( S ): A measure of the degree to which energy has been spread throughout a system or between a system and its surroundings. ∆𝑆 =

@ABC D Standard entropy of reaction ∆𝑺𝐫𝐱𝐧^ °^ = ∆𝑺𝐟^ °,^ 𝐩𝐫𝐨𝐝𝐮𝐜𝐭𝐬− ∆𝑺𝐟^ °,𝐫𝐞𝐚𝐜𝐭𝐚𝐧𝐭𝐬

Note: Entropy is maximized at equilibrium.

Gibbs Free Energy ( G )

Gibbs Free Energy ( G ): Derived from enthalpy and entropy.

D𝑮 = D𝐇 − 𝐓 D𝐒

Standard Gibbs free energy of reaction D𝑮𝐫𝐱𝐧^ °^ = ∆𝑮𝐟^ °,^ 𝐩𝐫𝐨𝐝𝐮𝐜𝐭𝐬− ∆𝑮𝐟^ °,𝐫𝐞𝐚𝐜𝐭𝐚𝐧𝐭𝐬

From equilibrium constant K eq ∆𝐺MNO^ °^ = −R 𝑇 ln (𝐾VW)

From reaction quotient Q ∆𝐺MNO = ∆𝐺MNO^ °^ + R 𝑇 ln (𝑄) ∆𝐺MNO = R 𝑇 ln (

Y ZB[^ ) D G < 0 : Spontaneous

D G = 0 : Equilibrium

D G > 0 : Non-spontaneous

Gibbs Free Energy ( G )

D𝑮 = D𝐇 − 𝐓 D𝐒

D H D S Outcome

+ + Spontaneous at HIGH temps

+ - Non-spontaneous at all temps

- + Spontaneous at all temps - - Spontaneous at LOW temps

Note: Temperature dependent when DH and DS have same sign.

General Chemistry 8: The Gas Phase

Ideal Gases

Ideal Gas: Theoretical gas whose molecules occupy negligible space and whose collisions are perfectly elastic. Gases behave ideally under reasonably temperatures and ¯pressures. STP: (^) 273 K (0°C), 1 atm 1 mol Gas: At STP 1 mol of gas = 22.4 L Units: 1 atm = 760 mmHg = 760 torr = 101. 3 kPa = 14. 7 psi

Ideal Gas Law

𝑷 𝑽 = 𝒏 𝐑 𝑻 R^ =^8.^314

=

?@ A

Density of Gas: r =

B C =^

DE FG

Combined Gas Law:

DHCH GH^ =^

DICI GI^ ( n^ is constant) V 2 = V 1 (D DHI ) (G GIH )

Avogadro’s Principle:

L C =^ k^ or^

LH CH^ =^

LI CI^ ( T^ and^ P^ are constant)

Boyle’s Law: PV = k or P 1 V 1 = P 2 V 2 ( n and T are constant)

Charles’s Law:

C G =^ k^ or^

CH GH^ =^

CI GI^ ( n^ and^ P^ are constant)

Gay-Lussac’s Law:

D G =^ k^ or^

DH GH^ =^

DI GI^ ( n^ and^ V^ are constant)

Other Gas Laws

Dalton’s Law: (total pressure from partial pressures)

P T = P A + P B + P C + …

Dalton’s Law: (partial pressure from total pressure)

P A = X A P T ( X = mol fraction)

Henry’s Law: [A]^ =^ k H x^ P A or^

[P]H DH^ =^

[P]I DI^ =^ k H

Kinetic Molecular Theory

Avg Kinetic Energy of a Gas :

𝐾𝐸 = S T 𝑚 𝑣T^ = W T 𝐾X 𝑇 𝐾X = 1. 38 × 10 [TW^ A=

T = molecules move FASTER molar mass = molecules move SLOWER Root-Mean- Square Speed: (^) 𝑢^>_ = `3R𝑇 𝑀

Diffusion: The spreading out of particles from [high] ® [low]

Effusion: The mvmt of gas from one compartment to another through a small opening under pressure Graham’s Law: bH bI^ =^ c

EI EH ¯molar mass = diffuse/effuse FASTER molar mass = diffuse/effuse SLOWER

Real Gases

Real gases deviate from ideal behavior at ¯temperature & pressure At Moderately P , ¯ V , or ¯ T :

Real gases will occupy less volume than predicted by the ideal gas law because the particles have intermolecular attractions. At Extremely P , ¯ V , or ¯ T :

Real gases will occupy more volume than predicted by the ideal gas law because the particles occupy physical space. Van der Waals Equation of State: d𝑃^ +^

𝑛T𝑎

𝑉T^

j (𝑉 − 𝑛𝑏)^ = 𝑛R𝑇

a corrects for attractive forces b corrects for volume of the particles themselves

Diatomic Gases

Exist as diatomic molecules, never a stand-alone atom. Includes H 2 , N 2 , O 2 , F 2 , Cl 2 , Br 2 , and I 2

Mnemonic: “ H ave N o F ear O f I ce C old B eer”

The 7 Diatomic Gases

General Chemistry 10: Acids and Bases

Arrhenius Acid: Produces H+^ (same definition as Brønsted acid) Arrhenius Base: Produces OH-

Brønsted-Lowry Acid: Donates H+^ (same definition as Arrhenius acid)

Brønsted-Lowry Base: Accepts H+

Lewis Acid: Accepts e-^ pair

Lewis Base: Donates e-^ pair

Note: All Arrhenius acids/bases are Brønsted-Lowry acids/bases, and all Brønsted-Lowry acid/bases are Lewis acids/bases; however, the converse of these statements is not necessarily true.

Amphoteric Species: Species that can behave as an acid or a base. Amphiprotic = amphoteric species that specifically can behave as a Brønsted- Lowry acid/base.

Polyprotic Acid: An acid with multiple ionizable H atoms.

Definitions

Properties

Equivalent: 1 mole of the species of interest.

Normality: Concentration of equivalents in solution.

Polyvalent: Can donate or accept multiple equivalents.

Example: 1 mol H 3 PO 4 yields 3 mol H+^. So, 2 M H 3 PO 4 = 6 N.

Polyvalence & Normality

Titrations

Water Dissociation Constant: 𝐾" = 10 &'(^ at 298 K 𝐾" = 𝐾) × 𝐾, pH and pOH: (^) pH = −log [H^4 ] [H^4 ] = 10 &^67 pOH = −log [OH&] pH + pOH = 14 p scale value approximation: (^) −log (𝐴 × 10 &=) p value ≈ −(𝐵 + 0. 𝐴) Strong Acids/Bases: Dissociate completely

Weak Acids/Bases: Do not completely dissociate

Acid Dissociation Constant: (^) 𝐾) = [^7 FGH][IJ] [7I] p𝐾)^ =^ −log^ (𝐾)) Base Dissociation Constant: (^) 𝐾, = [KH][G7J] [KG7] p𝐾,^ =^ −log^ (𝐾,) p𝐾) + p𝐾, = p𝐾" = 14

Conjugate Acid/Base Pairs: Strong acids & bases / weak conjugate Weak acids & bases / weak conjugate

Neutralization Reactions: Form salts and (sometimes) H 2 O

Half-Equivalence Point: (midpoint)

The midpoint of the buffering region, in which half the titrant has been protonated or deprotonated. [HA] = [A&] and pH = pK) and a buffer is formed.

Equivalence Point: The point at which equivalent amounts of acid and base have reacted. 𝑁' 𝑉' = 𝑁P 𝑉P pH at Equivalence Point: Strong acid + strong base, pH = 7 Weak acid + strong base, pH > 7 Weak base + strong acid, pH < 7 Weak acid + weak base, pH > or < 7 depending on the relative strength of the acid and base

Indicators: Weak acids or bases that display different colors in the protonated and deprotonated forms. The indicator’s p K a should be close the pH of the equivalence point.

Tests: Litmus : Acid = red; Base = blue; Neutral = purple Phenolphthalein : pH < 8.2 = colorless; pH > 8.2 = purple Methyl Orange : pH < 3.1 = red; pH > 4.4 = yellow Bromophenol Blue : pH < 6 = yellow; pH > 8 = blue

Endpoint: When indicator reaches full color. Polyvalent Acid/Base Titrations:

Multiple buffering regions and equivalence points.

Buffer: Weak acid + conjugate salt Weak base + conjugate salt

Buffering Capacity: The ability of a buffer to resist changes in pH. Maximum buffering capacity is within 1 pH point of the p K a. Henderson-Hasselbalch Equation:

pH = pK) + log

[IJ] [7I]

pOH = pK, + log [K

H] [7G7]

When [A - ] = [HA] at the half equivalence point, log(1) = 0, so pH = p K a

Buffers

Burette

Conical flask

Titrant (strong acid in this example)

Analyte / Titrand (weak base in this example)

Titration Setup

Midpoint pOH = pK (^) ,

Equivalence Point 𝑁' 𝑉' = 𝑁P 𝑉P

Titration Curve When titrating a weak base with a strong acid

General Chemistry 11: Oxidation-Reduction Reactions

• Any free element or diatomic species = 0
• Monatomic ion = the charge of the ion
• When in compounds, group 1A metals = +1; group 2A metals = +
• When in compounds, group 7A elements = -1, unless combined with an element of greater EN
• H = +1 unless it is paired with a less EN element, then = -
• O = -2 except in peroxides, when it = -1, or in compounds with more EN elements
• The sum of all oxidation numbers in a compound must = overall charge

Definitions

Oxidation # Rules

• Separate the two half-reactions
• Balance the atoms of each half-reaction. Start with all elements besides H and O. In acidic solution, balance H and O using water and H+^. In basic solution, balance H and O using water and OH-
• Balance the charges of each half-reaction by adding e-^ as necessary
• Multiply the half-reactions as necessary to obtain the same number of e-^ in both half-reactions
• Add the half-reactions, canceling out terms on both sides
• Confirm that the mass and charge are balanced

Balancing via Half-Reaction Method

Complete Ionic Equation: Accounts for all of the ions present in a reaction. Split all aqueous compounds into their relevant ions. Keep solid salts intact.

Net Ionic Equation: Ignores spectator ions Disproportionation Reactions: (dismutation)

A type of REDOX reaction in which one element is both oxidized and reduced, forming at least two molecules containing the element with different oxidation states

REDOX Titrations: Similar in methodology to acid-base titrations, however, these titrations follow transfer of charge

Potentiometric Titration: A form of REDOX titration in which a voltmeter measures the electromotive force of a solution. No indicator is used, and the equivalence point is determined by a sharp change in voltage

Net Ionic Equations

Oxidation: Loss of e-

Reduction: Gain of e-

With Respect to Oxygen Transfer:

Oxidation is GAIN of oxygen Reduction is LOSS of oxygen

Oxidizing Agent: Facilitates the oxidation of another compound. Is itself reduced

Reducing Agent: Facilitates the reduction of another compound. Is itself oxidized

Organic Chemistry 1: Nomenclature

Step 1: Find the parent chain, the longest carbon chain that contains the highest-priority functional group.

Step 2: Number the chain in such a way that the highest-priority functional group receives the lowest possible number.

Step 3: Name the substituents with a prefix. Multiples of the same type receive ( di -, tri -, tetra -, etc.). Step 4: Assign a number to each substituent depending on the carbon to which it is bonded.

Step 5: Alphabetize substituents and separate numbers from each other by commas and from words by hyphens.

IUPAC Naming Conventions

Alkane: Hydrocarbon with no double or triple bonds. Alkane = C)H(,)-,)

Naming: Alkanes are named according to the number of carbons present followed by the suffix – ane.

Alkene: Contains a double bond. Use suffix -ene.

Alkyne: Contains a triple bond. Use suffix –yne. Alcohol: Contains a –OH group. Use suffix –ol or prefix hydroxy-. Alcohols have higher priority than double or triple bonds.

Diol: Contains 2 hydroxyl groups. Geminal : If on same carbon Vicinal : If on adjacent carbons

Hydrocarbons and Alcohols

Aldehyde Ketone

Carbonyl Group: C=O. Aldehydes and ketones both have a carbonyl group. Aldehyde: Carbonyl group on terminal C.

Ketone: Carbonyl group on nonterminal C.

Aldehydes and Ketones

Carboxylic Acid

Carboxylic Acid: The highest priority functional group because it contains 3 bonds to oxygen.

Naming: Suffix –oic acid.

Ester Amide

Ester: Carboxylic Acid derivative where –OH is replaced with -OR.

Amide: Replace the –OH group of a carboxylic acid with an amino group that may or may not be substituted.

Carboxylic Acids & Derivatives

1 ° 2 ° 3 °

Alcohols:

Amines:

Primary, Secondary, and Tertiary

Organic Chemistry 2: Isomers

• Share only a molecular formula.
• Have different physical and chemical properties.

Structural Isomers

Chiral Center: Four different groups attached to a central carbon.

2 n^ Rule: 𝑛 = # of chiral centers # of stereoisomers = 23

Conformational Isomers

Anti Gauche Eclipsed

Differ by rotation around a single (s) bond

Cyclohexane Substituents:

Equatorial : In the plane of the molecule. Axial : Sticking up/down from the molecule’s plane.

Configurational Isomers

Enantiomers

Enantiomers: Nonsuperimposable mirror images. Opposite stereochemistry at every chiral carbon. Same chemical and physical properties, except for rotation of plane polarized light.

Optical Activity: The ability of a molecule to rotate plane-polarized light: d- or (+) = RIGHT, l- or (-) = LEFT. Racemic Mixture: (^) 50:50 mixture of two enantiomers. Not optically active because the rotations cancel out. Meso Compounds: Have an internal plane of symmetry, will also be optically inactive because the two sides of the molecule cancel each other out.

Diastereomers

Diastereomers: Stereoisomers that are NOT mirror image.

Cis-Trans: A subtype of diastereomers. They differ at some, but not all, chiral centers. Different chemical and physical properties.

Stereoisomers

Relative Configuration: Gives the stereochemistry of a compound in comparison to another compound. E.g. D and L.

Absolute Configuration: Gives the stereochemistry of a compound without having to compare to other compounds. E.g. S and R. Cahn-Ingold-Prelog Priority Rules:

Priority is given by looking at atoms connected to the chiral carbon or double-bonded carbons; whichever has the highest atomic # gets highest priority.

(Z) and (E) for Alkenes: ( Z ): Highest priority on same side. ( E ): Highest priority on opposite sides.

(R) and (S) for Stereocenters:

A stereocenter’s configuration is determined by putting the lowest priority group in the back and drawing a circle from group 1-2-3. ( R ): Clockwise ( S ): Counterclockwise

Fischer Projection: Vertical lines go to back of page (dashes); horizontal lines come out of the page (wedges).

Altering Fischer Projection:

Switching 1 pair of substituents inverts the stereochemistry; switching 2 pairs retains stereochemistry. Rotating entire diagram 90° inverts the stereochemistry; rotating 180° retains stereochemistry.

Relative & Absolute Configuration

Compounds with atoms connected in the same order but differing in 3D orientation.

Organic Chemistry 4: Analyzing Organic Reactions

Lewis Acid: e-^ acceptor. Has vacant orbitals or + polarized atoms.

Lewis Base: e-^ donor. Has a lone pair of e-^ , are often anions.

Brønsted-Lowry Acid: Proton donor

Brønsted-Lowry Base: (^) Proton acceptor

Amphoteric Molecules:

Can act as either acids or bases, depending on reaction conditions.

K a: Acid dissociation constant. A measure of acidity. It is the equilibrium constant corresponding to the dissociation of an acid, HA, into a proton and its conjugate base.

p K a: An indicator of acid strength. p K a decreases down the periodic table and increases with EN. p𝐾# = −log (𝐾#)

a -carbon: A carbon adjacent to a carbonyl.

a -hydrogen: Hydrogen connected to an a-carbon.

Acids and Bases

Oxidation Number: The charge an atom would have if all its bonds were completely ionic. Oxidation: Raises oxidation state. Assisted by oxidizing agents.

Oxidizing Agent: Accepts electrons and is reduced in the process.

Reduction: Lowers oxidation state. Assisted by reducing agents.

Reducing Agent: Donates electrons and is oxidized in the process.

REDOX Reactions

Nucleophiles: (^) “Nucleus-loving”. Contain lone pairs or p bonds. They have EN and often carry a NEG charge. Amino groups are common organic nucleophiles.

Nucleophilicity: A kinetic property. The nucleophile’s strength. Factors that affect nucleophilicity include charge, EN, steric hindrance, and the solvent.

Electrophiles: “Electron-loving”. Contain a + charge or are positively polarized. More positive compounds are more electrophilic.

Leaving Group: Molecular fragments that retain the electrons after heterolysis. The best LG can stabilize additional charge through resonance or induction. Weak bases make good LG.

SN1 Reactions: Unimolecular nucleophilic substitution. 2 steps. In the 1st step, the LG leaves, forming a carbocation. In the 2nd^ step, the nucleophile attacks the planar carbocation from either side, leading to a racemic mixture of products. Rate = 𝑘 [substrate] SN2 Reactions: Bimolecular nucleophilic substitution. 1 concerted step. The nucleophile attacks at the same time as the LG leaves. The nucleophile must perform a backside attack, which leads to inversion of stereochemistry. ( R ) and ( S ) is also changed if the nucleophile and LG have the same priority level. SN 2 prefers less-substituted carbons because steric hindrance inhibits the nucleophile from accessing the electrophilic substrate carbon. Rate = 𝑘 [nucleophile] [substrate]

Nucleophiles, Electrophiles and Leaving Groups

Both nucleophile-electrophile and REDOX reactions tend to act at the highest-priority (most oxidized) functional group. One can make use of steric hindrance properties to selectively target functional groups that might not primarily react, or to protect functional groups.

Chemoselectivity

Substrate Polar Protic Solvent

Polar Aprotic Solvent

Strong Small Base

Strong Bulky Base Methyl SN 2 SN 2 SN 2 SN 2

Primary SN 2 SN 2 SN 2 E

Secondary SN 1 / E1 SN 2 E2 E

Tertiary SN 1 / E1 SN 1 / E1 E2 E

Solvents

Polar Protic Polar Aprotic

Polar Protic solvents Acetic Acid, H 2 O, ROH, NH 3

Polar Aprotic solvents DMF, DMSO, Acetone, Ethyl Acetate

S N 1 SN 2 E1 E

Organic Chemistry 5: Alcohols

Alcohols: Have the general form ROH and are named with the suffix – ol. If they are NOT the highest priority, they are given the prefix hydroxy -

Phenols: Benzene ring with –OH groups attached. Named for the relative position of the –OH groups:

ortho meta para

• Alcohols can hydrogen bond, raising their boiling and melting points
• Phenols are more acidic than other alcohols because the aromatic ring can delocalize the charge of the conjugate base
• Electron-donating groups like alkyl groups decrease acidity because they destabilize negative charges. EWG, such as EN atoms and aromatic rings, increase acidity because they stabilize negative charges

Description & Properties

Quinones: Synthesized through oxidation of phenols. Quinones are resonance-stabilized electrophiles. Vitamin K 1 ( phylloquinone ) and Vitamin K 2 (the menaquinones ) are examples of biochemically relevant quinones

Quinone

Hydroxyquinones: Produced by oxidation of quinones, adding a variable number of hydroxyl gruops

Ubiquinone: Also called coenzyme Q. Another biologically active quinone that acts as an electron acceptor in Complexes I, II, and III of the electron transport chain. It is reduced to ubiquinol

Reactions of Phenols

Primary Alcohols:

Can be oxidized to aldehydes only by pyridinium chlorochromate (PCC); they will be oxidized all the way to carboxylic acids by any stronger oxidizing agents

Secondary Alcohols:

Can be oxidized to ketones by any common oxidizing agent

Alcohols can be converted to mesylates or tosylates to make them better leaving groups for nucleophilic substitution reactions Mesylates: Contain the functional group –SO 3 CH 3

Tosylates: Contain the functional group –SO 3 C 6 H 4 CH 3

Mesylate Tosylate

Aldehydes or ketones can be protected by converting them into acetals or ketals Acetal: (^) A 1° carbon with two –OR groups and an H atom

Ketal: A 2° carbon with two –OR groups

Acetal Ketal

Deprotection: The process of converting an acetal or ketal back to a carbonyl by catalytic acid

Reactions of Alcohols