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i

General Chemistry Biology

Behavioral Sciences Appendix

• 1 Atomic Structure 1 1 The Cell Contents
• 2 The Periodic Table 2 2 Reproduction
• 3 Bonding and Chemical Interactions 3 3 Embryogenesis and Development
• 4 Compounds and Stoichiometry 4 4 Nervous System
• 5 Chemical Kinetics 5 5 Endocrine System
• 6 Equilibrium 6 6 Respiratory System
• 7 Thermochemistry 7 7 Cardiovascular System
• 8 The Gas Phase 8 8 Immune system
• 9 Solutions 9 9 Digestive System
• 10 Acids and Bases 10 10 Kidney and Skin
• 11 Oxidation-Reduction Reactions 11 11 Muscular System
• 12 Electrochemistry 12 12 Genetics and Evolution
• 1 Nomenclature 13 1 Amino Acids, Peptides, and Proteins
• 2 Isomers 14 2 Enzymes
• 3 Bonding 15 3 Nonenzymatic Protein Function & Protein Analysis
• 4 Analyzing Organic Reactions 16 4 Carbohydrate Structure and Function
• 5 Alcohols 17 5 Lipid Structure and Function
• 6 Aldehydes and Ketones I 18 6 DNA and Biotechnology
• 7 Aldehydes and Ketones II 19 7 RNA and the Genetic Code
• 8 Carboxylic Acids 20 8 Biological Membranes
• 9 Carboxylic Acid Derivatives 21 9 Carbohydrate Metabolism I
• 10 N- and P-Containing Compounds 22 10 Carbohydrate Metabolism II
• 11 Spectroscopy 23 11 Lipid and Amino Acid Metabolism
• 12 Separations and Purifications 24 12 Bioenergetics and Regulation of Metabolism
• 1 Biology and Behavior 49 A Organic Chemistry Common Names ii
• 2 Sensation and Perception 50 B The Heart and Oxygen Transport
• 3 Learning and Memory 51 C Brain
• 4 Cognition, Consciousness, and Language 52 D Endocrine Organs and Hormones
• 5 Motivation, Emotion, and Stress 53 E Lab Techniques
• 6 Identity and Personality 54 F DNA and RNA
• 7 Psychological Disorders 55 G DNA Replication
• 8 Social Processes, Attitudes, and Behavior 56 H The Central Dogma
• 9 Social Interaction 57 I Amino Acids
• 10 Social Thinking 58 J Enzyme Inhibition
• 11 Social Structure and Demographics 59 K Metabolism Overview
• 12 Social Stratification 60 L Glycolysis
• M Gluconeogenesis
• N Citric Acid Cycle
• O Oxidative Phosphorylation
• 1 Kinematics and Dynamics 61 P More Metabolic Pathways
• 2 Work and Energy 62 Q Essential Equations
• 3 Thermodynamics
• 4 Fluids
• 5 Electrostatics and Magnetism
• 6 Circuits
• 7 Waves and Sound
• 8 Light and Optics
• 9 Atomic and Nuclear Phenomena
• 10 Mathematics
• 11 Design and Execution of Research
• 12 Data-Based and Statistical Reasoning

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.

» 2.0 2.5 3.0 3.5 4.

Atomic

Size

Only trend this direction

Cations < Neutral < Anions

0

Kr

Xe

Rare Earth Metal Rows

General Chemistry 3: Bonding and Chemical Interactions

Hybridization

e-^ Groups

Around

Central Atom

Bonded

Pairs

Lone

Pairs

Electronic

Geometry

Molecular

Shape

Bond

Angle

sp

2

1

0

1

Linear

Linear

Linear

180 °

sp^2

3

2

1

0

1

2

Trigonal Planar

Trig Planar

Bent

Linear

120 °

sp^3

4

3

2

1

0

1

2

3

Tetrahedral

Tetrahedral

Trig Pyramidal

Bent

Linear

109.5°

sp

d

5

4

3

2

0

1

2

3

Trigonal

Bipyramidal

Trigonal Bipyramidal

Seesaw

T-Shaped

Linear

90 °

&

120 °

sp^3 d^2

6

5

4

0

1

2

Octahedral

Octahedral

Square Pyramidal

Square Planar

90 °

Intermolecular Forces Formal Charge

Formal Charge = valence e^0 − dots − sticks Dots: Nonbonding e- Sticks: Pair of bonding electrons

Covalent Bonds

Valence Shell Electron Pair Repulsion Theory (VSEPR)

Hydrogen O-H, N-H, F-H

Dipole-Dipole

Strength London Dispersion

Sigma and Pi Bonds

1 s

1 s 1 p

1 s 2 p

Covalent Bond: Formed via the sharing of electrons between two elements of similar EN.

Bond Order: Refers to whether a covalent bond is a single, double, or triple bond. As bond order increases bond strength , bond energy , bond length ¯.

Nonpolar Bonds: (^) DEN < 0.5.

Polar Bonds: DEN is between 0.5 and 1.7.

Coordinate Covalent Bonds:

A single atom provides both bonding electrons. Most often found in Lewis acid-base chemistry.

Ionic Bonds

Ionic Bond: Formed via the transfer of one or more electrons from an element with a relatively low IE to an element with a relatively high electron affinity DEN > 1.7. Cation: POSITIVE +

Anion: NEGATIVE −

Crystalline Lattices: Large, organized arrays of ions.

Electronic Geometry: Bonded and lone pairs treated the same. Molecular Shape : Lone pairs take up less space than a bond to another atom.

Note: Van de Waals Forces is a general term that includes Dipole-Dipole forces and London Dispersion forces.

H-Bond acceptor

H-Bond donor

Bond Type According to D EN

(^0) Nonpolar covalent

0.5 (^) Polar covalent

1.7 (^) Ionic

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]@

2 𝑘

𝑀

𝑠

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

=

ln ( 2 ) 𝑘

𝑠

2 second order^ 𝑅 = 𝑘 [A]_^1 [A]

=

[A]@^ +^ 𝑘𝑡^ 𝑡^ _^ =^

𝑘 [A]@

𝑀 𝑠

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 6: Equilibrium

Equilibrium Constant

a A + b B ⇌ c C + d D

Equilibrium Constant ( K eq): (^) 𝐾"# =

[C](^ [D]+

[A]-^ [B]/

Reaction Quotient ( Q c ): (^) 𝑄 1 =

[C](^ [D]+

[A]-^ [B]/

Exclude pure solids and liquids

Reaction Quotient

Q < K eq D G < 0, reaction ®

Q = K eq D G = 0, equilibrium

Q > K eq D G > 0, reaction ¬

Le Châtelier’s Principle

If a stress is applied to a system, the system shifts to relieve that applied stress.

Example: Bicarbonate Buffer

CO2 (g) + H 2 O (l) ⇌ H 2 CO3 (aq) ⇌ H+ (aq) + HCO 3 - (aq)

¯pH Þ respiration to blow off CO 2

pH Þ ¯respiration, trapping CO 2

Kinetic ( E a) and Thermodynamic (D G ) Control

Kinetic Products: HIGHER in free energy than thermodynamic products and can form at lower temperatures. “Fast” products because they can form more quickly under such conditions.

Thermodynamic Products: LOWER in free energy than kinetic products, more stable. Slower but more spontaneous (more negative DG)

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=

(𝐾𝐸 ∝ 𝑇)

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 9: Solutions

Saturated solutions are in equilibrium at that particular temperature.

Solubility Product Constant:

Equilibrium expression for something that dissolves. For substance Aa Bb , 𝐾"# = [A]([B]* Ion Product: 𝐼𝑃 = [A]([B]* IP < K sp unsaturated IP = K sp saturated at equilibrium IP > K sp supersaturated, precipitate Formation or Stability Constant:

Kf. The equilibrium constant for complex formation. Usually much greater that K sp. Common Ion Effect:

¯solubility of a compound in a solution that already contains one of the ions in the compound. The presence of that ion shifts the dissolution reaction to the left, decreasing its dissociation. Chelation: When a central cation is bonded to the same ligand in multiple places. Chelation therapy sequesters toxic metals.

% by mass: (^) -.""-."" "0123506^ "01234 × 100%

Mole Fraction: 𝑋< = -014" 303.1 -014""

Molarity: (^) 𝑀 = (^) 1534>"-014" 0?^ "0123506"

Molality: 𝐶- =^

-014" " AB 0? "01C463 Can also just be a lowercase m

Normality: 𝑁 = (^) 1534>"#^ 0?^ 4F25C.1463" 0? "

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

Dilutions: M 1 V 1 = M 2 V 2

Colligative Properties: Physical properties of solutions that depend on the concentration of dissolved particles but not on their chemical identity. Raoult’s Law: (^) Vapor pressure depression. 𝑃< = 𝑋<𝑃<^ ° The presence of other solutes ¯evaporation rate of solvent, thus ¯ P vap. Boiling Point Elevation: ∆𝑇I = 𝑖 𝐾I 𝐶- 𝑖 = ionization factor 𝐾I = boiling point depression constant 𝐶- = molal concentration Freezing Point Depression: ∆𝑇? = 𝑖 𝐾? 𝐶- 𝐾? = freezing point depression constant Osmolarity: The number of individual particles in solution. Example: NaCl dissociates completely in water, so 1 M NaCl = 2 0"-01 1534>

Osmotic Pressure: “Sucking” pressure generated by solutions in which water is drawn into solution.

p = 𝑖 𝑀 R 𝑇

𝑖 = vanbt Hoff factor 𝑀 = molar concentration of solute 𝑅 = gas constant 𝑇 = temperature

Solution: Homogenous mixture. Solvent particles surround solute particles via electrostatic interactions. Solvation or Dissolution:

The process of dissolving a solute in solvent. Most dissolutions are endothermic, although dissolution of gas into liquid is exothermic. Solubility: Maximum amount of solute that can be dissolved in a solvent at a given temp. Molar Solubility: Molarity of the solute at saturation.

Complex Ions: Cation bonded to at least one ligand which is the e- pair donor. It is held together with coordinate covalent bonds. Formation of complex ions solubility. Solubility in Water: Polar molecules (with +/- charge) are attracted to water molecules and are hydrophilic. Nonpolar molecules are repelled by water and are hydrophobic.

Polar = Hydrophilic Nonpolar = Hydrophobic

Terminology

Concentration

Solutions Equilibria

Colligative Properties

Solubility Rules

Soluble Na+^ , K+^ , NH 4 + NO 3 -

Cl-^ , Br-^ , I- SO 4 2-

Except with Pb2+^ , Hg 2 2+, Ag+ Except with Ca2+^ , Sr2+^ , Ba2+^ , Pb2+^ , Hg 2 2+, Ag+

Insoluble S 2 - O 2 -

OH- CrO 4 2-

PO 4 3-^ & CO 3 2-

Except with Na+^ , K+^ , NH 4 +, Mg2+^ , Ca2+^ , Sr2+^ , Ba2+ Except with Na+^ , K+^ , Sr2+^ , Ba2+

Except with Na+^ , K+^ , Ca2+^ , Sr2+^ , Ba2+ Except with Na+^ , K+^ , Mg2+^ , NH 4 +

Except with Na+^ , K+^ , NH 4 +

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

General Chemistry 12: Electrochemistry

Electrochemical Cells

Galvanic Cell Electrolytic Cell

Reduction Potential: Quantifies the tendency for a species to gain e- and be reduced. More positive E red = greater tendency to be reduced.

Standard Reduction Potential: (^) 𝐸"#$^ °^. Calculated by comparison to the standard hydrogen electrode (SHE).

Standard Electromotive Force: 𝐸%#&&^ °^. The difference in standard reduction potential between the two half-cells.

Galvanic Cells: (^) +𝐸%#&&^ °

Electrolytic Cells: (^) - 𝐸%#&&^ °

Cell Potentials

Electromotive force and change in free energy always have OPPOSITE signs.

Type of Cell (^) 𝑬𝐜𝐞𝐥𝐥^ °^ D G °

Galvanic + -

Electrolytic - +

Concentration 0 0

𝐸%#&&^ °^ = 𝐸"#$^ °^ ,%-./0$#− 𝐸"#$^ ° ,-20$#

∆𝐺 ° = −𝑛 F 𝐸%#&&^ °

∆𝐺 ° = −R 𝑇 ln (𝐾>?)

∆𝐺 = ∆𝐺 ° + R 𝑇 ln (𝑄)

Faraday constant (F): 96,485 C

1 C = (^) FE

Emf & Thermodynamics

Describes the relationship between the concentration of species in a solution under nonstandard conditions and the emf.

When K eq > 1, then +𝐸%#&&^ °

When K eq < 1, then -𝐸%#&&^ °

When K eq = 1, then 𝐸%#&&^ °^ = 0

𝐸%#&& = 𝐸%#&&^ °^ − G IHJ ln (𝑄)

𝐸%#&& = 𝐸%#&&^ °^ − K.KMNO I log (𝑄)

Nernst Equation

Anode: Always the site of oxidation. It attracts anions.

Cathode: Always the site of reduction. It attracts cations.

Red Cat = Reduction at the Cathode

e-^ Flow Anode ® Cathode

Current Flow: (^) Cathode ® Anode

Galvanic Cells: (Voltaic)

House spontaneous reactions. -DG, +Emf, +𝐸%#&&^ ° Anode = NEG, Cathode = POS

Electrolytic Cells: (^) House non-spont reactions. +DG, -Emf, -𝐸%#&&^ ° Anode = POS, Cathode = NEG

Concentration Cells:

Specialized form of galvanic cell in which both electrodes are made of the same material. It is the concentration gradient between the two solutions that causes mvmt of charge. Rechargeable Batteries:

Can experience charging (electrolytic) and discharging (galvanic) states.

Lead-Acid: Discharging: Pb anode, PbO 2 cathode in a concentrated sulfuric acid solution. Low energy density.

Ni-Cd: Discharging: Cd anode, NiO(OH) cathode in a concentrated KOH solution. Higher energy density than lead-acid batteries.

NiMH: More common than Ni-Cd because they have higher energy density.

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 3: Bonding

Quantum Numbers: Describe the size, shape, orientation, and number of atomic orbitals in an element

Atomic Orbitals & Quantum Numbers

Bonding Orbitals: Created by head-to-head or tail-to-tail overlap of atomic orbitals of the same sign. ¯energy stable

Antibonding Orbitals: Created by head-to-head or tail-to-tail overlap of atomic orbitals of opposite signs. energy ¯stable

Single Bonds: (^) 1 s bond, contains 2 electrons

Double Bonds: 1 s + 1 p Pi bonds are created by sharing of electrons between two unhybridized p-orbitals that align side-by-side

Triple Bonds: 1 s + 2 p

Multiple bonds are less flexible than single bonds because rotation is not permitted in the presence of a p bond. Multiple bonds are shorter and stronger than single bonds, although individual p are weaker than s bonds

Molecular Orbitals

sp^3 : 25% s character and 75% p character Tetrahedral geometry with 109.5° bond angles

sp^2 : 33% s character and 67% p character Trigonal planar geometry with 120° bond angles

sp: 50% s character and 50% p character Linear geometry with 180° bond angles

Resonance: Describes the delocalization of electrons in molecules that have conjugated bonds

Conjugation: Occurs when single and multiple bonds alternate, creating a system of unhybridized p orbitals down the backbone of the molecule through which p electrons can delocalize

Hybridization

Quantum Number Name^ What it Labels^

Possible Values Notes

n Principal^ e^

• (^) energy level or shell number

1, 2, 3, … Except for d-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

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

Organic Chemistry 6: Aldehydes and Ketones I: Electrophilicity and Oxidation-Reduction

Aldehydes: Are terminal functional groups containing a carbonyl bonded to at least one hydrogen. Nomenclature: suffix –al. In rings, they are indicated by the suffix –carbaldehyde.

Ketones: Internal functional groups containing a carbonyl bonded to two alkyl chains. In nomenclature, they use the suffix –one and the prefix oxo- or keto-.

Carbonyl: A carbon-oxygen double bond. The reactivity of a carbonyl is dictated by the polarity of the double bond. The carbon has a d+^ so it is electrophilic. Carbonyl containing compounds have a BP than equivalent alkanes due to dipole interactions. Alcohols have BP than carbonyls due to hydrogen bonding.

Oxidation: Aldehydes and ketones are commonly produced by oxidation of primary and secondary alcohols, respectively. Weaker, anhydrous oxidizing agents like pyridinium chlorochromate (PCC) must be used for synthesizing aldehydes, or the reaction will continue oxidizing to a carboxylic acid.

1 ° Alcohol Aldehyde

Description and Properties

When a nucleophile attacks and forms a bond with a carbonyl carbon, electrons in the p bond are pushed to the oxygen atom. If there is no good leaving group (aldehydes and ketones), the carbonyl will remain open and is protonated to form an alcohol. If there is a good leaving group (carboxylic acid and derivatives), the carbonyl will reform and kick off the leaving group.

Hydration Rxns: Water adds to a carbonyl, forming a geminal diol.

Aldehyde or Gem -diol Ketone

Aldehyde + Alcohol: When one equivalent of alcohol reacts with an aldehyde, a hemiacetal is formed. When the same rxn occurs with a ketone, a hemiketal is formed.

When another equivalent of alcohol reacts with a hemiacetal (via nucleophilic substitution), an acetal is formed. When the same reaction occurs with a hemiketal, a ketal is formed.

Nitrogen + Carbonyl: (^) Nitrogen and nitrogen derivatives react with carbonyls to form imines , oximes, hydrazones, and semicarbazones. Imines can tautomerize to form enamines.

1 ° Amine Aldehyde or Imine Ketone

Imine Enamine HCN + Carbonyl: Hydrogen cyanide reacts with carbonyls to form cyanohydrins.

Nucleophilic Addition Reactions

Aldehydes: Aldehydes can be oxidized to carboxylic acids using an oxidizing agent like KMnO 4 , CrO 3 , Ag 2 O, or H 2 O 2. They can be reduced to primary alcohols via hydride reagents (LiAlH 4 , NaBH 4 ). Ketones: Ketones cannot be further oxidized , but can be reduced to secondary alcohols using the same hydride reagents.

Oxidation-Reduction Reactions

Oxidizing Agent Reactant Product

PCC 1 ° alcohol

Aldehyde

2 ° alcohol Ketone

KMnO 4 or H 2 Cr 2 O 4

1 ° alcohol Carboxylic Acid

2 ° alcohol Ketone Reducing Agent Reactant Product

NaBH 4

Aldehydes / Ketones 1 ° alcohol 2 ° alcohol

LiAlH 4 (LAH)

Aldehydes Ketones

Carboxylic Acid Ester

1 ° alcohol 2 ° alcohol

1 ° alcohol 2 ° alcohol

Common Oxidizing / Reducing Agents

Organic Chemistry 7: Aldehydes and Ketones II: Enolates

a -carbon: The carbon adjacent to the carbonyl is the a-carbon. The hydrogens attached to the a-carbon are the a -hydrogens.

a -hydrogens: Relatively acidic and can be removed by a strong base. The e-^ withdrawing O of the carbonyl weakens the C-H bonds on a-hydrogens. The enolate resulting from deprotonation can be stabilized by resonance with the carbonyl. Ketones: Ketones are less reactive toward nucleophiles because of steric hindrance and a-carbanion de-stabilization. The presence of an additional alkyl group crowds the transition step and increases energy. The alkyl group also donates e- density to the carbanion, making it less stable.

General Principles

Starts with an aldol addition to create an aldol and create a new C-C bond

Then it undergoes a dehydration to give a conjugated enone (α,β- unsaturated carbonyl) Aldol: Contains both aldehyde and an alcohol. “Ald – ol”

Aldol Nucleophile:

The nucleophile is the enolate formed from the deprotonation of the a-carbon.

Aldol Electrophile:

The electrophile is the aldehyde or ketone in the form of the keto tautomer.

Dehydration: After the aldol is formed, a dehydration reaction (loss of water molecule) occurs. This results in an a,b- unsaturated carbonyl.

Retro-Aldol Reactions:

Reverse of aldol reactions. Catalyzed by heat and base. Bond between a- and b-carbon is cleaved.

Aldol Condensation

Keto / Enol: Aldehydes and ketones exist in both keto form (more common) and enol form (less common).

Tautomers: Isomers that can be interconverted by moving a hydrogen and a double bond. Keto / Enol are tautomers.

Michael Addition: (^) An enolate attacks an a,b-unsaturated carbonyl, creating a bond.

Kinetic Enolate: Favored by fast, irreversible reactions at LOW TEMP, with strong, sterically hindered bases.

Thermodynamic Enolate:

Favored by slower, reversible reactions at HIGH TEMP with weaker, smaller bases.

Enamines: Tautomers of imines. Like enols, enamines are the less common tautomer.

Enolate Chemistry

Organic Chemistry 8: Carboxylic Acids

Amide Synthesis

Carboxylic acids contain a carbonyl and a hydroxyl group connected to the same carbon. They are always terminal groups.

Nomenclature: Suffix – oic acid. Salts are named with the suffix – oate , and dicarboxylic acids are –dioic acids

Physical Properties:

Carboxylic acids are polar and hydrogen bond well, resulting in high BP. They often exist as dimers in solution. Acidity: The acidity of a carb acid is enhanced by the resonance between its oxygen atoms. The acidity can be further enhanced by substituents that are electron-withdrawing, and decreased by substituents that are electron-donating

b -dicarboxylic Acids:

Like other 1,3-dicarbonyl compounds, they have an a- hydrogen that is also highly acidic

a-proton is the most acidic due to resonance

Description and Properties

Oxidation: (^) Carboxylic acids can be made by the oxidation of 1° alcohols or aldehydes or the oxidation of 1° or 2° alkyl groups using an oxidizing agent like KMnO 4 , Na 2 Cr 2 O 7 , K 2 Cr 2 O 7 , or CrO 3.

Nucleophilic Acyl Substitution:

A common reaction in carboxylic acids. Nucleophile attacks the electrophilic carbonyl carbon, opening the carbonyl and forming a tetrahedral intermediate. The carbonyl reforms, kicking off the L.G. Nucleophiles: Ammonia / Amine : Forms an amide. Amides are given the suffix –amide. Cyclic amides are called lactams. Alcohol : Forms an ester. Esters are given the suffix – oate. Cyclic esters are called lactones. Carboxylic Acid : Forms an anhydride. Both linear and cyclic anhydrides are given the suffix anyhydride.

Reduction: (^) Carboxylic acids can be reduced to a 1° alcohol with a strong reducing agent like LiAlH4. Aldehyde intermediates are formed, but are also reduced to 1° alcohols. NaBH 4 is not strong enough to reduce a carboxylic acid

Decarboxylation: b-dicarboxylic acids and other b-keto acids can undergo spontaneous decarboxylation when heated, losing a carbon as CO 2. This reaction proceeds via a six-membered cyclic intermediate

Saponification: Mixing long-chain carboxylic acids (fatty acids) with a strong base results in the formation of a salt we call soap. Soaps contain a hydrophilic carboxylate head and hydrophobic alkyl chain tail. They organize in hydrophilic environments to form micelles. A micelle dissolves nonpolar organic molecules in its interior, and can be solvated with water due to its exterior shell of hydrophilic groups.

Micelle : Polar heads, non-polar tails. The non-polar tails dissolve non-polar molecules such as grease

Reactions of Carboxylic Acids

Nucleophilic Acyl Substitution

Ester Synthesis

Anhydride Synthesis

Carboxylic Acid Synthesis via Oxidation

Reduction of Carboxylic Acid Yields a 1 ° Alcohol

Acid Halide Synthesis

Organic Chemistry 9: Carboxylic Acid Derivatives

Amides: The condensation product of carboxylic acid and ammonia or an amine. Amides are given the suffix –amide. The alkyl groups on a substituted amide are written at the beginning of the name with the prefix N-. Cyclic amides are called lactams , named with the Greek letter of the carbon forming the bond with the N.

N,N-Dimethylpropanamide b-Lactam

Esters: The condensation products of carboxylic acids with alcohols, i.e., a Fischer Esterification. Esters are given the suffix –oate. The esterifying group is written as a substituent, without a number. Cyclic esters are called lactones , named by the number of carbons in the ring and the Greek letter of the carbon forming the bond with the oxygen. Triacylglycerols include three ester bonds between glycerol and fatty acids.

Isopropyl butanoate b-Propiolactone

Anhydrides: The condensation dimers of carboxylic acids. Symmetric anhydrides are named for the parent carb acid, followed by anhydride. Asymmetric anhydrides are named by listing the parent carb acids alphabetically, followed by anhydride. Some cyclic anhydrides can be synthesized by heating dioic acids. Five- or six-membered rings are generally stable.

Ethanoic Ethanoic anhydride Succinic anhydride propanoic anhydride

Amides, Esters, and Anhydrides

In Nu-^ substitution reactions, reactivity is: acid chloride > anhydrides > esters > amides > carboxylate

Steric Hindrance: Describes when a reaction cannot proceed (or significantly slows) because substituents crowd the reactive site. Protecting groups , such as acetals, can be used to increase steric hindrance or otherwise decrease the reactivity of a particular portion of a molecule

Induction: (^) Refers to uneven distribution of charge across a s bond because of differences in EN. The more EN groups in a carbonyl-containing compound, the greater its reactivity Conjugation: Refers to the presence of alternating single and multiple bonds, which creates delocalized p electron clouds above and below the plane of the molecule. Electrons experience resonance through the unhybridized p-orbitals, increasing stability. Conjugated carbonyl-containing compounds are more reactive because they can stabilize their transition states.

Conjugation in Benzene Ring Strain: Increased strain in a molecule can make it more reactive. b-lactams are prone to hydrolysis because they have significant ring strain. Ring strain is due to torsional strain from eclipsing interactions and angle strain from compression bond angles below 109.5°

Reactivity Principles

All carboxylic acid derivatives can undergo nucleophilic substitution reactions. The rates at which they do so is determined by their relative reactivities. Cleavage: Anhydrides can be cleaved by the addition of a nucleophile. Addition of ammonia or an amine results in an amide and a carboxylic acid. Addition of an alcohol results in an ester and a carboxylic acid. Addition of water results in two carboxylic acids.

Transesterification: The exchange of one esterifying group for another on an ester. The attacking nucleophile is an alcohol.

Amides: Can be hydrolyzed to carboxylic acids under strongly acidic or basic conditions. The attacking nucleophile is water or the hydroxide anion.

Fischer Esterification^ Nucleophilic Acyl Substitution Reactions

Synthesis of an Anhydride via Carboxylic Acid Condensation

Organic Chemistry 10: Nitrogen- and Phosphorus-Containing Compounds

Amino Acid: The a-carbon of an amino acid is attached to four groups: an amino group, a carboxyl group, a hydrogen atom, and an R group. It is chiral in all amino acids except glycine.

All amino acids in eukaryotes are L-amino acids. They all have (S) stereochemistry except cysteine , which is (R).

Amphoteric: Amino acids are amphoteric, meaning they can act as acids or bases. Amino acids get their acidic characteristics from carboxylic acids and their basic characteristics from amino groups. In neutral solution, amino acids tend to exist as zwitterions (dipolar ions).

Aliphatic: Non-aromatic. Side chain contains only C and H. Gly, Ala, Val, Leu, Ile, Pro. Met can also be considered aliphatic. Peptide Bonds : Form by condensation reactions and can be cleaved hydrolytically. Resonance of peptide bonds restricts motion about the C-N bond, which takes on partial double bond character. A strong acid or base is needed to cleave a peptide bond. Formed when the N-terminus of an AA nucleophilically attacks the C-terminus of another AA.

Polypeptides: Made up of multiple amino acids linked by peptide bonds. Proteins are large, folded, functional polypeptides.

Amino Acids, Peptides, and Proteins

Biologically, amino acids are synthesized in many ways. In the lab, certain standardized mechanisms are used.

Strecker Synthesis:

Generates an amino acid from an aldehyde. An aldeyhyde is mixed with ammonium chloride (NH 4 Cl) and potassium cyanide. The ammonia attacks the carbonyl carbon, generating an imine. The imine is then attacked by the cyanide, generating an aminonitrile. The aminonitrile is hydrolyzed by two equivalents of water, generating an amino acid.

Gabriel Synthesis: Generates an amino acid from potassium phthalimide, diethyl bromomalonate, and an alkyl halide. Phthalimide attacks the diethyl bromomalonate, generating a phthalimidomalonic ester. The phthalimidomalonic ester attacks an alkyl halide, adding an alkyl group to the ester. The product is hydrolyzed, creating phthalic acid (with two carboxyl groups) and converting the esters into carboxylic acids. One carboxylic acid of the resulting 1,3-dicarbonyl is removed by decarboxylation.

Synthesis of a-Amino Acids

Phosphoric Acid: Sometimes referred to as a phosphate group or inorganic phosphate , denoted Pi. At physiological pH, inorganic phosphate includes molecules of both hydrogen phosphate (HPO 4 2-) and dihydrogen phosphate (H 2 PO 4 -^ ). Phosphoric Acid Structure:

Contains 3 hydrogens, each with a unique p K a. The wide variety in p K a values allows phosphoric acid to act as a buffer over a large range of pH values.

Phosphodiester Bonds:

Phosphorus is found in the backbone of DNA, which uses phosphodiester bonds. In forming these bonds, a pyrophosphate ( PPi , P 2 O 7 4-) is released. Pyrophosphate can then be hydrolyzed to two inorganic phosphates. Phosphate bonds are high energy because of large negative charges in adjacent phosphate groups and resonance stabilization of phosphates.

Organic Phosphates:

Carbon containing compounds that also have phosphate groups. The most notable examples are nucleotide triphosphates (such as ATP or GTP) and DNA.

Phosphorus-Containing Compounds

Gabriel Synthesis of an Amino Acid

Strecker Synthesis of an Amino Acid