MCAT Review Sheets PDF

MCAT REVIEW SHEETS Revised 2019 Please send questions or comments to: [email protected]  Contents General Chemistry Biology 1 Atomic Structure 1 1 The Cell...

MCAT REVIEW SHEETS Revised 2019 Please send questions or comments to: [email protected]  Contents General Chemistry Biology 1 Atomic Structure 1 1 The Cell 25 2 The Periodic Table 2 2 Reproduction 26 3 Bonding and Chemical Interactions 3 3 Embryogenesis and Development 27 4 Compounds and Stoichiometry 4 4 Nervous System 28 5 Chemical Kinetics 5 5 Endocrine System 29 6 Equilibrium 6 6 Respiratory System 30 7 Thermochemistry 7 7 Cardiovascular System 31 8 The Gas Phase 8 8 Immune system 32 9 Solutions 9 9 Digestive System 33 10 Acids and Bases 10 10 Kidney and Skin 34 11 Oxidation-Reduction Reactions 11 11 Muscular System 35 12 Electrochemistry 12 12 Genetics and Evolution 36 Organic Chemistry Biochemistry 1 Nomenclature 13 1 Amino Acids, Peptides, and Proteins 37 2 Isomers 14 2 Enzymes 38 3 Bonding 15 3 Nonenzymatic Protein Function & Protein Analysis 39 4 Analyzing Organic Reactions 16 4 Carbohydrate Structure and Function 40 5 Alcohols 17 5 Lipid Structure and Function 41 6 Aldehydes and Ketones I 18 6 DNA and Biotechnology 42 7 Aldehydes and Ketones II 19 7 RNA and the Genetic Code 43 8 Carboxylic Acids 20 8 Biological Membranes 44 9 Carboxylic Acid Derivatives 21 9 Carbohydrate Metabolism I 45 10 N- and P-Containing Compounds 22 10 Carbohydrate Metabolism II 46 11 Spectroscopy 23 11 Lipid and Amino Acid Metabolism 47 12 Separations and Purifications 24 12 Bioenergetics and Regulation of Metabolism 48 i Behavioral Sciences Appendix 1 Biology and Behavior 49 A Organic Chemistry Common Names 73 2 Sensation and Perception 50 B The Heart and Oxygen Transport 74 3 Learning and Memory 51 C Brain 75 4 Cognition, Consciousness, and Language 52 D Endocrine Organs and Hormones 76 5 Motivation, Emotion, and Stress 53 E Lab Techniques 77 6 Identity and Personality 54 F DNA and RNA 78 7 Psychological Disorders 55 G DNA Replication 79 8 Social Processes, Attitudes, and Behavior 56 H The Central Dogma 80 9 Social Interaction 57 I Amino Acids 81 10 Social Thinking 58 J Enzyme Inhibition 82 11 Social Structure and Demographics 59 K Metabolism Overview 83 12 Social Stratification 60 L Glycolysis 84 M Gluconeogenesis 85 N Citric Acid Cycle 86 Physics and Math O Oxidative Phosphorylation 87 1 Kinematics and Dynamics 61 P More Metabolic Pathways 88 2 Work and Energy 62 Q Essential Equations 89 3 Thermodynamics 63 4 Fluids 64 5 Electrostatics and Magnetism 65 6 Circuits 66 7 Waves and Sound 67 8 Light and Optics 68 9 Atomic and Nuclear Phenomena 69 10 Mathematics 70 11 Design and Execution of Research 71 12 Data-Based and Statistical Reasoning 72 ii General Chemistry 1: Atomic Structure A Z X A = Mass number = protons + neutrons Z = Atomic number = # of protons Note: Atomic Weight = weighted average Scientist Contributions Rutherford Model: 1911. Electrons surround a nucleus. Bohr Model: 1913. Described orbits in more detail. Farther orbits = ­Energy AHED Mnemonic Absorb light Photon emitted when n¯, absorbed when n­ Higher potential Heisenberg Uncertainty: It is impossible to know the momentum and Excited position simultaneously. Distant from nucleus Hund’s Rule: e- only double up in orbitals if all orbitals first have 1 e-. Pauli Exclusion Principle: Paired e- must be + " , − ". Diamagnetic vs. Paramagnetic # # Diamagnetic: All electrons are paired ­¯ REPELLED by an external magnetic field Constants Light Energy Paramagnetic: 1 or more unpaired electrons () ­ PULLED into an external magnetic field Avogadro’s Number: 6.022 × 10#F = 1 mol 𝐸= l 𝐸 =h𝑓 Follow Hund’s rule to build the atom’s electron configuration. If 1 or more Planck’s (h): 6.626 × 10 HFI J s 𝑓 = frequency orbitals have just a single electron, the atom is paramagnetic. If there are h = Planck 8 s constant no unpaired electrons, then the atom is diamagnetic. Speed of Light (c) 3.0 × 10K m c = speed of light s Examples: He = 1s2 = diamagnetic and will repel magnetic fields. C = 1s22s22p2 = paramagnetic and will be attracted to magnetic fields. Quantum Numbers Quantum Possible Name What it Labels Notes Number Values Atomic Orbitals on the Periodic Table 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 = s orbital 1 = p orbital 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 = 2n2 Maximum e- in subshell = 4l + 2 Free Radical: An atom or molecule with an unpaired electron. The Aufbau Principle 3D shapes of s, p, d, and f orbitals 1  General Chemistry 2: The Periodic Table Alkali Metals ds lloi Noble Gases Alkaline Earth Metals a et Halogens M Non-metals Post Transition Metals Transition Metals Rare Earth Metal Rows 8A Noble Gases have no affinity for e-. It Zeff IE EA Unchanged would take energy to force an e- on 0 them Pull between nucleus & valence e- Lose e- Gain e- 1st Ionization energies DHrxn < 0 when gaining e- but EA is reported as positive value Of the Noble Gases, only Kr and Xe have an EN Atomic EN Kr Xe Common Electronegativities Size H C N O F Force the atom exerts Exact 2.20 2.55 3.04 3.44 3.98 Only trend this direction on an e- in a bond » 2.0 2.5 3.0 3.5 4.0 Cations < Neutral < Anions 2 General Chemistry 3: Bonding and Chemical Interactions Bond Type According to DEN 0 Nonpolar 0.5 Polar 1.7 Ionic Covalent Bonds covalent covalent Covalent Bond: Formed via the sharing of electrons between two elements of similar EN. Ionic Bonds Bond Order: Refers to whether a covalent bond is a single, Ionic Bond: Formed via the transfer of one or more electrons double, or triple bond. As bond order increases from an element with a relatively low IE to an bond strength­, bond energy­, bond length¯. element with a relatively high electron affinity Nonpolar Bonds: DEN < 0.5. DEN > 1.7. Polar Bonds: DEN is between 0.5 and 1.7. Cation: POSITIVE + Anion: NEGATIVE − Coordinate A single atom provides both bonding electrons. Covalent Bonds: Most often found in Lewis acid-base chemistry. Crystalline Lattices: Large, organized arrays of ions. Intermolecular Forces Sigma and Pi Bonds Formal Charge Hydrogen O-H, N-H, F-H Formal Charge = valence e0 − dots − sticks 1s Strength Dipole-Dipole Dots: Nonbonding e- 1s 1p Sticks: Pair of bonding electrons London Dispersion 1s 2p Note: Van de Waals Forces is a general term that includes Dipole-Dipole forces and London Dispersion forces. H-Bond acceptor Valence Shell Electron Pair Repulsion Theory (VSEPR) H-Bond donor Electronic Geometry: Bonded and lone pairs treated the same. Molecular Shape: Lone pairs take up less space than a bond to another atom. Hybridization e- Groups Around Bonded Lone Electronic Molecular Bond Central Atom Pairs Pairs Geometry Shape Angle sp 2 2 1 0 1 Linear Linear Linear 180° 3 0 Trig Planar sp 2 3 2 1 1 2 Trigonal Planar Bent Linear 120° 4 0 Tetrahedral sp3 4 3 2 1 1 2 3 Tetrahedral Trig Pyramidal Bent Linear 109.5° 5 0 Trigonal Bipyramidal 5 4 1 Trigonal Seesaw 90° sp3d & 3 2 Bipyramidal T-Shaped 120° 2 3 Linear 6 0 Octahedral sp3d2 6 5 4 1 2 Octahedral Square Pyramidal Square Planar 90° 3  General Chemistry 4: Compounds and Stoichiometry Equivalents & Normality Naming Ions Equivalent Mass of an acid that yields 1 mole of or H+ Mass: mass of a base that reacts with 1 mole of H+. For elements (usually metals) that can Fe2+ Iron(II) form more than one positive ion, the Fe3+ Iron(III) !"#$% !$'' charge is indicated by a Roman numeral in GEW = Cu+ Copper(I) !"# () "% *+ parentheses following the name of the element Cu2+ Copper(II) Equivalents = !$'' ", -"!."/01 234 Older method: –ous and –ic to the atoms Fe2+ Ferrous with lesser and greater charge, Fe3+ Ferric Normality = 35 For acids, the # of equivalents respectively 6 (n) is the # of H+ available Cu+ Cuprous from a formula unit. Cu2+ Cupric Molarity = 0"%!$#789 !"# () "% *+ Monatomic anions drop the ending of the H- Hydride name and add –ide F- Fluoride O2- Oxide S2- Sulfide N3- Nitride P3- Phosphide Compound Formulas Oxyanions = polyatomic anions that NO3- Nitrate Empirical: Simplest whole-number ratio of atoms. contain oxygen. NO2 - Nitrite MORE Oxygen = –ate SO42- Sulfate Molecular: Multiple of empirical formula to show LESS Oxygen = –ite exact # of atoms of each element. SO32- Sulfite In extended series of oxyanions, prefixes ClO- Hypochlorite are also used. ClO2- Chlorite MORE Oxygen = Hyper- (per-) ClO3- Chlorate LESS Oxygen = Hypo- ClO4- Perchlorate Polyatomic anions that gain H+ to for HCO3- Hydrogen carbonate or bicarbonate anions of lower charge add the word HSO4 - Hydrogen sulfate or bisulfate Hydrogen or dihydrogen to the front. H2PO4- Dihydrogen phosphate Types of Reactions Combination: Two or more reactants forming one product 2H2 (g) + O2 (g) ® 2H2O (g) Acid Names Decomposition: Single reactant breaks down 2HgO (s) ® 2Hg (l) + O2 (g) -ic: Have one MORE oxygen than -ous. Combustion: Involves a fuel, usually a hydrocarbon, and O2 (g) -ous: Has one FEWER oxygen than -ic. Commonly forms CO2 and H2O CH4 (g) + 2O2 (g) ® CO2 (g) + H2O (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: Elements from two compounds swap places (metathesis) CaCl2 (aq) + 2AgNO3 (aq) ® Ca(NO3)2 (aq) + 2AgCl (s) Neutralization: A type of double-replacement reaction Acid + base ® salt + H2O HCl (aq) + NaOH (aq) ® NaCl (aq) + H2O (l) 4 General Chemistry 5: Chemical Kinetics Units of Rate m Order Rate Law Integrated Rate Law Half Life Constant 0 zeroth order 𝑅=𝑘 [A] = [A]@ − 𝑘 𝑡 [A]@ 𝑀 𝑡^ = _ 2𝑘 𝑠 1 first order 𝑅 = 𝑘 [A] [A] = [A]@ × 𝑒 &A 0 ln (2) 1 𝑡^ = _ 𝑘 𝑠 2 second order 𝑅 = 𝑘 [A]_ 1 1 1 1 = + 𝑘𝑡 𝑡^ = [A] [A]@ _ 𝑘 [A]@ 𝑀𝑠 [A] ln [A] 1 [A] Zeroth Order Reaction First Order Reaction Second Order Reaction 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 Types of Reactions reaction ceases to depend on substrate concentration. Combination: Two or more reactants forming one product. 2H2 (g) + O2 (g) ® 2H2O (g) Equations Decomposition: Single reactant breaks down. 2HgO (s) ® 2Hg (l) + O2 (g) Arrhenius: 𝑘 = 𝐴 × 𝑒 &' )* ( Combustion: Involves a fuel, usually a hydrocarbon, and O2 (g). Definition of Rate: For aA + bB ® cC + dD Commonly forms CO2 and H2O. D[-] D D D CH4 (g) + 2O2 (g) ® CO2 (g) + H2O (g) Rate = − /D0 =− 2D0 = 4D0 = 6D0 Single-Displacement: An atom or ion in a compound is replaced by Rate Law: rate = 𝑘 [A] = [B] ? another atom or ion. Radioactive Decay: [A]0 = [A]@ × 𝑒 A0 Cu (s) + AgNO3 (aq) ® Ag (s) + CuNO3 (aq) Double-Displacement: Elements from two compounds swap places. (metathesis) CaCl2 (aq) + 2AgNO3 (aq) ® Ca(NO3)2 (aq) + 2AgCl (s) Neutralization: A type of double-replacement reaction. Reaction Mechanisms Acid + base ® salt + H2O Overall Reaction: A2 + 2B ® 2AB HCl (aq) + NaOH (aq) ® NaCl (aq) + H2O (l) Step 1: A2 + B ® A2B slow Hydrolysis: Using water to break the bonds in a molecule. Step 2: A2B + B ® 2AB fast A2B is an intermediate Slow step is the rate determining step Arrhenius Equation Arrhenius: 𝑘 = 𝐴 × 𝑒 &' )* ( k = rate constant A = frequency factor Ea = activation energy G R = gas constant = 8.314 Gibbs Free Energy HIJ K T = temp in K ∆G = EO − EO PQR Trends: ­A Þ ­k −∆G = Exergonic ­T Þ ­k (Exponent gets closer to 0. Exponent becomes less negative) +∆G = Endergonic 5  General Chemistry 6: Equilibrium Equilibrium Constant Kinetic (Ea) and Thermodynamic (DG) Control aA + bB ⇌ cC + dD Kinetic Products: HIGHER in free energy than thermodynamic products and can form at lower temperatures. [C]( [D]+ “Fast” products because they can form more Equilibrium Constant (Keq): 𝐾"# = quickly under such conditions. [A]- [B]/ Thermodynamic Products: LOWER in free energy than kinetic products, [C]( [D]+ Reaction Quotient (Qc): 𝑄1 = more stable. Slower but more spontaneous [A]- [B]/ (more negative DG) Exclude pure solids and liquids Reaction Quotient Le Châtelier’s Principle Q < Keq DG < 0, reaction ® If a stress is applied to a system, the system shifts to relieve that applied stress. Q = Keq DG = 0, equilibrium Example: Bicarbonate Buffer CO2 (g) + H2O (l) ⇌ H2CO3 (aq) ⇌ H+ (aq) + HCO3- (aq) Q > Keq DG > 0, reaction ¬ ¯pH Þ ­respiration to blow off CO2 ­pH Þ ¯respiration, trapping CO2 6 General Chemistry 7: Thermochemistry Systems and Processes Temperature (T) and Heat (q) Isolated System: Exchange neither matter nor energy with Temperature (T): Scaled measure of average kinetic the environment. energy of a substance. Closed System: Can exchange energy but not matter with Celsius vs 0°C = 32°F Freezing Point H2O the environment. Fahrenheit: % 25°C = 75°F Room Temp Open system: Can exchange BOTH energy and matter ℉ = ( ℃ ) + 32 & 37°C = 98.6°F Body Temp with the environment. Heat (q): The transfer of energy that results Isothermal Process: Constant temperature. from differences of temperature. Hot Adiabatic Process: Exchange no heat with the environment. transfers to cold. Isobaric Process: Constant pressure. Isovolumetric: Constant volume. Enthalpy (H) (isochoric) Enthalpy (H): A measure of the potential energy of a system found in intermolecular attractions and chemical bonds. States and State Functions Phase Changes: Solid ® Liquid ® Gas: ENDOTHERMIC since gases have more heat energy than liquids and State Functions: Describe the physical properties of an equilibrium liquids have more heat energy than solids. state. Are pathway independent. Pressure, density, temp, volume, enthalpy, internal energy, Gibbs free Gas ® Liquid ® Solid: EXOTHERMIC since these energy, and entropy. reactions release heat. Standard Conditions: 298 K, 1 atm, 1 M Hess’s Law: Enthalpy changes are additive. Note that in gas law calculations, Standard ° Temperature and Pressure (STP) is 0°C, 1 atm. D𝐻-./ from heat of formations ∆𝑯𝐫𝐱𝐧 = ∆𝑯°𝐩𝐫𝐨𝐝𝐮𝐜𝐭𝐬 − ∆𝑯°𝐫𝐞𝐚𝐜𝐭𝐚𝐧𝐭𝐬 ° Fusion: Solid ® liquid ° Freezing: Liquid ® solid D𝐻-./ from bond dissociation energies ∆𝑯°𝐫𝐱𝐧 = ∆𝑯°𝐫𝐞𝐚𝐜𝐭𝐚𝐧𝐭𝐬 − ∆𝑯°𝐩𝐫𝐨𝐝𝐮𝐜𝐭𝐬 Vaporization: Liquid ® gas Sublimation: Solid ® gas Entropy (S) Deposition: Gas ® solid Entropy (S): A measure of the degree to which energy has been spread throughout a system or between a Triple Point: Point in phase diagram where all 3 phases exist. system and its surroundings. @ Supercritical Fluid: Density of gas = density of liquid, no distinction ∆𝑆 = ABC D between those two phases. Standard entropy of reaction ∆𝑺°𝐫𝐱𝐧 = ∆𝑺°𝐟,𝐩𝐫𝐨𝐝𝐮𝐜𝐭𝐬 − ∆𝑺°𝐟,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 D𝑮°𝐫𝐱𝐧 = ∆𝑮°𝐟,𝐩𝐫𝐨𝐝𝐮𝐜𝐭𝐬 − ∆𝑮°𝐟,𝐫𝐞𝐚𝐜𝐭𝐚𝐧𝐭𝐬 Gibbs Free Energy (G) From equilibrium constant Keq D𝑮 = D𝐇 − 𝐓 D𝐒 ° ∆𝐺MNO = −R 𝑇 ln (𝐾VW ) DH DS Outcome From reaction quotient Q ° + + Spontaneous at HIGH temps ∆𝐺MNO = ∆𝐺MNO + R 𝑇 ln (𝑄) Y ∆𝐺MNO = R 𝑇 ln (Z ) + - Non-spontaneous at all temps B[ DG < 0 : Spontaneous - + Spontaneous at all temps DG = 0 : Equilibrium - - Spontaneous at LOW temps DG > 0 : Non-spontaneous Note: Temperature dependent when DH and DS have same sign. 7  General Chemistry 8: The Gas Phase Ideal Gases Real Gases Ideal Gas: Theoretical gas whose molecules occupy negligible space Real gases deviate from ideal behavior at ¯temperature & ­pressure and whose collisions are perfectly elastic. Gases behave At Moderately ­P, ¯V, Real gases will occupy less volume than ideally under reasonably ­temperatures and ¯pressures. or ¯T: predicted by the ideal gas law because the STP: 273 K (0°C), 1 atm particles have intermolecular attractions. 1 mol Gas: At STP 1 mol of gas = 22.4 L At Extremely ­P, ¯V, Real gases will occupy more volume than Units: 1 atm = 760 mmHg = 760 torr = 101.3 kPa = 14.7 psi or ¯T: predicted by the ideal gas law because the particles occupy physical space. Van der Waals 𝑛T 𝑎 d𝑃 + T j (𝑉 − 𝑛𝑏) = 𝑛R𝑇 Equation of State: 𝑉 a corrects for attractive forces Ideal Gas Law b corrects for volume of the particles themselves = 𝑷𝑽=𝒏𝐑𝑻 R = 8.314 >?@ A B DE Density of Gas: r = C = FG Kinetic Molecular Theory DH CH DI CI = (n is constant) GH GI Combined Gas Law: DH GI Avg Kinetic 𝐾𝐸 = S 𝑚 𝑣 T = W 𝐾X 𝑇 𝐾X = 1.38 × 10[TW = V2 = V1( ) ( ) Energy of a Gas: T T A DI GH (𝐾𝐸 ∝ 𝑇) L LH L Avogadro’s Principle: = k or = CI (T and P are constant) C CH I ­T = molecules move FASTER ­molar mass = molecules move SLOWER Boyle’s Law: PV = k or P1V1 = P2V2 (n and T are constant) Root-Mean- C CH CI 3R𝑇 Charles’s Law: = k or =G (n and P are constant) Square Speed: 𝑢^>_ = ` G GH I 𝑀 D DH DI Gay-Lussac’s Law: = k or =G (n and V are constant) Diffusion: The spreading out of particles from [high] ® [low] G GH I Effusion: The mvmt of gas from one compartment to another through a small opening under pressure Other Gas Laws Graham’s Law: bH E = cEI bI H Dalton’s Law: PT = PA + PB + PC + … ¯molar mass = diffuse/effuse FASTER (total pressure from ­molar mass = diffuse/effuse SLOWER partial pressures) Dalton’s Law: PA = XAPT (X = mol fraction) (partial pressure from total pressure) [P]H [P]I Henry’s Law: [A] = kH x PA or DH = DI = kH Diatomic Gases Exist as diatomic molecules, never a stand-alone atom. Includes H2, N2, O2, F2, Cl2, Br2, and I2 Mnemonic: “Have No Fear Of Ice Cold Beer” The 7 Diatomic Gases 8 General Chemistry 9: Solutions Terminology Solubility Rules Solution: Homogenous mixture. Solvent particles surround Soluble solute particles via electrostatic interactions. Na+, K+, NH4+ Solvation or The process of dissolving a solute in solvent. Most Dissolution: dissolutions are endothermic, although dissolution of NO3- gas into liquid is exothermic. Cl-, Br-, I- Except with Pb2+, Hg22+, Ag+ Solubility: Maximum amount of solute that can be dissolved in a SO42- Except with Ca2+, Sr2+, Ba2+, Pb2+, Hg22+, Ag+ 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- Insoluble pair donor. It is held together with coordinate covalent S2- Except with Na+, K+, NH4+, Mg2+, Ca2+, Sr2+, Ba2+ bonds. Formation of complex ions ­solubility. O2- Except with Na+, K+, Sr2+, Ba2+ Solubility in Water: Polar molecules (with +/- charge) are attracted to water molecules and are hydrophilic. Nonpolar OH- Except with Na+, K+, Ca2+, Sr2+, Ba2+ molecules are repelled by water and are hydrophobic. CrO42- Except with Na+, K+, Mg2+, NH4+ Polar = Hydrophilic PO43- & CO32- Except with Na+, K+, NH4+ Nonpolar = Hydrophobic Concentration -."" "01234 % by mass: × 100% -."" "0123506 -014" "01234 Mole Fraction: 𝑋< = 303.1 -014" Molarity: 𝑀 = -014" "01234 1534>" 0? "0123506 Colligative Properties -014" "01234 Molality: 𝐶- = AB 0? "01C463 Can also just be a lowercase m Colligative Properties: Physical properties of solutions that depend on the concentration of dissolved particles but not For acids, the # of equivalents on their chemical identity. # 0? 4F25C.1463" Normality: 𝑁 = (n) is the # of H+ available 1534>" 0? "0123506 from a formula unit. Raoult’s Law: Vapor pressure depression. 𝑃< = 𝑋< 𝑃 IP = Ksp saturated at equilibrium Osmotic Pressure: “Sucking” pressure generated by solutions in IP > Ksp supersaturated, precipitate which water is drawn into solution. Formation or Kf. The equilibrium constant for complex formation. Stability Constant: Usually much greater that Ksp. p=𝑖𝑀R𝑇 𝑖 = vanb t Hoff factor Common Ion ¯solubility of a compound in a solution that already 𝑀 = molar concentration of solute Effect: contains one of the ions in the compound. The 𝑅 = gas constant presence of that ion shifts the dissolution reaction to 𝑇 = temperature the left, decreasing its dissociation. Chelation: When a central cation is bonded to the same ligand in multiple places. Chelation therapy sequesters toxic metals. 9  General Chemistry 10: Acids and Bases Definitions Polyvalence & Normality Arrhenius Acid: Produces H+ (same definition as Brønsted acid) Equivalent: 1 mole of the species of interest. Arrhenius Base: Produces OH- Normality: Concentration of equivalents in solution. Brønsted-Lowry Acid: Donates H+ (same definition as Arrhenius acid) Polyvalent: Can donate or accept multiple equivalents. Brønsted-Lowry Base: Accepts H+ Example: 1 mol H3PO4 yields 3 mol H+. So, 2 M H3PO4 = 6 N. Lewis Acid: Accepts e- pair Lewis Base: Donates e- pair Titrations Note: All Arrhenius acids/bases are Brønsted-Lowry acids/bases, and all Brønsted-Lowry acid/bases are Lewis acids/bases; however, the Half-Equivalence Point: The midpoint of the buffering region, in which half the converse of these statements is not necessarily true. (midpoint) titrant has been protonated or deprotonated. [HA] = [A& ] and pH = pK ) and a buffer is formed. Amphoteric Species: Species that can behave as an acid or a base. Amphiprotic = amphoteric species Equivalence Point: The point at which equivalent amounts of acid and base that specifically can behave as a Brønsted- have reacted. 𝑁' 𝑉' = 𝑁P 𝑉P Lowry acid/base. pH at Equivalence Point: Strong acid + strong base, pH = 7 Polyprotic Acid: An acid with multiple ionizable H atoms. 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 Properties Indicators: Weak acids or bases that display different colors in the protonated and deprotonated forms. The indicator’s Water Dissociation Constant: 𝐾" = 10&'( at 298 K pKa should be close the pH of the equivalence point. 𝐾" = 𝐾) × 𝐾, Tests: Litmus: Acid = red; Base = blue; Neutral = purple pH and pOH: pH = −log [H4 ] [H4 ] = 10&67 Phenolphthalein: pH < 8.2 = colorless; pH > 8.2 = purple pOH = −log [OH& ] Methyl Orange: pH < 3.1 = red; pH > 4.4 = yellow pH + pOH = 14 Bromophenol Blue: pH < 6 = yellow; pH > 8 = blue p scale value approximation: −log (𝐴 × 10&= ) Endpoint: When indicator reaches full color. p value ≈ −(𝐵 + 0. 𝐴) Polyvalent Acid/Base Multiple buffering regions and equivalence points. Strong Acids/Bases: Dissociate completely Titrations: Weak Acids/Bases: Do not completely dissociate Acid Dissociation Constant: 𝐾 = [7FGH ][IJ] p𝐾) = −log (𝐾) ) Titration Setup ) [7I] Base Dissociation Constant: 𝐾 = [KH][G7J] p𝐾, = −log (𝐾, ) , [KG7] Burette p𝐾) + p𝐾, = p𝐾" = 14 Conjugate Acid/Base Pairs: Strong acids & bases / weak conjugate Titrant (strong acid in this example) Weak acids & bases / weak conjugate Neutralization Reactions: Form salts and (sometimes) H2O Analyte / Titrand (weak base in this example) Conical flask Buffers Buffer: Weak acid + conjugate salt Weak base + conjugate salt Titration Curve When titrating a weak base with a strong acid Buffering Capacity: The ability of a buffer to resist changes in pH. Maximum buffering capacity is within 1 pH point of the pKa. Midpoint J Henderson-Hasselbalch pH = pK + log [I ] pOH = pK , ) [7I] Equation: [KH ] Equivalence Point pOH = pK , + log [7G7] 𝑁' 𝑉' = 𝑁P 𝑉P When [A-] = [HA] at the half equivalence point, log(1) = 0, so pH = pKa 10 General Chemistry 11: Oxidation-Reduction Reactions Definitions Balancing via Half-Reaction Method Oxidation: Loss of e- Separate the two half-reactions Balance the atoms of each half-reaction. Start with all elements besides H Reduction: Gain of e- and O. In acidic solution, balance H and O using water and H+. In basic With Respect to Oxidation is GAIN of oxygen solution, balance H and O using water and OH- Oxygen Transfer: Reduction is LOSS of oxygen 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 Oxidizing Agent: Facilitates the oxidation of another both half-reactions compound. Is itself reduced Add the half-reactions, canceling out terms on both sides Reducing Agent: Facilitates the reduction of another Confirm that the mass and charge are balanced compound. Is itself oxidized Oxidation # Rules Any free element or diatomic species = 0 Monatomic ion = the charge of the ion When in compounds, group 1A metals = +1; group 2A metals = +2 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 = -1 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 Net Ionic Equations 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: A type of REDOX reaction in which one element is both (dismutation) 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 11  General Chemistry 12: Electrochemistry Galvanic Cell Electrolytic Cell Electrochemical Cells Emf & Thermodynamics Anode: Always the site of oxidation. It attracts anions. Electromotive force and change in free energy always have OPPOSITE signs. Cathode: Always the site of reduction. It attracts cations. Type of Cell 𝑬°𝐜𝐞𝐥𝐥 DG° Red Cat = Reduction at the Cathode Galvanic + - e- Flow Anode ® Cathode Electrolytic - + Current Flow: Cathode ® Anode ° Concentration 0 0 Galvanic Cells: House spontaneous reactions. -DG, +Emf, +𝐸%#&& (Voltaic) Anode = NEG, Cathode = POS ° Electrolytic Cells: House non-spont reactions. +DG, -Emf, -𝐸%#&& ° ° ° 𝐸%#&& = 𝐸"#$,%-./0$# − 𝐸"#$,-20$# Anode = POS, Cathode = NEG ° ∆𝐺° = −𝑛 F 𝐸%#&& Concentration Specialized form of galvanic cell in which both electrodes are Cells: made of the same material. It is the concentration gradient ∆𝐺° = −R 𝑇 ln (𝐾>? ) between the two solutions that causes mvmt of charge. ∆𝐺 = ∆𝐺° + R 𝑇 ln (𝑄) Rechargeable Can experience charging (electrolytic) and discharging Batteries: (galvanic) states. Faraday constant (F): 96,485 C Lead-Acid: Discharging: Pb anode, PbO2 cathode in a concentrated E sulfuric acid solution. Low energy density. 1C= F 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. Nernst Equation Describes the relationship between the concentration of Cell Potentials species in a solution under nonstandard conditions and the emf. Reduction Potential: Quantifies the tendency for a species to gain e- ° When Keq > 1, then +𝐸%#&& and be reduced. More positive Ered = greater ° tendency to be reduced. When Keq < 1, then -𝐸%#&& ° Standard Reduction Potential: 𝐸"#$ °. Calculated by comparison to the standard When Keq = 1, then 𝐸%#&& =0 hydrogen electrode (SHE). ° GH 𝐸%#&& = 𝐸%#&& − ln (𝑄) ° IJ Standard Electromotive Force: 𝐸%#&&. The difference in standard reduction ° K.KMNO potential between the two half-cells. 𝐸%#&& = 𝐸%#&& − log (𝑄) I ° Galvanic Cells: +𝐸%#&& ° Electrolytic Cells: -𝐸%#&& 12 Organic Chemistry 1: Nomenclature IUPAC Naming Conventions 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. Hydrocarbons and Alcohols Carboxylic Acids & Derivatives 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. Carboxylic Acid Alkene: Contains a double bond. Use suffix -ene. Carboxylic Acid: The highest priority functional group because it Alkyne: Contains a triple bond. Use suffix –yne. contains 3 bonds to oxygen. Alcohol: Contains a –OH group. Use suffix –ol or prefix hydroxy-. Naming: Suffix –oic acid. Alcohols have higher priority than double or triple bonds. Diol: Contains 2 hydroxyl groups. Geminal: If on same carbon Vicinal: If on adjacent carbons Ester Amide Aldehydes and Ketones 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 Aldehyde Ketone substituted. Carbonyl Group: C=O. Aldehydes and ketones both have a carbonyl group. Aldehyde: Carbonyl group on terminal C. Ketone: Carbonyl group on nonterminal C. Primary, Secondary, and Tertiary 1° 2° 3° Alcohols: Amines: 13  Organic Chemistry 2: Isomers Structural Isomers Relative & Absolute Configuration Share only a molecular formula. Relative Configuration: Gives the stereochemistry of a compound in Have different physical and chemical properties. 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. Compounds with atoms connected in the Stereoisomers same order but differing in 3D orientation. Cahn-Ingold-Prelog Priority is given by looking at atoms connected to Priority Rules: the chiral carbon or double-bonded carbons; Chiral Center: Four different groups attached to a central carbon. whichever has the highest atomic # gets highest priority. 2n Rule: 𝑛 = # of chiral centers # of stereoisomers = 23 (Z) and (E) for Alkenes: (Z): Highest priority on same side. Conformational Isomers (E): Highest priority on opposite sides. (R) and (S) for A stereocenter’s configuration is determined by Stereocenters: 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); Anti Gauche Eclipsed horizontal lines come out of the page (wedges). Differ by rotation around a single (s) bond Cyclohexane Equatorial: In the plane of the molecule. Substituents: Axial: Sticking up/down from the molecule’s plane. Configurational Isomers Altering Fischer Switching 1 pair of substituents inverts the Projection: stereochemistry; switching 2 pairs retains stereochemistry. Rotating entire diagram 90° inverts the stereochemistry; rotating 180° retains stereochemistry. 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. 14 Organic Chemistry 3: Bonding Atomic Orbitals & Quantum Numbers Hybridization Quantum Numbers: Describe the size, shape, orientation, and number sp3: 25% s character and 75% p character of atomic orbitals in an element Tetrahedral geometry with 109.5° bond angles Quantum Possible sp2: 33% s character and 67% p character Name What it Labels Notes Number Values Trigonal planar geometry with 120° bond angles n Principal e- energy level or shell number 1, 2, 3, … Except for d-orbitals, the shell # matches the row of the periodic table sp: 50% s character and 50% p character l Azimuthal 3D shape of orbital 0, 1, 2, …, n-1 0 = s orbital Linear geometry with 180° bond angles 1 = p orbital 2 = d orbital Resonance: Describes the delocalization of electrons in 3 = f orbital 4 = g orbital molecules that have conjugated bonds ml Magnetic Orbital sub-type Integers –l ® +l Conjugation: Occurs when single and multiple bonds alternate, " " creating a system of unhybridized p orbitals down ms Spin Electron spin + ,− # # the backbone of the molecule through which p electrons can delocalize Maximum e- in terms of n = 2n2 Maximum e- in subshell = 4l + 2 Molecular Orbitals 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 15  Organic Chemistry 4: Analyzing Organic Reactions Acids and Bases REDOX Reactions Lewis Acid: e- acceptor. Has vacant orbitals or + polarized atoms. Oxidation Number: The charge an atom would have if all its bonds were completely ionic. Lewis Base: e- donor. Has a lone pair of e-, are often anions. Oxidation: Raises oxidation state. Assisted by oxidizing agents. Brønsted-Lowry Acid: Proton donor Oxidizing Agent: Accepts electrons and is reduced in the process. Brønsted-Lowry Base: Proton acceptor Reduction: Lowers oxidation state. Assisted by reducing agents. Amphoteric Can act as either acids or bases, depending on Molecules: reaction conditions. Reducing Agent: Donates electrons and is oxidized in the process. Ka: 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. Chemoselectivity pKa: An indicator of acid strength. pKa decreases down the Both nucleophile-electrophile and REDOX reactions tend to periodic table and increases with EN. act at the highest-priority (most oxidized) functional group. p𝐾# = −log (𝐾# ) One can make use of steric hindrance properties to a-carbon: A carbon adjacent to a carbonyl. selectively target functional groups that might not primarily react, or to protect functional groups. a-hydrogen: Hydrogen connected to an a-carbon. SN 1 SN2 E1 E2 Solvents Nucleophiles, Electrophiles and Leaving Groups Polar Protic Polar Aprotic Nucleophiles: “Nucleus-loving”. Contain lone pairs or p bonds. They have Polar Protic solvents Polar Aprotic solvents ­EN and often carry a NEG charge. Amino groups are Acetic Acid, H2O, DMF, DMSO, common organic nucleophiles. ROH, NH3 Acetone, Ethyl Acetate 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. Substrate Polar Protic Polar Aprotic Strong Small Strong Bulky Leaving Group: Molecular fragments that retain the electrons after heterolysis. The best LG can stabilize additional charge Solvent Solvent Base Base through resonance or induction. Weak bases make good LG. Methyl SN1 Reactions: Unimolecular nucleophilic substitution. 2 steps. In the 1st SN2 SN2 SN2 SN2 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. Primary Rate = 𝑘 [substrate] SN2 SN2 SN2 E2 SN2 Reactions: Bimolecular nucleophilic substitution. 1 concerted step. The Secondary nucleophile attacks at the same time as the LG leaves. The nucleophile must perform a backside attack, which leads to SN1 / E1 SN2 E2 E2 inversion of stereochemistry. (R) and (S) is also changed if the nucleophile and LG have the same priority level. SN2 prefers less-substituted carbons because steric hindrance Tertiary inhibits the nucleophile from accessing the electrophilic SN1 / E1 SN1 / E1 E2 E2 substrate carbon. Rate = 𝑘 [nucleophile] [substrate] 16 Organic Chemistry 5: Alcohols Description & Properties Reactions of Alcohols Alcohols: Have the general form ROH and are named with the suffix –ol. Primary Can be oxidized to aldehydes only by pyridinium If they are NOT the highest priority, they are given the prefix Alcohols: chlorochromate (PCC); they will be oxidized all the way to hydroxy- carboxylic acids by any stronger oxidizing agents Phenols: Benzene ring with –OH groups attached. Named for the relative Secondary Can be oxidized to ketones by any common oxidizing agent position of the –OH groups: Alcohols: Alcohols can be converted to mesylates or tosylates to make them better leaving groups for nucleophilic substitution reactions Mesylates: Contain the functional group –SO3CH3 ortho meta Tosylates: Contain the functional group –SO3C6H4CH3 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 Mesylate Tosylate Electron-donating groups like alkyl groups decrease acidity Aldehydes or ketones can be protected by converting them into acetals or because they destabilize negative charges. EWG, such as EN ketals atoms and aromatic rings, increase acidity because they stabilize negative charges Acetal: A 1° carbon with two –OR groups and an H atom Ketal: A 2° carbon with two –OR groups Acetal Ketal Reactions of Phenols Deprotection: The process of converting an acetal or ketal back to a Quinones: Synthesized through oxidation of phenols. Quinones carbonyl by catalytic acid are resonance-stabilized electrophiles. Vitamin K1 (phylloquinone) and Vitamin K2 (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 17  Organic Chemistry 6: Aldehydes and Ketones I: Electrophilicity and Oxidation-Reduction Description and Properties Nucleophilic Addition Reactions Aldehydes: Are terminal functional groups containing a carbonyl bonded When a nucleophile attacks and forms a bond with a carbonyl carbon, to at least one hydrogen. Nomenclature: suffix –al. In rings, electrons in the p bond are pushed to the oxygen atom. If there is no good they are indicated by the suffix –carbaldehyde. leaving group (aldehydes and ketones), the carbonyl will remain open and is protonated to form an alcohol. If there is a good leaving group Ketones: Internal functional groups containing a carbonyl bonded to (carboxylic acid and derivatives), the carbonyl will reform and kick off the two alkyl chains. In nomenclature, they use the suffix –one leaving group. and the prefix oxo- or keto-. Hydration Rxns: Water adds to a carbonyl, forming a geminal diol. 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. Aldehyde or Gem-diol Alcohols have ­BP than carbonyls due to hydrogen bonding. Ketone Oxidation: Aldehydes and ketones are commonly produced by oxidation Aldehyde + Alcohol: When one equivalent of alcohol reacts with an of primary and secondary alcohols, respectively. Weaker, aldehyde, a hemiacetal is formed. When the same anhydrous oxidizing agents like pyridinium chlorochromate rxn occurs with a ketone, a hemiketal is formed. (PCC) must be used for synthesizing aldehydes, or the reaction will continue oxidizing to a carboxylic acid. 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. 1° Alcohol Aldehyde Oxidation-Reduction Reactions Aldehydes: Aldehydes can be oxidized to carboxylic acids using an oxidizing agent like KMnO4, CrO3, Ag2O, or H2O2. They can be reduced to primary alcohols via hydride reagents (LiAlH4, NaBH4). Ketones: Ketones cannot be further oxidized, but can be reduced to secondary alcohols using the same hydride reagents. Nitrogen + Carbonyl: Nitrogen and nitrogen derivatives react with carbonyls to form imines, oximes, hydrazones, and semicarbazones. Imines can tautomerize to form Common Oxidizing / Reducing Agents