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MCAT > Gen Chem > Flashcards

Gen Chem Flashcards

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MCAT flashcards
1
Q

effective nuclear charge (Zeff)

A

a measure of the net positive charge experienced by the outermost electrons (how close valence e- are to the nucleus), increases from left to right for elements in the same period

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2
Q

Principal quantum number

A

increases by one each time as one moves down the elements of a given group

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3
Q

Atomic radius

A

decreases from left to right across a period, increases down a group

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4
Q

Ionization energy

A

the energy required to remove an electron from a gaseous species, increases from left to right across a period and from bottom to top in a group

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5
Q

Electron affinity

A

the energy dissipated by a gaseous species when it gains an electron, increases across a period from left to right and decreases in a group from top to bottom

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6
Q

Electronegativity

A

a measure of the attractive force that an atom will exert on an electron in a chemical bond, increases across a period from left to right and decreases in a group from top to bottom (trends are the same as ionization energy)

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7
Q

Isomers

A

molecules w/ the same molecular formulas but different structures

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8
Q

Stereoisomers

A

have the same chemical formulas and the same atomic connectivity (the same structural backbone), they differ in how these atoms are arranged in space (their wedge-and-dash pattern)

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9
Q

Conformational isomers

A

differ in rotation around single (σ) bonds,

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10
Q

Anti Staggered conformation

A

substituents are at the maximum distance from each other (no overlap of atoms along line of sight), also called anti coformation

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11
Q

Gauche staggered conformation

A

the two largest groups are 60 degrees apart

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12
Q

Eclipsed conformation

A

when the two groups are 120 degrees apart and overlap w/ hydrogens on the adjacent carbon, when two groups atoms and/or groups directly overlap each other with 0 degrees separation (totally eclipsed)

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13
Q

Cis

A

both groups located on the same side

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14
Q

Trans

A

groups are on opposite sides

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15
Q

Configurational isomers

A

can only change from one form to another by breaking and reforming covalent bonds, 2 categories: enantiomers and diastereomers

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16
Q

Chiral

A

mirror image cannot be superimposed on the original object; the molecule lacks an internal plane of symmetry (ex: left hand cannot fit in right-handed glove even though they are essentially identical)

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17
Q

Enantiomers

A

two molecules that are non superimposable mirror images of each other, have the same connectivity but opposite configurations at every chiral center in the molecule

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18
Q

Diastereomers

A

molecules are chiral and share the same connectivity but are not mirror images of each other, have different physical properties

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19
Q

Meso compound

A

a molecule w/ chiral centers that has an internal plane of symmetry (two halves that are mirror images)

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20
Q

Cahn-Ingold-Prelog priority rules

A

priority is assigned based on the atom bonded to the double-bonded carbons: the higher the atomic number, the higher the priority

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21
Q

S

A

when assigning priority the circle is drawn counterclockwise (lowest priority group projects into the page- dashes, #4 doesn’t count when determining the order– true designation is the opposite of what you find when initially assigning priority)

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22
Q

R

A

when assigning priority the circle is drawn clockwise

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23
Q

Incomplete octet

A

elements that are stable w/ fewer that 8 electrons in their valence shell and include hydrogen (2), helium (2), lithium (2), beryllium (4), and boron (6)

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24
Q

Expanded octet

A

any element in period 3 and greater can hold more than 8 electrons, including phosphorus (10), sulfur (12), chlorine (14), and many others

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25
Q

Odd numbers of electrons

A

any molecule w/ an odd number of valence electrons cannot distribute those electrons to give eight to each atom; for example, nitric oxide (NO) has eleven valence electrons

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26
Q

Ionic bonding

A

one or more e- from an atom w/ low ionization energy, typically a metal, are transferred to an atom w/ high electron affinity, typically a nonmetal
–atom that loses e- (cation)
–atom that gains e- (anion)
–difference in electronegativity must be greater than 1.7 on the Pauling scale

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27
Q

Covalent bonding

A

an electron pair is shared btwn 2 atoms, typically nonmetals, that have relatively similar values of electronegativity
–if electron pair is shared equally = nonpolar
–if electron pair is shared unequally = polar

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28
Q

Ionic compound properties

A

-have high melting and boiling points
-many dissolve readily in water and other polar solvents
-in the molten or aqueous state they are good conductors of electricity
-in the solid state they ionic constituents of the compound form a crystalline lattice

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29
Q

Covalent compound properties

A

-have lower melting and boiling points
-bc they do not break down into constituent ions, they are poor conductors of electricity in the liquid state or in aqueous solutions
-the # of shared electron pairs btwn 2 atoms = bond order
-bond length= avg distance btwn 2 nuclei of atoms in a bond (decreases as # of shared e- pairs increases)
-bond energy= the energy required to break a bond by separating its components (the greater the # of pairs of e- shared, the more energy required)

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30
Q

Lewis acid

A

any compound that will accept a lone pair of electrons

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31
Q

Lewis base

A

any compound that will donate a pair of electrons

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32
Q

Formal electrons

A

= valence electrons - dots (lone electron) - sticks (bond)

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33
Q

Electronic geometry

A

describes the spatial arrangement of all pairs of electrons around the central atom, including both the bonding and the lone pairs

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34
Q

Molecular geometry

A

describes the spatial arrangement of only the bonding pairs of electrons

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35
Q

Hydrogen bonds

A
  • H bonds to three highly electronegative atoms– nitrogen, oxygen, or fluorine
    -substances that display H bonding have unusually high boiling points
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36
Q

Formula for determining the number of moles of a sample substance

A

Moles = Mass of sample (g)/ Molar mass (g/mol)

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37
Q

Percent composition formula

A

Percent composition = mass of element in formula/ molar mass x 100%

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38
Q

Combination reaction

A

has two or more reactants forming one produce (*they have more reactants than products)
A + B -> C

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39
Q

Decomposition reactions

A

a single reactant breaks down into two or more products, usually the result of heating , high-frequency radiation, or electrolysis (*have more products than reactants)
A -> B + C

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40
Q

Combustion reaction

A

a special type of reaction that involves a fuel- usually a hydrocarbon- and an oxidant (normally oxygen), commonly these reactants form the two products of carbon dioxide and water (*involves oxidation of a fuel)

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41
Q

Percent yield

A

Percent yield = actual yield/ theoretical yield x 100%

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42
Q

Rate-determining step

A

the slowest step in any proposed mechanism
-the rate of the whole reaction is only as fast as the rate-determining step

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43
Q

The collision theory of chemical kinetics

A

the rate of a reaction is proportional to the number of collisions per second between the reacting molecules
-suggests that not all collisions result in a chem rxn
-an effective collision occurs only if the molecules collide with each other in the correct orientation and with sufficient energy to break their existing bonds and form new ones

44
Q

Collision theory equations

A
  • rate = Z (# of collisions) x f (fraction of collisions that are effective)
  • Arrhenius equation: k = Ae x (-Ea/RT) where k is the rate constant, A is the frequency factor, Ea is the activation energy, R is the ideal gas constant, and T is the temp
45
Q

Free energy change of the reaction (ΔGrxn)

A

the difference between the green energy of the products and the free energy of the reactants
+ΔG = exergonic = energy absorbed
-ΔG = endergonic = energy given off

46
Q

rate law

A

rate = k[A] ^x [B]^y
–the values of k, x, and y must be determined experimentally for a given rxn at a given temp

47
Q

Zero-order reaction

A

one in which the rate of formation of product C is independent of changes in concentrations of any of the reactants
- rate = k[A]^0[B]^0 = k, where k has the units M/s
-it is possible to change the rate by changing the temp or by the addition of a catalyst which lowers the activation energy thereby increasing the value of K
-results in a linear graph

48
Q

First-order reaction

A

has a rate that is directly proportional to only one reactant, such that doubling the concentration of that reactant results in a doubling of the rate of formation of the product
- rate = k[A]^1 or rate = k[B]^1, where k has the units s^-1
-ex: process of radioactive decay
-results in a nonlinear graph

49
Q

Second-order reaction

A

has a rate that is proportional to either the concentrations of two reactants or to the square of the concentration of a single reactant
- rate = k[A]^1[B]^1 or rate = K[A]^2 or rate = k[B]^2, where k has the units of M^-1x^-1
-suggest a physical collision between two reactant molecules
-results in a nonlinear graph

50
Q

Reaction quotient

A

Q<Keq: ΔG<0, reaction proceeds in forward direction
Q=keq: ΔG=0, reaction is in dynamic equilibrium
Q>Keq: ΔG>0, reaction proceeds in the reverse direction

51
Q

Le Chatelier’s Principle

A

states that if a stress is applied to a system, the system shifts to relieve that applied stress
-change in concentration: the system will always react in the direction away from the added species or toward the removed species
-change in pressure (and volume): if the pressure is increased the system will react in the direction that produces fewer moles of gas, and if the pressure is decreased the system will react in the direction that produces more moles of gas
-change in temp: results in a change in Keq, if a rxn is endothermic heat functions as a reactant, if a rxn is exothermic heat functions as a product (the equilibrium shifts in the direction that consumes energy)

52
Q

First law of thermodynamics

A

ΔU (change in internal energy) = Q (heat added to the system) - W (work done by the system)

53
Q

Isothermal process

A

the system’s temperature is constant
-implies that the total internal energy of the system (U) is constant, ΔU= 0 and the first law simplifies to Q = W

54
Q

Adibatic process

A

no heat is exchanged between the system and the environment, the thermal energy of the system is constant throughout the process
-when Q = 0, the first law of thermodynamics simplifies to ΔU = -W (the work done on the system)

55
Q

Isobaric process

A

the pressure of the system is constant
-do not alter the first law of thermodynamics

56
Q

Isovolumetric (isochoric) process

A

experience no change in volume
-bc the gas neither expands nor compresses, no work is performed in the process
-the first law of thermodynamics simplifies to ΔU = Q

57
Q

Standard condition

A

for measuring enthalpy, entropy, and Gibbs free energy changes of a reaction
— 25 degrees celsius (298 K), 1 atm pressure, and 1 M concentrations
-used for kinetics, equilibrium, and thermodynamics problems

58
Q

Standard temperature and pressure (STP)

A

the temperature is 0 degrees celsius (273 K) and pressure is 1 atm
-used for ideal gas calculations

59
Q

Fusion or melting

A

the transition from solid to liquid

60
Q

Solidification, crystallization, or freezing

A

the transition from liquid to solid

61
Q

Sublimation

A

solid goes directly into the gas phase

62
Q

Deposition

A

from the gaseous to the solid phase

63
Q

Triple point

A

the point at which the three phase boundaries (solid-liquid, liquid-gas, and solid-gas) meet

64
Q

Critical point

A

where the phase boundary between the liquid and gas phases terminates (there is no distinction between the two phases)

65
Q

Endothermic

A

processes in which the system absorbs heat (ΔQ >0)

66
Q

Exothermic

A

process in which the system releases heat (ΔQ <0)

67
Q

The equation for heat transfer

A

q = mcΔT (m= mass, c= specific heat of the substance)

68
Q

Enthalpy

A

heat changes at constant pressure

69
Q

Hess’s law

A

enthalpy changes of reactions are additive
-applies to any state function, including entropy and Gibbs free energy

70
Q

The second law of thermodynamics

A

states that energy spontaneously disperses from being localized to becoming spread out if it is not hindered from doing so

71
Q

Entropy

A

is a measure of the spontaneous dispersal of energy at a specific temperature: how much energy is spread out, or how widely spread out energy becomes in a process
ΔS = Q rev (heat that is gained or lost in a reversible process) /T (temp in kelvins), units are J/ mol x K

72
Q

The change in Gibbs free energy

A

the maximum amount of energy released by a process– occurring at constant temperature and pressure–that is available to do work
ΔG = ΔΗ - ΤΔS

73
Q

Exergonic

A

movement toward the equilibrium position, when a system releases energy (ΔG <0)

74
Q

Endergonic

A

movement away from the equilibrium position, (ΔG>0)

75
Q

ΔG

A

-if negative, the reaction in spontaneous
-if positive, the reaction is non spontaneous
-if zero, the system is in a state of equilibrium; ΔH = TΔS

76
Q

Variables of gaseous samples

A

-pressure, volume, temperature, and number of moles (n)
1 atm = 760 mmHg = 760 torr = 101.325 kPa

77
Q

Ideal gas

A

represents a hypothetical gas with molecules that have no intermolecular forces and occupy no volume

78
Q

The ideal gas law

A

shows the relationship among four variables that define a sample of gas:
PV = nRT
-R represents the ideal gas constant, which has a value of 8.21 x 10^-2 L x atm / mol x K, may also be given as 8.314 J/ K x mol which is derived when the SI units of pascal

79
Q

Combined gas law

A

can be used to relate changes in temperature, volume, and pressure of a gas:
P1V1/T1 = P2V2/T2

80
Q

Avogadro’s principle

A

states that all gases at a constant temperature and pressure occupy volumes that are directly proportional to the number of moles of gas present
n/V = k or n1/V1 = n2/V2

81
Q

Boyle’s law

A

for a given gaseous sample held at constant temperature (isothermal conditions), the volume of the gas is inversely proportional to its pressure:
PV = k or P1V1 = P2V2

82
Q

Charles’s Law

A

at constant pressure, the volume of a gas is proportional to its absolute temperature, expressed in kelvin
V/T = k or V1/T1 = V2/T2

83
Q

Gay-Lussac’s law

A

is complementary to Charles’s law, it utilizes the same derivation from the ideal gas, but it relates pressure to temperature instead
P/T = k or P1/T1 = P2/T2

84
Q

Dalton’s law of partial pressures

A

states that the total pressure of a gaseous mixture is equal to the sum of the partial pressures of the individuals
PT= PA + PB + PC +…
-The partial pressure of a gas is related to its mole fraction: PA = XAPT, XA = moles of gas A/ total moles of gas

85
Q

Solvation

A

the electrostatic interaction between solute and solvent molecules
-exothermic = new interactions are stronger than original ones
-endothermic= new interactions are weaker than original ones

86
Q

General solubility rules in aqueous solutions

A
  1. all salts containing ammonium and alkali metal cations are water-soluble
  2. all salts containing nitrate and acetate anions are water-soluble
  3. Halides excluding fluorides, are water-soluble, with the exceptions of those formed with Ag+, Pb2+, and Hg 2+
  4. All salts of the sulfate ion are water-soluble, with the exceptions of those formed with Ca2+, Sr2+, and Ba2+
  5. all metal oxides are insoluble, with the exception of those formed with the alkali metals, ammonium, and CaO, SrO, and BaO, all of which hydrolyze to form solutions of the corresponding metal hydroxides
  6. All hydroxides are insoluble, with the exception of those formed with the alkali metals, ammonium, and Ca2+, Sr2+, and Ba 2+
  7. All carbonates, phosphates, sulfides, and sulfites are insoluble with the exception of those formed with the alkali metals and ammonium
87
Q

Percent composition by mass

A

mass of solute/ mass of solution x 100%

88
Q

mole fraction (X)

A

XA = moles of A/total moles of all species
*the sum of the moles fractions in a system will always equal 1

89
Q

molarity (M)

A

moles of solute/liters of solution

90
Q

colligative properties

A

physical properties of solutions that are dependent on the concentration of dissolved particles but not on the chemical identity of the dissolved particles
-include: vapor pressure depression, boiling point elevation, freezing point depression, and osmotic pressure

91
Q

Raoult’s law

A

accounts for vapor pressure depression caused by solutes in solution
-as solute is added to a solvent, the vapor pressure of the solvent decreases proportionately
-as more solute is dissolved into solvent, the vapor pressure of the solvent decreases
PA = XAPA (PA = vapor pressure of solvent A when solutes are present, XA= the mole fraction the solvent A in the solution, PA is the vapor pressure of solvent A in its pure state)

92
Q

The extent to which the boiling point of a solution is raised relative to that of the pure solvent is given by the formula

A

ΔTb = iKbm (where Tb is the increase in boiling point, I is the Van’t Hoff factor, Kb is a proportionality constant characteristic of a particular solvent, and m is the molality of the solution)

93
Q

The formula for calculating the freezing point depression for a solution is

A

ΔTf = iKfm (Tf is the freezing point depression)

94
Q

Osmotic pressure

A

a sucking pressure generated by solutions in which water is drawn into a solution, it is the amount of pressure that must be applied to counteract this attraction of water molecules for the solution
- π = iMRT

95
Q

Arrhenius acids & bases

A

-acid: will dissociate to form an excess of H+ in solution
-base: will dissociate to form an excess of OH- in solution
—*these behaviors are generally limited to aqueous acids and bases
—acids contain H at the beginning of their formula, bases contain OH at the end of their formula

96
Q

Bronsted-Lowry acids & bases

A

-acid: a species that donates hydrogen ions (H+)
-base: a species that accepts hydrogen ions
-not limited to aqueous solutions

97
Q

Lewis acids & bases

A

-acid: an electron pair acceptor
-base: an election pair donor

98
Q

pH

A

pH = -log[H+] = log 1/[H+]

99
Q

pOH

A

-log[OH-] = log 1/[OH-]

100
Q

Henderson-Hasselbach equation

A

used to estimate the pH or pOH of a buffer solution
pH = pka + log [A-]/[HA]

101
Q

dehydrogenases

A

enzymes that catalyze oxidations in which electrons are lost in the form of hydrogen (dehydrogenation)

102
Q

Oxidizing agents

A

*almost all contain oxygen or another strongly electronegative element (such as a halogen)
- 02, H202, F2, Cl2, Br2, I2, H2SO4, HNO3, NaClO, KMn04, Cr03, Na2Cr2O7, pyridinium chlorochromate (PCC), NAD+, FADH

103
Q

Reducing agents

A

often contain metal ions or hydrides (H+)
-CO, C, B2H6, Sn2+ and other pure metals, Hydrazine, Zn(Hg), Lindlar’s catalyst, NABH4, LiAlH4, NADH, FADH2

104
Q
A
105
Q
A

MCAT (6 decks)

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