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MCAT Prep > General Chemistry > Flashcards

General Chemistry Flashcards

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MCAT flashcards
1
Q

Principle quantum number

A

(n) – larger the value = higher energy level and radius of electron shell
max # of e- in a shell = 2n^2

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2
Q

Azimuthal Quantum number

A 
(l) angular momentum 
l=0 --- s
l=1 --- p
l=2 --- d
l=3 --- f
max # of electrons within subshell = 4l + 2
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3
Q

Hund’s Rule

A

within electron sub shell, e- will only begin pairing up once all orbitals are half-filled

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4
Q

Paragmagnetic

A

unpaired e- orient to align with magnetic field

–meaning that a magnetic field will cause parallel spins in unpaired electrons and cause an attraction

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5
Q

Diamagnetic

A

paired electrons will be slightly repelled by a magnetic field

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6
Q

Ionization Energies (Periodic Table)

A

Increasing moving to the right and up

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7
Q

Electron Affinity (Periodic Table)

A

Increasing moving to the right and up (just like ionization energy)

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8
Q

Atomic and Ionic Radii (Periodic Table)

A

increases moving to the left and down

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9
Q

Alkali Metals (Periodic Table)

A

Group 1

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10
Q

Alkaline Earth Metals (Periodic Table)

A

Group 2

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11
Q

Halogens (Periodic Table)

A

Group 17

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12
Q

Noble Gases (Periodic Table)

A

Group 18

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13
Q

Stable with incomplete octet (5)

A 
Hydrogen H (2)
Helium He (2) 
Lithium Li (2)
Beryllium Be (4)
Boron B (6)
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14
Q

Stable with expanded octet

A 
elements in group 3 (or greater)
Phosphorus P (10) 
Sulfur S (12)
Chlorine Cl (14)
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15
Q

Ionic Bonds

A

one low ionization with atom of high electron affinity
typically between metal (become cation) and nonmetal (becomes anion)
Physical properties:
very high melting and boiling points
dissolved readily in water and other polar solvents
great conductors

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16
Q

covalent bonding

A

e- pair shared between 2 atoms, how they are shared determines polarity
Typically between two metals
Physical Properties:
weak intermolecular interactions
tend to have lower boiling and melting points

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17
Q

Coordinate Covalent Bonds

A

both shared e- originate from the same atom

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18
Q

Lewis Acid

A

any compound that will accept lone pair of e

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19
Q

lewis Base

A

any compound that will donate a lone pair of e

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20
Q

electronic vs. molecular geometry

A

electronic: spatial arrangement of electron pairs around central atom
molecular: spatial arrangement of only bonding pairs (coordination # determines molecular geometry)

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21
Q

Intermolecular Forces (weakest to strongest) 3

A

Dispersion forces then dipole-dipole interactions then H bonding

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22
Q

Dispersion Forces

A

attractive and repulsive interactions (rapidly)

Important key element for why noble gases can liquify

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23
Q

Dipole-Dipole interactions

A

oppositely charged ends orient to be closer together

longer duration of attraction

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24
Q

Hydrogen Bonding

A

strong
no sharing or transferring of electrons (no actual bond)
Only exists in HIGHLY electronegative atoms
high boiling points

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25
Q

Combination Reactions

A

2 or more reactants form 1 product

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26
Q

Decomposition Reactions

A

one reactant forms 2 or more products

usually due to heating, high freq. radiation, electrolysis

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27
Q

Combustion reaction

A 
involves oxidation (hydrocarbon fuels, or S, or sugars) 
products: CO2, H2O is almost always present
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28
Q

Single Displacement Reaction

A

1 atom or ion in a compound is replaced with another

RedOx Rxns

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29
Q

Double Displacement Reaction

A

elements of 2 compounds swap to form 2 new compounds

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30
Q

Neutralization Reactions

A

type of double displacement

acid reacts with base to produce a salt

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31
Q

Three types of systems

A

isolated, closed, open

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32
Q

Isolated System

A

cannot exchange energy (heat or work) OR matter with surroundings

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33
Q

Closed System

A

can exchange energy (heat or work), but NOT matter with surroundings

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34
Q

Open System

A

Can exchange BOTH energy (heat and work) and matter with surroundings

35
Q

First law of thermodynamics

A

deltaU = Q - W

36
Q

Four types of Processes

A

Isothermal Process
Adiabatic Process
Isobaric Process
Isovolumetric Process

37
Q

Isothermal Process

A

system’s temperature is constant = U is constant
U is constant (deltaU=0)
Q = W

38
Q

Adiabatic Process

A

no heat exchange, thermal energy is constant
Temperature is not held constant
Q = 0, deltaU = -W

39
Q

Isobaric Process

A

Pressure is held constant

40
Q

Isovolumetric Process

A

(Isochoric Process)
Volume is constant, no work is being done
deltaU = Q

41
Q

State Functions

A

P, density, T, V, enthalpy (H), Internal energy (U), Gibbs Free Energy (G), Entropy (S)
independent of path taken, but they relate to eachother

42
Q

standard conditions versus standard Temperature and Pressure

A

Standard conditions = 25C (298K), I atm P, I [M] = used for kinetics, equilibrium, and thermodynamics
STP = 0C (273K), I atm P = used for ideal gas calculations

43
Q

deltaQ > 0

A

Endothermic Reaction

44
Q

deltaQ < 0

A

Exothermic Reaction

45
Q

Bond formation (endothermic or exothermic?)

A

Exothermic process

46
Q

Bond dissociation (endothermic or exothermic?)

A

Endothermic Process

47
Q

Solvation

A

“Dissolution”

breaking of intermolecular interactions between solute molecules and solvent molecules to form new interactions together

48
Q

Solvation: Exothermic reaction

A

when new interactions are STRONGER than original ones

Favored at low temperature

49
Q

Solvation: Endothermic reaction

A

when new interactions are WEAKER than original ones

favorite at high temperatures

50
Q

Solvation: Ideal Solution

A

when overall strengths between original and new interactions are approximately equal

51
Q

Two prominent solubility rules

A
  1. All salts with ammonium (NH4+) and alkali metals (group 1) cations are water soluble
  2. All salts with nitrate (NO3-) and acetate (CH3COO-) anions are water soluble
52
Q

Ion Product (IP)

A

allows you to determine where the system is with respect to equilibrium position

53
Q

IP < Ksp

A

not at equilibrium
“unsaturated”
Dissolution is thermodynamically favored

54
Q

IP > Ksp

A

Beyond equilibrium
“supersaturated”
Thermodynamically unstable

55
Q

IP = Ksp

A

at equilibrium
saturated solution
“molar solubility”

56
Q

Kf constant

A

formation/stability constant

significantly larger than Ksp

57
Q

Ka constant

A

acid dissociation constant
=Products/Reactants
=higher value is a weaker acid (higher dissociation in more reactive acids)

58
Q

Kb constant

A

base dissociation constant

similarly to Ka

59
Q

Q < Keq , deltaG < 0

A

reaction proceeds in the forward direction

60
Q

Q = Keq , deltaG = 0

A

reaction @ dynamic equilibrium

61
Q

Q > Keq , deltaG > 0

A

reaction proceeds in the reverse direction

62
Q

Le Chatelier’s Principle

A

3 main types of stress

Concentration, pressure/volume, temperature

63
Q

Le Chatelier’s Principle: increase [R], decrease [P]

A

increase [R], decrease [P] = shift reaction to the right

vice versa will shift to the left

64
Q

Le Chatelier’s Principle: increase P, decrease V

A

increase P, decrease V = shift in direction of fewer mols of gas
vice versa will shift towards more mols of gas

65
Q

Le Chatelier’s Principle: increasing/decreasing T of Endothermic Reaction

A

increasing T will shift rxn to the right

decreasing T will shift rxn to the left

66
Q

Le Chatelier’s Principle: increasing/decreasing T of Exothermic Reaction

A

increasing T will shift rxn to the left

decreasing T will shift rxn to the right

67
Q

Strong Bases

A

NaOH Sodium Hydroxide

KOH Potassium Hydroxide

68
Q

Strong Acids

A 
HCl: hydrochloric acid
HBr: hydrobromic acid
HI: hydroiodic acid
H2SO4: sulfuric acid
HNO3: nitric acid
HClO4: perchloric acid
69
Q

Electrochemical Cells

A
  • any cell with red-ox reactions
  • anode: site of oxidation
  • cathode: site of reduction
  • e flow anode to cathode
  • current flows cathode to anode
70
Q

Galvanic (voltaic) Cell

A

-spontaneous reaction (deltaG<0) w/ positive electromotive force
-anode (-) charge
-cathode (+) charge
(electromotive force (+) and deltaG (-) )

71
Q

Electrolytic Cells

A

-nonspontaneous reactions (deltaG>0) w/ negative electromotive force
-anode (+) charge
-cathode (-) charge
(electromotive force (-) and deltaG (+) )

72
Q

3 type of rechargeable battery cell systems

A
  1. lead-acid battery
  2. nickel-cadmium batteries
  3. nickel-metal hydride (NiMH)
73
Q

Lead-Acid battery

A

discharging: Pb anode and PbO2 cathode concentrated in H2SO4
charging: PbSO4- electrodes are dissociated to restore original Pb and PbO2 electrodes
- low energy density

74
Q

Nickel-Cadmium Battery

A

discharging: Cd anode, NiO(OH) cathode in concentrated KOH
charging: Ni(OH)2 and Cd(OH)2 electrodes are dissociated
- high energy density than lead-acid

75
Q

Nickel-Metal Hydride Battery

A

-replaced Ni-Cd batteries
-higher energy density
surge content: above average content @ start of discharge phase

76
Q

concentration cell

A

specialized galvanic cell
-both electrodes are made of the same material
-concentration gradient causes movement of charge
(electromotive force (0) and deltaG (0) )

77
Q

relationship between electromotive force and Keq constant

A

Keq > 1 = E cell is +
Keq < 1 = E cell is -
Keq = 1 = E cell is 0

78
Q

Pauli Exclusion Principle

A

no 2 e- can possess the same set of 4 quantum numbers

79
Q

Aufbau Principle

A

e- will fill lower energy orbitals before filling higher energy orbitals

80
Q

Pauli Exclusion Principle

A

no 2 e- can possess the same set of 4 quantum numbers

81
Q

Aufbau Principle

A

e- will fill lower energy orbitals before filling higher energy orbitals

82
Q

Nucleophiles

A

negative charge
long pairs of e- or pi bonds
amino groups are common

83
Q

Electrophiles

A

positive charge

alcohols, aldehydes, ketones, carboxylic acids

84
Q

Oxidation of:
primary alcohol
aldehyde
secondary alcohol

A

primary = aldehyde, or carboxylic acid
aldehyde = carboxylic acid
secondary alcohol = ketone
(tertiary alcohol cannot be oxidized)

MCAT Prep (11 decks)

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