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Malika Gen Chem MCAT > Additional IMPORTANT Concepts > Flashcards

Additional IMPORTANT Concepts Flashcards

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1
Q

Oxidation state of oxygen and ONLY exception to its normal oxidation #

A

-2 is normal oxidation state
-1 is oxidation ONLY as peroxide

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2
Q

Oxidation state of nitrogen

A

varies depending on what it is bound to

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3
Q

Oxidation state of gases in their lone state

A

ALWAYS 0 when they are by themselves.
N2, O2, F2, H2, etc

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4
Q

Disproportionation reaction

A

A redox reaction in which one element is both oxidized and reduced. The products include at least two different molecules which both contain the element (that has different oxidation states).

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5
Q

Single-replacement reaction

A

Reaction between a compound & a free element where the compound now is bonded to the free element

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6
Q

Double-replacement reaction

A

When two compounds swap ions to form two new compounds

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7
Q

Combination reaction

A

When two elements/compounds combine to form one compound

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8
Q

Oxidation-reduction reaction

A

Look out for atoms:
-That are reduced (oxidation state decreases & they gain more bonds to H, less bonds to O)
-Atoms that are oxidized (gain more double bonds to O or any halogen)

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9
Q

Out of the following ions, which one is smallest?
Na+
K+
Cl-

A

Na+ because first you compare the number of energy levels each ion has. K+ has more energy levels than Cl- and Na+, therefore it is the biggest. Then, since you are dealing with ions you need to compare ionic radius.
Na+ has a more positive nucleus which pulls the electron cloud inward, making it a smaller atom. Cl- has a less positive nucleus, so electron cloud stays in place.

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10
Q

Calculating formal charge of an atom

A

= valence electrons – (nonbonding electrons + 1/2 of bonding electrons)

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11
Q

Why is H2O not included in an equilibrium expression (Keq)?

A

Generally, water is the solvent and it is much greater in concentration than all the other species in solution and it’s concentration doesn’t change. Remains at 55.5, so it is omitted.

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12
Q

How does a catalyst affect Keq?

A

It doesn’t affect the position of equilibrium because it speeds up the rate of the forward and reverse reaction equally.

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13
Q

Question about reaction w/ coordination complex. Pay attention to how many moles of C2O4 is required in place of H20

A
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14
Q

Difference between attractive ion-dipole interaction and repulsive ion-dipole interaction

A

Ion-dipole interaction is attraction between a cation/anion and a polar molecule

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15
Q

Equation for max # of electrons in a shell

A

2n^2
n is the # of shells

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16
Q

Properties of ionic compounds

A

-In a solid state, they form crystalline lattice structures
-They have high melting and boiling points
-They dissolve in water and polar solvents readily

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17
Q

What is the electron geometry of an atom that has three bonds connected to it? What are the bond angles?

A

Trigonal planar; 120

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18
Q

What is the electron geometry of an atom connected to one lone pair and 2 bonds?

A

Bent

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19
Q

What is the electron geometry of an atom that is connected to one lone pair and 3 bonds?

A

Trigonal pyramidal

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20
Q

What is the electron geometry of an atom that is connected to 2 lone pairs and 2 bonds?

A

Bent

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21
Q

What is the electron geometry of an atom that is connected to 5 bonds?

A

Trigonal byramid

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22
Q

What is the electron geometry of an atom that is connected to one lone pair and 4 bonds?

A

Sawhorse (seesaw)

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23
Q

What is the electron domain geometry of an atom that is connected to two lone pairs, and three bonds?

A

T shape

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24
Q

What is the electron geometry of an atom that is connected to three lone pairs, and two bonds?

A

Linear

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25
Q

What is the electron pair geometry for an item that is connected to 6 bonds? Name the rest of the geometries of an atom connected to six substituents in order of 1 lone pair up to 4 lone pairs.

A

Octahedral; square pyramid, square planar, T-shape, linear

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26
Q

What factors affect lattice energy?

A

The more ionic a compound is, the higher lattice energy it has.
-Higher charge and smaller ion size (cations in the top of groups)

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27
Q

How do you calculate bond order of a molecule which has resonance?

A

bonds to central atom (a double bond counts as 2 separate bonds) / # of atoms bonded to central atom

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28
Q

As the intermolecular forces in a compound increases, how is the vapor pressure affected?

A

Vapor pressure is the amount of particles that have escaped into the gas phase above a liquid at equilibrium. If the IMFs are stronger, than less molecules can escape from the liquid, which causes a lower vapor pressure.

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29
Q

Compare the overlap of sigma vs pi bonds.

A

Sigma bonds have head to head overlap of orbitals.
Pi bonds have side to side/lateral overlap of orbitals.

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30
Q

Differentiate between intermolecular and intramolecular forces. Which ones are stronger?

A

Intermolecular forces are the attractions between various molecules in a substance which determine the physical properties of a substance.
Intramolecular forces are the attractions within one molecule (e.g ionic bonds, covalent bonds).
Intramolecular forces are stronger than intermolecular forces.

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31
Q

Electron configuration of charged ions (+ or - charged)

A

First write the ground state electron configuration of the atom.
Then, when writing the electron configuration of the ion, always remove/add electrons from the highest energy level.

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32
Q

Ground state Electron configuration exceptions to memorize (special atoms)

A

Cr: [Ar] 4s1,3d5
Cu: [Ar] 4s1,3d10
Mo: [Kr] 5s1, 4d5
Ag: [Kr] 5s1, 4d10
Pd: [Kr] 4d10

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33
Q

Diamagnetic vs. paramagnetic

A

Diamagnetic means that all the electrons are paired; produces its own magnetic field in the opposite direction.
Paramagnetic means there are one or more unpaired electrons, pulled into an external magnetic field.

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34
Q

Which elements can have expanded octets?

A

Sulfur, phosphorus, chlorine, silicon

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35
Q

Half-life question: How much sample was present prior to decaying?

A
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36
Q

Half-life question:

A
  1. Find # of half-lifes by halving until you get to final number.
  2. Divide the time it takes by # of half-lives.
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37
Q

Equation for fraction of substance remaining after “n” number of half-lives

A

1/2^n
n is the # of half-lives

Ex: If the half life is 2 mins, then in 10 minutes what is the fraction of sample that decays?
In 10 minutes, the sample is able to go through 5 half lives (2x5=10). Therefore, (1/2^5 = 1/32 fraction remaining. 31/32 is the amount of SAMPLE THAT DECAYS.

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38
Q

Reversible reactions have a ________ ∆G.
Irreversible reactions have a _______ ∆G.

A

small positive ∆G
large negative ∆G

39
Q

Calculating molecular formula question:

A
  1. Divide the masses given for each element by their MOLAR MASS (found in periodic table)
  2. Obtain # moles for each element
  3. Divide each element’s moles by the smallest molar amount (in this case 1.5) to get the correct ratio for each element
  4. C2H2NO is the empirical formula
  5. To get the molecular formula obtain the molar mass of the the empirical formula and divide the molar mass given in the QUESTION by the empirical molar mass
  6. Use this # to multiple each element ratio to get the correct molecular formula
40
Q

AAMC FL5 #10

A
41
Q

AAMC FL5 #15
Approximately how many moles of Kr+ are contained in the laser tube (11 cm^3) at 0°C and 1 atm?

A

*Use conversion 1 mol = 22.4 L of gas at STP
1 cm^3 = 1 mL
(1 mol/2.24 x 10^4 mL) x (11 mL) units cancel & you’re left w/ moles

42
Q

AAMC FL#16

A

Recognize:
-Radiation of 605 nm (longer wavelength) corresponds to lower f and lower E
-In the graph, a radiation of 604 nm is emitted & shown
-605 nm can’t be absorbed to produce 604 nm emitted because the frequency absorbed is lower than the frequency emitted

Thus, energy absorbed is always higher than energy emitted –> λ absorbed is lower than λ emitted

43
Q

If Keq > 1, reaction is (endergonic/exergonic).

A

exergonic

44
Q

pKa + pKb = 14
pH + pOH = 14 = pKw

A
45
Q

How to calculate Kb or Ka when given the pKa or pKb?

Calculating pKa/pKb when given Ka/Kb?

A

Ka = 1x10^(-pKa)
Kb = 1x10^(-pKb)

If [Ka] = n x 10 ⁻ᵉ, then pKa = {e-1}.{10-n}

46
Q

Shortcut equation to calculate pH or pOH when given the [H+] or [OH-] concentration of a strong acid/base

A

If [H⁺] = n x 10 ⁻ᵉ, then pH = {e-1}.{10-n}

Example [H+] = 5 x 10^-7
pH = {7-1}.{10-5} = 6.5

47
Q

Define Kw and how to calculate it

A

It is the water dissociation constant which is equal to the concentration of hydroxide and hydronium ions.
Kw = 10^-14 = Ka x Kb

48
Q

What is the pKa of a species in solution?

A

The pKa is the point at which half of the species in the solution are deprotonated.
If the pH of a solution is equal to 3, which is also the pKa of the carboxy group of the amino acid in solution, half of the carboxy group is deprotonated while the other half remains protonated.

49
Q

A weak acid has a Ka value of?
A strong acid has a Ka value of?

A

Weak acid Ka < 1
Strong acid Ka > 1

50
Q

Tip: when dealing with diprotic bases such as Ca(OH)2, don’t forget to multiply the concentration of calculated OH by 2.

A
51
Q

Calculating the H+ concentration of a weak acid without using a ICE table

A
52
Q

What is the normality of a 2 M H3PO4 solution?

A

H3 = 3 H+ protons
3 x 2 = 6 N

53
Q

Define equivalence point. Identifying on a titration curve.

A

The point at which equivalent amounts of acid and base have reacted. Use the Henderson-Hasselbalch equation to calculate the pKa in this region.
The middle of the vertical region on a titration curve.
pH = pI.

54
Q

Henderson-Hasselbalch EQUATION

A
55
Q

Define half equivalence point and identify on a titration curve.

A

The point where half of the acid/base has been neutralized by the titrant. Therefore [HA] = [A-]
Horizontal region on a titration curve. In this “buffer region” the pKa = pH/ pKb = pOH

56
Q

Define a buffer

A

Solutions which resist changes in pH; consists of a mixture of a weak acid/base conjugate pair
Ex: CH3COOH & CH3COONa

57
Q

Ideal condition of a buffer

A

Must have a pKa value within 1 pH unit (either above or below) of the analyte you wish to maintain

For pH = 6.3
Buffer must be either 5.3 up to 7.3

58
Q

approximations of pH at equivalence point for various titration scenarios:
Strong acid + strong base
Weak acid + strong base
Weak base + strong acid

A

Strong acid + strong base, pH = 7
Weak acid + strong base, pH > 7
Weak base + strong acid, pH < 7

59
Q

Henderson-Hasselbach example #1
Carbonic acid has a pKa of 6.1. If a buffer contains 0.012 M of H2CO3 and 1.2 M of HCO3- what is the pH of the buffer system?

A
  1. Starting pKa listed in question = 6.1.
  2. Figure out if there is more acid or base in the buffer. More acid means starting pKa will go down; more base means it will go up.
  3. Figure out the decimal pt. difference between the concentrations of acid & base
  4. If difference is 2, pKa will go up by 2, so final pH is 8.1.
60
Q

Henderson-Hasselbach example #2
Carbonic acid has a pKa of 6.1 and buffer created from this acid and its conjugate has a pH of 5.1. What is the ratio of acid to base in this solution?

A
  1. Compare pH of buffer to original pKa. Find the difference. This case is only 1 pH unit which corresponds to 10.
    2 pH difference = 100 and so on
  2. If pH is less, means there is a higher amt. of acid in buffer than base
  3. Ratio of acid:base is 10:1.
61
Q

Henderson-Hasselbach example #3
If the pH of an acetate buffer is 2.8 and the ratio of CH3COOH to CH3COO- is 100 to 1. Then what is the pKa of acetic acid?

A
  1. pH after buffer was added= 2.8
  2. Which concentration is higher in the buffer, acid or base? Here, it’s acid. Means that the buffer caused the pH to drop to 2.8. But what was the original pKa?
  3. Ratio is 100:1, so the original pKa has to be 2 units higher.
62
Q

Question:
You are given pKa and [HA] weak acid concentration. Asked to find pH.

A
  1. Find Ka from the pKa. Ka = 1x10^-pKa
  2. Set up expression:
    Ka = x^2/[HA]
  3. Solve for x = [H3O+]
  4. Plug in:
    pH = -log[H3O+]
63
Q

Question:
You are given pH and [HA] weak acid concentration. Asked to find pKa.

A
  1. Find [H3O+] from pH=3 –> 1x10^-3 M
  2. This is your x.
    Set up expression:
    Ka = x^2/[HA]
  3. Solve for Ka
  4. Solve for pKa = -logKa
64
Q

Ionic Equation vs. Net Ionic Equation

A

Ionic equation - includes all spectator ions
Net ionic equation - no spectator ions

65
Q

Finding solubility product constant: Ksp (the value for the limit of solubility of a compound)

A

*Polyatomic ions such as PO4 are considered simply “x”
*Solids & pure liquids AREN’T included
*Ksp is temperature dependent

66
Q

Low solubility of a compound correlates with low ionization in solution –> which can correlate with low Ka or Kb (weak A/B)

A
67
Q

Define chelation

A

In complex ions this occurs when the central cation is bonded to the same ligand in multiple places.
Often ligands are large and organic.

68
Q

Define molality

A

moles of solute/kg of solvent
*Independent of temperature unlike molarity

69
Q

Define saturation point of a solution (that is when a solution becomes saturated)

A

The equilibrium point of a solution where rate of dissolution = rate of precipitation

70
Q

How to determine which compound is more soluble given 2 Ksp values

A

When comparing 2 compounds with same # of ions:
-Higher Ksp = more soluble
Comparing 2 compounds w/ diff # of ions:
-Calculate x, larger x = more souble

71
Q

How do these two factors affect solubility:
1. Complex Ion formation
2. Common Ion effect

A
  1. Increases solubility
  2. Decreases solubility
72
Q

Colligative Properties

A

Physical properties that are dependent on the concentration of dissolved particles, but not on the chemical identity of them

73
Q

Calculating solubility with a common ion present

A

*There are 2 moles of Cl in CaCl2, so multiple M by 2 to get 0.06

74
Q

Different indicator tests and corresponding values indicating pH change

A

The chosen indicator should have a range in which the equivalence point pH is within

75
Q

True or false:
An atom with a higher electronegativity, also has a higher reduction potential.

A

True:
Yes because it wants to attract/gain electrons.

76
Q

True or false:
When comparing reduction potentials, the more positive a reduction potential for a half reaction, the more likely the atom will get reduced.

A

True

77
Q

Oxidation potential = −Reduction potential

A
78
Q

Calculating standard cell potential (Ecell - the amount of energy a battery can provide)

A

E˚cell = Eoxidation + Ereduction

E˚cell = E˚red,cathode − E˚red,anode

79
Q

Equation relating ∆G˚to E˚cell

A

∆G˚= −nFE˚cell

n is # of moles of e- transferred
F is Faraday’s constant; 96,485 C/mol e-

80
Q

Different types of Cells and respective values

A

*E˚cell is the same as Emf (electron motive force)

81
Q

Define the anode and cathode in an electrochemical cell

A

Anode: Site of oxidation. Only attracts anions.
Cathode: Site of reduction. Only attracts cations.

82
Q

In an electrochemical cell, e- flow from:

A

Anode to cathode

83
Q

In an electrochemical cell, current flows from:

A

cathode to anode

84
Q

Writing half reactions

A

*Oxidized species: e- are on the right of equation
Reduced species: e- on left of equation

85
Q

Comparing electrolytic vs. galvanic cells

A

Electrolytic - ∆G>0, driven by current from battery:
anode is + electrode, cathode is − electrode; current flows in a nonspontaneous direction (- to +)

Galvanic - ∆G<0, cell produces electric current itself:
anode is −, cathode is +

86
Q

1 coulomb is equal to?

A

1Joule/1Volt

87
Q

In a spontaneous reaction, current will flow from a _____ to ______ potential. Electrons will flow from negative to positive potential.

A

positive to negative

88
Q

What is a concentration cell?

A

A type of galvanic cell in which electrons flow from a lower concentration half-cell into a higher concentration half-cell. The cell has the same electrodes on both sides.
Oxidation occurs on the lower concentration side (anode), while reduction occurs at the higher concentration side (cathode).
E˚cell > 0

89
Q

Define electrolysis.

A

The operation of an electrolytic cell which runs a nonspontaneous electrochemical decomposition reaction.

90
Q

Rechargeable batteries:
When a battery is charging, it operates as an _________ cell.
When discharging, the battery operates as a ___________ cell.

A

electrolytic, galvanic

91
Q

Compare differences in setups of rechargeable batteries

A

Lead acid:electrolyte solution is sulfuric acid has low energy density, but high E˚cell

Ni-Cd: electrolyte solution is concentrated KOH high energy density, but lower E˚cell

92
Q

Explain a direct redox reaction setup (could be present on MCAT passage)
FL 1 #53

A

A metal strip of (Al, Zn, etc) is placed in a solution of ions of (another transition metal such as Pb).

  1. A reaction will occur if the strip is a stronger reducing agent. It will reduce the solution of ions and cause the solution to deposit as metal onto the strip. The strip itself will be oxidized.
  2. How do we know reduction has occurred? If there is a deposit of Pb metal onto the strip.
93
Q

Nerst Equation for Ecell

A
94
Q

When comparing polyprotic acids, with multiple Ka’s remember that:

A

Ka1 > Ka2 > Ka3, etc..
H3PO4 will dissociate more (have higher Ka) than H2PO4-
The higher the negative charge value on the acid, the more unstable it is, and less likely it will dissociate

Malika Gen Chem MCAT (3 decks)

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