General Chemistry Mcat

Front 1

Describe the Bohr model of the atom

What makes up the nucleus?

Emitted and Absorbed energy directions?

Back 1

The nucleus is made up of protons and neutrons and the e-'s are found within the e- cloud

Front 2

1) Describe what the 12 signifies and what the 6 signifies.

2) Explain how e-'s come into play?

3) What is an isotope?

Back 2

1) The 12 is the mass #, which is protons + neutrons
The 6 is the atomic number, which never changes unless the element changes. This is only protons.
2) The e-'s = the amt of protons in the neutral state
3) Another variation of an element that differs by weight: 12C, 13C, 14C

Front 3

1) How much does a proton weigh?
2) How much does a neutron weigh?
3) How much does an electron weigh?

Back 3

1) 1.67 x 10^-27 kg
2) 1.67 x 10^-27 kg
3) very close to 0

Front 4

1) Which is the prevalent form of Cl? 35Cl or 37Cl
2) How much more prevalent is that form?


atomic mass: 35.5 g/mol

Back 4

1) 35Cl
2) 75% more prevalent (3:1 ratio)

Front 5

What are the 5 major groups you should know from the periodic table?

Describe their reactivity

Back 5

Alkali Metals: Form very strong bases
Alkaline Earth Metals: React very strongly with H2O
Transition Metals: Often form colored compounds
Halogens: Reactivity increases as you go up this column, F is the most reactive
Noble Gases: Not very reactive

Front 6

Describe the location of metals and nonmetals

Back 6

Metals tot he left of staircase, nonmetals to the right

Front 7

Describe the amount of orbitals for s, p, and d

How many e-'s can be held by s, p, and d?

Back 7

s --> 1 & 2e-
p --> 3 & 6e-
d --> 5 & 10e-

Front 8

1) Paramagnetic?

2) Diamagnetic?

Back 8

1) Ion/element attracted to a magnetic field (unpaired e-)
2) Slight deflection away from magnetic field (all paired e-)

Front 9

How many valence electrons are present?
- S[Ne]3s^2 3p^4
-Cr[Ar]4s^1 3d^5
-Br[Ar]4s^2 3d^10 4p^5
-Cu[Ar]4s^1 3d^10

Back 9

-6
-6
-7(remember once the d orbital is filled it doesn't count)
-1ish

Front 10

Explain the significance of the quantum numbers:
n -->
l -->
m(l) -->
m(s) -->

Back 10

n = principal quantum #/shell # [1....infinity]
l = Azimuthal #, the subshell (s,p,d,etc) 0=s, 1=p, 2=d [0...(n-1)]
m(l) = Magnetic # - specific order [-l ...+l]
m(s) = spin# [+1/2 or - 1/2]

Front 11

What is the strong nuclear force/nuclear binding energy?

Back 11

The force that holds protons together within the nucleus even though they repel one another.

Front 12

Describe the 5 routes of nuclear decay

Back 12

Make real index card of this, a problem, and the graph

Front 13

What are two main characteristics of a stable nucleus?

Back 13

1) Even # of neutrons and/or protons
2) N/Z =1for Z<20, the closer to 1 the better

Front 14

1) The rate at which decay occurs is proportional to _____
2) Solve this problem: If the 1/2 life is 10yrs of a 100g substance, how much remains after 40 years?

Back 14

1) how much you have
2) 6.25g

Front 15

1) Why does a nucleus always weigh less than it's separate parts?
2) What is the most stable nucleus in the universe? WHat can we determine based on this?

Back 15

1) Some mass turns into energy, E=mc^2
2) 56_26Fe the closer atoms are to this atom, the higher nuclear bonding energy they have
(Ex: 54C vs 238U --> 54C is higher)

Front 16

_Periodic Trends_
1) Atomic radius (Ex)
2) Ionic radius (Ex)
3) Electronegativity (Def)
4) Ionization Energy (Def)
5) Electron Affinity (Def)

Back 16

1) Smaller towards F (Ex. Mg has 1 more e- in the valence shell than Na and therefore feels a stronger attraction towards the nucleus)
2) Ex. --> Mg+2 is smaller b/c it looses an entire valence shell in comparison to Mg. The more positive the smaller; the more negative the larger
3) Increases towards F (how tightly an atom holds onto e-'s)
4) Increases towards He (E required to remove an e-)
5) Becomes more negative/exothermic to the right (E change when gaining an e-)

Front 17

Types of bonds:
1) Covalent
2) Ionic
3) Metallic

Back 17

1) non-metal & non-metal
-covalent network: diamond(C) & quartz(SiO2)
-Molecular-forms molecules
2) non-metal & metal (salt)
-can't move or conduct electricity unless in liquid
-high m.p. & b.p.
3) metal & metal
-conductive (ex. metal bench)
-ductile (can be easily drawn into wires)
-malleable (can be pounded into sheets)
-luster (shine)

Front 18

Compare Lewis, Arrhenius, and Bronsted-Lowry acids and bases

Back 18

Lewis - (covers all acids & bases) base: e- donor, acid: e- acceptor
Bronsted-Lowry - (the second best) base: H+ acceptor, acid: H+ donor
Arrhenius - (worst) base: H+ acceptor in H2O, acid: H+ donor in H2O

Front 19

Difference between ionic and covalent bonds

Back 19

Ionic: Atom gives away electron, both become ions, the + and - charges attract (metal and nonmetal)
Covalent: Electrons are shared between two atoms (nonmetal and nonmetal)

Front 20

Describe the electron domain geometry and angle value for the 5 most popular types

Back 20

Linear 180
Trigonal Planar 120
Tetrahedral 109.5
Trigonal Bipyramidal 90 & 180 & 120
Octahedral 90

Front 21

What is the electron domain geometry and molecular geometry for NH3?

H2O?

Back 21

EDG: Tetrahedral
MG: trigonal pyramidal

EDG: Tetrahedral
MG: Bent

Front 22

For the e- domains, what would the hybrid of the orbitals be?
2:
4:
6:

Back 22

2: sp
4: sp3
6: sp3d2

Front 23

Describe polar vs non-polar bonds

Are these polar or non-polar?
1) HCl
2) Cl2
3) CO2
4) BF3
5) NF3

Back 23

If the dipoles within the molecule cancel each other out, it is non-polar. If they dont, it's polar

1) polar
2) nonpolar
3) nonpolar
4) nonpolar
5) polar

Front 24

Intermolecular Forces:
1) H-bonding
2) Dipole-Dipole
3) London Dispersion (Van der Waals)
4) Ion - dipole

Back 24

All are significantly weaker than ionic and covalent bonds

1) F-H, O-H, N-H --> strongest intermolecular force
2) H+-Cl-----H+-Cl-
3) Nonpolar molecules ONLY have london dispersion. Bigger molecules have more
4) Ex. - Na surrounded by water molecules, partial neg. on O attracted to + charge on Na -->close in strength to H-bond

Front 25

If a molecule has a lot of I.M. forces, it is:

1) Big/Small
2) has __M.P.
3) has __B.P.
4) has __vapor pressure
5) has __viscosity
6) has __surface tension

Back 25

1. big
2. higher
3. higher
4. lower
5. higher
6. higher

Front 26

1) Which has the highest B.P.?
a: CH4
b: CH3OH
c: CH3F

2) Which has the highest B.P.?
a: NaCl
b: CH4
c: CH3OH
d: CH3F

Back 26

1) B
2) A

Front 27

What is it called when:
1. solid --> liquid
2. liquid --> gas
3. solid --> gas
4. gas --> liquid
5. liquid --> solid
6. gas --> solid

Back 27

1. fusion
2. vaporization
3. sublimation
4. condensation
5. crystallization
6. deposition

Front 28

1. Which phase changes are endothermic?
2. What would the change in H and S be for endo?
3. Which phase changes are exothermic?
4. What would the change in H and S be for exo?

Back 28

1. fusion, vaporization, sublimation
2. ΔH > 0, ΔS > 0
3. condensation, crystallization, deposition
4. ΔH < 0, ΔS < 0

Front 29

1. What is the change in H?
2. Describe Gibbs Free Energy
3. At eq, Gibbs = ?

Back 29

1. The change in Enthalpy - the amount of Energy within a system
2. G takes enthalpy and entropy into account in order to determine the spontaneity of a rxn. The formula for Gibbs is ΔG = ΔH - TΔS. A rxn is spontaneous when ΔG is negative
3. 0

Front 30

Describe phase changes based on the addition of heat and the change in tempture.

What is the formula for heat?

Does ΔT change for when switched from K to C?

Back 30

When heat is added at a temperature below the freezing point, the temperature increases until it reaches the phase change pt (solid to liquid). As heat increases, the temperature does not change during the phase change. When the phase change pt pases, temp increases again until it reaches the next phase change.....temp remains constat throughout, then increases again when past that pt

q = mcΔT (MCAT*)

No

Front 31

Describe the phase diagram of CO2

Where is the triple pt, critical pt, and normal B.P. located?

What is the B.P., triple point, and critical pt?

Back 31

-Mention solid, liquid, gas locations and the axis

B.P. - Pt at which the vapor pressure of H2O equals the atmospheric pressure
Triple Point - Pt at which all three phases exist
Critical Point - Pt at which you cannot differentiate between liquid and gas

Front 32

What are the 2 ideal gas assumptions?

What is gas at STP?

What 4 units (w/values) can P be converted into?

Back 32

a: No intermolecular attractions
b: Gases have no volume. They take up a lot of space, but the majority of the space is empty. (only at low P)

P = 1atm, T = 273K, V = 22.4L for 1 mol of gas

1 atm = 760mmHg = 760torr = 100,000Pa

Front 33

1. What is the ideal gas law equation?
2. What is the combined gas law equation?
3. If gas is in a container with a piston and the piston is pushed down, decreasing the volume, P ___ , T ___ , & Intermolecular forces ___
4. What is the value of R?

Back 33

1. PV = nRT
2. P1V1/n1T1 = P2V2/n2T2
3. P increases, T increases, & Intermolecular forces increase
4) 0.08206 L*atm/mol*K or 8.314 J/mol*K

Front 34

1. What is Avogadro's Law?
2. What is Boyle's Law?
3. What is Charle's Law?

Back 34

1. V is proportional to n, 6.02 x 10^23 particles w/in a mol
2. V is proportional to 1/P, As P increases - V decreases
3. V is proportional to T (Why? T is measured by how excited atoms are - as V decreases there is less space for the molecules to act in the excited state. Hence, as V decreases, T decreases)

Front 35

Equations for density and molar mass?

Describe Dalton's Law od Partial Pressures.

Back 35

d = m/V
mu = m/n

Ptotal exerted by the mixture of non-reactive gases is equal to the sum of the partial pressures of individual gases:
Pa + Pb + Pc..... = Ptotal
Pa = Xa(mol fraction) * Ptotal

Front 36

Why, in Grahm's law of diffusion, does a lighter gas move at a higher velocity - equation based question?

Back 36

Grahm's law explains how gases move through a narrow slit. The lighter gas moves faster because the K.E.'s of the gases are equal to one another and smaller mass divides to give a larger number.

Ex. 1/2mv^2 O2 = 1/2mv^2 H2

Front 37

1. If you mix two liquids, how do you determine which is the solvent and which is the solute?
2. Saturated?
3. Unsaturated?
4. Supersaturated?

Back 37

1. The solvent is the more prevalent liquid
2. The max amt of solute has been added to the solution
3. The max amt of solute has not been added to the solution
4. More than the max amt of solute has been added to the solvent - this is done by using heat

Front 38

1. Molarity?
2. Molality?

Back 38

1. M = moles of solvent/Liters of solution (solvent & solute)
2. m = moles solute/kg of solvent

(these two are similar in water b/c 1L = 1kg, but for other substances, this is not the same)
(they differ the most when the concentration increases)

Front 39

What are 2 solubility rules?

Back 39

1) Soluble = ALL group 1 metals, NO3-, NH4+, ClO4-, C2H3O2- salts
2) Insoluble = Most Ag+, Pb2+, Hg2+2 salts (unless they're with something from 1)

Front 40

1) Solids are more at ___________ temps
2) Gases are less soluble at _____________ temps
3) Gases are more soluble in a liquid at _________ pressure

Back 40

1) higher
2) higher
3) higher (higher pressure forces the gases down into the solution)

Front 41

1) What would make a freezing pt even lower?
2) What is the van't hoff factor for:
a. NaCl
b. Na2SO4

Back 41

1) If irregularities are present.
2) a. i = 2 b. i = 3

Front 42

1) Osmosis?
2) Osmotic Pressure?
3) Describe a basic reaction coordinate diagram

Back 42

1) The shift of water from high P to low P
2) The pressure it takes to keep water from entering

Front 43

1) Why do most chemical rxns slow down overtime?
2) What is the rate law equation for:
2NO(g) + Cl2(g) ---> 2NOCl(g)
3) The rate law is determined from the (slow/fast) rxn
4) What is a catalyst?

Back 43

1) Because it's rate is proportional to reactant conc.
2) Rate = k[NO]^2[Cl2]
3) slow
4) Speeds up a rxn by lowering the Ea by providing a diff. mechanism (pathway) for the rxn to occur

Front 44

1) What is the Arrhenius equation?
2) What is the definition of equilibrium?
3) If Keq >> 1, this means:
4) If Keq << 1, this means:
5) If Keq = 1, this means

Back 44

1) K = Ae^(-Ea/RT)
2) forward rate = reverse rate
3) there are more products than reactants at eq
4) more reactants than products at eq
5) similar amts of products and reactants at eq

Front 45

1) If Q > K, the result would be:
2) If Q < K, the result would be:
3) If Q = K, the result would be:

Back 45

1) shift left
2) shift right
3) at equilibrium

Front 46

1) What does the equilibrium constant & rxn quotient not include from a rxn?
2) What happened to Keq when we:
a: reverse the rxn
b: multiply everything by 3

Back 46

1) Solids & liquids
2) a: inverts it 1/Keq b: cubes it

Front 47

Le Chateliers Principle Example:
C(s) + O2(g) <---> 2CO(g)
1) If O2 is added to the rxn, what happens?
2) If C is added, what happens?
3) If CO is removed, what happens?
4) If P increases, what happens?
5) If the inert gas Ar is initially added to the rxn, what happens?
6) What is the only thing that changes K?

Back 47

1) shifts right
2) nothing
3) shifts right
4) shifts left, shifts to sode with less moles of gas
5) the pressure increases due to more molecules present, but doesn't change the moles or V of the other species. There's no shift - doesn't participate in rxn
6) T. When concentration changes, Q changes, but not K

Front 48

What are 7 strong acids?

Back 48

H2SO4
HClO4
HI
HBr
HCl
HNO3
HClO3

Front 49

What are the strong bases?

Back 49

Group 1 metal hydroxides
Ba(OH)2
Sr(OH)2
Ca(OH)2

Front 50

1) Describe the binary acid trend
2) What are two oxoacid trends?

Back 50

1) Acidity increases to the right and down (larger bonds have weaker bonds and are stronger acids)
2) the more oxygens, the more acidic --- the more electronegative heteroatom, the more acidic

Front 51

What are the equation for pH, pOH, H+, OH-?

Back 51

pH = -log[H+]
pOH = -log[OH]
[H+] = 10^-pH
[OH] = 10^-pOH
pH + pOH = 14
[H+][OH-] = 1x10^-14

Front 52

What are the 3 laws of thermodynamics (in order)?

Back 52

1) Energy cannot be created or destroyed
2) Entropy increases in spontaneous processes
3) A perfect crystal at 0K has zero entropy

Front 53

What is the change in E equation?

What is the work equation?

Back 53

ΔE = q(heat) + w(work)

w = -PΔV

Front 54

1) Temperature wise --> expanding gases ____a____, and compressing gases _____b_____
2) Isobaric?
3) Isochoric?
4) Isothermal?
5) Adiabatic?

Back 54

1) a: cool b: heat up
(only apply to non-ideal gases)
2) ΔP = 0 ---> w = 0
3) ΔV = 0 ---> w = 0
4) ΔT = 0 ---> E = 0, E is proportional to T (q = mcΔT)
5) q = 0

Front 55

1) In this rxn:
2C(s) + O2(g) ---> 2CO (g)
is entropy - or +?
2) What two things aren't state functions?
3) What is the ΔH of a bond forming and one breaking?
4) What is the ΔH formula for bond enthalpies?
5) What is the ΔH formula for enthalpies of formation?

Back 55

1) Entropy is positive b/c there is a 1 to 2 reactant product ratio (s is not included)
2) q and w (the pathway matters)
3) Forming bond: ΔH < 0
Breaking bond: ΔH >0
4) ΔHrxn = bonds broken(reactants) - bonds formed(products)
5) ΔHrxn = sum of ΔH of products - reactants

Front 56

1) Describe 2 aspects of formation rxns
2) Does the universe prefer exo or endo rxns? why?
3) Using Gibbs eqn, how would you adjust a system that has -ΔH and -ΔS to make it spontaneous? +ΔH and +ΔS?

Back 56

1) a: forms 1 mole of product
b: all reactants are elements in their std states
2) Exo b/c the system becomes more stable
3) Lower T for negative and Increase T for positives

Front 57

1) What 3 things fully dissociate?
2) What are the formulas for an acid or base dissociating into water? (used for when ions don't fully dissociate)
3) What are 2 eqns to be used instead of ice tables?
4) Kw = ?
5) pKa + pKb = ?

Back 57

1) Strong acids & bases, salts, group 1 & 2's
2) Ka = [H3O+][A-]/[HA] Kb = [OH-][HA]/[A-]
3) [H+] = sqrt(Ka[HA]) [OH-] = sqrt(Kb[A-])
(HA + H2O <---> H+ + A-)
4) [H+][OH-] = KaKb = 1 * 10^-4 @ 25C
5) 14

Front 58

1) The conjugate bases and acids are considered:
2) Why are salts more soluble in acidic solutions?

Back 58

1) Neutral or negligible
2) The H+ can bind to a product and create, example: HF from F-, this decreases the amount of product and causes a shift left

Front 59

1) What is used to calculate the pH of a buffer?
2) What do we want to happen in this eqn - to have max buffering capacity?
3) The buffer range is:
4) How much more acidic is 4 from 7?

Back 59

1) pH = pKa + log[A-]/[HA]
2) have the [A-]/[HA] ration = 1, so that pH = pKa
3) +/- 1 from pKa
4) every pH unit is a power of 10, so 4 is 10^3 times more acidic

Front 60

Titrations:
1) Equation?
2) Where is the eq pt for:
a: SA/SB
b:SA/WB
c:WA/SB
3) What does the 1/2 eq pt signify?
4) What does the eq pt signify?

Back 60

1. nMV(1) = nMV(2)
2. a: 7 b: <7 c: >7
3. [HA] = [A-] & pH = pKa of acid being titrated/pOH=pKb
4. pH can be determined & molA = molB
(the eqn "MaVa = MbVb" can be used)

Front 61

Oxidation States:
1. Elements in their elemental form =
2. Group 1 elements =
3. Group 2 elements =
4. H = ____, unless _____________ (then it = ____)
5. Ag = ____, Zn = ______, Al = _______, Cd = _______

Back 61

1. 0
2. +1
3. +2
4. +1, unless with a metal, then it = -1
5. Al=+3, Zn = +2, Cd = +2, Ag=+1

Front 62

Electrochemical cells:
1. An anode is always the site of __________
2. A cathode is always the site of __________
3. Electrons flow from ________ to __________
4. Anions flow to the ________, cations to the _________
5. For non-spontaneous rxns, the anode has a ___ charge and the cathode a ___ charge. For spontaneous rxns?

Back 62

1. oxidation (An Ox)
2. reduction (Red Cat)
3. anode to cathode
4. anions to anode, cations to cathode
5. spontaneous=cathode+ and anode-,
non-spontaneous is opposite

Front 63

Describe the Galvanic(voltaic) cell and the electrolytic cell

Back 63

Galvanic is a spontaneous process that produces electricity, while electrolytic is non-spontaneous and consumes electricity
(Red Cat/An Ox)

Front 64

1. Ecell =
2. What is the molar conc. of reactants and products in a std rxn?
3. The nonstandard Ecell eqn =
4. Describe what ΔG, Ecell, and K would equal for a spontaneous rxn

Back 64

1. Ecell = Ereduced + Eoxidized
2. reactants & products = both have 1 molar conc.
3. Ecell = E - 0.0592/n * logQ
(It's the Nernst eqn --> Ecell = E0cell - (RT/nF)lnQ)
4. ΔG < 0, Ecell > 0, K >Q (if 1 is true, they're all true)

Front 65

Describe electrolysis

What are the 2 quantitative eqns to solve for grams and moles of product?

Back 65

Uses a power supply to overcome non-spontaneity and produce elements.
1) Molton (salts)
2) Aqueous (salts dissolved in sol.)

Two eqns:
(Amps)(Ts)(MW)/(n)(F) = g product
(Amps)(Ts)/(n)(F) = moles product

look at notes - red notebook