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DAT BOOT CAMP > Week 1 > Flashcards

Week 1 Flashcards

1
Q

Atomic # or Z

A

number of protons (retrieve from periodic table)

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2
Q

Mass #

A

N + Z (neutrons + protons)

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3
Q

isotopes

A

same # of protons but different # neutrons

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4
Q

alkali metals

A

group 1 - basic, react with H2O - explosive, exothermic and more reactive as you move down the group

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5
Q

alkaline earth metals

A

group 2 - basic, but not as reactive as alkaline metals

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6
Q

noble gases

A

group 8 - full octet, not very reactive

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7
Q

halogens

A

group 7 - single electron acceptor

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8
Q

transition metals

A

form bright colored compounds, i.e. copper = green/blue

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9
Q

exceptions to noble gas configurations

A

Cr, Mo, Cu, Ag, Au (because more stable if shells are half full or full)

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10
Q

s,p,d,f orbitals

A

s = 2e, p= 6e, d = 10e, f = 14e

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11
Q

paramagnetic

A

attracted to magnetic field (unpaired electrons)

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12
Q

diamagnetic

A

slight deflection of magnetic field (all electrons are paired)
*remember electrons like to fill each shell across before pairing up

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13
Q

Quantum #: n

A

principal - shell - [1….infinity]

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14
Q

Quantum #: l

A

azimuthal - subshell (type of orbital) - [0…(n-1)]

s = 0, p =1, d=2, f=3

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15
Q

Quantum #ml

A

magnetic - specific orbital (orientation in space), [-l…+l]

i.e. if l =2; [-2,-1,0,1,2] (typical d orbital has 5 orbitals)

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16
Q

Quantum #: ms

A

up or down - spin - (-1/2) or (+1/2)

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17
Q

Quantum # format

A

[n, l, ml, ms]

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18
Q

What forces hold the nucleus together?

A

strong nuclear force or nuclear binding energy

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19
Q

nuclear symbols

A

alpha (same as He), p (proton), n (neutron), beta (electron or positron), gamma (gamma ray light)

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20
Q

Stable Nucleus features

A
  • even # protons and/or neutrons
  • N/Z ratio ~ 1 for Z= 20 (up to Ca)
  • daughter (product) of parent nuclear rxn is more stable due to release in energy
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21
Q

Alpha decay

A

reduces mass # of large nuclei (Z>83) - product = 4/2 alpha (like helium but not electrons)

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22
Q

Beta (-) decay/emmission

A

converts neutrons to protons (occurs when N/Z ratio is too high) product = electron or 0/-1 beta

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23
Q

beta (+) decay (positron emmission)

A

converts proton to neutrons (occurs when N/Z ratio is too low) product = 0/+1 beta

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24
Q

electron capture

A

converts protons to neutrons (occurs when N/Z ratio is too low) (reactant side = 0/-1 electron)

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25
Q

gamma decay

A

product is a photon (emitted when coming from excited state to ground state)

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26
Q

Most stable nucleus

A

iron-56 (56-Fe) —atoms closer to this is more likely to be more stable

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27
Q

Periodic Trends

A
  1. Atomic radius - down + left of the table; radius gets bigger because z(eff) is smaller
    z(eff) = # of protons - shielding/core electrons
  2. Ionic Radius - cations are usually smaller because they lose electrons, anions gain electrons and protons can’t hold as close
  • non metals form anions
  • metals form cations
  1. 1st ionization energy = energy req’d to remove 1 e-
    (trend is opposite of atomic radius – up and right harder to remove e-)
    *big jumps in ionization energy indicates all valence electrons have been removed and moving into core electrons
    **EXCEPTIONS: Be + B and N + O
    These are slightly lower because filled/half-filled (making them more stable (Be over B and N over O)
  2. Electron affinity - energy in gaining an electron (exothermic) - increasing trend from left to right (no vertical trend, but 3rd row is more negative than others - Cl)
    * **EXCEPTIONS: groups in Be, N, noble gases (endothermic due to filled shell/half-filled shell
  3. Electronegativity - associated with sharing an electron (shows how close it can pull a shared e- in a covalent bond)
    Trend: up and to the right (fluorine)
    Q: which is most polar? look for most electronegative atom
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28
Q

Covalent bond

A

non metal with non metal (covalent network solids - diamond) or molecular compounds (low mp/bp)

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29
Q

Ionic bond

A

metal with non metal (crystalline, high mp/bp)

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30
Q

Metallic

A

metal with metal

  • conductive (electric + thermal)
  • ductile (draw them into wires)
  • malleable (pound them to sheets)
  • Luster (shiney)
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31
Q

Lewis base

A

ligand, anything with lone pair of e-

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32
Q

Lewis acid

A

usually a metal ion

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33
Q

Octet Rule

A

Exceptions:

  1. 2H, 4Be, 6B, 6Al
  2. 3rd row and lower can go over octet rule
  3. odd # e- overall (i.e. NO)
34
Q

Formal Charge

A

Normal # of valence electrons - (each line + dots)

lowest formal charge = preferred species

35
Q

Electron Domain = 2

A

sp, 180 degrees, linear, 0 nonbonding e, linear

36
Q

Electron Domain = 3

A

sp2, 120 degrees, trigonal planar, 1 nonbonding e, trigonal planar

1 non-bonding e- = bent

37
Q

Electron Domain = 4

A

sp3, 109.5 degrees, tetrahedral, 0 nonbonding e- pair, tetrahedral

1 non-bonding e- = trigonal bipyramidal

2 non-bonding e- = bent

38
Q

Electron Domain = 5

A

sp3d, 90/120 degrees, trigonal bipyramidal, 0 nonbonding e- pair, trigonal bipyramidal

1 nonbonding e- pair = see-saw
2 nonbonding e- pair = T-shaped
3 nonbonding e- pair = linear

39
Q

Electron Domain = 6

A

sp3d2, 90 degrees, octahedral, 0 nonbonding e- pair, octahedral

1 nonbonding e- pair = square pyramidal
2 nonbonding e- pair = square planar

40
Q

sigma

A

single bonds

41
Q

pi bond

A 
double bond (has 1 sigma and 1 pi bonds)
triple bond (has 1 sigma and 2 pi bonds)
42
Q

Intermolecular forces - Hydrogen bonds

A

strongest over dipole-dipole and london, needs F-H, O-H, or N-H “FON”

43
Q

Intermolecular forces - London Dispersion Forces

A

Van der Waals - all molecules have them, larger ones have MORE, nonpolar molecules have the most

44
Q

intermolecular forces - dipole-dipole

A

partial + and partial - weak attractive forces (i.e. HCl with HCl) between polar molecules

45
Q

intermolecular forces - ion-dipole

A

stronger than dipole-dipole, similar to hydrogen bond, occurs between something that is completely charged to a partially charged atom

i.e. Na+ in water

46
Q

Increase in intermolecular forces

A

increase MP, BP, viscosity, surface tension

decrease vapor pressure (the measure of a liquids vapor that is in equilibrium above it)

47
Q

solid to liquid

A

fusion “melting”; delta H>0; delta S>0 (more random), endothermic

48
Q

liquid to gas

A

vaporization “boiling”; delta H>0; delta S>0 (more random), endothermic

49
Q

solid to gas

A

sublimation (i.e. dry ice); delta H>0; delta S>0 (more random), endothermic

50
Q

gas to liquid

A

condensation; delta H

51
Q

liquid to solid

A

crystallization; delta H

52
Q

gas to solid

A

deposition; delta H

53
Q

Calorimetry

A

q = mc(delta)T
m= mass in grams
c = specific heat of substance
delta T = change in temperature (kelvin or C)

**During phase changes, cant use mcdeltaT, must use ndeltaH (n = mols)

54
Q

Triple Point

A

solid, liquid, gas phases are in equilibrium

55
Q

critical point

A

when beyond this point, there is no phase change between liquid and gas (substance is in “super critical fluid” state)

56
Q

boiling point

A

when vapor pressure equals the atmospheric pressure pushing down on it

57
Q

normal atmospheric pressure

A

1 atm = 760 torr

58
Q

phase diagram of water

A

different because solid H20 is less dense than liquid H20 at 1 atm

59
Q

Phase diagram of CO2

A

At 1 atm, there is no liquid phase– must increase pressure to get liquid; CO2 will continue to sublime at 1 atm (Solid to gas)

60
Q

Ideal Gas Assumptions

A
  • gas molecules have no volume (good at low pressures)
  • no intermolecular attractive forces (all collisions are elastic) good at high temperatures because gas molecules are moving so fast there is no time for feel other forces of H-bonds, dipole-dipole, or london
61
Q

STP for ideal gas

A

1 atm, 273 K, 22.4L for 1 mol gas

one mol of any gas takes up the same volume

62
Q

Ideal Gas Law

A

PV = nRT

P= pressure (atm)
V = volume (L)
n = mol
R = given constant (0.08206 for gases)
T = Kelvin
63
Q

Boyle’s Law

A

V proportional to 1/P

64
Q

Charles’ Law

A

V proportional to T

increase T = increase Volume = decrease in density (hence a balloon rises)

65
Q

Avogadro’s Law

A

V proportional to mols

66
Q

Combined Gas Law

A

P1V1/n1T1 = P2V2/n2T2

use this when given two parts

understand proportions of the Laws’s – will help deduce answer without calculator

67
Q

gas density

A

d = mass/volume

68
Q

two main equations for molar mass

A

d = m/v and molar mass = m/n (g/mol)

69
Q

Dalton’s Law of Partial Pressures

A

Pa + Pb + Pc … = Ptotal

Pa = (Chi) x Ptotal

(Chi) = mol fraction

70
Q

Graham’s Law of Effusion

A

when gas escapes through a narrow opening

rate2/rate1 = sqrt(molar mass1/molar mass2)

**simply at the same temp, the lighter gas will move at a higher velocity and escape faster

71
Q

saturated

A

max amt of dissolved solute in solvent

72
Q

unsaturated

A

not the max amt dissolved solute in solvent

73
Q

supersaturated

A

more than max solute dissolved by heating solvent

74
Q

Molarity

A

mol of solute/Liter of soln

75
Q

Molality

A

mol of solute/Kg of solvent

76
Q

mol fraction

A

Mol(a)/Mol(total)

77
Q

Solubility Rules (MEMORIZE)

A
  • ALL group 1 metal salts, NO3-, NH4+, ClO4-, C2H3O2- (acetate) salts are SOLUBLE in water
  • Most Ag+, Pb2+, Hg2 (2+) salts are INSOLUBLE in water unless it is with any salts in RULE #1
78
Q

Phase solubility

A
  • solids are more soluble at higher temp
  • gases are less soluble at higher temp
  • gases are more soluble in a liquid at higher pressures
    (i. e. if we compressed air into a beaker, it pushes the molecules down into the liquid that were originally in the air)
79
Q

Colligative Properties: Freezing point depression

A

deltaT(f) = -i Kf m

i = van't Hoff # (how many ions can a molecule dissociate into)
Kf = constant (degree C/molality)
m = molality
Tf = change in freezing temp

i.e. NaCl = Na+ and Cl- = 2 = i
Na2SO4 = 3 = i

80
Q

Colligative Properties: Boiling point elevation

A

delta Tb = i Kb m

i = van't Hoff # (how many ions can a molecule dissociate into)
Kb = constant (degree C/molality)
m = molality
Tb = change in boiling temp

i.e. NaCl = Na+ and Cl- = 2 = i
Na2SO4 = 3 = i

81
Q

Colligative Properties: Vapor Pressure Depression

A

Rauolt’s Law - basically pure water vapor pressure is 22 mm Hg, the vapor pressure for (8 mol H2O and 2 mol CH3OH) is just 8/10 mol ratio times the pure water vapor pressure 22 mm Hg)

82
Q

Colligative Properties: Osmosis

A

Osmostic pressure (pi) = iMRT

i = van't Hoff factor
M = molarity mol/volume (L)
R = constant
T = temp (K)

DAT BOOT CAMP (3 decks)

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