CHE Final Exam Flashcards

1
Q

Alpha and beta particles

A

Positively and negatively charged particles, respectively

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2
Q

Gamma rays

A

Neutral charge, high-energy radiation

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3
Q

Nuclear theory

A

Supported by Rutherford’s experiment:

nucleus = most of atom’s mass w/protons & neutrons

empty space = most of atom’s space w/electrons

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4
Q

Polyatomic molecules in nature

A

P4
S4
Se4
O3

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5
Q

Noble metals

A

Ag, Pt, Au

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6
Q

Transition metals with predictable charge (3)

A

Ag (1B = 1+)
Zn (2B = 2+)
Cd (2B = 2+)

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7
Q

1 prefix (HC)

A

meth-

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8
Q

2 prefix (HC)

A

eth-

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9
Q

3 prefix (HC)

A

prop-

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10
Q

4 prefix (HC)

A

but-

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11
Q

HC suffixes (single, double, & triple bonds)

A

-ane, -ene, -yne

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12
Q

Alkanes

A

Hydrocarbons with only single bonds (-anes)

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13
Q

Percent yield

A

% yield = ( actual / theoretical ) * 100%

*NOT PERCENT ERROR!

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14
Q

Dilution formula

A

M1V1 = M2V2

M = molarity, V = volume (liters)

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15
Q

Dissociation

A

Ions of a salt separate when dissolved

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16
Q

Ionization

A

Process of ion formation in solution

(acids, bases)

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17
Q

Arrhenius acids & bases

A

Acids produce H+ (H3O+) in water

Bases produce OH- in water

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18
Q

Brønsted acids & bases

A

Acid = proton donor

Base = proton acceptor

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19
Q

Monoprotic, diprotic, and tripotic acids

A

Each unit of acid yields one, two, and H+ ion(s) respectively

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20
Q

Neutralization reaction result

A

salt + water + heat

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21
Q

Gas-evolving reactions & compounds (4)

A

Acid + salt –> salt + gas + water

Sulfides

Bicarbonates

Bisulfites

Ammonium

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22
Q

Oxidation half-reaction

A

Shows loss of electrons (OIL)

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23
Q

Reduction half-reaction

A

Shows the gain of electrons (RIG)

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24
Q

Oxidizing agents

A

Are reduced (RIG)

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25
Q

Reduction agents

A

Are oxidized (OIL)

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26
Q

Disproportionation reaction

A

Same element is simultaneously oxidized & reduced

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27
Q

Modified dilution equation

A

MacidVacid(# of H+) = MbaseVbase(# of OH-)

28
Q

Pressure unit conversion

A

760 mmHg = 760 torr = 1 atm = 101,325 Pa

29
Q

Molar mass of gas formula

A

M = dRT/P

molar mass = (density) * R * (Kelvin) / (atm)

D = PM/RT
= pressure * molar mass / R * temp

30
Q

Average KE

A

Constant tempterature: gases have same average KE

31
Q

Root mean square velocity

A

urms = (3RT/M)1/2

R = J/mol K
T = Kelvin
M = kg/mol
32
Q

Graham’s law of effusion

A

ra/rb = (M<span>b</span>/M<span>a</span>)1/2

33
Q

Real gases

A

Behavior differse at high pressure/low temperature

34
Q

Energy unit conversion (4 equalities)

A

1 J = 1 kg*m2/s2

1 cal = 4.184 J

1 Cal = 1kcal

1 L*atm = 101.3 J

35
Q

Internal energy & change in internal energy

A

E (internal energy) = KE + PE

ΔE = Eproducts - Ereactants

ΔE = q + w

*+w work done ON system, -w work done BY system

ΔEsystem = -ΔEsurroundings

36
Q

Pressure-volume work

A

w = -PΔV

work = negative of external pressure * change in volume

37
Q

Enthalpy/change in enthalpy formula

A

Enthalpy = H

H = E + PV
(internal energy + pressure*volume)

ΔH = ΔE + PΔV

ΔH = Hproducts - Hreactants

38
Q

Heat capacity

A

q = C * Δt

Amount of heat to raise temperature by 1°C

C = J/°C or J/K

System absorbs heat = temperature increasse

39
Q

Specific heat capacity

A

q = m * Cs * Δt

Measure of substance’s ability to absorb heat/amount of heat required to raise 1 g of substance by 1°C

40
Q

Bomb calorimeter

A

Constant-volume calorimetry

ΔErxn = qv = qrxn = -qcal

qcal = Ccal * ΔT = -qrxn

41
Q

Coffee-cup calorimetry

A

Constant-pressure calorimetry

ΔHrxn = qp = qrxn = -qsoln

qsoln = msoln * Cs, soln * ΔT = -qrxn

42
Q

Standard enthalpy of formation

A

ΔHf°

Heat change when one mole is formed from elements at 1 atm

Most stable = 0

Stable liquids: Hg, Br

Stable gases: H2, N2, O2, F2, Cl2, noble gases

Graphite is stable, diamond is not

S8 rhombic is stable

43
Q

Standard enthalpy of a reaction

A

The enthalpy of a reaction carried out at 1 atm

ΔHrxn° = ΣnproductsΔHf° - ΣnreactantsΔHf°

44
Q

Hess’ law

A

ΔHrxn° = ΔH1° + ΔH2°

45
Q

Frequency formula

A

v = c / λ

46
Q

Electromagnetic spectrum

A

Low energy to high:

radio
microwave
infared
visible light
ultraviolet
X-ray
gamma ray

47
Q

Photoelectric effect

A

Many metals emit electrons when light shines on surface

48
Q

Number of photons formula

A

Number of photons = Epulse / Ephoton

*Ephoton = hc/λ

49
Q

Energy & frequency combined formula

A

v = c / λ

E = hv

E = hc/λ

50
Q

Wavelength formula

A

λ = h/mv

Planck’s constant / mass * velocity

51
Q

Energy of electron orbital change formula

A

ΔEH atom = -R (1/n2final - 1/n2initial)

R = Joules

52
Q

1 Hz = ? s-1

A

1 Hz = 1 s-1

53
Q

Pauli exclusion principle

A

No two electrons can have same four quantum numbers

Orbital = 2 electrons max

54
Q

Aufbau principle

A

Lower energy orbitals fill before higher energy orbitals to minimize energy of atom

55
Q

Hund’s rule

A

Electrons first occupy orbitals of equal energy singly with parallel spins

56
Q

Transition elements with irregular electron configurations (10)

A
  1. Cr (Chromium) 24
  2. Cu (Copper) 29
  3. Nb (Niobium) 41
  4. Mo (Molybdenur) 42
  5. Ru (Rutherium) 44
  6. Rh (Rhodium) 45
  7. Pd (Palladium) 46
  8. Ag (Silver) 47
  9. Pt (Platinum) 78
  10. Au (Gold) 79
57
Q

Van der Waals radius

A

Nonbonding radius of an atom

58
Q

Cation v. anion radius size

A

Cations < anions

Except Rb+ and Cs+ are larger than F- and O2-

59
Q

Periodic trends in first ionization energy

A

The larger the Zeff, the more energy it takes to remove it.
*IE1 increases to the right
**Except: Group 2A to 3A & Group 5A to 6A

The farther the electron is from the nucleus, the less energy it takes to remove it
*IE1 decreases toward the bottom

60
Q

Electron affinity

A

X (g) + e- → X- (g) + EA

The energy released (-kJ/mol) when a neutral gaseous atom gains an electron

The more energy released when electron is gained, the more negative the EA (-kJ/mol)

*INCREASES TO THE RIGHT

61
Q

Diagonal relationship

A
  1. Li & Mg
  2. Be & Al (Al is a metal like 2A)
  3. B & Si (metalloids)

Similar because of charge density (ion charge / volume)

62
Q

Properties of oxides across a period (3)

A
  1. Metals/Groups 1A/2A form basic oxides
  2. Nonmetals form acidic oxides
  3. Al form amphoteric oxides (both basic/acidic)
63
Q

Formal charge (4)

A

Charge an atom would have if all bonding electrons were shared equally

FC of an atom = Valence e- - nonbonding e- - (1/2)(bonding e-)

  1. Sum of all formal charges in neutral molecule = 0
  2. Sum of all formal charges in ion = ion charge
  3. Small (or zero) formal charges on individual atoms are better than large ones
  4. When formal charge cannot be avoided, negative formal charge should reside on the most electronegative atom
64
Q

The Clausius-Clapeyron Equation

A

ln P2/P1 = (-ΔHvap/R) (1/T2-T1)

R = 8.314 J/mol K

T = in Kelvin

65
Q

Equation for heat involved in completion of a phase change

A

q = nΔHtransition

heat = # of mols * heat of [transition]

66
Q

Equation for heat involved in temperature change

A

q = (m) (Cs) (ΔT)