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Chemistry 121 > Midterm II Equations & Terms > Flashcards

Midterm II Equations & Terms Flashcards

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1
Q

stoichiometric amounts

A

when compounds are mixed in the exact amounts needed for all molecules to be “used up” completely

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2
Q

limiting reagent (reactant)

A

compound that is completely used up

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3
Q

% yield

A

[(actual yield)/(theoretical yield)] x 100%

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4
Q

Water collective properties

A

1) many different forms of ice 2) high heat capacity 3) high boiling point 4) expands upon cooling 5) dissolves many different substances

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5
Q

What kind of molecule is water?

A

polar covalent (unequal electron sharing)

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6
Q

hydration

A

when a solid dissolves in water, the H (partial positive charge) orients itself toward a negative charge and vice versa with the partial negative charge

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7
Q

when a salt dissolves…

A

the cations and anions separate into individual ions

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8
Q

solute

A

substance dissolved

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9
Q

solvent

A

water (dissolved into)

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10
Q

strong electrolytes…

A

conduct electrical current strongly (ionize completely!)

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11
Q

weak electrolytes…

A

conduct electrical current weakly (ionize partially!)

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12
Q

non-electrolytes…

A

do not conduct electrical current (no ionization!)

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13
Q

what are strong electrolytes

A

substances that ionize completely when dissolved in water (ex. soluble salts, strong acids, strong bases)

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14
Q

generally soluble ions

A

Li+, Na+, K+, NH4^+, NO3^-, C2H3O2^-, Cl-, Br-, I-, SO4^2-

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15
Q

Cl-, Br-, I- solubility EXCEPTIONS

A

Ag+, Hg2^2+, Pb^2+ make compound INSOLUBLE

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16
Q

SO4^2- solubility EXCEPTIONS

A

Sr^2+, Ba^2+, Pb^2+, Ag+, Ca^2+ make compound INSOLUBLE

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17
Q

generally insoluble compounds

A

OH-, S^2-, CO3^2-, PO4^3-

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18
Q

OH- and S^2- insolubility EXCEPTIONS

A

Li+, Na+, K+, NH4^+, Ca^2+, Sr^2+, Ba^2+ make compound SOLUBLE

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19
Q

CO3^2- and PO4^3- insolubility EXCEPTIONS

A

Li+, Na+, K+, NH4^+ make compound SOLUBLE

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20
Q

strong acid list

A

HCl, HBr, HI, HNO3, H2SO4, HClO4

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21
Q

strong base list

A

NaOH, LiOH, KOH, Ca(OH)2, Ba(OH)2

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22
Q

weak acid

A

any acid that dissociates (ionizes) to only a slight extent in aqueous solution (ex. HC2H3O2, acetic acid)

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23
Q

non-electrolyte

A

substance that dissolves in water but doesn’t produce any ions

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24
Q

Molarity (M)

A

(moles of solute)/(liters of solution)

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25
Q

standard solution

A

solution whose concentration is accurately known

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26
Q

dilution equation

A

M1V1 = M2V2 (remember, DILUTION ONLY!)

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27
Q

molecular equation (ME)

A

shows reactants, products, and physical states

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28
Q

complete ionic equation (CIE)

A

shows everything as ions, INCLUDES “spectator” ions that do not participate in reaction

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29
Q

net ionic equation (NIE)

A

includes only the ions that participate in the reaction

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30
Q

when diluting, remember…

A

AAA (always add acid)

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31
Q

Bronsted-Lowry Definition

A

acid=proton (H+) donor, base=proton acceptor

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32
Q

the reaction is likely acid base if…

A

the products of the reaction include water and some ionic compound

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33
Q

volumetric analysis

A

technique for determining the amount of a certain substance by doing a titration

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34
Q

titration

A

carefully adding a solution of known concentration (titrant) to a solution of unknown concentration (analyte)

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35
Q

equivalence point (stoichiometric point)

A

when enough of the titrant has been added to react exactly with the analyte

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36
Q

indicator

A

a substance that’s added to the analyte that changes color at or very near the equivalence point (usually a weak acid/base)

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37
Q

endpoint

A

the point where the indicator actually changes color

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38
Q

successful titration

A

1) exact reaction b/w titrant and analyte must be known 2) equivalence point must be marked accurately 3) volume of titrant required to reach equivalence point must be known

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39
Q

acid-base titration

A

analyte and titrant are acid/base

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40
Q

gas-evolution reaction

A

when 2 aqueous solutions mix and produce a new gaseous substance that bubbles out of the solution

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41
Q

oxidation-reduction reactions (redox)

A

reactions in which one or more electrons are transferred

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42
Q

assigning oxidation state/# rules

A

1) oxidation state of pure element is zero 2) oxidation state of any monoatomic ion is equal to the charge on the ion 3) sum of oxidation states on a polyatomic ion equal the charge on the ion (0, -1, +3, etc) 4) fluorine in a molecule is always -1 5) hydrogen in a molecule is usually +1 6) oxygen is usually -2

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43
Q

oxidation

A

an increase in oxidation state (LOSS of electrons)

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44
Q

reduction

A

a decrease in oxidation state (GAIN of electrons)

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45
Q

Evangelista Torricelli

A

invented the mercury barometer (at sea level, barometer is at 760mmHg)

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46
Q

manometer

A

measures pressure of sample gas, heigh difference in sides indicates pressure of gas relative to atmospheric pressure

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47
Q

pressure conversions

A

1 atm = 760mmHg = 760 torr = 101,325 Pa = 29.92 inHg = 14.7 psi

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48
Q

Boyle’s law (Robert Boyle)

A

PV=k; pressure and volume are inversely proportional, P1V1=P2V2

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49
Q

Charles’s law (Jacques Charles)

A

V=bT; volume of gas is directly proportional to the temperature, T is in Kelvin, (V1/T1)=(V2/T2)

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50
Q

Avogadro’s law (Lorenzo Avogadro)

A

V=an; for a given gas at a constant temperature and pressure the volume is directly proportional to the number of moles of gas, (V1/n1)=(V2/n2)

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51
Q

Ideal gas law

A

PV=nRT, R = 0.08206 Latm/Kmol

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52
Q

molar mass of a gas

A

dRT/P, d=density

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53
Q

ideal conditions for a gas

A

1 atm, 273.15K, 22.4L

54
Q

Dalton’s law of partial pressures

A

P(total)= P1+P2+P3+…+Pn

55
Q

assumptions of ideal gas law (simplifications)

A

1) volume of the individual gas particles must not be important 2) forces among the particles must not be important

56
Q

Mole fraction (X)

A

the ratio of the number of moles of a given component in a mixture to the total number of moles in the mixture

57
Q

mole fraction equations

A

X1 = n1/n(total), X1 = P1/P(total)

58
Q

pressure/mole fraction equation

A

P1=X1(P(total)), n1 = X1(n(total))

59
Q

vapor pressure of water

A

pressure due to water vapor over liquid water; if calculating pressure of sample over water, make sure to subtract this from total pressure

60
Q

Kinetic Molecular Theory (KMT)

A

1) particles are so small compared with the distances b/w the, that the volume of individual particles can be assumed to be negligible (0) 2) particles are in constant motion, collisions of the particles with the walls of the container are the cause of the pressure exerted by the gas 3) particles are assumed to exert no forces on each other 4) average kinetic energy of a collection of gas particles is assumed to be directly proportional the the temperature in K of the gas

61
Q

Pressure and volume relation (Boyle)

A

if volume DECREASES then pressure INCREASES (inversely proportional) (all else constant)

62
Q

Pressure and temperature relation

A

if temperature INCREASES then pressure INCREASES (directly proportional) (all else constant)

63
Q

Volume and temperature relation (Charles)

A

if volume DECREASES then temperature DECREASES (all else constant)

64
Q

Volume and number of moles relation (Avogadro)

A

if number of moles INCREASES then the volume must INCREASE (all else constant)

65
Q

Kinetic energy equation

A

KE = (1/2)(mv^2)

66
Q

AVERAGE kinetic energy equation

A

KEavg=(2/3)(RT), if temperature increases so does energy (direct proportion)

67
Q

the lighter a gas is…

A

the larger its range of velocity, and the higher its peak velocity is, see how H2 has HIGH velocity

68
Q

as temperature increases, molecular velocity…

A

has a wider range/higher average

69
Q

diffusion

A

describes the mixing of gases; rate of diffusion is rate of mixture

70
Q

effusion

A

describes the passage of a gas through a tiny orifice into an evacuated chamber; rate of effusion is the rate of transfer into the chamber

71
Q

Grahams law of effusion (Thomas Graham)

A

the rate of effusion for a gas is inversely proportional to the square root of the mass of the gas

72
Q

effusion rate equation

A

(rate of effusion for gas 1= M2^1/2) / (rate of effusion for gas 2 = M1^1/2), [(mm2)^1/2]/[(mm1)^1/2]

73
Q

van der Waals equation (Johannes Diderik van der Waals)

A

P+(an^2/V^2)=nRT

74
Q

correction for intermolecular forces

A

a(n/V)^2

75
Q

correction for particle volume

A

nb (used as V-nb)

76
Q

energy

A

the capacity to do work or produce heat

77
Q

law of conservation of energy

A

energy can be converted from one form to another but neither created nor destroyed

78
Q

potential energy

A

energy due to composition or position

79
Q

kinetic energy

A

energy due to the motion of the object

80
Q

heat

A

the transfer of energy between two objects due to a temperature difference

81
Q

temperature

A

a property reflecting the motions of particles in a substance

82
Q

work

A

a force acting over a distance (W = F x d), (W = -PΔV)

83
Q

system

A

the part of the universe on which we wish to focus attention

84
Q

surroundings

A

everything else in the universe (other than the system)

85
Q

exothermic

A

when a reaction gives off heat (E flows FROM the system)

86
Q

endothermic

A

when a reaction absorbs heat (E flows INTO the system)

87
Q

change in potential energy (ΔPE)

A

the energy required to break the bonds in the reactants minus the energy released when the bonds in the products are formed

88
Q

energy of a system equation

A

ΔE = q + w, q=heat and w=work

89
Q

enthalpy

A

H = E + PV

90
Q

calorimeter

A

the device used experimentally to determine the heat associated with a chemical reaction

91
Q

calorimetry

A

the science of measuring heat

92
Q

specific heat capacity

A

heat capacity of a substance in grams (J/°Cg or J/Kg)

93
Q

molar heat capacity

A

heat capacity of a subtance in moles (J/°Cmol or J/Kmol)

94
Q

wavelength (λ)

A

the distance b/w two consecutive peaks or troughs in a wave, unit is m or nm

95
Q

frequency (ν)

A

the number of waves (cycles) per second that pass a given point in space, unit is Hertz or 1/s

96
Q

speed of light equation

A

c = (ν)(λ)

97
Q

speed of light

A

c = 2.9979 x 10^8 m/s

98
Q

electromagnetic spectrum (in DECREASING wavelength)

A

radio, microwave, infrafred, visible, ultraviolet, X-ray, Gamma ray

99
Q

red light wavelength

A

625nm-750nm

100
Q

orange light wavelength

A

590nm-625nm

101
Q

yellow light wavelength

A

565nm-590nm

102
Q

green light wavelength

A

500nm-565nm

103
Q

blue light wavelength

A

450nm-500nm (lighter blue has longer wavelength)

104
Q

violet light wavelength

A

380nm-450nm

105
Q

Planck’s constant

A

h=6.62610^-34 Js

106
Q

Planck’s equation

A

ΔE=hν (planck’s constant times frequency)

107
Q

quantum

A

quantized packets, only was energy can be absorbed or emitted

108
Q

quantum equation/energy of photon equation (Einstein)

A

E(photon)= hc/λ, also used for ΔE for an electron

109
Q

de Broglie equation

A

λ = h/mv (v is velocity)

110
Q

1 J equals…

A

(1 kg*m^2)/(s^2)

111
Q

emission spectrum

A

range of electromagnetic radiation of various wavelengths

112
Q

quantum model (Niels Henrik David Bohr)

A

electron in a hydrogen atom moves around the nucleus only in certain allowed orbits

113
Q

Bohr constant

A

E = -2.178 x 10^-18 J

114
Q

Bohr electron transition equation

A

ΔE = -2.178 x 10^-18 J [(1/nf^2) - (1/ni^2)]

115
Q

problems with Bohr model

A

1) only works with things that work as a circle (2D things that can be approximated as a circular thing) 2) only works on ONE electron systems

116
Q

Schrodinger equation

A

Hψ=Eψ; H is the Hamiltonian operator/multiplier, ψ is the wavefunction (cartesian coordinate representation of the orbital), E is the summed energy of the atom

117
Q

orbital

A

probability of finding an electron in a given area of space around the nucleus in all 3 dimensions

118
Q

Heisenberg uncertainty principle

A

there is a fundamental limitation to just how precisely we can know both the position and the momentum of a particle at any given time

119
Q

Heisenberg equation

A

ΔxΔ(mv) = h/4pi, or ΔxΔp = h/4pi; Δx is changing position of electron and Δmv is the changing momentum; the MORE we know about position, the LESS we know about velocity and vice versa

120
Q

ψ^2

A

probability density/distribution

121
Q

orbital size defintion

A

size of an orbital is the radius of the sphere that encloses 90% of the total electron probability

122
Q

principal quantum number (n)

A

has integer values; related to the SIZE and ENERGY of the orbital. 1s,2p,3d,4f, etc

123
Q

angular momentum quantum number (l)

A

has integer values from 0 to n-1 for each n value; is related to the SHAPE of the atomic orbital, sometimes called subshells (ex. l=0, s orbital)

124
Q

magnetic quantum number (mₗ)

A

has integer values between -l and l INCLUDING zero; tells us about the relative orientation in space of the orbital

125
Q

spin quantum number (mₛ)

A

either -1/2 or 1/2; indicates orientation of electron spin

126
Q

node

A

area of low electron probability (n-1 gives number of nodes)

127
Q

Pauli exclusion principle

A

in a given atom, no two electrons can have the same set of four quantum numbers; two electrons in same orbital must have OPPOSITE spins

128
Q

a given orbital…

A

can only hold 2 electrons and the occupied orbital must contain electrons of opposite spin

129
Q

shielding

A

the outermost electrons have other electrons between themselves and the nucleus leading to a net smaller attraction to the nucleus for the outermost electrons

130
Q

electron approximation

A

we approximate by saying an electron is a field of charge that is the net result of the nuclear attraction and the average repulsions of all other electrons

131
Q

types of energy in complex atoms

A

1) the kinetic energy of electrons moving around the nucleus 2) the potential energy of nuclear/electron attractions 3) the potential energy of electron/electron repulsions

Chemistry 121 (6 decks)

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