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Ch > H > Flashcards

H Flashcards

1
Q

Under conditions of given temperature and pressure a substance will be
found in one of these three states depending on:

A

Intermolecular forces V(r)
vs.
Kinetic energy 1/2mv²

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2
Q

phase

A

a form of matter that is uniform throughout in chemical composition and physical state.

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3
Q

as kinetic energy incerases

A

solid –> liquid –> gas

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4
Q

Kinetic energy and temperature

A

average kinetic energy of molecules is proportional to Temperature

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5
Q

avg kinetic energy formula

A

v¯ = √(8RT/πM)

M = molecular weight
R = ideal gas constant
T = temp (in K?)
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6
Q

molecules being separated and brought together

A
  • attractive forces bring molecules together

- kinetic energy separates molecules

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7
Q

phase transition

A

the spontaneous conversion of one phase to another phase. Occurs at a characteristic temperature (transition
temperature) for a given pressure.

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8
Q

Enthalphy of phase changes: pure substances

DHₜᵣₐₙₛᵢₜᵢₒₙ
note: says transition

A

For a pure substance each phase transition has a specific enthalpy change per mole

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9
Q

vapor pressure

A

pressure exerted by the vapor at equilibrium with its condensed phase at a given temperature

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10
Q

solid-liquid equilibria

A

when kinetic energy is high enough molecules at the surface start moving freely out of their positions in the solid.

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11
Q

solid-vapor equilibria

A

when kinetic energy is high enough molecules at the surface start moving freely out of their positions in the
solid.

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12
Q

phase diagrams of pure substances

A

check slides 1

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13
Q

vapor pressure and liquids

A

Vapour Pressure is the pressure that exists above a liquid in a sealed container

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14
Q

THE CLAUSIUS-CLAPEYRON EQUATION

A

ln(p) = △Hᵥₐₚ/RT + constant

Two point clausius-clapeyron:
ln(p₂/P₁) = -△Hᵥₐₚ/R x (1/T₂ - 1/T₁)

P = pressure
T = temp (K) 
△H = enthalpy of vaporisation
R = gas constant
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15
Q

solution

A

= a homogeneous mixture, with no boundaries separating its components

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16
Q

types of solution

A
  • heterogeneous eg. sand in water

- homogeneous eg. salt dissolved in water

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17
Q

solute

A

minority component in a solution.

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18
Q

solvent

A

majority component in a solution

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19
Q

concentration

A

amount of solute dissolved in solvent

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20
Q

Molarity formula

A

M = moles of solute/L of soln

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21
Q

Molality formula

A

m = moles of solute/mass (kg) of solvent

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22
Q

Molar fraction

A

x = moles of solute/(moles of solute + moles of solvent)

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23
Q

Molar fractions added up

A

molar fractions add up to 1

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24
Q

solubility

A

maximum amount of solute that can dissolve in a solvent at a given temperature

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25
Q

when will a solute spontaneously dissolve

A

solute will spontaneously dissolve if the formation of a solution leads to
the lowering of the total Gibbs energy (△G)

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26
Q

when spont solution will and wont form

A

if △G < 0 spont solution formation

if △G > 0 solution will not form

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27
Q 
relating the equilibrium
constant for the solubilization
reaction to the Standard Gibb’s
Free Energy change.
(formula)
A

check pg 9 on slide 2

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28
Q

what is entropy related to

A

related to the number of ways in which a system can distribute
its energy.

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29
Q

entropy of solution vs entropy of solute and solvent

A

solution usually has higher entropy than the pure solute and pure solvent:

-more possible interactions when solute and solvent are
mixed than when pure.
-greater number of microstates due to greater number of different interactions present in solution

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30
Q

entropy and kinetic energy

A

The more freedom of motion particles have, the more ways they can
distribute their kinetic energy…….hence

S𝓰ₐₛ > Sₗᵢᵩᵤᵢ𝒹 > Sₛₒₗᵢ𝒹

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31
Q

△Sᵥₐₚ
△S𝒻ᵤₛ

formulas

A

△Sᵥₐₚ = S𝓰ₐₛ - Sₗᵢᵩᵤᵢ𝒹 > 0

△S𝒻ᵤₛ = Sₗᵢᵩᵤᵢ𝒹 - Sₛₒₗᵢ𝒹 > 0

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32
Q

△H°ₛₒₗᵤₜᵢₒₙ

A

△H°ₛₒₗᵤₜᵢₒₙ = △H°(lattice) + △H°(hydration)

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33
Q

probability molecule of a solvent (A) will go into gas phase

A

probability that a molecule of A will go into the gas phase is proportional to the probability of finding A at the liquid surface

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34
Q

probability that a molecule of solvent (A) will go into the liquid phase

A

probability that a molecule of A will go into the liquid phase is proportional to the probability of finding A in the gas phase.

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35
Q

rate of vaporization

A

rate of vaporization = k*xₐ

  • probably means multiply lol
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36
Q

rate of condensation

A

rate of condensation = k’*pₐ

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37
Q

Raoult’s Law

A

At equilibrium condensation and vaporization rates must be equal

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38
Q

Raoult’s law formula

A

pₐ = (xₐ)(p*ₐ)

pₐ = vapor pressure of solvent
xₐ = mole fraction of A
p*ₐ = vapor pressure of pure solvent A
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39
Q

Raoult’s Law diagram

A

check pg.5 of slide 3

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40
Q

ideal solutions

A

solutions that obey Raoult’s law throughout the composition range from pure A to pure B.

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41
Q

vapor pressure and molar fraction

A

At very low concentrations the vapor pressure of the solute is proportional to the molar fraction but with a constant different from p*A

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42
Q

Henry’s Law

A

S𝓰ₐₛ = P𝓰ₐₛ x Kₕ

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43
Q

ideal-dilute solutions

A

solutions in which the solvent obeys Raoult’s law and the solute obeys Henry’s law.

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44
Q

pressure and gas solubility

A

Pressure has a major effect on the solubility of gases in solution

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45
Q

colligative property

A

A colligative property of a solution depends not on the nature of the chemical species dissolved in the solution but only on the amount of solute dissolved in the solution

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46
Q

examples of colligative properties

A
  • vapor pressure lowering
  • boiling point elevation
  • freezing point depression
  • osmosis
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47
Q

vapor pressure lowering

A
  • Raoult’s Law can be used to estimate change in vapor pressure of solvent A for ideal dilute solution of solute B
  • Assumed that only solvent obeys Raoult’s law. Lowering of vapor pressure does not depend on type of solute.
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48
Q

vapor pressure lowering equation

A

pₐ - pₐ = (xᵦ)(pₐ)

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49
Q

boiling point elevation

A

△T = (Kᵦ)(mᵦ)

mᵦ = molality of B
Kᵦ = ebullioscopic cosntant or boiling point elevation constant
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50
Q

when boiling occurs

A

when the vapor pressure of the solvent is equal to that of the atmosphere

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51
Q

Adding a solute to solvent and effects on phase boundaries

A

-eg. adding NaCl to water leads to an inc in boiling point of water

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52
Q

why add ice to roads

A

addition of salt to the roads lowers the freezing point of water, this means that colder temperatures are required to form ice

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53
Q

Freezing point depression

A

△T = (K𝒻)(mᵦ)

Kf = cryoscopic constant

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54
Q

osmotic pressure (π)

A

π = hydrostatic pressure due to △h

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55
Q

osmosis

A

π = nᵦ/Vₛₒₗᵤₜᵢₒₙ x RT = cRT

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56
Q

hydrostatic pressure

A

a colligative property

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57
Q

applications of osmosis

A

responsible for movement of nutrients and water across cell walls.

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58
Q

colligative properties and ionic solutes

A

In this case, remember that the concentrations to use in the colligative property formula is that of TOTAL number of chemical species:

△T = i(Kᵦ)(mᵦ)
△T = i(K𝒻)(mᵦ)
π = icRT
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59
Q

i

A

van’t hoff factor, ideally equal to # of particles

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60
Q

what constitutes 1 phase

A

two miscible gases/liquids or solids

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61
Q

components

A

minimum number of independent species necessary to define the composition of ALL the phases in the system.

eg. sugar soln has 2 components (sugar + water) but a single phase

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62
Q

total pressure in vapor phase at equilibrium equation

A

p = (pₐ - pᵦ)(xₐ) + p*ᵦ

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63
Q

phase changes for one component systems vs multi component systems

A

one comp systems - phase changes at well defined temps

multi comp - phase changes over a range of temps

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64
Q

use of binary phase diagram

A

determine

  • phases present
  • composition of phases
  • relative fractions of phases
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65
Q

efficiency of fractionating column

A

expressed as the number of theoretical plates.

66
Q

azeotrope

A

check slides 5 pg 12

max and min in T vs Xₐ diagrams

67
Q

Eutectic mixture/

composition

A

mixture composition that melts at lowest temp in a binary diagram.

68
Q

Eutectic temperature

A

Temp. at which E melts/freezes.

69
Q

Eutectic point

A

point defined by Tₑ and xₑ

note e = E, just cant subscript it

70
Q

notes on equilibrium

A
  • system must be closed
  • equilibrium reached spontaneously
  • system shows no macroscopic evidence of change
  • equilibrium is dynamic, ie. microscopic change takes place
  • same equil conditions obtained at given temp
71
Q

at equilibrium,

A

rate𝒻𝓌𝒹 = rateᵣₑᵥ

72
Q

reaction quotient: reaction shifted towards products

A

Q increases

73
Q

reaction quotient: reaction shifted towards reactants

A

Q decreases

74
Q

Law of Mass Action

A

At equilibrium,

Q = Kₑᵩ = [C]ᶜ[D]ᵈ/[A]ᵃ[B]ᵇ

75
Q

comparing Q and K

A

Q < K
[reactant] large wrt [product] –> inc prod formation until equil reached

Q > K
[product] large wrt. [reactant] –> inc reactant to reach equil

Q = K
system at equilib

76
Q

second law of thermodynamics

A

a spontaneous process is one that results in an increase of entropy in the universe

△S(universe) = △S(system) + △S(surroundings)

under standard conditions:
△S°(universe) = △S°(system) + △S°(surroundings)

77
Q

△S°

A

△S°(system) = ΣS°(products) - ΣS°(reactants)

78
Q

Δ𝐻°

A

Δ𝐻° = ΣΔH𝒻°(prod) - ΣΔH𝒻°(react)

79
Q

gibbs free energy: when is equilibriumr eached?

A

when minimum G is reached:

check slides 6 pg 19

80
Q

writing equilibrium constants: aqueous solutions

A

Kₑᵩ = [C]ᶜ[D]ᵈ/[A]ᵃ[B]ᵇ

81
Q

writing equilibrium constants: gases

A

Kₑᵩ = pₚᵣₒ𝒹/[(pᵣₑₐ𝒸ₜ)(pᵣₑₐ𝒸ₜ)]

p = partial pressure??

82
Q

writing equilibrium constants: pure solids and liquids

A

K𝓌 = [A]ᵃ[B]ᵇB]/1

eg. Water dissociation:
H2O H+ + OH-

Kw = [H+][OH-]/1

83
Q

manipulating equilibrium constants: inverting a reaction

A

K₁

K₂ = 1/K₁

84
Q

manipulating equilibrium constants: multiplying a reaction by a constant

A

K₁

K₂ = K₁ⁿ

85
Q

manipulating equilibrium constants: combining reactions and associated equilibrium constants

A

K₁
K₂

Kₑᵩ = K₁K₂

86
Q

eg of manipulating equil constants

A

pg 7 on slides 7

87
Q

Le Chatelier’s Principle

A

when chem system at equil is disturbed, it reattains equil by undergoing a net reaction that reduces effect of disturbance

88
Q

sources of disturbance to a system

A
  • change in reactant or product
  • change in pressure
  • change in temp

disturbance shifts equil position to right or left

89
Q

Equilibrium: change in concentration

A

if [R] inc, system will consume R and product more product (P)

If [P] increases sytem will counteract by converting P into R

90
Q

Equilibrium: change in pressure + ways to change pressure

A

-negligible in solids/liquids

WAYS TO CHANGE PRESSURE:

  • changing conc of gaseous component: effect on Q
  • Adding inert gas: gas does not appear in Q - will have no effect on Q
  • change volume of reaction vessel
91
Q

Equilibrium: changing volume of reaction vessel

A

Inc volume of container - all partial pressures become smaller

Due to pV=nRT, increasing V decreases P

92
Q

effects on equil

A

check chemistry hardback for summary on effects on equil

93
Q

equilibrium: changing temp

A

Inc T: rate of both process increases, but rate of endo process increased more

value of K changes

94
Q

Only thing that changes value of K

A

temperature

95
Q

Van’t Hoff Equation (expression for effect of T on Kₑᵩ)

A

ln(k₂/k₁) = △Hᵣₓₙ/R x (1/T₂ - 1/T₁)

96
Q

slope of Van’t Hoff plot

A

plot is:
X-axis: 1/T
Y-axis: ln(K)

slope = -△Hᵣₓₙ/R

97
Q

equilibria very shifted to right

A

Many reactions that are usually presented as going to completion are actually equilibria…. only VERY shifted towards products!

98
Q

partition constant

A

Equilibrium constant for the distribution of a solute between two immiscible phases. It is a measure of differential solubility.
The two phases usually chosen are a polar (e.g. water) and apolar (e.g. octanol).

99
Q

Partition constant

A

Kₒ𝓌 = [S]ₒ/[S]ₐᵩ

100
Q

rate

A

A rate is a change in some variable with respect to time

101
Q

chemical kinetics

A

study of reaction rates.

102
Q

reaction rate

A

change of reactant or product concentration vs. time.

103
Q

Reaction rate formula

A

rate = -△[R]/△t = △[P]/△t

  • rate is in mol/dm³/s
  • negative sign indicates reactant
[R] = reactant concentration
[P] = product concentration
104
Q

graph of [R[ and [P] over time

A

check slides 9 pg 9

105
Q

for a generic reaction, expression for the overall reaction rate is

A

aA + bB –> cC + dD

rate = -1/a △[A]/△t = -1/b △[B]/△t = +1/c △[C]/△t = +1/d △[D]/△t

check slides 9 pg 10 for graph and slope formula

106
Q

monitoring reaction rates

A
  • manometric methods

- conductometric methods

107
Q

manometric methods

A
  • used for gas phase

- products/reactants - can follow reaction rate by monitoring change in pressure

108
Q

conductometric methods

A
  • used for ionic dissociation

- soln becomes conductive + can monitor extent of conductivity

109
Q

measuring rate of a reaction

A
  • average rate
  • instantaneous rate
  • initial rate
110
Q

average reaction rate

A

over the course of a chosen time interval (△t = t(final) - t(initial))

-find slope of graph over chosen time period

111
Q

instantaneous rate

A
  • rate of reaction at specific moment int ime in reaction
  • slope of line at tangent to curve

△y/△x

112
Q

initial reaction rate

A

-instantaneous rate the moment reactants are mixed, ie. rate at t=0

113
Q

Rate law

A

equation that expresses the rate of reaction in terms of molar concentrations of the species in the overall chemical equation for the reaction at some time.

114
Q

reaction order

A

the power to in the rate law

-not the stoichiometric coefficients in balanced chem equation, they are the values of reaction orders determined from experimental observation

115
Q

zeroth order

A

rate of the reaction is not dependant on reactant concentration

116
Q

first order

A

rate of the reaction doubles when the concentration is doubled

117
Q

second order

A

rate of the reaction quadruples when concentration doubles

118
Q

rate (to do with orders)

A

rate = k[A]ⁿ[B]ᵐ

119
Q

integrated first order rate law

A

ln[A]ₜ ] -kt + ln[A]₀

[A]ₜ = conc of A at time t
[A]₀ = initial conc of A (@ t=0)
120
Q

integrated second rate law

A

1/[A]ₜ ] kt + 1/[A]₀

[A]ₜ = conc of A at time t
[A]₀ = initial conc of A (@ t=0)
121
Q

integrated rate law for 0th order reaction

A

[A]ₜ = -kt + [A]₀

122
Q

units of rate constant

A

-change depending on order of reaction

reaction order units
0 mol dm⁻³ s⁻¹
1 s⁻¹
2 dm³ mol⁻¹ s⁻¹

123
Q

half life of a reaction

A

time taken for initial conc of a reactant to fall to 1/2 its initial value

[A]ₜ₁/₂ = [A]₀/2

124
Q

half life formula

A

t₁/₂ = (ln(2))/k

125
Q

Arrhenius equation

A

k = Ae⁻ᴱᵃ/ᴿᵀ

Eₐ = activation energy
A = pre-exponential factor (units depend on reaction order as for k)
126
Q

activation energy

A

represents an energy barrier that colliding molecules must overcome to react

127
Q

activation energy graphs for endo and exo thermic graphs

A

revise

128
Q

effects of diff things on activation energy

A

revise

129
Q

catalysts

A

materials that are not consumed in the reaction but they can alter the rate of reaction

130
Q

enzymes

A

examples of naturally occurring catalysts used in biological reactions

131
Q

Temp effect on activation energy

A

As T increases more collisions reach the activation energy required for the reaction to occur and the rate increases

132
Q

catalysts effect on actiavation energy

A

act to lower the activation energy resulting in a larger value of k

-does not improve the yield in this case but does increase the rate of reaction

133
Q

Solubility product (Kₛₚ)

A

equil constnt for a solid and its dissolved product in a saturated solution

(diff to solubility)

134
Q

Kₛₚ

A

eg.

Kₛₚ = [Na+][Cl-]

135
Q

Qₛₚ

A

in a saturated solution Qₛₚ attains a special value called Kₛₚ

Qc = Qₛₚ

Qc = [Na+][Cl-]/NaCl]

136
Q

common ion effect

A

tells us if we add a second salt which contains a common ion, equil will adjust by decreasing conc of added ions.

137
Q

current

A

flow of charge with time

I = dq/dt

138
Q

electrolyte

A

yield charged species in soln

eg. ionic salts are strong electrolytes, completely dissociates into cations and anions when dissolved in water

139
Q

non-electrolyte

A

do not yield charged species in soln

140
Q

K𝓌

A

K𝓌 = [H+][OH-] = 10⁻¹⁴ @ 298K

K𝓌 = [H3O+][OH-]

[H3O+] = [OH-]

141
Q

hydronium ion

A

water also dissociates into ions, the H+ proton can’t exist as free H+ in solution, so it binds to another water molecule and forms hydronium ion - H3O+

142
Q

arrhenius acid

A

An acid is a substance which increases concentration of the hydronium ion (H3O+) when dissolved in water

143
Q

arrhenius base

A

A base is a substance which increases concentration of the hydroxide ion (OH-) when dissolved in water

144
Q

bronsted-lowry acid

A

acids are proton donors;

145
Q

bronsted-lowry base

A

bases are proton acceptors

146
Q

Lewis acid

A

acid accepts a pair of electrons

147
Q

Lewis base

A

base donates a pair of electrons

148
Q

pH formula

A

pH = -log[H3O+]

POH - -log[OH-]

pK𝓌 = pH + pOH = 14
@ 25°C

149
Q

acids and bases

A

substances that change the concentration of hydronium and hydroxide ions in water

150
Q

strong acid

A

fully dissociates in water, i.e. almost every molecule breaks up
to form H+ ions – Large Ka

151
Q

weak acid

A

partially dissociates in water – Small Ka

152
Q

strong base

A

fully dissociates in water, i.e. almost every molecule breaks up
to form OH- ions

153
Q

weak base

A

partially dissociates in water

154
Q

Kₐ Acid Dissociation Constant

A

-indicates strength of acid

If Kₐ&raquo_space; 1, acid fully dissociates
If Kₐ &laquo_space;1, only slightly dissociated

155
Q

Kₐ formula

A

Kₐ = [A-][H+]/[AH]

156
Q

conjugate acid-base pair

A

An acid and a base which differ only by the presence or absence of a proton

157
Q

amphoteric

A

substances that can both donate and accept a proton

eg. water

158
Q

Measuring pH

A

indicators
universal pH paper
electronic pH meter

159
Q

polyprotic acids

A

if an acid can donate more than 1 proton

160
Q

1st, 2nd, 3rd dissociation etc

A

Ka₁ > Ka₂ > Ka₃

161
Q

oxidation of iron

A
  • can be considered as acid-base reaction
  • under acidic conditions Fe reacts with gaseous oxygen to prod iron oxide
  • lower pH : greater rust formation

Ch (3 decks)

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