midterm 1 Flashcards
1
Q
define electromagnetic radiation (= electromagnetic energy or radiant energy)
A
energy propagated by perpendicular electric and magnetic fields that increase and decrease in intensity as they move through space as waves.
2
Q
define wavelength ( λ)
A
distance between identical points on adjacent waves, such as from the crest/trough of one wave to the crest/trough of the next
3
Q
what are the units for wavelength?
A
usually metres or nanometers, nm= 10^-9 metres or pico metres pm=10^-12
4
Q
define frequency
A
the number of complete waves per second that pass through a given point
5
Q
what are the units for frequency?
A
s^-1 = 1Hz
6
Q
define the speed of a wave
A
the distance it moves per unit time (m/s)
7
Q
state the equation relating, frequency, wavelength, and speed of light
A
v λ = c
8
Q
state the 7 parts of the electromagnetic spectrum
A
highest frequency:
gamma rays
x rays
ultraviolet
visible
infrared
microwave
radio waves
9
Q
define the amplitude
A
the height of the crest of the wave or the depth of the trough
10
Q
for an electromagnetic wave, the amplitude is related to the
A
intensity of the radiation, or its brightness in the case of visible light
11
Q
dimmer light
A
lower amplitude, less intense
12
Q
brighter light
A
higher amplitude, more intense
13
Q
what is the similarity between all electromagnetic waves?
A
they travel at the same speed through a vacuum
14
Q
state the 2 distinctions between energy and matter
A
- refraction and dispersion
- diffraction and interference
15
Q
refraction
A
- when a light wave passes from one medium into another at an angle other, the speed of the wave changes
- if the wave strikes the boundary between media at an angle other than 90, the change in speed causes a change in direction, and the wave continues at a different angle
16
Q
dispersion
A
white light separates (disperses) into its component colours when it passes through a prism or another reframing object because each incoming wave is refracted at a slightly different angle
17
Q
how does a particle of matter contrast to a wave of light in terms of refraction?
A
particle of matter does not undergo refraction, it just slows down gradually along a curved path (after entering water)
18
Q
diffraction (inc diagram)
A
when a wave strikes the edge of an object and bends around it. if the wave passes through a slit as wide as its wavelength, it bends around both edges of the slit and forms a semi-circular wave on the other side of the opening
p294
19
Q
how does a particle of matter contrast to a wave of light in terms of diffraction?
A
when you throw a collection of particles at a small opening, only some go through the opening.
20
Q
interference (inc diagram)
A
when waves of light pass through two adjacent slits, the nearby emerging circular waves interact through the process of interference to form a diffraction pattern (p295)
21
Q
describe the two possible outcomes for a diffraction pattern
A
- in phase: if the crests of the waves coincide, the interfere constructively - the amplitudes add together to form a brighter region in the diffraction pattern
- out of phase: if crests coincide with troughs, the interfere destructively - the amplitudes cancel to form a darker region
22
Q
how does a particle of matter contrast to a wave of light in terms of interference ?
A
particles passing through adjacent openings continue straight paths, some colliding and moving at different angles
23
Q
what did scientists observe that caused them to question their view of energy?
A
- blackbody radiation -> quantum theory of energy
- the photoelectric effect -> photon theory of light
24
Q
explain blackbody radiation
A
- when an object is heated, it emits radiation of various wavelengths, including visible light
- the light of maximum intensity that is emitted shifts to shorter wavelengths as the temperature increases (red by cooler, blue by hotter)
25
why was the blackbody radiation observed confusing?
the classical wave model could not explain the relationship between the energy given off by a hot object and the wavelength of the energy emitted
26
define a blackbody
an idealised object that absorbs all the radiation incident on it, eg a hollow cube with a small hole in one wall approximates one.
27
what did Planck theorise?
- an atom changes its energy state by emitting/absorbing one or more quanta (energy packet of fixed quantity)
- the energy of the emitted/absorbed radiation is equal to the difference in the atom's energy states.
28
planck's equation and explain a more popular alternative
ΔE = E(emitted/absorbed radiation) = Δnhv
atoms can change energy only by integer multiples of hv, so the smallest energy change occurs when Δn=1:
ΔE = hv
or ΔE = hc/λ
29
h
Planc's constant
30
v
frequency of radiation
31
E
energy of radiation
32
energy is directly proportional to ------- and indirectly proportional to --------
frequency; wavelength
33
describe the photoelectric effect
when monochromatic light of sufficient frequency shines on a metal plate, a current flows (p297)
34
two confusing features that led scientists to the idea of the photoelectric effect
presence of a threshold frequency:
- wave theory associates the energy of light with its amplitude (intensity) not frequency (color)
- it thus predicts that an electron would break free when it absorbed enough energy from light of any color
- for current to flow, the light shining on the metal must have a minimum frequency (diff metals = diff frequencies)
absence of a time lag:
- wave theory predicts that with dim light there would be a time lag before the current flows as the electrons would have to absorb enough energy to break free
- current flows the moment light of the min frequency shines on the metal, regardless of the light's intensity
35
describe photon theory
Einstein proposed that light is particulate, quantised into photons - tiny bundles of energy.
E(photon) = hv = ΔE(atom)
36
how does photon theory explain the two features of the photoelectric effect?
threshold frequency:
- the intensity (brightness) of a beam of light is related to the number of photons, but not to the energy of each.
- a photon of a certain minimum energy must be absorbed to free an electron from the surface
absence of time lag:
- an electron breaks free when it absorbs a photon of enough energy
- current is weak in dim light cos fewer photons of enough energy can free fewer electrons per unit time, but some current flows as soon as light of sufficient energy/freq strikes the metal
37
above the minimum frequency of light, the kinetic energy of the ejected electron increases with....
light frequency
38
increased light intensity increases ...... but not...
the number of ejected electrons; their kinetic energy
39
give a roadmap of new ideas of light and matter, from before 1900 to after
before 1900:
light: wave character - eg wavelength, frequency, diffraction
matter: particulate character - eg mass, position
experiments:
- photoelectric effect
- diffraction by electrons
- atomic line spectra
wave-particle duality:
light: both wave and particulate character
matter: both particulate and wave character
40
describe how the emission of light from atoms was confusing for scientists
the passage of electricity through gas of atoms causes atoms to emit light.
- scientists expected to observe continuous light (of all wavelengths)
- instead, they only observed discrete frequencies in the visible part of the spectrum
41
describe the Schrodinger model of the atom
H^ψ = Eψ (wave equation)
- has many solutions psi
42
what does ψ psi mean?
'wavefunction' or 'orbital'
= a mathematical function describing the shape of a wave
43
what does ψ^2 mean?
the probability of finding the electron at any point about the nucleus
44
each orbital is described by 3 quantum numbers, and electrons are described by a quantum number:
1. principal quantum number (n) - size of the orbital
2. angular momentum quantum number (l) - shape of the orbital
3. magnetic quantum number (ml or m) - orientation of the orbital
4. electron spin quantum number (ms or s) - +1/2 or -1/2, up or down
45
principal quantum number (n)
size of orbital
n = 1, 2, 3
46
angular momentum quantum number (l)
shape of orbital
l = 0 (s), 1 (p) ... n-1
47
magnetic quantum number (ml or m)
orientation of the orbital
m = -l, -l+1, ..., l-1, l
48
describe s orbitals
spherical, l=0
49
1s orbital
- radial probability distribution plot is highest slightly out from the nucleus
50
2s orbital
- has two regions of higher electron density
- radial probability distribution of the more distant region is higher than that of the closer one
- between the two regions is a spherical node
51
why is the radial probability distribution of the more distant 2s region higher than that of the closer one?
because the sum of psi^2 for it is taken over a much larger volume
52
what happens at the node?
the probability of finding the electron drops to 0
53
what is the result of the 2s orbital being larger than the 1s?
an electron in the 2s spends more time farther from the nucleus than it does when it occupies the 1s
54
3s orbital
- three regions of high electron density and two nodes
- highest radial probability is at the greatest distance from the nucleus
55
describe p orbitals
l=1
- has two regions (lobes) of high probability, one on either side of the nucleus
- nucleus lies at the nodal plane of this dumbbell shaped orbital
56
what do the three possible ml values of p orbitals -1,0,+1 refer to?
three mutually perpendicular orientations associated with the x, y, and z axis
57
a d orbital has any one of ------ orientations
five
58
describe d orbitals
- 4 of the 5 d orbitals have 4 lobes with two mutually perpendicular nodal planes between them and the nucleus at the junction of the lobes
- 3 of these orbitals lie in the xy, xz, and yz planes, with their lobes between the axes, and are called dxy, dxz, and dyz orbitals.
- a fourth, the dx^2-y^2 orbital, also lies in the xy plane, but its lobes are along the axes
- the 5th d orbital, the dz^2, has two major lobes along the z axis, and a donut shaped region girdling the centre
59
radial probability distribution for s orbitals
60
radial probability distribution for a p orbital
61
radial probability distribution for a d orbital
62
describe f orbitals
l=3; seven orbitals with seven orientations
63
what is the function of the Rydberg equation?
predicts the position and wavelength of any line in a given series
64
state the Rydberg equation
1/ λ = R ( 1 /n2,1 – 1 /n2,2 )
- λ is the wavelength of the line, n1 and n2 are positive integers with n2>n1, and R is the Rydberg constant
65
what does the value n1 determine?
the region of the electromagnetic spectrum in which the lines occur
66
n1= 1
Rydberg equation
(n2=2, 3, 4...) for the lines in the ultraviolet series
67
n1 = 2
Rydberg equation
(n2 = 3, 4, 5...) for the lines the visible series
68
n1 = 3
Rydberg equation
(n2= 4, 5, 6...) for the lines in the infrared series
69
describe main features of the Bohr model for the hydrogen atom
- the H atom has only certain energy levels called stationary states - electrons cannot be between states
- the lower the quantum number value (n), the smaller the radius of the orbit, the closer the electron is to the nucleus, and the lower the energy level
- E(photon) = ΔE(atom) = E(final) - E(initial)
70
when does the electron of a H atom move to an outer (higher energy) orbit?
if an H atom absorbs a photon whose energy equals the difference between lower and higher energy levels
71
when does an atom emit a photon?
if an H atom in a higher energy level returns to a lower energy level, the atom emits a photon whose energy equals the difference between the two levels
72
Why is an atomic spectrum not continuous?
an atom's energy is not continuous, but rather has only certain states
73
when an electron drops from outer orbits to the n=1 orbit,
photons with the energy of UV radiation are emitted, creating the ultraviolet series of lines
74
when an electron drops to the n=2 orbit,
photons of lower energy visible radiation are emitted, creating the visible series of lines
75
when electrons drop to the n=3 orbit,
photons of even lower energy infrared radiation are emitted, resulting in the infrared series of lines
76
what is the main limitation of the Bohr model?
fails for atoms with more than one electron - the electron-electron repulsions and additional nucleus-electron attractions that are present create much more complex interactions. also, electrons do not move in fixed, defined orbits.
77
Bohr's work leads to an equation for calculating the energy levels of an atom:
E = -2.18 x 10^-18 (Z^2/n^2)
78
how to find the difference in energy between two levels mathematically using Bohr's equation
ΔE = E(final) - E(initial) = -2.18 x 10^-18 (1/n^2(final) - 1/n^2(initial))
79
how to find the wavelength of a spectral line mathematically using Bohr's equation
sub ΔE into Bohr's model to derive Rydberg's equation (p302)
80
the wavelength of any particle of mass m moving at speed u can be calculated using the equation
de Broglie wavelength eq:
λ = h/mu
81
what does de Broglie's wavelength imply?
matter behaves as though it moves in a move. An object's wavelength is inversely proportional to its mass.
82
give the application for De Broglie's equation but for photons
λ = h/p
where p is the momentum and h Planck's constant
83
describe Heisenberg's uncertainty principle
For a particle with constant mass m,
Δx * mΔu >/ h/4pi
where Δx is the uncertainty in position, Δu is the uncertainty in speed, and h is Planck's constant
84
define the Pauli exclusion principle
requires each electron in an atom to have a unique set of four quantum numbers; therefore, an orbital can hold no more than two electrons and their spins must be paired (opposite)
85
give 3 ways in which electrostatic interactions determine sub level energies
1. greater nuclear charge lowers sub level energy, making electrons harder to remove
2. electron-electron repulsions raise sub level energy, making electrons easier to remove. repulsions shield electrons from the full nuclear charge, reducing it to an effective nuclear charge, Z(eff). Inner electrons shield outer electrons very effectively.
3. penetration makes an electron harder to remove because nuclear attraction increases and shielding decreases. therefore, an energy level is split into sub levels with the energy order s
86
define penetration and give 2 examples of its effects
the ability of an electron to be closer to the nucleus
- increases nuclear attraction for a 2s electron over that for a 2p electron
- decreases the shielding of a 2s electron by the 1s electrons
87
define the Aufbau principle
electrons will fill orbitals with lowest energies first
88
Hund's rule
orbitals of equal energy become half-filled, with electron spins parallel, before any pairing of spins occurs
89
what are considered valence electrons for main group elements vs transition elements?
- MG: in the outer (highest energy level) only
- T: (n-1)d electrons are also considered valence electrons
90
describe how to figure out which sub shell will fill first in a many-electron atom
draw the orbitals out
arrows diagonal to the left
91
Z(eff) =
Z(actual) - electron shielding
92
draw out the s, p, d, and f blocks
s d p and lanthanides + actinides are f
93
the elements of a group have similar
outer electron configurations and similar chemical behaviour
94
describe two factors affecting atomic size
1. changes in n; as the principal quantum number (n) increases, the probability that outer electrons spend most of their time farther from the nucleus increases as well; atomic size increases.
2. changes in Z(eff): as effective nuclear charge increases, outer electrons are pulled closer to nucleus; atomic size decreases
95
atomic radius generally ----- down a group
increases
96
atomic radius generally ----- across a period
decreases
97
define atomic size
half the distance between nuclei of adjacent atoms
98
what is the trend in atomic size for a transition series?
size remains relatively constant
99
define first ionisation energy
energy required to remove a mole of electrons from a mole of gaseous atoms or ions
100
relationship between first ionisation energy and atomic size
inversely related
101
why do successive ionisation energies of an element show a very large increase after all valence electrons have been removed?
the first inner (core) electron is in an orbital of much lower energy so is held very tightly
102
define electron afinity
the energy involved in adding a mole of electrons to a mole of gaseous atoms or ions
103
the 1s22s22px2 state of a carbon atom is
excited
104
describe sub shells for a many-electron atom
not degenerate (due shielding)
105
as n increases, the size of the orbital and number of nodes ------
increase
106
define shielding and effective nuclear charge
Shielding occurs when inner electrons protect or shield outer electrons from the full nuclear attractive force. The
effective nuclear charge is the nuclear charge an electron actually experiences. As the number of inner electrons
increases, shielding increases, and the effective nuclear charge decreases
107
what is penetration?
Penetration occurs when the probability distribution of an orbital is large near the nucleus, which results in an increase of the overall attraction of the nucleus for the electron, lowering its energy.
108
define ionic bonding
the electrostatic forces of attraction between oppositely charged ions
109
Why do bonds form?
Because the molecule has a lower energy (i.e. is more stable) than its separated atoms
110
Why do ionic solids have a non-directional bond?
because oppositely charged ions are attracted to each other in all directions
111
define covalent bonding
the electrostatic forces of attractions between the shared pairs of electrons and positively charged nuclei of non-metals
112
define electronegativity
the ability of an atom in a molecule to attract electrons toward itself
113
how does a polar covalent bonds rise?
the difference in electronegativity of the atoms leads to an unsymmetrical electron distribution
114
non polar covalent
- electronically symmetrical
- electronegativity difference = 0-0.4
115
polar covalent
- partial charges
- electronegativity difference<2.0
116
ionic bond
- full charges
- electronegativity difference>2
117
define a lewis structure
a representation of covalently bonded molecules
118
describe VSEPR
Valence Shell Electron Pair Repulsion Model
- observed shapes of molecule arise from each electron group around an atom arranging themselves as far away as possible from other electron groups to minimise repulsions between them
119
electron group
single bond, multiple (double/triple) bond or lone pair
120
five electron group arrangements of minimum energy seen in large majority of molecules
1. linear (2ED)
2. trigonal planar (3ED)
3. tetrahedral (4ED)
4. trigonal bipyramidal (5ED)
5. octahedral (6ED)
121
describe linear molecules (BP only)
- no of bond pairs
- bond angle
- 3 examples
- shape
- 2 bond pairs
- 180
- BeCl2, CO2 and all diatomic molecules
- straight line
122
describe trigonal planar molecules
- no of bond pairs
- bond angle
- 1 example
- shape
- 3 bond pairs
- 120
- BCl3
- flat peace-sign
123
describe tetrahedral molecules
- no of bond pairs
- bond angles
- 1 examples
- shape
- 4 bond pairs
- 109.5
- methane, CH4
- Eiffel Tower
124
describe trigonal bipyramidal molecules
- no of bond pairs
- bond angles
- 1 example
- shape
- 5 bond pairs
- 90 and 120
- PF5
- fidget spinner shot by an arrow
125
describe octahedral molecules
- no of bond pairs
- bond angles
- 1 example
- shape
- 6 bond pairs
- 90
- SF6
- christian cross that has been shot by an arrow
126
why does a lone pair have a greater repulsive effect than a bonding pair?
- lone pair electrons are localised to an atom, so they are closer to each other
127
state the hierarchy of repulsion
lp to lp> lp to bp> bp to bp
128
name the 7 types of lone pair inclusive molecule shapes
- trigonal pyramidal molecules
- v shaped/bent molecules
- square planar molecules
- square pyramidal molecules
- seesaw molecules
- t/arrow shaped molecules
- linear molecules
(two very silly Swiss singers terrify Lav)
129
describe trigonal pyramidal molecules
- which bpo shape are they like
- no of bp, lp, ed
- bond angles
- example
- shape
- like tetrahedral molecules but without top
- 3 bp, 1 lp, 4 ed
- 107
- NH3
- beheaded Eiffel tower
130
describe v shaped/bent molecules
- which bpo shape are they like
- no of bp, lp, ed
- bond angles
- example
type 1 (SO2 DB):
- like trigonal planar but without top
- 2 bp, 1 lp, 3ed
- 104.5
type 2 (H2O):
- like type 1 but extra pair
- 2 bp, 2 lp, 4 ed
- 104.5
131
describe square planar molecules
- which bpo shape are they like
- no of bp, lp, ed
- bond angles
- example
- shape
- like octahedral but without both vertical bits
- 4 bp, 2 lp, 6 ed
- 90
- XeF4
- cross laid on its side + 2lp above and below
132
describe square pyramidal
- which bpo shape are they like
- no of bp, lp, ed
- bond angles
- example
- shape
- like octahedral but without bottom vertical bit
- 5 bp, 1 lp, 6 ed
- 85-87.5
- BrF5
- cross laid on its side + stick up + lp below
133
describe see saw molecules
- which bpo shape are they like
- no of bp, lp, ed
- bond angles
- example
- like trigonal bipyramidal but without left side bit
- 4 bp, 1 lp, 5 ed
- 87.5-90 (equatorial-axial), 117 (e-e)
- SF4
134
describe t/arrow shaped molecules
- which bpo shape are they like
- no of bp, lp, ed
- bond angles
- example
- like trigonal bipyramidal but only side 3 + 2 lp
- 3 bp, 2 lp, 5 ed
- 87.5-90
- XeOF2
135
describe linear molecules (WITH lone pairs)
- which bpo shape are they like
- no of bp, lp, ed
- bond angles
- example
- shape
- like trigonal bipyramidal but only 2 vertical bits
- 2 bp, 3 lp, 5 ed
- 180
- I3-
- straight line
136
dipole moment
a measure of the separation of charge in a molecule arising from the unequal sharing of electrons in polar bonds
137
when a dipole is permanent, the molecule is
polar
138
distinguish between intermolecular and intramolecular forces
intermolecular forces are much weaker than intramolecular forces
139
4 types of intermolecular forces
1. ion-dipole
2. dipole-dipole
3. hydrogen bonds
4. London dispersion forces
140
what are intermolecular forces important in?
state of matter, solubility, boiling point, melting point
141
describe an ion-dipole interaction
- interaction between fully charged ion and partial charges of a polar molecule
- the energy of attraction increases with the charge of the ion and decreases with the square of the distance between the ion and dipole
142
describe dipole dipole interactions
polar molecules attract one another when they orient with unlike charges close together
143
describe hydrogen bonding
the dipole-dipole interaction that arises when a hydrogen atom is bonded to Nitrogen, Fluorine or Oxygen atoms with a lone pair of electrons
144
describe dispersion forces
London forces are caused by temporary dipoles which arise in atoms due to uneven distribution of electrons
145
define what it means for an atom's electron cloud to be polarisable
- susceptible to distortion by a neighbouring charge
- increases with the number of electrons which increases with molecular mass
146
macroscopic level of gases
properties
147
molecular level of gases
- structure
- dynamics (motion)
- intermolecular forces
148
equation and definition for pressure
P=F/A
= force exerted per area
149
SI units for pressure
Pa = N/m^2
kPa=10^3 Pa
150
1atm =
101.3 kPa, 760mmHg (Torr)
151
1 bar =
100.0 kPa
152
equation for pressure with density, g, and h
p x h x g, where g = 9.8m/s^2
153
what does the manometer lead to?
an equalisation of pressure:
P(atm) = P(Hg) = p(Hg)gh(Hg)
154
if the Hg in a manometer was replaced with water, would the height of the water be greater or less than Hg?
greater, due to water's lower density (P constant, p goes down, g constant, so h goes up)
155
describe kinetic molecular theory - the ideal gas
1. gases made of tiny particles moving completely randomly
2. total volume of particles very small compared to size of container
3. particles do not interact with each other
4. particle collisions are elastic (no energy lost)
5. kinetic energy (KE) increases with temperature
156
describe (5) KE increases with temperature
for a large collection of molecules:
- at a given temperature, all gases have the same distribution of kinetic energy
- each molecule: KE = 1/2mv^2
157
what is the effect of temperature on average kinetic energy?
the average kinetic energy increases with temperature
158
state 2 equations relating kinetic energy with temperature
KE(avg) = 3RT/2
Temperature is in Kelvin
KE is in J/mol
KE (avg) = 3RT/2Na
for a single gas molecule
159
units for R
J/molK
160
3RT/2Na =
1/2mu^2
161
state the 2 equations relating u, R, T, m, Na, and M
162
how is kinetic energy dependent on the mass of the particles?
at the same temperature, more massive molecules move slower
163
what does pressure result from?
gas particles colliding with container walls
164
Boyle's law
at constant T and fixed n, volume is inversely proportional to pressure, or V=k/P
165
low Pext = Pgas
-> high V
166
high Pext = Pgas
-> low V
167
Charles Law
at same P and fixed n, volume is proportional to temperature, or V= kT
168
Avogadro's law
V=kn
169
what equation describes an ideal gas?
PV = nRT
P - atm
V - L
T - K
170
P1V1/n1T1 =
P2V2/n2T2
171
mole fraction of gas A
Xa = na/ntotal = Pa/Ptotal
172
law of partial pressures
Pa + Pb + Pc +.... = PTotal
173
define effusion
escape of a pas through a hole into a vacuum
174
define diffusion
movement of one gas through another
175
crystalline solids
- well-ordered matter within the solid
- arrangement of atoms in the solid repeats itself
176
amorphous solids
don't have extensive ordering of particles
177
types of crystalline solids
- molecular solids
- covalent network solids
- metallic solids
- ionic solids
178
molecular solids
molecules held together by intermolecular forces, with relatively low melting points (eg ice or benzoic acid)
179
covalent network solids
extended structures of atoms held together by covalent bonds with very high melting points
180
allotropes
different structural forms of an element
181
metallic solids
- metallic bonding between atoms
- metal atoms as cations in sea of delocalised electrons
- high electrical conductivity
- malleable, ductile
182
ionic solids
held together by electrostatic attraction between cations and anions, with high melting points
183
describe how most liquids are molecular
intermolecular forces keep particles close together but not strong enough to keep particles from moving past each other
184
define surface tension
amount of energy required to expand a liquid surface
185
how does surface tension work?
at top - intermolecular attractions from below mean that the net attraction is down, causing surface to contract.
in the middle - there are intermolecular attractions in all directions; no net attraction
186
how does strength of forces between particles in a liquid affect surface tension?
the stronger the forces between particles in a liquid, the greater the surface tension
187
define capillary action
the rising of a liquid in a narrow space against the pull of gravity
188
cohesive molecules
between molecules
189
adhesive forces
between molecules and container walls
190
how does water rise in a capillary tube?
- strong H-bonding interactions between the water and glass (SiO2) pull the water up
- gravity and cohesive forces pull the liquid down
191
diagram for water in a tube when adhesive>cohesive and vice versa
192
gas -> solid
deposition
193
solid -> gas
sublimation
194
sold -> liquid
fusion (melting)
195
liquid -> gas
vaporisation
196
gas to liquid
condensation
197
liquid to solid
freezing
198
attractive forces in solids, liquids, gases
solid: many attractive forces
liquid: some attractive forces
gas: no attractive forces
199
breaking attractive forces
requires (absorbs) energy: endothermic, so ∆H is positive
- adding heat increases temperature, increases KE, and overcomes attractive forces
200
forming attractive forces
releases energy: exothermic ∆H is negative
201
graph for phase changes
moving from one state of matter to another
202
how does ∆H of vaporisation relate to ∆H of fusion?
∆H(vap)>∆H(fus) for all substances
203
what type of process is a phase change?
a spontaneous process
204
spontaneous
∆G (change in free energy) is negative
205
equation for ∆G
∆G=∆H-T∆S
both ∆H and ∆S determine whether a process occurs spontaneously
206
∆H =
change in enthalpy due to phase change
207
∆S =
change in entropy due to phase change
208
entropy, S
the greater the degree of randomness or disorder in a system, the greater its entropy
209
∆S =
S (final) - S (initial)
210
for a reaction to be favourable, ∆S must be
positive
211
solid -> liquid -> gas requires
heat and increases entropy
212
describe gas-liquid equilibrium and draw a graph for it
as equilibrium between Rate(vap) and Rate(cond) is achieved, gas pressure reaches a constant value
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vapour pressure
the pressure of gas in equilibrium (co-existing) with its liquid (or solid) at a specified temperature
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vapour pressure depends on
temperature and the strength of intermolecular forces
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effect of boiling point on vapour pressure
- the temperature at which vapour pressure of a liquid equals the external pressure
- enough kinetic energy to vaporise against external pressure; water can form vapour bubbles within
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normal boiling point
boiling point at 1atm
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effect of intermolecular forces on vapour pressure
greater intermolecular forces:
- fewer molecules escape liquid
- more KE required to overcome
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Colligative properties
solution properties that depend on the concentration of a solute, not its identity
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freezing point depression
solutes lower freezing point of the solvent
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boiling point elevation
solutes raise boiling point of the solvent
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vapour pressure lowering
solutes decrease vapour pressure of the solvent
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osmotic pressure
pressure that must be applied to a solution to prevent osmosis from a sample of a pure solvent
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how do non-volatile solutes act on the vapour pressure of the solvent?
they decrease it
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Raoul's law for non-volatile solute:
P = XP^0
P - vapour pressure of solvent in solution
X = mole fraction of the solvent
P^0 = vapour pressure of pure solvent
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volatile
something that can evaporate
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what is the difference between the effect of a volatile and non-volatile solute?
with a volatile solute, both the solvent and solute contribute to vapour pressure
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P total =
P sub 1 + P sub 2
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solutions that obey Raoul's law are called
ideal solutions
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osmosis
net flow of solvent through a semi permeable membrane from a dilute solution to more concentrated solution
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semi-permeable membrane only allows
solvent to go through, not solvent
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osmotic pressure
pressure that must be applied to the solution to prevent osmosis from a sample of pure solvent
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how do you signify osmotic pressure?
pi
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equation for osmotic pressure
pi = MRT
pi - osmotic pressure
M - total molarity of all solutes
R - gas constant
T - temperature (K)
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how do you know which value of R to use?
whichever one the units cancel out for
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dissociation of 0.1M NaCl
dissociation of 0.1M CaCl2
dissociation of 0.1M glucose
state the osmotic pressure for all
Na+ and Cl-
Ca2+ and 2Cl-
stays same
0.200M, 0.100, 0.300M
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state van't Hoff's factor
i = moles of particles in solution/moles of solute dissolved
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define a solution
one or more substances (solutes) mixed at the molecular level (dissolved) in a medium (solvent, usually liquid)
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two properties of a solution
- uniform mixing
- constant composition throughout
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percent A by mass =
mass of solute A/mass of solution x 100
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mole fraction (Xa) =
moles of solute A/moles of solution
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parts per million (ppm) A =
mass of solute A/mass of solution x 10^6
or
moles of solute A/moles of solution x 10^6
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what needs to occur for substances to form a solution?
the solute-solvent interaction has to overcome both solute-solute interactions and solvent-solvent interactions
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substances with ----- -------- ----- form solutions
similar intermolecular forces
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formula for solubility
solubility = k P
P - partial pressure of gas over the solution
k - Henry's Law constant - depends on gas and solvent (usually H2O)
