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CHM135 > UNIT 1 > Flashcards

UNIT 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.

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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

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3
Q

what are the units for wavelength?

A

usually metres or nanometers, nm= 10^-9 metres or pico metres pm=10^-12

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4
Q

define frequency

A

the number of complete waves per second that pass through a given point

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5
Q

what are the units for frequency?

A

s^-1 = 1Hz

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6
Q

define the speed of a wave

A

the distance it moves per unit time (m/s)

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7
Q

state the equation relating, frequency, wavelength, and speed of light

A

v λ = c

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8
Q

state the 7 parts of the electromagnetic spectrum

A

highest frequency:
gamma rays
x rays
ultraviolet
visible
infrared
microwave
radio waves

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9
Q

define the amplitude

A

the height of the crest of the wave or the depth of the trough

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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

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11
Q

dimmer light

A

lower amplitude, less intense

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12
Q

brighter light

A

higher amplitude, more intense

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13
Q

what is the similarity between all electromagnetic waves?

A

they travel at the same speed through a vacuum

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14
Q

state the 2 distinctions between energy and matter

A
  1. refraction and dispersion
  2. diffraction and interference
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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
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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

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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)

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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

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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.

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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)

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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
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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

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23
Q

what did scientists observe that caused them to question their view of energy?

A
  1. blackbody radiation -> quantum theory of energy
  2. the photoelectric effect -> photon theory of light
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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)
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25
Q

why was the blackbody radiation observed confusing?

A

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

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26
Q

define a blackbody

A

an idealised object that absorbs all the radiation incident on it, eg a hollow cube with a small hole in one wall approximates one.

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27
Q

what did Planck theorise?

A
  • 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.
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28
Q

planck’s equation and explain a more popular alternative

A

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

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29
Q

h

A

Planc’s constant

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30
Q

v

A

frequency of radiation

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31
Q

E

A

energy of radiation

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32
Q

energy is directly proportional to ——- and indirectly proportional to ——–

A

frequency; wavelength

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33
Q

describe the photoelectric effect

A

when monochromatic light of sufficient frequency shines on a metal plate, a current flows (p297)

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34
Q

two confusing features that led scientists to the idea of the photoelectric effect

A

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

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35
Q

describe photon theory

A

Einstein proposed that light is particulate, quantised into photons - tiny bundles of energy.

E(photon) = hv = ΔE(atom)

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36
Q

how does photon theory explain the two features of the photoelectric effect?

A

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

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37
Q

above the minimum frequency of light, the kinetic energy of the ejected electron increases with….

A

light frequency

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38
Q

increased light intensity increases …… but not…

A

the number of ejected electrons; their kinetic energy

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39
Q

give a roadmap of new ideas of light and matter, from before 1900 to after

A

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

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40
Q

describe how the emission of light from atoms was confusing for scientists

A

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

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41
Q

describe the Schrodinger model of the atom

A

H^ψ = Eψ (wave equation)
- has many solutions psi

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42
Q

what does ψ psi mean?

A

‘wavefunction’ or ‘orbital’
= a mathematical function describing the shape of a wave

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43
Q

what does ψ^2 mean?

A

the probability of finding the electron at any point about the nucleus

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44
Q

each orbital is described by 3 quantum numbers, and electrons are described by a quantum number:

A
  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
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45
Q

principal quantum number (n)

A

size of orbital
n = 1, 2, 3

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46
Q

angular momentum quantum number (l)

A

shape of orbital
l = 0 (s), 1 (p) … n-1

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47
Q

magnetic quantum number (ml or m)

A

orientation of the orbital
m = -l, -l+1, …, l-1, l

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48
Q

describe s orbitals

A

spherical, l=0

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49
Q

1s orbital

A
  • radial probability distribution plot is highest slightly out from the nucleus
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50
Q

2s orbital

A
  • 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
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51
Q

why is the radial probability distribution of the more distant 2s region higher than that of the closer one?

A

because the sum of psi^2 for it is taken over a much larger volume

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52
Q

what happens at the node?

A

the probability of finding the electron drops to 0

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53
Q

what is the result of the 2s orbital being larger than the 1s?

A

an electron in the 2s spends more time farther from the nucleus than it does when it occupies the 1s

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54
Q

3s orbital

A
  • three regions of high electron density and two nodes
  • highest radial probability is at the greatest distance from the nucleus
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55
Q

describe p orbitals

A

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

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56
Q

what do the three possible ml values of p orbitals -1,0,+1 refer to?

A

three mutually perpendicular orientations associated with the x, y, and z axis

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57
Q

a d orbital has any one of —— orientations

A

five

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58
Q

describe d orbitals

A
  • 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
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59
Q

radial probability distribution for s orbitals

A
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60
Q

radial probability distribution for a p orbital

A
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61
Q

radial probability distribution for a d orbital

A
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62
Q

describe f orbitals

A

l=3; seven orbitals with seven orientations

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63
Q

what is the function of the Rydberg equation?

A

predicts the position and wavelength of any line in a given series

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64
Q

state the Rydberg equation

A

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

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65
Q

what does the value n1 determine?

A

the region of the electromagnetic spectrum in which the lines occur

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66
Q

n1= 1
Rydberg equation

A

(n2=2, 3, 4…) for the lines in the ultraviolet series

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67
Q

n1 = 2
Rydberg equation

A

(n2 = 3, 4, 5…) for the lines the visible series

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68
Q

n1 = 3
Rydberg equation

A

(n2= 4, 5, 6…) for the lines in the infrared series

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69
Q

describe main features of the Bohr model for the hydrogen atom

A
  • 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)
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70
Q

when does the electron of a H atom move to an outer (higher energy) orbit?

A

if an H atom absorbs a photon whose energy equals the difference between lower and higher energy levels

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71
Q

when does an atom emit a photon?

A

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

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72
Q

Why is an atomic spectrum not continuous?

A

an atom’s energy is not continuous, but rather has only certain states

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73
Q

when an electron drops from outer orbits to the n=1 orbit,

A

photons with the energy of UV radiation are emitted, creating the ultraviolet series of lines

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74
Q

when an electron drops to the n=2 orbit,

A

photons of lower energy visible radiation are emitted, creating the visible series of lines

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75
Q

when electrons drop to the n=3 orbit,

A

photons of even lower energy infrared radiation are emitted, resulting in the infrared series of lines

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76
Q

what is the main limitation of the Bohr model?

A

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.

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77
Q

Bohr’s work leads to an equation for calculating the energy levels of an atom:

A

E = -2.18 x 10^-18 (Z^2/n^2)

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78
Q

how to find the difference in energy between two levels mathematically using Bohr’s equation

A

ΔE = E(final) - E(initial) = -2.18 x 10^-18 (1/n^2(final) - 1/n^2(initial))

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79
Q

how to find the wavelength of a spectral line mathematically using Bohr’s equation

A

sub ΔE into Bohr’s model to derive Rydberg’s equation (p302)

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80
Q

the wavelength of any particle of mass m moving at speed u can be calculated using the equation

A

de Broglie wavelength eq:
λ = h/mu

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81
Q

what does de Broglie’s wavelength imply?

A

matter behaves as though it moves in a move. An object’s wavelength is inversely proportional to its mass.

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82
Q

give the application for De Broglie’s equation but for photons

A

λ = h/p

where p is the momentum and h Planck’s constant

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83
Q

describe Heisenberg’s uncertainty principle

A

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

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84
Q

define the Pauli exclusion principle

A

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)

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85
Q

give 3 ways in which electrostatic interactions determine sub level energies

A
  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<p<d<f
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86
Q

define penetration and give 2 examples of its effects

A

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

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87
Q

define the Aufbau principle

A

electrons will fill orbitals with lowest energies first

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88
Q

Hund’s rule

A

orbitals of equal energy become half-filled, with electron spins parallel, before any pairing of spins occurs

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89
Q

what are considered valence electrons for main group elements vs transition elements?

A
  • MG: in the outer (highest energy level) only
  • T: (n-1)d electrons are also considered valence electrons
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90
Q

describe how to figure out which sub shell will fill first in a many-electron atom

A

draw the orbitals out
arrows diagonal to the left

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91
Q

Z(eff) =

A

Z(actual) - electron shielding

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92
Q

draw out the s, p, d, and f blocks

A

s d p and lanthanides + actinides are f

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93
Q

the elements of a group have similar

A

outer electron configurations and similar chemical behaviour

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94
Q

describe two factors affecting atomic size

A
  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
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95
Q

atomic radius generally —– down a group

A

increases

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96
Q

atomic radius generally —– across a period

A

decreases

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97
Q

define atomic size

A

half the distance between nuclei of adjacent atoms

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98
Q

what is the trend in atomic size for a transition series?

A

size remains relatively constant

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99
Q

define first ionisation energy

A

energy required to remove a mole of electrons from a mole of gaseous atoms or ions

100
Q

relationship between first ionisation energy and atomic size

A

inversely related

101
Q

why do successive ionisation energies of an element show a very large increase after all valence electrons have been removed?

A

the first inner (core) electron is in an orbital of much lower energy so is held very tightly

102
Q

define electron afinity

A

the energy involved in adding a mole of electrons to a mole of gaseous atoms or ions

103
Q

the 1s22s22px2 state of a carbon atom is

A

excited

104
Q

describe sub shells for a many-electron atom

A

not degenerate (due shielding)

105
Q

as n increases, the size of the orbital and number of nodes ——

A

increase

106
Q

define shielding and effective nuclear charge

A

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
Q

what is penetration?

A

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
Q

define ionic bonding

A

the electrostatic forces of attraction between oppositely charged ions

109
Q

Why do bonds form?

A

Because the molecule has a lower energy (i.e. is more stable) than its separated atoms

110
Q

Why do ionic solids have a non-directional bond?

A

because oppositely charged ions are attracted to each other in all directions

111
Q

define covalent bonding

A

the electrostatic forces of attractions between the shared pairs of electrons and positively charged nuclei of non-metals

112
Q

define electronegativity

A

the ability of an atom in a molecule to attract electrons toward itself

113
Q

how does a polar covalent bonds rise?

A

the difference in electronegativity of the atoms leads to an unsymmetrical electron distribution

114
Q

non polar covalent

A
  • electronically symmetrical
  • electronegativity difference = 0-0.4
115
Q

polar covalent

A
  • partial charges
  • electronegativity difference<2.0
116
Q

ionic bond

A
  • full charges
  • electronegativity difference>2
117
Q

define a lewis structure

A

a representation of covalently bonded molecules

118
Q

describe VSEPR

A

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
Q

electron group

A

single bond, multiple (double/triple) bond or lone pair

120
Q

five electron group arrangements of minimum energy seen in large majority of molecules

A
  1. linear (2ED)
  2. trigonal planar (3ED)
  3. tetrahedral (4ED)
  4. trigonal bipyramidal (5ED)
  5. octahedral (6ED)
121
Q

describe linear molecules (BP only)
- no of bond pairs
- bond angle
- 3 examples
- shape

A
  • 2 bond pairs
  • 180
  • BeCl2, CO2 and all diatomic molecules
  • straight line
122
Q

describe trigonal planar molecules
- no of bond pairs
- bond angle
- 1 example
- shape

A
  • 3 bond pairs
  • 120
  • BCl3
  • flat peace-sign
123
Q

describe tetrahedral molecules
- no of bond pairs
- bond angles
- 1 examples
- shape

A
  • 4 bond pairs
  • 109.5
  • methane, CH4
  • Eiffel Tower
124
Q

describe trigonal bipyramidal molecules
- no of bond pairs
- bond angles
- 1 example
- shape

A
  • 5 bond pairs
  • 90 and 120
  • PF5
  • fidget spinner shot by an arrow
125
Q

describe octahedral molecules
- no of bond pairs
- bond angles
- 1 example
- shape

A
  • 6 bond pairs
  • 90
  • SF6
  • christian cross that has been shot by an arrow
126
Q

why does a lone pair have a greater repulsive effect than a bonding pair?

A
  • lone pair electrons are localised to an atom, so they are closer to each other
127
Q

state the hierarchy of repulsion

A

lp to lp> lp to bp> bp to bp

128
Q

name the 7 types of lone pair inclusive molecule shapes

A
  • 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
Q

describe trigonal pyramidal molecules
- which bpo shape are they like
- no of bp, lp, ed
- bond angles
- example
- shape

A
  • like tetrahedral molecules but without top
  • 3 bp, 1 lp, 4 ed
  • 107
  • NH3
  • beheaded Eiffel tower
130
Q

describe v shaped/bent molecules
- which bpo shape are they like
- no of bp, lp, ed
- bond angles
- example

A

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
Q

describe square planar molecules
- which bpo shape are they like
- no of bp, lp, ed
- bond angles
- example
- shape

A
  • 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
Q

describe square pyramidal
- which bpo shape are they like
- no of bp, lp, ed
- bond angles
- example
- shape

A
  • 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
Q

describe see saw molecules
- which bpo shape are they like
- no of bp, lp, ed
- bond angles
- example

A
  • like trigonal bipyramidal but without left side bit
  • 4 bp, 1 lp, 5 ed
  • 87.5-90 (equatorial-axial), 117 (e-e)
  • SF4
134
Q

describe t/arrow shaped molecules
- which bpo shape are they like
- no of bp, lp, ed
- bond angles
- example

A
  • like trigonal bipyramidal but only side 3 + 2 lp
  • 3 bp, 2 lp, 5 ed
  • 87.5-90
  • XeOF2
135
Q

describe linear molecules (WITH lone pairs)
- which bpo shape are they like
- no of bp, lp, ed
- bond angles
- example
- shape

A
  • like trigonal bipyramidal but only 2 vertical bits
  • 2 bp, 3 lp, 5 ed
  • 180
  • I3-
  • straight line
136
Q

dipole moment

A

a measure of the separation of charge in a molecule arising from the unequal sharing of electrons in polar bonds

137
Q

when a dipole is permanent, the molecule is

A

polar

138
Q

distinguish between intermolecular and intramolecular forces

A

intermolecular forces are much weaker than intramolecular forces

139
Q

4 types of intermolecular forces

A
  1. ion-dipole
  2. dipole-dipole
  3. hydrogen bonds
  4. London dispersion forces
140
Q

what are intermolecular forces important in?

A

state of matter, solubility, boiling point, melting point

141
Q

describe an ion-dipole interaction

A
  • 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
Q

describe dipole dipole interactions

A

polar molecules attract one another when they orient with unlike charges close together

143
Q

describe hydrogen bonding

A

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
Q

describe dispersion forces

A

London forces are caused by temporary dipoles which arise in atoms due to uneven distribution of electrons

145
Q

define what it means for an atom’s electron cloud to be polarisable

A
  • susceptible to distortion by a neighbouring charge
  • increases with the number of electrons which increases with molecular mass
146
Q

macroscopic level of gases

A

properties

147
Q

molecular level of gases

A
  • structure
  • dynamics (motion)
  • intermolecular forces
148
Q

equation and definition for pressure

A

P=F/A
= force exerted per area

149
Q

SI units for pressure

A

Pa = N/m^2
kPa=10^3 Pa

150
Q

1atm =

A

101.3 kPa, 760mmHg (Torr)

151
Q

1 bar =

A

100.0 kPa

152
Q

equation for pressure with density, g, and h

A

p x h x g, where g = 9.8m/s^2

153
Q

what does the manometer lead to?

A

an equalisation of pressure:
P(atm) = P(Hg) = p(Hg)gh(Hg)

154
Q

if the Hg in a manometer was replaced with water, would the height of the water be greater or less than Hg?

A

greater, due to water’s lower density (P constant, p goes down, g constant, so h goes up)

155
Q

describe kinetic molecular theory - the ideal gas

A
  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
Q

describe (5) KE increases with temperature

A

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
Q

what is the effect of temperature on average kinetic energy?

A

the average kinetic energy increases with temperature

158
Q

state 2 equations relating kinetic energy with temperature

A

KE(avg) = 3RT/2
Temperature is in Kelvin
KE is in J/mol

KE (avg) = 3RT/2Na
for a single gas molecule

159
Q

units for R

A

J/molK

160
Q

3RT/2Na =

A

1/2mu^2

161
Q

state the 2 equations relating u, R, T, m, Na, and M

A
162
Q

how is kinetic energy dependent on the mass of the particles?

A

at the same temperature, more massive molecules move slower

163
Q

what does pressure result from?

A

gas particles colliding with container walls

164
Q

Boyle’s law

A

at constant T and fixed n, volume is inversely proportional to pressure, or V=k/P

165
Q

low Pext = Pgas

A

-> high V

166
Q

high Pext = Pgas

A

-> low V

167
Q

Charles Law

A

at same P and fixed n, volume is proportional to temperature, or V= kT

168
Q

Avogadro’s law

A

V=kn

169
Q

what equation describes an ideal gas?

A

PV = nRT
P - atm
V - L
T - K

170
Q

P1V1/n1T1 =

A

P2V2/n2T2

171
Q

mole fraction of gas A

A

Xa = na/ntotal = Pa/Ptotal

172
Q

law of partial pressures

A

Pa + Pb + Pc +…. = PTotal
PA = XA × Ptotal

173
Q

define effusion

A

escape of a pas through a hole into a vacuum

174
Q

define diffusion

A

movement of one gas through another

175
Q

how does a barometer work?

A

The mercury column in the mercury barometer stays up due to the force exerted by the atmosphere on the mercury
in the outer reservoir just balancing the gravitational force on the mercury in the tube. Its height adjusts according
to the air pressure on the reservoir. The column of mercury is shorter on a mountaintop as there is less atmosphere
to exert a force on the mercury reservoir. On a mountaintop, the air pressure is less, so the height of mercury it
balances in the barometer is shorter than at sea level where there is more air pressure

176
Q

Colligative properties

A

solution properties that depend on the concentration of a solute, not its identity

177
Q

freezing point depression

A

solutes lower freezing point of the solvent

178
Q

boiling point elevation

A

solutes raise boiling point of the solvent

179
Q

vapour pressure lowering

A

solutes decrease vapour pressure of the solvent

180
Q

osmotic pressure

A

pressure that must be applied to a solution to prevent osmosis from a sample of a pure solvent

181
Q

how do non-volatile solutes act on the vapour pressure of the solvent?

A

they decrease it

182
Q

Raoul’s law for non-volatile solute:

A

P = XP^0
P - vapour pressure of solvent in solution
X = mole fraction of the solvent
P^0 = vapour pressure of pure solvent

183
Q

volatile

A

something that can evaporate

184
Q

what is the difference between the effect of a volatile and non-volatile solute?

A

with a volatile solute, both the solvent and solute contribute to vapour pressure

185
Q

P total =

A

P sub 1 + P sub 2

186
Q

solutions that obey Raoul’s law are called

A

ideal solutions

187
Q

osmosis

A

net flow of solvent through a semi permeable membrane from a dilute solution to more concentrated solution

188
Q

semi-permeable membrane only allows

A

solvent to go through, not solvent

189
Q

osmotic pressure

A

pressure that must be applied to the solution to prevent osmosis from a sample of pure solvent

190
Q

how do you signify osmotic pressure?

A

pi

191
Q

equation for osmotic pressure

A

pi = MRT
pi - osmotic pressure
M - total molarity of all solutes
R - gas constant
T - temperature (K)

192
Q

how do you know which value of R to use?

A

whichever one the units cancel out for

193
Q

dissociation of 0.1M NaCl
dissociation of 0.1M CaCl2
dissociation of 0.1M glucose
state the osmotic pressure for all

A

Na+ and Cl-
Ca2+ and 2Cl-
stays same
0.200M, 0.100, 0.300M

194
Q

state van’t Hoff’s factor

A

i = moles of particles in solution/moles of solute dissolved

195
Q

define a solution

A

one or more substances (solutes) mixed at the molecular level (dissolved) in a medium (solvent, usually liquid)

196
Q

two properties of a solution

A
  • uniform mixing
  • constant composition throughout
197
Q

percent A by mass =

A

mass of solute A/mass of solution x 100

198
Q

mole fraction (Xa) =

A

moles of solute A/moles of solution

199
Q

parts per million (ppm) A =

A

mass of solute A/mass of solution x 10^6

or

moles of solute A/moles of solution x 10^6

200
Q

what needs to occur for substances to form a solution?

A

the solute-solvent interaction has to overcome both solute-solute interactions and solvent-solvent interactions

201
Q

substances with —– ——– —– form solutions

A

similar intermolecular forces

202
Q

formula for solubility

A

solubility = k P

P - partial pressure of gas over the solution
k - Henry’s Law constant - depends on gas and solvent (usually H2O)

203
Q

crystalline solids

A
  • well-ordered matter within the solid
  • arrangement of atoms in the solid repeats itself
204
Q

amorphous solids

A

don’t have extensive ordering of particles

205
Q

types of crystalline solids

A
  • molecular solids
  • covalent network solids
  • metallic solids
  • ionic solids
206
Q

molecular solids

A

molecules held together by intermolecular forces, with relatively low melting points (eg ice or benzoic acid)

207
Q

covalent network solids

A

extended structures of atoms held together by covalent bonds with very high melting points

208
Q

allotropes

A

different structural forms of an element

209
Q

metallic solids

A
  • metallic bonding between atoms
  • metal atoms as cations in sea of delocalised electrons
  • high electrical conductivity
  • malleable, ductile
210
Q

ionic solids

A

held together by electrostatic attraction between cations and anions, with high melting points

211
Q

describe how most liquids are molecular

A

intermolecular forces keep particles close together but not strong enough to keep particles from moving past each other

212
Q

define surface tension

A

amount of energy required to expand a liquid surface

213
Q

how does surface tension work?

A

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

214
Q

how does strength of forces between particles in a liquid affect surface tension?

A

the stronger the forces between particles in a liquid, the greater the surface tension

215
Q

define capillary action

A

the rising of a liquid in a narrow space against the pull of gravity

216
Q

cohesive molecules

A

between molecules

217
Q

adhesive forces

A

between molecules and container walls

218
Q

how does water rise in a capillary tube?

A
  • strong H-bonding interactions between the water and glass (SiO2) pull the water up
  • gravity and cohesive forces pull the liquid down
219
Q

diagram for water in a tube when adhesive>cohesive and vice versa

A
220
Q

gas -> solid

A

deposition

221
Q

solid -> gas

A

sublimation

222
Q

sold -> liquid

A

fusion (melting)

223
Q

liquid -> gas

A

vaporisation

224
Q

gas to liquid

A

condensation

225
Q

liquid to solid

A

freezing

226
Q

attractive forces in solids, liquids, gases

A

solid: many attractive forces
liquid: some attractive forces
gas: no attractive forces

227
Q

breaking attractive forces

A

requires (absorbs) energy: endothermic, so ∆H is positive
- adding heat increases temperature, increases KE, and overcomes attractive forces

228
Q

forming attractive forces

A

releases energy: exothermic ∆H is negative

229
Q

graph for phase changes

A

moving from one state of matter to another

230
Q

how does ∆H of vaporisation relate to ∆H of fusion?

A

∆H(vap)>∆H(fus) for all substances

231
Q

what type of process is a phase change?

A

a spontaneous process

232
Q

spontaneous

A

∆G (change in free energy) is negative

233
Q

equation for ∆G

A

∆G=∆H-T∆S
both ∆H and ∆S determine whether a process occurs spontaneously

234
Q

∆H =

A

change in enthalpy due to phase change

235
Q

∆S =

A

change in entropy due to phase change

236
Q

entropy, S

A

the greater the degree of randomness or disorder in a system, the greater its entropy

237
Q

∆S =

A

S (final) - S (initial)

238
Q

for a reaction to be favourable, ∆S must be

A

positive

239
Q

solid -> liquid -> gas requires

A

heat and increases entropy

240
Q

describe gas-liquid equilibrium and draw a graph for it

A

as equilibrium between Rate(vap) and Rate(cond) is achieved, gas pressure reaches a constant value

241
Q

vapour pressure

A

the pressure of gas in equilibrium (co-existing) with its liquid (or solid) at a specified temperature

242
Q

vapour pressure depends on

A

temperature and the strength of intermolecular forces

243
Q

effect of boiling point on vapour pressure

A
  • 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
244
Q

normal boiling point

A

boiling point at 1atm

245
Q

effect of intermolecular forces on vapour pressure

A

greater intermolecular forces:
- fewer molecules escape liquid
- more KE required to overcome

246
Q

why Hvap is greater than H fus

A

In the change from solid to liquid, the kinetic energy of the molecules must
increase only enough to partially offset the intermolecular attractions between molecules. In the change from
liquid to gas, the kinetic energy of the molecules must increase enough to overcome the intermolecular forces

247
Q

does Xe have a lower or higher boiling point than Iodine

A

Xe would have a lower boiling point than iodine. Both are nonpolar with dispersion forces, but the forces
between xenon atoms would be weaker than those between iodine molecules since the iodine molecules are more
polarizable because of their larger size

CHM135 (23 decks)

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