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CONDENSED Flashcards

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1
Q

is energy released or absorbed in sublimation

A

absorbd

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2
Q

is energy released or absorbed in deposition

A

released

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3
Q

density =

A

mass / volume

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4
Q

what is the resultant volume

A

the final volume

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5
Q

the number of particels in a mole of a substance is numerically equal to the

A

Avogadro’s constant, 6.02 × 10^23 mol-1

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6
Q

what is molar mass numerically equal to

A

relative atomic mass

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7
Q

percentage composition formula

A

mass of element in compound / molar mass of compound x 100%

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8
Q

molar volume of a gas

A

This states that one mole of a gas at STP occupies a volume of 22.7 dm3 (22700 cm3 or 0.0227 m3).

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9
Q

relationship between the amount of a gas (in mol) and its volume

A

amount in mol = vol (dm^3) / molar vol (22.7dm^3)

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10
Q

an ideal gas is

A

a gas that exhibigts the five postulates of the kinetic molecular theory, as well as obeying gas laws.

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11
Q

what are the gas laws

A

Boyle’s law, Charles law, Gay-Lussac’s Law

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12
Q

what temperature scale must you use for gas laws

A

kelvin

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13
Q

Boyles Law

A

at constant temperature, the pressure and volume of a fixed mass of an ideal gas are inversely proportional to each other.

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14
Q

Charles Law

A

at constant pressure, the volume of a fixed mass of an ideal gas, is directly proportional to its absolute temp in kelvin

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15
Q

Gay-Lussac’s law

A

at constant volume the pressure of a fixed mass of an ideal gas is directly proportional to its absolute temperature (in kelvin).

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16
Q

Combined gas law:

A

PV/T = k
P1V1/T1=P2V2/T2

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17
Q

ideal gas equation

A

PV = nRT

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18
Q

cm cubed to m cubed

A

divide by 10^6

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19
Q

convert dm cubed to m cubed

A

divide by 10^3

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20
Q

with real gases, which of the assumptions made about ideal gases no longer apply under certain conditions.

A

At very high pressure the gas particles are closer together. Under these conditions, the actual volume of the particles becomes significant.
At low temperatures, the particles move less rapidly (have lower average kinetic energy). This means that there is a greater opportunity for intermolecular forces between the particles to have an effect.

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21
Q

when is the greatest deviation from ideal behaviour seen

A

when the gas is subjected to a low temperature and a high pressure

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22
Q

why is ideal behaviour deviated from at high pressure

A

spaces between particles lessen so intermolecular forces operate and particles become attracted to eachother. additionally, the non zero volume of the particles becomes significant in that the volume of a real gas at high pressure is higher than expected and PV/RT > 1.

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23
Q

which gases show the most ideal behaviour

A

low molar mass and weakest intermolecular forces

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24
Q

what is a standard solution

A

a solution with an accurately known concentration

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25
Q

what is a primary standard solution

A

it has:
High purity (99.9 %).
High molar mass.
Low reactivity.
Does not change composition in contact with air.

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26
Q

what is a secondary standard solution

A

a solution that has been standardised against a primary standard solution.

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27
Q

describe a titration

A

A titration is a method of volumetric analysis in which the concentration of a solution can be determined. In a titration, a burette is filled with a standard solution of known concentration (the titrant). A carefully measured volume of the solution with the unknown concentration (the analyte) is placed in a conical flask below the burette (Figure 2). An indicator is used to determine the end-point of the titration. In acid–base titrations, for example, a suitable indicator is added to the solution in the conical flask and the volume of titrant required to reach the end-point of the titration is recorded. Acid–base titrations are covered in more detail in section 8.2.2.

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28
Q

isotopes have the same chemical reactions because

A

they have the same number of electrons

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29
Q

isotopes have differnt physical properties becaues

A

they have more neutrons, so more mass

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30
Q

what is the mass spectrometer used for

A

to determine the realtive atomic masses of elements. it can also be used to determine the structure of organic compounds

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31
Q

how does mass spec work

A
  1. sample vapourised
  2. bombarded with high energy electrons
  3. cations produced
  4. cations accelerated in a electric field
  5. cations reach the detector where they produce a mass spec.

the cations are deflected in a magnetic field depending on their mass to charge ration (m/z). ions with a higher mass to charge ratio are deflected less in the mag
field than ion with alower m/z.

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32
Q

what is the shape of an s orbital

A

sphere

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33
Q

the main energy levels are split into…

A

sub levels which are assigned a number and the letter s,p,d or f

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34
Q

what is the shape of a p orbital

A

dumbbell

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35
Q

what does the pauli exclusion prinicple state

A

two electrons cannot have the same quantum number. two electrons can only occupy the same atomic orbital if they have opposite spins.

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36
Q

what is heisenberg’s uncertainty principle

A

it is not possible to know, at the same time, the exact position and momentum of an electron

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37
Q

n=1

A

s 2

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38
Q

n=2

A

s 2
p 6

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39
Q

n=3

A

s 2
p 6
d 10
f 14

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40
Q

aufbau principle

A

electrons fill atomic orbitals of lowest energy first

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41
Q

which sub level has the lowest energy

A

the 1s sub level

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42
Q

are s orbitals or p lower energy

A

s

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43
Q

degenerate orbitals

A

equal energy

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44
Q

there is an overlap between the 3d and 4s sub levels

A

this means the 4s sub level is of lower energy and fills before the 3d sub level

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45
Q

chromium electron configuration

A

1s2 2s2 2p6 3s2 3p6 4s1 3d5 or
1s2 2s2 2p6 3s2 3p6 3d5 4s1

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46
Q

copper electron configuration

A

1s2 2s2 2p6 3s2 3p6 4s1 3d10 or 1s2 2s2 2p6 3s2 3p6 3d10 4s1

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47
Q

hunds rule

A

electrons fill orbitals in the saem sub level singly before pairing up.

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48
Q

as freq increases

A

wavelength decreases

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49
Q

energy increases alongside

A

frequency

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50
Q

higher energy =

A

higher frequency = shorter wavelength.

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51
Q

lower energy =

A

lower frequency = longer wavelength

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52
Q

differences between spectra

A
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53
Q

when electrons are excited

A

they jump to higher energy levels

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54
Q

rel between energy and frequency formula

A

E = hv

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55
Q

Electron transitions to the n = 1 energy level

A

UV

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56
Q

Electron transitions to the n = 2 energy level

A

visible light

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57
Q

Electron transitions to the n = 3 energy level

A

IR

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58
Q

what is the highest energy end of each series of spectral lines known as

A

convergence limit

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59
Q

Plancks constant, h

A

6.6.3 x 10^-34 J/s

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60
Q

ground state

A

The ionisation of a hydrogen atom in its ground state corresponds to the electron transition from n = 1 to n = ∞. At this point, the electron is no longer attracted to the nucleus and the atom has been ionised.

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61
Q

what is the first ionisation energy of an element

A

the energy required to remove one mole of electrons from one mole of gaseous atoms to produce one mole of gaseous 1+ ions.

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62
Q

why are ionisation energies always positive (endothermic)

A

energy must be added to overcome the electrostatic attraction between the nucleus and valence electrosn

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63
Q

what is second ionisation energy

A

If an additional mole of electrons is removed from one mole of gaseous 1+ ions

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64
Q

why should we expect the ionisation energies to increase progressively as we remove negatively charged electrons from increasingly positive ions

A

it results in a stronger electrostatic attraction between the nucleus and the remaining electrons.

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65
Q

s block

A

groups 1 and 2

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66
Q

d block

A

groups 3 to 12

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67
Q

p block

A

13 to 18

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68
Q

f block

A

bottom of the periodic table

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69
Q

what is a transition element

A

defined as an element that has an incomplete d sub-level in its atom or one or more of its ions.

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70
Q

properties of transition elements

A

They have variable oxidation states.
They form coloured compounds.
The elements or their compounds show catalytic activity.
They form complex ions in solution.
The metals and their complexes show magnetic properties.

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71
Q

why is scandium the only element that cant have a +2 oxidation state

A

has only one unparied electron

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72
Q

what is the reason for these variable oxidation states is

A

the closeness in energy of the 3d and 4s sub-levels.

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73
Q

how to find the coordination number

A

number of ions in contact with other ions

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74
Q

what does the mp of an ionic compound depend on

A

the ionic charge and the ionic radius of its component ions

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75
Q

which exist as diatomic molecules

A

halogens
hydrogen
oxygen
nitrogen

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76
Q

what is different about diatomic oxygen and nirtogen

A

atoms are bonded via double and tripke bonds

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77
Q

coordiante covalent bond

A

both bonding electrons come from one atom

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78
Q

what is a dimer

A

a larger molecule composed of two identical smaller molecules and can be linekd by coordinate covalent bodns or by hyrdrogen bodns

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79
Q

what is bodn order

A

the number of bonds between a pair of atomso

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80
Q

over 1.8 units

A

ionic

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81
Q

1.8 units plus diff in electronegativity

A

ionic

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82
Q

0.5-1.7 units diff in electronegativity

A

polar covalent

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83
Q

0.1-0.4 units difference in electronegativity

A

non polar or weakly polar covalent

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84
Q

0 units electronegativity

A

pure covalent

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85
Q

when do pure covalent bodns occur

A

between atoms that have no differene in electronegativity, such as those in molecular oxygen chlorine nitgoen

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86
Q

when do polar covalent bonds occur

A

between atoms that have a difference in electronegativity of between 0.5 and 1.7 units

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87
Q

what causes a bond dipole

A

unequal sharing of electrons in a covalent bond

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88
Q

what does the octet rule state

A

the most stable arrangement is 8 electrons

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89
Q

exceptions to the octet rule

A

Hydrogen is stable with only two electrons in its outer shell.

Atoms such as boron, beryllium and aluminium (in compounds) are stable with fewer than eight electrons in their outer shell.

Atoms in period three and higher, such as sulfur, can form expanded octets with up to twelve electrons in their valence shell.

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90
Q

benefit of delocalised electrons

A

give greater stability to a molecule or polyatomic ion

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91
Q

silicon and silicon dioxide

A

giant covalent structure
tetrahedral
109.5 degree bond angle
strong covalent bonds
poor conductor of electricity

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92
Q

what are the allotropes of carbon

A

diamond
graphite
fullerenes

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93
Q

graphite

A

layered structure
planar sheets of hexagonally arranged carbon atoms
The layers are held together by relatively weak London dispersion forces.
Each carbon atom has an electron which becomes delocalised across the plane. The presence of delocalised electrons explains the ability of graphite to conduct electricity along the plane of the crystal when a voltage is applied.

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94
Q

fullerene

A

bond angle is 120
poor conductor of electricity
rings of 5 and 6 carbon atoms

fullerene C60 is by definition a simple molecular substance even though it contains so many bonded carbon atoms
The structure is made up of carbon atoms bonded together in 20 hexagons (six-carbon rings) and 12 pentagons (five-carbon rings), known as a truncated icosahedron.

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95
Q

graphene

A

Its tensile strength is 1000 times greater than steel.

It behaves as a semi-metal, so it is very suitable for electronic devices.

Adding only 1% content of graphene to plastics could allow those plastics to conduct electricity.

graphene is the most chemically reactive. This is because of the reactive edges of the structure, where there are carbon atoms with unoccupied (‘dangling’) bonds. This reactivity could be used in important ways. For example, membranes of graphene oxide have been shown to be preferentially permeable to water, which could be useful in desalination and water purification.

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96
Q

how many atoms is each carbon directly bonded to in these allotropes
diamond
graphite
fullerene

A

4
3
3

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97
Q

temporary dipole

A

caused by changes in electron density within an atom or molecule

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98
Q

waht does the strength of london forces depend on

A

The ease with which the electrons in an atom or molecule form a temporary or induced dipole (their polarisability).

The surface area of the molecule.

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99
Q

when does polarisability increase

A

alongside molar mass of a molecule. so does melting point and strength of London dispersion forces

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100
Q

what does the strength of the metallic bnd depend on

A

charge ont he metal ion and ionic radius of the metal ion. they also affect density

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101
Q

how is a sigma bond formed

A

direct head-on (axial) overlap of atomic orbitals
The two 1s atomic orbitals overlap head-on, forming a sigma bond.

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102
Q

pi bond formation

A

formed by the sideways overlap of two unhybridised p orbitals

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103
Q

double bond

A

one sigma one pi

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104
Q

which is stronger sigma or pi and why

A

sigma

The extra strength of the sigma bond comes from the greater overlap of the atomic orbitals in the bond. In a pi bond, the atomic orbitals cannot overlap as much which results in a weaker bond.

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105
Q

purpose of formal charge

A

used ot determin ewhich lewis structure is the preferred one when there is more than one possibility

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106
Q

formal charge equation

A

V - 1/2B - NB

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107
Q

the sum of formal charges in a neutral molecule must equal

A

zero

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108
Q

CFC’s

A

Chlorofluorocarbons (CFCs) are highly stable compounds. This stability was initially seen as an advantage but it meant that CFC molecules released into the lower atmosphere could remain intact and reach the upper atmosphere. Here, when exposed to UV radiation, compounds such as trichlorofluoromethane (CCl3F) decompose to produce chlorine free radicals (Cl*), as shown in the equation below.

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109
Q

standard conditions (Ɵ)

A

A pressure of 100 kPa.

A temperature of 25°C (298 K).

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110
Q

what is the enthalpy change of neutralisation

A

the enthalpy change when an acid and base react together to form one mole of water

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111
Q

waht is the molar enthalpy of combustion, or the standard enthalpy of combustion (ΔHƟc),

A

the enthalpy change when one mole of a substance is burned completely in oxygen under standard conditions.

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112
Q

Percentage error =

A

((experimental value - theoretical value) ÷ theoretical value) × 100

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113
Q

You should be aware of the limitations of calculating enthalpy changes in a school laboratory. These include but are not limited to:

A

Heat loss to the surroundings and heat absorbed by the calorimeter

Incomplete combustion of the fuel

Assumptions made about the specific heat capacity and density of aqueous solutions.

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114
Q

what does hess law state

A

the total enthalpy change in a chemical reaction is independent of the route by which the chemical reaction takes place, as long as the initial and final conditions are the same.

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115
Q

what is the standard enthalpy of formation

A

the enthalpy change when one mole of a compound is formed from the elements in their standard states under standard conditions.

116
Q

Enthalpy of formation values are useful in that they indicate

A

he stability of compounds in relation to their elements

117
Q

does hgaving a double bond increase or decrease reactivity and why

A

incease

This region of high electron density is the site of chemical reactivity within the molecule. This means that alkenes undergo addition reactions that take place across the carbon-carbon double bond.

118
Q

enthalpy change of a reaction using combustion

A

ΔH⦵ = ΣΔH⦵c (reactants) − ΣΔH⦵c (products)

119
Q

bond breaking is

A

endothermic (releases energy)

120
Q

bond making is

A

exothermic (absorbs energy)

121
Q

bond enthalpy aka

A

bond dissociation energy

122
Q

bond enthalpy definition

A

It is defined as the energy required to break one mole of chemical bonds in the gaseous state

123
Q

average bond enthalpy definitio n

A

when one mole of bonds are broken in the gaseous state averaged for the same bond in similar compounds

124
Q

The enthalpy change for a reaction can be calculated using average bond enthalpy values.

A

ΔH = ΣE(bonds broken) − ΣE(bonds formed)

125
Q

which has a higher wavelength oxygen or ozone

A

ozone

oxygen
λ < 242 nm
higher energy radiation of shorter wavelength
ozone
λ < 330 nm
lower energy radiation of longer wavelength

126
Q

benefit of the decomposition of oxygen

A

The two reactions in this cycle depend on different wavelengths of UV radiation, and the effect is to remove the higher energy radiation (λ < 242 nm), so that only the longer wavelength, which is less damaging radiation, reaches the Earth’s surface.

127
Q

thes strong double covalent bond in oxygen is disrupted by

A

the suns high energy UV-C raditation to form atoms which are free radicals since they have an unpaired electron. Such oxygen radicals can then react with an oxygen molecule to form ozone. The bonds in ozone, being weaker, can then be broken by the less energetic UV-B radiation (of longer wavelength) to reform oxygen and an oxygen free-radical.

128
Q

the surface of the earth is protected by these breakdown of ozone reactions from the damaging effects of

A

UV-B and UV-C radiation

129
Q

The lattice enthalpy (ΔH⦵lat) is

A

the enthalpy change when one mole of a solid ionic compound breaks down to form gaseous ions under standard conditions.

130
Q

The magnitude of the lattice enthalpy of an ionic compound can be thought of as

A

a measure of the strength of the ionic bonds between the ions.

131
Q

The magnitude of the lattice enthalpy depends on two factors:

A

the charge of the ions (ionic charge)

the size of the ions (ionic radii).

132
Q

the lattice enthalpy is directly proportional to

A

the product of the ionic charges and inversely proportional to the distance between the nuclei of the ions.

133
Q

lattice enthalpy increases alongside

A

ionic charge increasing or ionic radii decreasing

134
Q

lattice enthalpy

A

The enthalpy change that occurs when one mole of an ionic solid is broken down into its gaseous ions.

135
Q

purpose of born harber

A

to calculate lattice enthalpy

136
Q

enthalpy of formtation

A

the enthalpy change when one mole of a compound is formed from its elements in their standard states under standard conditions.

137
Q

enthalpy of atomisation

A

the enthalpy change when one mole of gaseous atoms is formed from an element in its standard state.

138
Q

Bond dissociation energy, E

A

the energy required to break one mole of bonds in the gaseous state.

139
Q

Ionisation energy, ΔH⦵IE:

A

the energy required to remove one mole of electrons from one mole of gaseous atoms.

140
Q

Electron affinity, ΔH⦵EA:

A

the energy released when one mole of electrons are added to one mole of gaseous atoms

141
Q

steps in born harber

A

Enthalpy of formation, ΔH⦵f

Enthalpy of atomisation, ΔH⦵at

Bond dissociation energy, E

Ionisation energy, ΔH⦵IE

Electron affinity, ΔH⦵EA

142
Q

why do stable ionic compounds tpicall have large negative values for enthalpy of formation

A

because as more energy is released, the compound becomes more stable

143
Q

waht is enthalpy of hydration

A

the enthalpy change when one mole of gaseous ions dissolves in water to form a solution of infinite dilution. It can be represented by the general equation:

X+ (g) → X+ (aq)

144
Q

what is a sol of infinite dilution

A

A solution of infinite dilution is a solution that has a large excess of water and the addition of more water would not cause any more heat to be released or absorbed.

145
Q

The enthalpy change of solution, ΔH⦵sol,

A

the enthalpy change when one mole of solute dissolves to form a solution of infinite dilution.

146
Q

enthalpy change of sol is the sum of

A

lattic and hydration enthalpies

147
Q

if a substnace has a high positive value fo the enthalpy of solution

A

generally insoluble

148
Q

factors that affect enthalpy of hydration

A

ionic radius and charge on ion

149
Q

the magnitude of enthalpy of hydration decreases whilst

A

ionic radius increases.
This is due to the weaker ion-dipole forces produced as the size of the ion increases.
charge decreases

150
Q

the charge density of an ion increases with

A

greater ionic charge and a decrease in ionic radius

151
Q

what does a higher charge density result in

A

a stronger ion-dipole force between the ion and the water molecule, and a greater, more negative value for the enthalpy of hydration.

152
Q

substances with a high positive (endothermic) value of ΔH⦵sol, are

A

less soluble

153
Q

is melting a positive or negative entropty change?

A

increase, positive tnropy

154
Q

is deposition a positive or negative entropty change?

A

decrease, negative entropy

155
Q

is sublimation a positive or negative entropty change?

A

increase, positive entropy

156
Q

dissolving a solute to form a solution positive or negative entropty change?

A

increase, positive entropy

157
Q

spontaneous meaning

A

A spontaneous process occurs without the addition of energy, other than that required to overcome the initial energy barrier (also known as the activation energy).

158
Q

For a spontaneous process, the total entropy of the system and surroundings (ΔStotal) must increase. This can be represented in equation form as:

A

ΔStotal = ΔSsystem + ΔSsurroundings ≥ 0

159
Q

how can negative entropy reactions be spontaneous

A

the entropy of the surroundings increases to a much greater extent, which gives an overall increase in entropy for the process.

160
Q

For the surroundings, the entropy change is given by the following relationship:

A

ΔSsurroundings=−ΔH/T

161
Q

entropy change for the universe

A

ΔSuniverse=ΔSsystem−ΔH/T

162
Q

gibbs free energy change formula

A

ΔG=ΔH−TΔS

163
Q

what is gibbs free energy

A

the energy associated with a chemical reaction that can be used to do work.

164
Q

This example shows us that there are three factors to be considered when determining the spontaneity of a reaction:

A

the sign of the ΔH

the sign of the ΔS

the temperature at which the reaction takes place.

165
Q

if gibbs is negative

A

reaction is spontaneous

166
Q

The Gibbs free energy of formation, ΔG⦵f, is defined as

A

the change in free energy when one mole of a compound is formed from its elements in their standard states under standard conditions.

167
Q

gibbs free energy change formula

A

ΔG⦵ = ΣΔG⦵f (products) − ΣΔG⦵f (reactants)

168
Q

requirements fo r a chemical reaction to take place

A

correct orientation
sufficent energy

169
Q

what is the transitions tate

A

the highest energy state on a reaction coordinate; it indicates a point at which new bonds are being formed at the same time as old bonds are being broken.

170
Q

what is the maxwell boltzmann distribution

A

a graph where the x axis is kinetic energy and the y axis fraction of particles

171
Q

the area beneath the curve in the MB curve is directly proportional to

A

directly proportional to the number of molecules having a value of kinetic energy in that range.

172
Q

as temp increases in the mb curve,

A

the particles gain ke and the curve flattens out

173
Q

how is the maxwell boltzmann distribution helpful

A

in understanding how changes in temperature and the use of a catalyst will affect the rate of a reaction.

174
Q

what factors affect rate of reaction

A

temperature
concentration
pressure
surface area

175
Q

when does particle size affect rate of reaction

A

when the solid is being reacted directly, not when it is dissolved in solution.

176
Q

The factor by which the concentration of a reactant affects the rate of a reaction is known as

A

its order of reaction

177
Q

what is the only way the order of a reaction with respect to a particular reactance can be determined

A

through experimental data

178
Q

what does the rate expression give

A

the relationship between reactant concentrations and the rate of reaction

179
Q

what is k

A

the rate constant

180
Q

what is the overall oder of reaction

A

the sum of indiviual orders of reactions

181
Q

zeroth order

A

Rate is independent of [A]; any change in the concentration of A does not affect the rate of reaction

182
Q

first order

A

Rate is directly proportional to [A]; if the concentration of A is doubled, the rate also doubles

183
Q

the rate expression gives the rel between the conc of the reactants and the

A

overall rate of reaction

184
Q

k is what dependent

A

temperature

185
Q

zero order units

A

mol dm-3 s-1

186
Q

first order units

A

s -1

187
Q

second order units

A

mol-1 dm3 s-1

188
Q

third order units

A

mol-2 dm6 s-1

189
Q

CT graphs

A

0, straight line
1, exp curve
2, exp curve

190
Q

RC graphs

A

0, horizontal straight line
1, straight line
2, parabolic curbe

191
Q

what is half life

A

the time it takes for the concentration of a reactant to decrease by half

192
Q

a reaction that takes place in a single step is known as a

A

elementary reaction

193
Q

what are the steps in a non elementary reaction called

A

elementary steps

194
Q

slowest elementary step is the

A

rds

195
Q

the elementary steps must add together to give

A

the overall balanced equation for the reaction

196
Q

a substance in multiple steps but not the overall equation is known as

A

a reaction intermediate

197
Q

what does molecularity tell us

A

the number of reactant particels in an elementary step

198
Q

If only one particle is involved in the elementary step, it is

A

unimolecular

199
Q

If two reactant particles are involved,

A

it is bimolecular

200
Q

if three reactant particles are involved,

A

termolecular

201
Q

the rds has the highest

A

activation energy as its the slowest

202
Q

why do catalysts appear in the elementary steps of a reaction mechanism and in the rate-determining step but not in the overall balanced equation for the reaction

A

they are chemically unchanged and can be reused

203
Q

what value constant does a catalyst change

A

increasing the value of the rate constant k

204
Q

rate constant k is temp dependent, this relatiponsh[ is shown by the arrhenius equation:

A

k=Ae ^−Ea/RT

205
Q

what is a in the arrhenius equation

A

the pre-exponential factor or the frequency factor.

206
Q

what is Ea in the arrhenius equation

A

activation energy (J/mol^-1)

207
Q

what is r in the arrhenius equation

A

universal gas constant (8.31 J K^1mol^-1)

208
Q

what is e in the arrhenius equation

A

eulers number

209
Q

what does the arrhenius constant take into account with a

A

the frequency of collisions and the probability that they have the correct orientation (or geometry)

210
Q

The expression (e−Ea /RT) (the exponential factor) is the fraction of molecules that have

A

sufficient kinetic energy to react at a certain temperature

211
Q

modified version of arrhenus equation (to find ativation energy)

A

lnk=−Ea/RT +lnA

212
Q

lnk =−Ea/RT +lnA is a form of y = mx+c

A

therefore, a graph of ln k against 1/T gives a straight line (Figure 1). The gradient of the line is equal to –Ea/R (where R is the universal gas constant with a value of 8.31 J K−1 mol−1). The intercept on the y-axis is equal to ln A. Once we have determined the activation energy for the reaction, we can calculate the Arrhenius constant by substituting the activation energy value into the Arrhenius equation.

213
Q

Gradient=

A

−Ea/R

214
Q

−Ea=

A

gradient×R

215
Q

the rate constant, k, for a reaction increases ???? with increasing temperature

A

exponentially

216
Q

4 key characteristics of a reaction at equilibrium

A
  1. The forward and reverse reactions occur at the same rate.
  2. The concentrations of reactants and products at equilibrium are constant.
  3. Equilibrium requires a closed system.
  4. There is no change in macroscopic properties at equilibrium.
217
Q

what is Kc

A

a constant for a given temperature, therefore, it is temperature-dependent. A change in temperature would result in a change in the value of Kc for a reaction (this is discussed in more detail in a later section).

218
Q

high Kc

A

reaction almost at completion

219
Q

low Kc (less than 10^-10)

A

reaction barely proceeded

220
Q

what value of Kc means reactants are predominate at eq

A

0.01

221
Q

what value of Kc means there are equal amounts of reactants and products

A

1

222
Q

what value of Kc means products are predominate at eq

A

100

223
Q

what is the value of Kc when the reactioon has almost gone to completion

A

over 10^10

224
Q

if a reaction is reversed, what happens to the kc

A

it will be the reciprocal of that value

225
Q

if the reaction coefficents are doubled, the Kc is

A

squared

226
Q

if the reaction coefficients are halved, the kc is

A

square rooted

227
Q

what is the reaction quotient Q

A

a measure of the relative amounts of reactants and products for a reaction that is not yet at equilibrium. The reaction quotient is useful for predicting in which direction a reaction will proceed to reach equilibrium.

228
Q

what is the different in reaction quotient expressiona dn equilibrum constant expression (q and kc)

A

kc uses equilibrium values

229
Q

Q vs Kc

A

Q less than Kc, reaction exceeded eq
Q = Kc, reaction at eq
Q bigger than Kc, reaction not at eq yet

230
Q

Le Châtelier’s principle states that:

A

‘When a system at equilibrium is subjected to a change, the system will respond to minimise the effect of the change.’

231
Q

The chemical process that produces ammonia gas involves the following reaction:

A

N2 (g) + 3H2 (g) ⇌ 2NH3 (g) ΔH = –92 kJ mol–1

232
Q

the haber process: This means that increasing the pressure will shift the equilibrium position to the right, yielding more what

A

ammonia

233
Q

catalyst of haber process

A

iron

234
Q

what does contact process create

A

sulfuric acid

235
Q

main uses of contact process

A

Fertilisers.
Paints and pigments.
Detergents and soaps.
Dyestuffs.

236
Q

pressure and temp in the contact process

A

low temp and high pressure according to le chatlier, but this is uneconomic

237
Q

Kc involving the forward and reverse reaction

A

Kc = Kf / Kr

238
Q

If the rate constant for the forward reaction (kf1) is greater than that of the reverse reaction (kr1), the equilibrium constant will be

A

large, and the forward reaction is favoured

239
Q

if the rate constant for the backward reaction (kr2) is greater than that of the forward reaction (kf2), the equalibrium constant will be

A

small and the reverse reaction is favoured

240
Q

You can see that the more negative the value of the standard Gibbs free energy change, the

A

larger the value of the equilibrium constant, K. This means that the more spontaneous a reaction is, the further the position of equilibrium lies to the right

241
Q

if gibbs is negative, ln k is positive and Kc is less than 1, waht is the pos of e?

A

Lies to the right – favouring the products

242
Q

the BL theory defines acids and bases as

A

proton donors and proton acceptors

243
Q

amphiprotic species

A

able to act as both a bronsted lowry acid and a bronsted lowry base depending on what its reacting with

244
Q

amphiprotic only applies to what theory

A

the BL theory

245
Q

amphoteric refers to

A

substances that can act as any acid or base (not restricted to Bl theory)

246
Q

metal + acid →

A

salt + hydrogen

247
Q

metal oxide + acid →

A

salt and water

248
Q

metal hydroxide + acid →

A

salt and water

249
Q

what is a neutralisation reaction

A

an acid reacts with a base or alkali to produce a salt and water

250
Q

ingredients in most antacids

A

The active ingredients of most antacid tablets are metal carbonates or hydrogen carbonates, NaHCO3, CaCO3, MgCO3, for instance, or insoluble metal hydroxides such as Mg(OH)2 or Al(OH)3. These react with excess stomach acid in neutralisation reactions to relieve the symptoms of heartburn.

251
Q

purpose of an acid base titration

A

to determine the unknown concentration of an acidic or basic solution using a solution of known concentration

252
Q

thermometric titration

A

when heat is released during a titration.

253
Q

pH formula

A

pH = −log[H+(aq)]

254
Q

change in one pH unit is equal to

A

ten times the change in hydrogen ion conc

255
Q

acid base indicator is

A

a weak acid or weak base in which the dissociated and undissociated forms have different colours

256
Q

litmus comes in two colours

A

red and blue

257
Q

in an acidic sol, blue litmus urns to

A

red

258
Q

in an alkaline sol red litmus turns to

A

blue

259
Q

universal indicator

A

red in acid
purple in alkali

260
Q

pH probe

A

more accurate method of measuring pH

261
Q

Water molecules do dissociate however, but only to a very small extent; this is known as the auto-ionisation of water. This reaction can be represented by the following equation:

A

H2O (l) ⇌ H+ (aq) + OH– (aq)

262
Q

The equilibrium constant expression (Kc) for the auto ionisatin of water is

A

Kc=[H+][OH−]/[H2O]

The position of equilibrium for the dissociation of H2O lies very far to the left, so the concentration of the water is effectively constant. Therefore, we can write a new expression which is known as the ionic product of water (Kw):

Kw=[H+][OH−]

263
Q

ionic product of water

A

Kw=[H+][OH−]

264
Q

Kw

A

1 x 10^-14

265
Q

The strength of an acid refers to its

A

degree of dissociation (or ionisation) in aqueous solution.

266
Q

three common strong bases

A

sodium hydroxide, potassium hydroxide and barium hydroxide

metaly hydroxides of group 1

267
Q

for the dissociation of weak bases the eq lies very far to the

A

left
Kc is very small

268
Q

remember to use an equilibrium sign (⇌) for

A

weak acids and bases.

269
Q

equimolar solutions

A

ones that have equal concentrations

270
Q

active metals are those

A

above hydrogen in the activity series

271
Q

which have higher elec conductivity

A

strong bases

272
Q

Natural, unpolluted rainwater is acidic, having a pH of approximately

A

5.6

273
Q

acid rain pH

A

less than 5

274
Q

how is acid rain formed

A

when acidic gases such as sulfur dioxide (SO2) and the oxides of nitrogen (NOx) dissolve in the water in the atmosphere to produce sulfuric acid or nitric acid.

275
Q

how can nitrogen monoxide be formed

A

during lightning storms by the reaction of nitrogen and oxygen, the two predominant atmospheric gases. The second is by the reactions that take place in internal combustion engines. Nitrogen reacts directly with oxygen at high temperatures to produce nitrogen monoxide

276
Q

benefit of a catalytic converter

A

reduces the level of these pollution emissions. This reaction also takes place in the jet engines of aircraft.

277
Q

how is sulfur dioxide formed

A

occurs naturally in the atmosphere as it is released during volcanic eruptions

278
Q

sulfur dioxide photochemical oxidation in the atmosphere

A

2SO2 (g) + O2 (g) → 2SO3 (g)

279
Q

sulfur dioxide photochemical oxidation in the atmosphere

A

2SO2 (g) + O2 (g) → 2SO3 (g)

280
Q

what does sulfur dioxide form with water

A

Sulfur dioxide and sulfur trioxide react with water in the atmosphere to form sulfurous and sulfuric acid:

SO2 (g) + H2O (l) → H2SO3 (aq)

SO3 (g) + H2O (l) → H2SO4 (aq)

281
Q

what is dry deposition

A

acidic particles and gases fall to the ground via dust and smoke in the absence of precipitation. This form of deposition can be washed into streams, lakes and rivers, causing harm to biological systems.

282
Q

which are the stronger acids

A

These four acids (nitric, nitrous, sulfuric and sulfurous acids) are all much stronger than carbonic acid, H2CO3.

283
Q

environmental impact of acid deposition

A

prevent fish eggs from hatching
damage plants
harms biodiversity

284
Q

how is acid rain detrimental on statues

A

calcium carbonate can react with sulfuric acid and nitric acid

CaCO3 (s) + H2SO4 (aq) → CaSO4 (aq) + H2O (l) + CO2 (g)

CaCO3 (s) + 2HNO3 (aq) → Ca(NO3)2 (aq) + CO2 (g) + H2O (l)

285
Q

acid rain on humans

A

can adversely affect the mucous membranes and lungs, causing irritation and possibly exacerbating the symptoms for people with asthma and other respiratory conditions.

286
Q

how to reduce acid deopisition

A

pre combustion methods: removing the sulfur before the coal is combusted
post comubstin methods: removing the sulfur oxides from exhaust gases once they have been formed by reacting with a base.

287
Q

hydrodesulfurisation

A

chemistry (13 decks)

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