physical chem Flashcards
1
Q
why do elements in groups have similar chemical properties
A
similar outer shell electron configuration
2
Q
periodicity definition
A
repeating pattern across different periods
3
Q
order of blocks in the periodic table
A
s,d,p,f
4
Q
what happens to atomic radius across a period
A
- atomic radius decreases
- as there is an increased number of protons, which creates more nuclear attraction
- similar shielding
5
Q
first ionisation energy definition
A
the energy needed to remove an electron from each atom in one mole of gaseous atoms
6
Q
example equation for first ionisation energy
A
H(g) → H+(g) + e-
7
Q
which factors affect ionisation energy
A
- nuclear attraction
- atomic radius
- shielding
8
Q
why are successive ionisation energies always larger
A
- the ion formed is smaller than the atom (smaller atomic radius)
- the proton to electron ratio in the 2+ ion is greater than in the 1+ ion
- nuclear attraction is therefore stronger
9
Q
if there is a big jump in ionisation energy between IE 2 and 3, what does this mean
A
- the element must be in group 2
- as the 3rd electron is being removed from an electron shell closer to the nucleus and therefore experiences more nuclear attraction
- there must also be less shielding
10
Q
why do 1st ionisation energies decrease down a group
A
the outer electrons are found in shells further from the nucleus and are more shielded so the attraction of the nucleus becomes smaller
11
Q
Why is there a general increase in first ionisation energy across a period?
A
the electrons are being added to the same
shell which has the same distance from the nucleus and same shielding
effect. The number of protons increases, however, making the effective attraction of the nucleus greater.
12
Q
why is there a small drop in IE from group 2 to 3
A
- group 2 has its outer electrons in the 3s sub shell, whereas group 3 is starting to fill a 3p sub shell
- The electrons in the 3p subshell are slightly easier to remove because
the 3p electrons are higher in energy and are also slightly shielded by the 3s electrons
13
Q
why is there a small drop in IE from group 5 to 6
A
in group 6 r there are 4 electrons in the 3p sub shell and the 4th is starting to doubly
fill the first 3p orbital.
When the second electron is added to a 3p orbital there is a slight repulsion between
the two negatively charged electrons which makes the second electron easier to
remove.
14
Q
metallic bonding definition
A
the electrostatic force of attraction between
the positive metal ions and the delocalised electrons
15
Q
what are the three main factors that affect the strength of metallic bonding
A
- Number of protons/ Strength of nuclear attraction.
(The more protons the stronger the bond) - Number of delocalised electrons per atom (the outer shell electrons are delocalised)
The more delocalised electrons the stronger the bond - Size of ion.
The smaller the ion, the stronger the bond
16
Q
why has Mg got stronger metallic bonding than Na
A
Mg has stronger metallic bonding than Na and hence a higher melting point.
The Metallic bonding gets stronger because in Mg
there are more electrons in the outer shell that are released to the sea of electrons.
The Mg ion is also smaller and has one more proton. There is therefore a stronger electrostatic attraction
between the positive metal ions and the delocalised electrons and higher energy is needed to break bonds
17
Q
explain the structure of diamond
A
MACROMOLECULAR
- Tetrahedral arrangement of
carbon atoms.
- 4 covalent bonds per atom
- high melting points because of strong covalent forces in the giant structure. It takes a lot of energy to break the many strong covalent bonds
18
Q
explain grpahite sutreucue
A
MACROMOLECULAR
- Planar arrangement of carbon atoms in layers.
- 3 covalent bonds
per atom in each layer. 4th outer electron per atom is delocalised.
- Delocalised electrons between layers.
- high melting points because of strong covalent forces in the giant structure. It takes a lot of energy to break the many strong covalent bonds
19
Q
which structures have covalent bonding
A
MACROMOLECULAR
20
Q
which structures have metallic bonding
A
GIANT METALLIC LATTICES
21
Q
properties of macromolecular substances
A
bp and mp: high due to many covalent bonds which take a lot of energy to break
solubility: insoluble
conductivity when solid: diamond and sand are poor because electrons cant move, whereas graphite is good as free electrons between layers
conductivity when molten: poor
22
Q
properties of giant metallic substances
A
mp and bp points: high due to strong electrostatic forces between positive ions and sea of deloc electrons
solubility: insoluble
conductivity: good - deloc electrons can move through structure
23
Q
explain the general trend of mp and bp across period one
A
For Na, Mg, Al- Metallic bonding : strong bonding – gets stronger the more electrons there are in the outer shell that
are released to the sea of electrons. A smaller positive
centre also makes the bonding stronger. High energy is needed to break bonds.
Si is Macromolecular: many strong covalent bonds
between atoms high energy needed to break covalent
bonds– very high mp +bp
Cl2 (g), S8 (s), P4 (S)- simple Molecular : weak London forces between molecules, so little energy is needed to break them – low mp+ bp
S8 has a higher mp than P4 because it has more electrons
(S8 =128)(P4=60) so has stronger London forces
Ar is monoatomic weak London forces between atoms
24
Q
what happens to atomic radius down group 2
A
- increases
- atoms have more shells of electrons
25
melting point down group 2
- decreases
- metallic bonding weakens as the atomic size increases
- distance between the + ions and deloc electrons increases
- therefore, electrostatic
attractive forces between the positive ions and the
delocalized electrons weaken
26
what happens to first and second IE's down group 2
- decrease
- The outermost electrons are held more weakly because they are successively further from the nucleus in additional
shells
- In addition, the outer shell electrons become more shielded from the attraction of the nucleus by the repulsive force of inner shell electron
27
reactivity down group 2
- increases
- As the atomic radii increase there is more shielding.
- The nuclear attraction decreases and it is easier to remove outer electrons.
- Cations form more
easily
28
group 2 metals reactions with oxygen
The group 2 metals will burn in oxygen.
Mg burns with a bright white flame
2Mg + O2 → 2MgO
- Mg will also react slowly with oxygen without a flame.
- Mg ribbon will often have a thin layer of magnesium oxide on it formed by reaction with oxygen
- this is cleaned off by emery paper before doing reactions with Mg ribbon as Mg and MgO react at different rates
29
Mg + steam
produces magnesium oxide and hydrogen
the Mg would burn with a bright white flame
30
Mg + warm water
produces magnesium hydroxide and hydrogen
this is a slower reaction than with steam and produces no flame
31
group 2 metals and cold water
- react with increasing vigour down the group to form hydroxides
eg : Ca + 2 H2O (l) → Ca(OH)2
(aq) + H2(g)
32
what are the observations for group 2 metals reacting with cold water
what are the observations for group 2 metals reacting with cold water
33
group 2 metals reacting with acid
- react with acids with increasing vigour down the group to form a salt and hydrogen
34
why does barium react slowly with H2SO4
the insoluble barium
sulfate produced will cover the surface of the metal and act as a barrier to further attack.
35
group 2 oxides + water
- forms hydroxides
36
what is magnesium hydroxide used for
s used in medicine (in suspension as milk of magnesia) to neutralise excess acid in the stomach and to treat constipation.
Mg(OH)2 + 2HCl → MgCl2 + 2H2O
It is safe to use as it is weakly alkaline.
37
what is calcium hydroxide used for
- It is used in agriculture to neutralise acidic soils.
- If too much calcium hydroxide is added to the
soil, excess will result in soils becoming too alkaline to sustain crop growth
38
appearance of the halogens
Fluorine (F2): very pale yellow gas. It is highly reactive
Chlorine : (Cl2) greenish, reactive gas, poisonous in high concentrations
Bromine (Br2) : red liquid, that gives off dense brown/orange poisonous fumes
Iodine (I2) : shiny grey solid sublimes to purple gas.
39
mp and bp down the halogens
increases down the group
As the molecules become larger they have more electrons and so have larger induced dipole-dipole forces
(London forces) between the molecules. As the intermolecular forces get larger more energy has to be put into break these intermolecular forces. This increases the melting and boiling points
40
reactivity of halogens down the group
- decreases
- as the atoms get bigger with more shielding so they less
easily attract and accept electrons. They therefore form -1 ions less easily down the group
41
reactivity of halogens explanations
- Chlorine is more reactive than bromine because it will
gain an electron and form a negative ion more easily than
bromine.
- This is because an atom of chlorine is smaller than bromine and the outermost shell of chlorine is less shielded than bromine so the electron to be gained is
attracted more strongly to the nucleus in chlorine than
bromine.
42
halogen displacement equations
Cl2(aq) + 2Br – (aq) → 2Cl - (aq) + Br2(aq)
Cl2(aq) + 2I – (aq) → 2Cl – (aq) + I2(aq)
Br2(aq) + 2I – (aq) → 2Br – (aq) + I2(aq)
43
disproportionation definition
the name for a reaction where
an element simultaneously oxidises and reduces.
44
chlorine + water equation
how is it disproportionation?
what will happen is UI is added?
Cl2(g) + H2O(l) →HClO(aq) + HCl (aq)
Chlorine is both simultaneously reducing and oxidising. It changes from 0 in Cl2 to -1 in HCl and +1 in HClO
If some universal indicator is added to the solution it will first turn red due to the acidity of both reaction products. It will then turn
colourless as the HClO bleaches the colour.
45
what are the benefits and negatives of chlorine being used in water treatment?
- kills bacteria
- toxic and irritates the respiratory system
46
chlorine + cold dilute NaOH
Cl2(aq) + 2NaOH(aq) → NaCl (aq) + NaClO (aq) + H2O(l)
- the colour of the halogen solution will fade to colourless
- The mixture of NaCl and NaClO (sodium chlorate (I)) is used as Bleach and to disinfect/ kill bacteria
47
chlorine + hot NaOH
3Cl2 + 6NaOH → NaClO3 + 5NaCl + 3H2O
- sodium chlorate (V) is formed
48
reaction of halide ions with silver nitrate
- used for?
- why is nitric acid added?
- colours?
- how can ammonia be used to help differentiate
used as a test to identify which halide ion is present. The test solution is made acidic with nitric acid, and then Silver nitrate solution is added drop wise
The role of nitric acid is to react with any carbonates
present to prevent formation of the precipitate
Ag2CO3 - This would mask the desired observations
- Fluorides produce no precipitate
- Chlorides produce a white precipitate
Ag+(aq) + Cl- (aq) →AgCl(s)
- Bromides produce a cream precipitate
Ag+ (aq) + Br- (aq) → AgBr(s)
- Iodides produce a pale yellow precipitate
Ag+ (aq) + I- (aq) →AgI(s)
The silver halide precipitates can be treated with ammonia
solution to help differentiate between them if the colours look similar:
Silver chloride dissolves in dilute ammonia to form a
complex ion
AgCl(s) + 2NH3(aq) → [Ag(NH3)2]+(aq) + Cl- (aq)
Silver bromide dissolves in concentrated ammonia to form a complex ion
AgBr(s) + 2NH3(aq) →[Ag(NH3)2]+(aq) + Br - (aq)
Silver iodide does not react with ammonia – it is too insoluble
49
is △H positive or negative in exothermic and endothermic reactions?
In an exothermic reaction the
∆H is negative
In an endothermic reaction
the ∆H is positive
50
activation energy definition
the minimum energy which particles need to collide to start a reaction
51
standard conditions
* 100 kPa pressure
* 298 K (room temperature or 25oC)
* Solutions at 1mol dm-3
* all substances should have their normal state at 298K
52
enthalpy change of reaction definition
the enthalpy change when the number of moles of reactants as specified in the
balanced equation react together
53
standard enthalpy change of formation definition
the enthalpy change when 1 mole of the compound is formed from its elements under standard conditions (298K and 100kpa), all
reactants and products being in their standard states
54
standard enthalpy change of combustion definition
the enthalpy change that occurs when one mole of a substance is combusted completely in oxygen under standard conditions (298K and 100kPa), all reactants and products being in their
standard states
55
enthalpy change of neutralisation definition
the enthalpy change when solutions of an acid and an alkali react together under standard conditions to produce 1 mole of water.
56
errors in calorimetry experiments
Energy losses from calorimeter
* Incomplete combustion of fuel
* Incomplete transfer of energy
* Evaporation of fuel after weighing
* Heat capacity of calorimeter not included
* Measurements not carried out under standard conditions as H2O is gas, not liquid, in this experiment
57
mean bond enthalpy
the enthalpy change when
one mole of bonds of (gaseous covalent) bonds is broken (averaged over different molecules)
58
hess' law definition
total enthalpy change for a reaction is independent of the route by which the chemical changes takes place
59
effect of increasing conc and pressure on rate of reaction
At higher concentrations(and pressures) there are more
particles per unit volume and so the particles collide
with a greater frequency and there will be a higher
frequency of effective collisions.
60
rate of reaction definition
the change in concentration of a substance in unit time
Its usual unit is mol dm-3s-1
61
catalysts definition
Catalysts increase reaction rates without getting used up.
Explanation: They do this by providing an alternative route or mechanism with a lower activation energy so
more molecules have energy above activation energy.
62
heterogeneous catalysts vs homogenous catalysts
A heterogeneous catalyst is in a different phase from the reactants
- Heterogeneous catalysts are usually solids whereas the reactants are gaseous or in solution.
- The reaction occurs at the surface of the catalyst.
A homogeneous catalyst is in the same phase as the reactants
- When catalysts and reactants are in the same phase, the reaction proceeds through an intermediate species.
63
benefits of catalysts
- lower temps and pressures can be used
- This can save energy costs as there is reduced energy demand for providing high temperature and less electrical pumping costs for producing pressure.
- This can mean fewer CO2 emissions from burning
of fossil fuels.
- Catalysts can enable different reactions to be
used, with better atom economy and with
reduced waste, or fewer undesired products or
less use of hazardous solvents and reactants.
- Catalysts are often enzymes, generating very specific products, and operating effectively close to room temperatures
and pressures.
64
effect of increasing temp of rate of reaction
At higher temperatures the energy of the particles
increases. They collide more frequently and more
often with energy greater than the activation energy.
More collisions result in a reaction
65
effect of increasing surface area on rate of reaction
causes collisions to occur more frequently between the reactant particles and this increases the rate of the reaction.
66
effect of catalysts on rate of reaction
If the activation energy is lower, more particles will have energy > EA, so there will be a higher frequency of effective collisions. The reaction will be faster
67
what are the conditions of dynamic equilibrium
- forward and backward reactions are occurring at equal rates.
- The concentrations of reactants and products stays
constant and the reaction is continuous.
68
effect of temp on equilibrium
If temperature is increased the equilibrium will
shift to oppose this and move in the endothermic direction to try to reduce the
temperature by absorbing heat.
69
what affect does temp have on yield
Low temperatures may give a higher yield of product but will also result in slow rates of
reaction. Often a compromise temperature is used that gives a reasonable yield and rate
70
effect of pressure on equilibrium
Increasing pressure will cause the equilibrium to shift
towards the side with fewer moles of gas to oppose
the change and thereby reduce the pressure.
If the number of moles of gas is the same on both
sides of the equation then changing pressure will have
no effect on the position of equilibrium
71
what affect does pressure have on yield
Increasing pressure may give a higher yield of product and will produce a faster rate. Industrially high
pressures are expensive to produce (high electrical energy costs for pumping the gases to make a high
pressure) and the equipment is expensive (to contain the high pressures)
72
effect of conc on equilibrium
- equilibrium shifts to opposite direction to oppose this
73
effect of catalysts on equilibrium
A catalyst has no effect on the position of equilibrium, but it will speed up the rate at which the equilibrium is achieved.
It does not effect the position of equilibrium because it speeds up the rates of the forward and backward reactions by the same amount.
74
haber process
- equation
- effect of temp and pressure
N2 + 3H2 →2NH3 △H = -ve exo
Low temp gives good yield but slow rate: compromise temp used
High pressure gives good yield and high rate:
too high a pressure would lead to too high energy costs for pumps to produce the pressure
75
contact process
- equations
- temp and pressure effect
Stage 1 S (s) + O2 (g) → SO2(g)
Stage 2 SO2 (g) + ½ O2 (g) → SO3 (g)
△H = -98 kJ mol-1
Low temp gives good yield but slow rate: compromise
moderate temp used
High pressure gives slightly better yield and high rate: too
high a pressure would lead to too high energy costs for
pumps to produce the pressure
76
production of methanol from CO
- equation
- effects of temp and pressure
CO (g) + 2H2(g) →CH3OH (g) △H = -ve exo
Low temp gives good yield but slow rate:
compromise temp used
High pressure gives good yield and high rate: too
high a pressure would lead to too high energy
costs for pumps to produce the pressure
77
hydration of ethene to produce ethanol
- equation
- effects of temp and pressure
CH2=CH2 (g) + H2O (g) →CH3CH2OH(l) △H = -ve
Low temp gives good yield but slow rate: compromise
temp used.
High pressure gives good yield and high rate: too high a
pressure would lead to too high energy costs for pumps
to produce the pressure.
High pressure also leads to unwanted polymerisation of
ethene to poly(ethene).
78
what affects Kc or Kp
ONLY TEMP
79
what does a large Kc indicate
- a large amount of products
- equilibrium has favoured the products
80
how to work out total order for a reaction
add all individual orders together
81
effect of temp on rate constant
Increasing the temperature increases the value of the
rate constant
82
acid definition
substance that can donate a proton
83
base definition
a substance that can accept a proton
84
acid + metal
salt + hydrogen
85
acid + alkali
salt + water
86
acid + carbonate
salt + water + co2
87
strong acids definition
completely dissociate
88
In all aqueous solutions and pure water the following equilibrium occurs....
H2O (l) →H+(aq) + OH-(aq)
89
weak acids definition
slightly dissociate when dissolved in water, giving an equilibrium mixture
90
buffer solution definition
one where the pH does not change significantly if small amounts of acid or alkali are added to it.
91
how is an acidic buffer solution made
from a weak acid and a salt
of that weak acid ( made from reacting the weak acid with a
strong base)
92
how is a basic buffer solution made
A basic buffer solution is made from a weak
base and a salt of that weak base ( made from reacting the weak base with a strong
acid)
93
how can the salt content be added to a buffer
a salt solution
could be added to the acid or some solid salt added. A buffer can also be made by partially neutralising a weak acid with alkali and
therefore producing a mixture of salt and acid
94
how do buffer solutions work
If small amounts of acid is added to the buffer then the equilibrium will shift in the direction to oppose this removing the H+ ions added
If small amounts of alkali is added to the buffer. The OHions will react with H+
ions to form water. The Equilibrium will then shift to the right to produce more H+ ions. Overall the concentration of H+ ions and pH remains constant
95
buffering action in blood
A carbonic acid– hydrogencarbonate equilibrium acts as a buffer in the control of blood pH
The H2CO3
/HCO3 – buffer is present in
blood plasma, maintaining a pH between
7.35 and 7.45.
H2CO3 (aq) ⇌ H+(aq) + HCO3–(aq)
Adding alkali reacts with H+ with the equation
H+ + OH- → H2O
so the above equilibrium would shift right forming new H+ and more HCO3-
Adding acid shifts the above equilibrium left.
The reaction is
H+ + HCO3- → H2CO3
96
enthalpy change of formation definition
the energy transferred when
1 mole of the compound is formed from its elements under standard conditions (298K and 100kpa), all reactants and products being in their standard states
97
enthalpy of atomisation definition
the enthalpy
change when 1 mole of gaseous atoms is formed
from the element in its standard state
98
1st ionisation enthalpy definition
2nd ionisation enthalpy definition
The first ionisation enthalpy is the enthalpy change
required to remove 1 mole of electrons from 1 mole
of gaseous atoms to form 1 mole of gaseous ions
with a +1 charge
The second ionisation enthalpy is the enthalpy
change to remove 1 mole of electrons from one mole
of gaseous 1+ ions to produces one mole of gaseous 2+ ions.
99
first electron affinity definition
second electron affinity definition
The first electron affinity is the enthalpy change that
occurs when 1 mole of gaseous atoms gain 1 mole of
electrons to form 1 mole of gaseous ions with a –1
charge
- The first electron affinity is exothermic for atoms that
normally form negative ions because the ion is more
stable than the atom and there is an attraction
between the nucleus and the electron
The second electron affinity is the enthalpy change
when one mole of gaseous 1- ions gains one
electron per ion to produce gaseous 2- ions
- The second electron affinity for oxygen is
endothermic because it take energy to overcome
the repulsive force between the negative ion and
the electron
100
lattice enthalpy definition
the standard enthalpy change
when 1 mole of an ionic crystal lattice is formed from
its constituent ions in gaseous form.
101
enthalpy of hydration definition
Enthalpy change when one mole of gaseous ions
become aqueous ions .
This always gives out energy (exothermic, -ve)
because bonds are made between the ions and the
water molecules.
102
enthalpy of solution definition
the standard enthalpy
change when one mole of an ionic solid dissolves in
a large enough amount of water to ensure that the
dissolved ions are well separated and do not
interact with one another.
103
what does strength of lattice enthalpy formation depend on
1. The sizes of the ions:
The larger the ions, the less negative the enthalpies of lattice formation (i.e. a weaker lattice). As the ions are larger the charges become further apart and so have a weaker attractive force between them.
2. The charges on the ion:
The bigger the charge of the ion, the greater the attraction
between the ions so the stronger the lattice enthalpy (more negative values).
104
entropy definition
measure of disorder within a system
105
a increase in entropy will occur if...
-there is a change of state from solid or liquid to gas
- there is a significant increase in number of molecules
between products and reactants.
106
advantages of fuel cells
- less pollution and less CO2
- greater efficiency
107
state the charge carriers that transfer current
- through the wire
- through the solution
- electrons
- ions
108
hydrogen electrode equilibrium
H2 (g) → 2H+ (aq) + 2e- (reversible)
109
why is Zn not a transition metal
Zn can only form a +2 ion. In this ion the Zn2+ has a complete d orbital and so does not meet the criteria of having an incomplete d orbital in one of its ions.
110
why is Sc not a transition metal
Sc can only form a +3 ion. In this ion the Sc3+ has an empty d orbital and so does not meet the criteria of having an incomplete d orbital in one of its ions.
111
transition metals
d-block elements that form at least one ion with a partially filled d-orbital
112
typical properties of transition metals
-The existence of more than one oxidation state for each element in its compounds
- the formation of coloured ions
- the catalytic behaviour of the elements and their compounds and their importance in the manufacture of chemicals by industry
113
complex definition
a central metal ion surrounded by ligands
114
ligand definition
an atom, ion or molecule which can donate a lone electron pair
115
co-ordinate bonding definition
when the shared pair of electrons in the covalent bond come from only one of the bonding atoms.
116
co-ordination number
the number of co-ordinate bonds formed to a central metal ion.
117
unidente definition
forms one coordinate bond per ligand
118
how to make a standard solution as accurate as possible
Use of a weighing boat to measure the solid into; [1 mark]
Method of rinsing the solid from the weighing boat using distilled water, to ensure no solid residue is left behind; [1 mark]
Transferring the solid (from the weighing boat) to a beaker, rather than straight into the volumetric flask; [1 mark]
Adding a minimum amount of distilled water (to the beaker) to dissolve the product before transferring to volumetric flask; [1 mark]
Transfer (from beaker) to volumetric flask using a funnel; [1 mark]
Rinse all pieces of equipment - weighing boat, beaker, funnel - with distilled water when transferring to volumetric flask, to ensure that all solid has been transferred; [1 mark]
Use a (teat) pipette to make up to the mark on volumetric flask; [1 mark]
Ensure that the bottom of the meniscus sits exactly on the graduation mark of the volumetric flask; [1 mark]
Inverting volumetric flask multiple times would be better than shaking
119
what can barium chloride be used for
to test for sulfate ions
120
write an equation for the dissociation of lactic acid (CH3CH(OH)COOH)
CH3CH(OH)COOH (aq) ⇌ CH3CH(OH)COO- (aq) + H+ (aq)
121
why would a substance not appear in the ionic equation
it is a spectator ion
122
what affect does removing the funnel during a titration have
Make the reading appear less; [1 mark]
(Because) liquid trapped between the funnel and the side of the burette will flow down into the burette;
123
explain the bonding present in silicon
- giant macromolecular structure
- strong covalent bonds between the Si atoms
124
when does shielding increase
down the group
125
explain why sulfur has a larger melting point than phosphorus
Both sulfur and phosphorus contain weak induced dipole-dipole forces
The strength of the induced dipole-dipole forces increases with the number of electrons
Sulfur has more electrons than phosphorus
126
trend in ph as the group 2 elements react with water
ph increases due to increased dissociation therefore more OH- ions in the solution
127
state how a group 2 element forms an ionic bond with non-metals
loses two outer electrons to non metals
128
describe trend in solubility of group 2 metal sulfates down the group
decreases
129
describe trend in solubility of group 2 metal hydroxides down the group
increases
130
give a medical use of barium sulfate and state why it is safe for use
- taken before an x ray
- safe as it is insoluble
131
ionic equation for neutralisation
H+(aq) + OH-(aq) = H20(l)
132
electronegativity down the halogens
- decreases down the group
- the atoms get larger and the distance between the nucleus and bonding electrons increase
- more shielding
133
dissociation equation of chloric acid in water
HClO(aq) + H2O(l) → ClO-(aq) + H3O+ (aq) REVERSIBLE
134
why is adding chlorine to drinking water safe even though cl2 is a toxic substance
- only added in small volumes
- thus not harmful to humans at these levels
135
reducing agent
species that can donate electrons
136
draw boltzmann distrbution
- with higher temp
- higher `Ea
137
what is k
the rate constant
138
draw conc rate graphs for zero, first and second order
139
draw conc/time graphs for zero first and second order
140
how can you tell order of a reaction from half lives
constant half life = 1st order
rapidly increasing half life = 2nd order
141
as Ea gets smaller rate constant gets....
bigger
- more particles have enough energy to react when they collide
142
rearrange arrhenius equation to get activation energy
Ea = (lnA-lnK) x RT
143
effect of temp on position of equilibrium and Kc
if temp is increased reaction shifts to oppose change in the backwards direction.
144
kw expression
[H+(aq)] [OH-(aq)]
145
ka expression
Ka = [H+] [A-] / [HA]
146
assumptions made when simplifying Ka expression
1. [H+] = [A-]
2. as amount of dissociation is small we assume that the initial concentration of the undissociated acid has remained constant
147
methyl orange ph colour change point
3-4.5
- use for strong acid!
148
phenolphthalein ph colour change point
8.2-10
- use for weak acid
149
weak acid weak base indicator
CANT USE ONE as no vertical part of the curve
ph meter used instead
150
what acid should be used for a manganate titration
if the manganate is In the burette what will be the colour change
- dilute sulfuric acid
- colourless to purple
151
why is starch added in the iodine thiosulfate reaction
what si the colour change
- to emphasise the colour change
- the colour of iodine would fae gradually during the reaction and the colour fading would not be clear
- blue black to white
152
what colour does a mixture containing bromine ions turn when cyclohexane is added
brown
153
why is bromine a liquid at room temp but fluorine and chlorine are gases
bromine atoms are bigger/ have a bigger atomic radius
induced dipole dipole forces are stronger
so more energy required to pull the molecules apart
154
draw a reaction profile for an endothermic and an exothermic reaction
155
give the formula for calculating standard enthalpy change of reaction using bond energies
Σ bond energies broken - Σ bond energies made
156
why is the enthalpy of formation of hydrogen zero
it is an element and so will not have an enthalpy change of formation value
157
how to improve calorimetry experiments
- use a lid
- insulation of the sides of the calorimeter
158
what two things do particles need for a successful collision
- sufficient energy
- correct orientation
159
comment how the area under a boltzmann curve changes with temperature
does not change at all as the total number of particles remains constant
160
why do boltzmann graphs always start at the origin
there are no molecules with no energy
161
closed system definition
a system where none of the reactants or products can escape
162
half life definition
time taken for half the reactant to be used up
163
if initial concentration of a substance is halved what is the effect on half life
no effect as half life is independant of initial concentration
164
in a clock reaction the time taken for a visual change to occur is timed and the initial rate is taken to be proportional to 1/t
what assumption is made so that initial rate can be taken to be proportional to 1/t?
There is no significant change in the rate of reaction during the time for the visual change to occur
165
mole fraction definition
The fraction of the total number of moles that each chemical in a reaction is responsible for
166
why is enthalpy of atomisation always positive
- energy is taken in to break the bonds
- so the reaction is endothermic
167
as temperature increases, what happens to the ph value of pure water
decreases
168
explain why two acids which both have the same concentration can have different ph values
- one acid could be a strong acid, the other could be weak
- ph depends on the dissociation of H+ ions: weak acids only partially dissociate, strong acids fully dissociate
169
assumptions made when using Ka expression for buffer solutions
[HA] = the concentration of the weak acid used
[A-] = the concentration of ([A-] in) the salt of the acid being used
[H+] is negligible/ too small to account for/ minimal dissociation of the acid occurs
170
predict the effect on the pH of the buffer when the concentration of the salt is increased
ph will increase
as H+ will decrease
171
give an example of a diprotic acid and explain the term dibasic
h2so4
releases two protons
172
units of Kw
mol2dm-6
173
what happens to Kw as temp increases
- Kw increases as bonds are being broken during dissociation
- bond breaking is an endothermic process
- therefore, equilibrium moves to the right hand side to oppose the change in temperature
174
why does the kw expression not include the concentration of water
- water only dissociates a very small amount
- concentration of water is effectively constant
175
considering the expression for Ka, explain why the value of Ka is larger in hydrofluoric acid than in ethanoic acid
- HF acid is a stronger acid so dissociated more
- ka = [H+] [A-]/[HA]
- numerator is larger in a stronger acid
176
explain how ethanoic acid acts as a buffer when small amounts of alkali are added
- OH- ions are added and equilibrium moves to the right
- conc of H+ ions remains constant
177
what adjustment needs to be made in the potassium manganate titration
- burette readings taken from top not bottom of meniscus
- as the bottom of meniscus is not visible
178
difference between primary and secondary cells
primary cells and non rechargeable
secondary cells are rechargeable
179
two difficulties of disposal of lead acid batteries
- heavy
- lead is toxic
180
why is platinum used as an electrode
it is inert AND conducts electricity
181
how does a salt bridge provide an electrical connection between two solutions
- has mobile ions that can carry charge
182
how is an electrical current produces in a fuel cell
why do fuel cells not need to be recharged
Hydrogen (electrode) produces electrons
Oxygen (electrode) accepts electrons;
there is a constant supply of hydrogen
183
disadvantages of fuel cells
- limited lifetime
- use of toxic chemicals in production
- not carbon neutral as co2 is released when manufactured
184
write ionic equations for
- reaction between manganate and iron
- reaction between iodine and thiosulfate
185
Use ideas about the enthalpy changes that take place during bond breaking and bond
making to explain why some reactions are exothermic.
- bond breaking absorbs energy
- bond making releases energy
- more energy released than absorbed
186
state le chateliers principle
the position of equilibrium will shift so as to minimise the effect of any change in conditions
187
It is very difficult to determine the standard enthalpy change of formation of
hexane directly. Suggest a reason why
many different hydrocarbons would form
188
State one important manufacturing process in which hydrogen is used.
production of margarine
ammonia
haber process
189
The standard enthalpy change of combustion of butane is –2877 kJ mol–1. What
does standard mean in this context?
reaction carried out at 298K and 100Kpa
190
Predict and explain whether a barium ion is larger, smaller or the same size as a
barium atom.
smaller as there is less shielding/a shell has been lost
191
Suggest two ways in which the use of catalysts helps chemical companies to make
their processes more sustainable.
- catalyst not used up in reaction
- reactions take place at lower temp with lower Ea so less co2 emitted into atmosphere
192
Suggest why the entropy of water is zero at 0 K.
particles are in perfect order
193
The student had added the exact amount of calcium required to react with the
hydrochloric acid used. After carrying out the experiment, the student accidentally
added some more calcium. The student was surprised that the extra calcium still
reacted.
Explain this observation
Calcium reacts with water producing
hydrogen and calcium hydroxide
194
Describe and explain one anomalous property of water which results from
hydrogen bonding.
Liquid H2O is denser than solid
In solid state H2O molecules are held apart by hydrogen bonds
195
Use your understanding of electronic structure to explain why chlorine is a
stronger oxidising agent than iodine.
Cl atom is smaller/has less shells
electron to be captured will be attracted more
