Periodic Table and Energy Flashcards
1
Q
what is activation energy
A
minimum amount of energy required for a reaction to take place
2
Q
what is average bond enthalpy
A
the energy required to break one mole of gaseous bonds. Actual bond enthalpies may differ from the average as the average bond enthalpy considers a particular bond in a range of molecules.
3
Q
what does endothermic mean
A
a reaction that takes in energy
ΔH is ALWAYS positive
more energy is required to break the bonds than is released forming them
4
Q
What is Enthalpy
A
value that represents heat content in a system
5
Q
what is enthalpy change of combustion
A
when 1 mol of a substances completely combusts with oxygen under standard conditions
6
Q
what is enthalpy change of formation
A
enthalpy change when 1 mol of a substance is formed from its elements under standard conditions
7
Q
what is enthalpy change of neutralisation
A
The enthalpy change that takes place when
one mole of water is formed from a neutralisation reaction
8
Q
what is enthalpy change (ΔH)
A
the change in the heat content of a system during a reaction. This can be determined from experimental results
9
Q
What is the equation for enthalpy change and what does each symbol stand for
A
q = MCΔT
q = heat change of surroundings
m = mass
c = specific heat capacity
ΔT = change in temperature
10
Q
what does exothermic mean
A
a reaction that gives out energy
ΔH is ALWAYS NEGATIVE
more energy is released by bonds forming than is when bonds are broke
11
Q
What is enthalpy change of reaction
A
the enthalpy change that is associated with a
particular chemical equation.
12
Q
What is Hess’s Law
A
the enthalpy change of a reaction is independent of the route it takes
13
Q
what are standard conditions?
A
1atm and 298K
14
Q
what is standard state
A
physical state of a substance under standard conditions
15
Q
how do you calculate percentage error
A
% error = resolution of equipment/measured value x 100
16
Q
when given ΔHf, what way do the arrows face on your enthalpy profile diagram
A
upwards
17
Q
when given ΔHc what way do the arrows face on your enthalpy profile diagram
A
downwards
18
Q
how do do you find ΔHr for ΔHf
A
-ΣΔHf (reactants) + ΣΔHf(products)
19
Q
how do do you find ΔHr for ΔHc
A
ΣΔHc (reactants) - ΣΔHc (products)
flip the sign on the products as you go against the arrow
20
Q
what is a system in a chemical reaction
A
the atoms and bonds involved in the reaction
21
Q
which way does the arrow for activation energy point on a reaction profile
A
always points up
22
Q
what are the advantages of a bomb calorimeter
A
minimises heat loss
pure oxygen is used which ensures complete combustion
23
Q
what are the reasons that experimental methods for enthalpy determination might not be accurate
A
heat loss to the surroundings
not in standard conditions
may not have a complete reaction
24
Q
why is using bond enthalpies not be as accurate as using standard enthalpy of combustion/formation
A
bond enthalpies are a mean average for the same bond across different molecules whereas standard enthalpy apply to just that molecule therefore they are more accurate
25
what is the equation to calculate rate
rate = change in concentration/time
26
what is the unit for rate of reaction
mol dm ^-3 s ^-1
27
what must particles do to react
collide with enough energy (activation energy) and correct orientation
28
how many collisions result in a reaction
not many
29
what are factors affecting rate of reaction
- temperature
- pressure
- concentration
- surface area
- catalyst
30
how does an increase in temperature increase rate
leads to an increase in rate
higher proportion of particles have more energy than activation energy
more successful collisions --> faster rate
31
what effect does an increase in pressure/concentration have on rate
leads to an increase in rate as there are more particles in a given volume -> more frequent successful collisions = faster rate reaction rate
32
how do you calculate rate from a concentration time graph
draw tangent
work out gradient of tangent using
change in y/change in x
32
what variables can be monitored to calculate rate
conc of reactant/product
gas volume of product
mass of substances formed
33
definition of a catalyst
substance that increases rate of reaction by lowering activation energy but is not used up in the reaction
34
how does a catalyst work
provide another reaction pathway that has a lower activation energy
so particles have more energy - collide more often successfully - increased reaction rate
35
what is dynamic equilibrium
occurs when rate of forwards reaction equals rate of backwards reaction and concentration of both remain constant in a closed system
36
what factors affect position of equilibrium
temperature
pressure
concentration
37
what is Le Chatelier's principle
if a system is disturbed, the position of equilibrium shifts to oppose/reduce the change
38
if a forwards reaction is endothermic
what effect would increasing temperature have on equilibrium
equilibrium would shift right - endothermic side - to lower the temperature
39
why is high pressure good for production of some chemicals
increase in pressure means more particles in a given volume - more successful collisions - faster rate
40
does a catalyst ahve an effect on position of equilibrium
no - effects both sides equally
41
what is a heterogenous catalyst
a catalyst that is a different state to reactants
42
what is a homogenous catalyst
a catalyst that is the same state as reactants
43
what is periodicity
the repeating trends in physical and chemical properties
44
what is first ionisation energy
the energy required to remove one electron from each atom in one mol of gaseous elements to form 1 mol of gaseous 1+ ions
45
what are factors affecting first ionisation energy
atomic radius
nuclear charge
electron shielding
46
explain the trend in Period 3 and explain why there is a dip between Mg and Al and between P and S
first ionisation energy increases across period 3
because of - increased nuclear charge, decreased atomic radius and same electron shielding.
so more energy is required to remove the electron
dip at Al
- outer electron is in a 3p orbital which is higher energy than the 3s - less energy required to remove electron
Dips at S
- one 3p orbital contains 2 electrons - repulsion between the paired electrons - less energy required to remove one
47
why does IE decrease between group 2 to 3
because in group 3 the outermost electrons are in p orbitals whereas in group 2 they are in S orbital so the electrons are easier to remove in group 3
48
what is the order of electron shells
1s2s2p3s3p4s3d4p
49
what do successive ionisation energies allow predictions to be made about
number of electrons in outer shell
the group of the element in the periodic table
identity of an element
50
For the first 20 elements, what are the two trends in first ionisation energies
a general increase in ionisation energy across each period
a sharp decrease in in the first ionisation between the end of one period and the start of the next
51
why is there a decrease in first ionisation energies down a group
although the nuclear charge increases, its effect is outweighed by the increased radius and, to a lesser extent the increased shielding
52
why is there a fall in ionisation energy from beryllium to boron
start of filling the 2p sub shell
the 2p sub shell in boron has a higher energy than the 2s shell in beryllium and is therefore easier to remove an electron from boron
53
explain the fall in ionisation energies from nitrogen to oxygen
both have electrons in the p shell
in oxygen, the paired electrons in one of the 2p orbitals repel each other making it easier to remove an electron from an oxygen than a nitrogen so IE is lower
54
what are the general trends across a period
across period 2, the 2s sub-shell fills with 2 electrons, followed by the 2p subshell with 6
across period 3, same pattern with filling the 3s and 3p shells
across period 4, although the 3d sub-shell is involved, the highest shell number is n=4
only 4s and 4p are filled
55
what change happens across each period
turns from metals to non-metals
56
write the equation for the first ionisation energy of magnesium
Mg (g) --> Mg+ (g) + e-
57
what are the properties of giant metallic lattices
high mp and bp
good electrical conductors
malleability
ductility
58
what is a ductile metal
the metal can be stretched
59
what is a malleable metal
metal that can be shaped into different forms
61
what are Group 2 elements
metals
62
what charge forms on Group 2 elements
+2
63
what are the properties of Group 2 elements
relatively high mp and bps
low densities
form colourless compounds
64
what is second ionisation energy
energy required to remove 1 electron from each ion in one mole of gaseous 1+ ions to form 1 mol of gaseous 2+ ions
65
what do Group 2 elements form with oxygen
metal oxides with ionic bonding
66
what is an example of a Group 2 reaction with oxygen
2Ca (s) + 2O (g) --> 2CaO (s)
Calcium is being oxidised from 0 to +2
Oxygen is reduced from 0 to -2
67
what do Group 2 elements form with water
metal hydroxides and hydrogen
68
what is an example of a Group 2 reaction with water
Mg(s) + 2H2O(l) --> Mg(OH)2 + H2
69
what is the trend of hydroxide solubility down Group 2
solubility of hydroxides increase down the group as there are more OH- ions in the solution
pH increases
70
what do Group 2 elements form with dilute acid
salts and hydrogen
71
what is an example of a Group 2 reaction with dilute acid
Mg(s) + 2HCl --> MgCl2 + H2
Mg + 2H2SO4 --> MgSO4 + H2
Mg + 2HNO3 --> MgNO3 + H2
72
what is are Group 2 elements also known as other than metals
reducing agents as they reduce another species
(they are themselves oxidised)
73
what is the trend in reactivity down Group 2 and why
reactivity increases because the atomic radius increases, leading to weaker nuclear attraction for the outer electrons making them easier to lose
74
A student adds magnesium metal to an aqueous solution of ethanoic acid, CH3COOH.
Write the overall equation for this reaction and explain, in terms of oxidation numbers, which element has been oxidised and which element has been reduced.
Mg + 2CH3COOH --> Mg(CH3COO)2 + H2
Mg is oxidised from 0 to +2
H is reduced from +1 to 0
75
The reaction of barium with bromine is more vigorous than the reaction of calcium with bromine.
Explain why.
Ba has greater atomic radius than Ca
Less nuclear attraction in Ba
76
A student plans to prepare a solution of Ba(OH)2 from barium by two different reaction routes.
Outline 2 reaction routes for preparing a solution of Ba(OH)2 from barium in the laboratory.
Include relevant equations.
Route 1
Add water to Ba
Ba + 2H2O --> Ba(OH)2 + H2
Route 2
2Ba + O2 --> 2BaO
BaO + H2O --> Ba(OH)2
77
Write the equation for the reaction of strontium with water.
Sr + 2H2O --> Sr(OH)2 + H2
78
Describe two observations which would be different if the student had used calcium in place of strontium when reacting with water
less vigorous effervescence
dissolves slower
less alkaline
precipitate forms/ less soluble
79
Write the equation for the reaction of calcium oxide with water
CaO + H2O --> Ca(OH)2
80
Calcium nitride reacts with water to form a solution containing two alkaline compounds.
Write an equation for this reaction.
Ca3N2 + 6H2O --> 3Ca(OH)2 + 2NH3
81
Barium chloride can be prepared from barium hydroxide in a neutralisation reaction
Ba(OH)2 + 2HCl --> 2BaCl + 2H2O
82
The reactivity of the Group 2 elements Mg–Ba increases down the group.
Explain why.
atomic radius increases
nuclear attraction decreases
ionisation energy decreases
83
Identify a compound of calcium that could be used to convert a soil pH from 5.8 to 7.5.
calcium hydroxide / calcium oxide
84
Write the equation for the reaction of magnesium with water.
Mg + 2H2O --> Mg(OH)2 + H2
85
Explain why the first ionisation energy of strontium is less than the first ionisation energy of calcium.
larger atomic radius
increased charge outweighed by increased distance
less nuclear attraction
86
putting a small piece of calcium to a large excess
of water.
Describe what the student would observe and write the equation for the reaction.
Effervescence
the solid dissolved
87
what kind of reaction is calcium carbonate to calcium oxide
thermal decomposition
88
Predict the structure and bonding of Ba3N2
Giant Ionic
89
what group are the halogens
group 7
90
why does boiling point increase down Group 7
more electrons
stronger london forces
more energy required to break intermolecular forces
91
what are the states
F2
Cl2
Br2
I2
at room temperature and pressure
F2 and Cl2 are gas
Br2 is a liquid
I2 is a solid
92
what is the trend in oxidising power down Group 7
weakens as
increased atomic radius
increased no. shells and shielding
93
what is the trend in atomic radius in Group 7
increases down a group because of an increase in shielding
94
what is the trend in electronegativity in Group 7
decreases down the group as there is weaker nuclear attraction because of increased shells
95
what is the trend in melting and boiling points in Group 7
increases down a group because there are more electrons going down the group which cause stronger london forces
96
what colour is
Cl2
Br2
I2
in water solution
pale green
orange
brown
97
what colour is
Cl2
Br2
I2
in solution of cyclohexane
pale green
orange
violet
98
What can Cl2 be displaced by and what is the colour change
neither Br2 or I2
no change
99
What can Br2 be displaced by and what is the colour change
Cl2
goes orange as Br2 molecules are formed
Cl2(aq) + 2Br-(aq) --> 2Cl-(aq) + Br2(aq)
no reaction with I2
100
What can I2 be displaced by and what is the colour change
Violet from I2 formation
Cl2(aq) + 2I(aq)- --> 2Cl- (aq) + I2 (aq)
Violet from I2 formation
Br2(aq) + 2I-(aq) --> 2Br- (aq) + I2 (aq)
101
write the equation of the redox reaction between chlorine and sodium bromide
Cl2 + 2naBr --> 2NaCl + Br2
102
why won't iodine displace chlorine
it is less reactive
103
write the equation process that gives you chlorate ions
1. Cl2 + H2O --> HClO + HCl
2. HClO + H2O --> ClO- + H3O
104
what are the benefits of using Cl2 in water
kills lots of disease-causing micro-organisms
105
what are the risks of using Cl2 in water chlorination
toxic
is an irritant in small concentrations and fatal in large concentrations
reacts with hydrocarbons to form carcinogens
however the risks are outweighed by the benefits
106
how do you test the anion CO3-
add dilute HNO3
observe effervescence
collect carbon dioxide gas
bubble gas through limewater which would go cloudy
CO3- + 2H+ --> CO2+ H2O
107
how do you test for SO4(2-)
add solution of BaNO3 (if not testing gfor halides then use BaCl2)
white precipitate forms
Ba (2+) + SO4 (2-) --> BaSO4
108
how do you test for X- (halides)
dilute HNO3
add AgNO3
the ppt forming:
Cl- = white
Br- = cream
I- = yellow
109
how do you further distinguish chlorine from the other halides after HNO3 and AgNO3
add dilute NH3 which will make the precipitate dissolve
doesn't dissolve the ppt in I and Br
110
how can you further distinguish bromine from the other halides after HNO3 and AgNO3
add concentrated NH3 - ppt will dissolve
111
how do you test for NH4+
add OH- ions
warm mixture
NH3 gas is produced
will turn damp red litmus paper blue
112
what is a disproportionation reaction
when the same element is reduced and oxidised in the same reaction
113
Sodium chloride - orange
Sodium Iodide - violet
Explain how the student’s results provide evidence for the trend in reactivity of the halogens down group 7 and write an ionic equation for any reaction that takes place.
Use your chemical knowledge to explain the trend in reactivity.
orange = bromine
violet = iodine
Br2 + 2I- --> 2Br - + I2
reactivity decreases down the group and oxidising power decreases
greater atomic radius
less nuclear attraction down the group
114
why does reactivity decrease down Group 7
more shielding
greater atomic radius
nuclear attraction decreases
115
Chlorine is used in water treatment.
State one benefit and one risk of using chlorine in water treatment.
kills bacteria
toxic/forms carcinogens/forms chlorinated hydrocarbons
116
Write the systematic name for NaBrO3.
sodium bromate (V)
117
Explain the trend in boiling points of the halogens
number of electrons increases down the group
London forces present
London forces get stronger down the group and so they require more energy to break
118
Iodine reacts with water as shown below.
I2 + H2O ⇌ HI + HIO
Using oxidation numbers, explain why this reaction is a disproportionation.
disproportionation = oxidation and reduction of same element
iodine has been oxidised and reduced
oxidation from 0 to +1
reduction from 0 to -1
119
Explain why iodine is less reactive than bromine.
iodine has a larger atomic radius
iodine has greater shielding/ more shells
iodine has weaker nuclear attraction
120
Bromine disproportionates when it reacts with potassium hydroxide solution.
Suggest an equation for this reaction.
Br2 + KOH --> KBr + KBrO + H2O
121
What is the systematic name for KCl O4?
potassium chlorate (VII)
122
Precipitation reactions can be used to distinguish between halide ions.
State the reagent needed for these precipitation reactions.
AgNO3
123
Write the electron configuration of a bromide ion, in terms of sub‐shells.
4p6
124
State the solutions that need to be added together in order to prove the trend in reactivity
of the halogens, using the minimum number of test-tube experiments.
* Describe the colour seen in the organic solvent at the end of each test-tube experiment.
* Write an ionic equation for one reaction that takes place.
add aqueous barium chloride to bromine water
- colour goes orange
add aqueous calcium chloride to bromine water
- colour goes orange
add aqueous magnesium bromide to aqueous iodine
- colour goes violet
Br2 + 2I- --> 2Br- + I2
125
Chlorine reacts with aqueous sodium hydroxide to form bleach.
Write the equation and state the conditions for this reaction.
2NaOH + Cl2 --> NaClO + NaCl + H2O
cold and dilute NaOH
126
Outline a simple practical test that would confirm the presence of chloride ions in the lower layer, and
give the expected result
add a few drops of AgNO3
127
Name the apparatus that could be used to separate the two liquid layers present at the end of the
experiment
separating funnel
128
A student bubbles chlorine gas through aqueous potassium iodide. A reaction takes place.
state observations and give the ionic equation
cream ppt
Ag+ + Br- --> AgBr
129
Chlorine gas can be added to a cold, dilute alkaline solution to form bleach.
Write the equation for this reaction.
Cl2 + 2NaOH → NaClO + NaCl + H2O
