periodic table and energy Flashcards
1
Q
what is the periodic table? (3)
A
the arrangement of elements by:
1) increasing atomic (proton) number
2) periods showing repeating trends in physical and chemical properties (periodicity)
3) in groups having similar chemical properties
2
Q
what is the definition of first ionisation energy?
A
energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous + ions.
3
Q
what does ionisation energy measure?
A
measures how easily an atom loses electrons to form positive ions
4
Q
what factors affect ionisation energy? (3)
A
1) no. of electron shells/ atomic radius
2) nuclear charge
3) electron shielding (inner electron shells blocking charge from the nucleus)
5
Q
what is the first ionisation energy equation for sodium?
A
Na (g) –> Na+(g) + e-
6
Q
what is the general rule in successive ionisation energies?
A
no. of electrons = ionisation energies
7
Q
every time an electron is removed… (3) (successive ionisation energy)
A
1) there is a greater nuclear attraction
2) atomic radius is slightly smaller
3) successive ionisation energies will increase
8
Q
how can you tell the electronic structure of an element from data which gives ionisation number and energies? (2)
A
1) name where the jump in energy occurs
2) say how many electrons are in its outer shell from that
9
Q
what do trends in 1st ionisation energy provide evidence for?
A
the existence of shells and sub shells
10
Q
what is the trend in first ionisation energies down a group? (3)
A
1) first ionisation energy decreases down a group
2) atomic radius increases down a group
3) so shielding increases meaning it is easier for the electron to be lost as attraction is weaker between the nucleus and electron
11
Q
what is the trend in first ionisation energies across period 2? (3)
A
1) general increasing trend along the period
2) due to decreasing atomic radius and increasing nuclear charge meaning outer electrons are held more strongly
3) which requires more energy to overcome
12
Q
why is boron an exception to the trend in first ionisation energy across period 2? (2)
A
1) Boron has a lower first ionisation energy than expected due to the energy difference between 2s and 2p sub shells / new sub shell shielding
2) the electron is being removed from a higher energy level that is further from the nucleus so the electron is held less strongly.
13
Q
why is oxygen an exception to the trend in first ionisation energy across period 2? (3)
A
1) oxygen has a lower first ionisation energy than expected by the general trend due to repulsion within the 2p orbital
2) two electrons with opposite spins are placed in the same orbital
3) allows electron to be removed from oxygen more easily as the repulsion is destabilising
14
Q
what is the trend in first ionisation energies across period 3? (3)
A
1) general increasing trend
2) due to decreasing atomic radius and increasing nuclear charge
3) meaning outer electrons are held more strongly
15
Q
why is aluminium an exception to the trend in first ionisation energy across period 3? (3)
A
1) aluminium has a lower first ionisation energy due to the energy difference between the 3s and 3p sub shells
2) the electron is being removed at a higher energy level which is further away from the nucleus
3) held less strongly so requires less energy to remove
16
Q
why is sulfur an exception to the trend in first ionisation energy across period 3? (3)
A
1) lower first ionisation energy than expected due to the electron repulsion in the 3p sub shell
2) electrons with opposite spins held in the same orbital
3) allows electron to be removed from sulfur more easily as the repulsion is destabilising
17
Q
what is the general trend for first ionisation energy down a group? (3)
A
1) first ionisation energy decreases down a group
2) more electron shells lead to increasing atomic radius so weaker nuclear attraction
3) shielding also increases down a group which further reduces attraction
18
Q
what is the general trend for first ionisation energy across a period? (4)
A
1) first ionisation energy increases across a period
2) due to a decreasing atomic radius due to extra nuclear charge
3) greater nuclear attraction to overcome
4) more nuclear charge because of an extra proton
19
Q
what 3 things affect first ionisation energy? describe the trend in these across a group and down a period (3)
A
1) atomic radius (increases down a group, decreases across a period)
2) nuclear charge (increases across a period, decreases down a group)
3) shielding (constant across a period, increases down a group)
20
Q
why does shielding remain constant across a period?
A
the number of energy levels is the same
21
Q
why does atomic radius decrease across a period? (2)
A
1) elements gain an extra proton for the same number of electron shells, which increases nuclear charge
2) attracting the electron closer which decreases atomic radius
22
Q
how are atomic radius and nuclear charge linked?
A
A higher nuclear charge causes greater attractions to the electrons, pulling the electron shells closer to the nucleus which results in a smaller atomic radius.
23
Q
what is the periodic trend in electron configurations across Periods 2 and 3? (2)
A
1) across period 2, the 2s sub shell fills with 2 electrons and then the 2p subs shells fill with 6 electrons
2) in period 3 the pattern repeats itself with the 3s and 3p sub-shells.
24
Q
How are elements classified into SPDF block elements?
A
classification is based on the name of orbitals which receive the last electron.
25
what can successive ionisation data be used for? (3)
1) predict or confirm the simple electronic configuration of elements (the number of ionisation energies shows the number of electrons)
2) confirm the number of electrons in the outer shell of an element
3) deduce the group an element belongs to in the periodic table
26
what is electron shielding?
when the inner shells create a barrier that blocks the attractive forces
27
what is metallic bonding?
strong electrostatic attraction between cations and delocalised electrons
28
what is ionisation energy measured in?
kJmol-1
29
when do successive ionisation energies occur?
when further electrons are removed after first ionisation energy
30
why does successive ionisation energy require more energy? (2)
1) as electrons are removed, nuclear attraction between the positive nucleus and the negative outer shell electron increases
2) more energy is needed to overcome this attraction, so ionisation energy increases
31
what is the structure of metals? (3)
1) metal loses all of its outer shell electrons causing a sea of delocalised electrons
2) lattice of cations
3) strong electrostatic attraction between the oppositely charged particles
32
what should you remember when drawing a diagram of metallic structure?
when asked to draw a specific metal, add the charge to the cation
e.g
Ca would be 2+
33
what are the properties of metals? (3)
1) conductor of electricity as delocalised electrons are able to move and carry charge through their structure
2) high melting and boiling points due to strong electrostatic attraction between cations and electrons which require a lot of energy to break
3) insoluble as metallic bonds are too strong to be overcome
34
which 3 elements form giant covalent structures?
Boron, Carbon and Silicon
35
what are the properties of giant covalent structures? (3)
1) high melting and boiling points due to a high proportion of strong covalent bonds which require a lot of energy to break
2) insoluble as covalent bonds are too strong to be overcome
3) non conductors of electricity as there are no delocalised electrons which can move and carry charge.
36
what is the periodic trend in melting points across period 2? (4)
1) increases between Li and Be as Be forms 2+ ions so there is a stronger metallic bonds as there are more free electrons
2) increases between Be to B to C as bonding switches to giant covalent structures
3) decreases from C to N, O and F as they have simple molecular structure
4) decreases most at Ne as it is a group 0 element
37
what are the simple molecular structures in period 2?
N2, O2, F2, Ne
38
what are the simple molecular structures in period 3?
P4, S8, Cl2, Ar
39
what is a giant covalent lattice?
a solid giant covalent lattice is a network of atoms bonded by many strong covalent bonds
40
what are the properties of giant covalent structures? (3)
1) high melting and boiling points due to each atom having multiple strong covalent bonds
2) insoluble as covalent bonds are too strong to be overcome
3) non conductors of electricity as there are no delocalised electrons which can move and carry charge
41
what are the allotropes of carbon?
diamond, graphite and graphene
42
what are the features of diamond? (3)
1) each carbon makes 4 covalent bonds
2) causes a rigid/ hard structure
3) non conductor of electricity as there are no delocalised electrons which can move and carry charge
43
what are the features of graphite? (3)
1) each carbon makes 3 covalent bonds in flat sheets
2) one delocalised electron per atom of carbon which can move between layers so it can conduct electricity
3) weak London forces between molecules allow layers to slide over each other easily
44
what are the features of graphene? (4)
1) 2D sheets of graphite
2) form hexagonal carbon rings
3) strong, rigid, light structure
4) conductor of electricity due to only 3 covalent bonds
45
what causes the variation in melting point in period 2?
due to different bond strength and structure
46
what is the melting point like for Li and Be in period 2? (3)
1) melting point increases from Li to Be because Be has a greater positive charge of 2+ when forming ions
2) more free electrons in Be lattice
3) electrostatic forces are greater than that in lithium
47
what is the melting point like for B and C in period 2? (2)
1) B and C form giant covalent lattices with very strong covalent bonds
2) covalent bonds require a lot of energy to break, causing high melting points
48
what is the melting point like for N2 O2, F2 and Ne in period 2? (2)
1) simple covalent molecules with weak London forces
2) relatively similar low melting and boiling point as not much energy is required to overcome these forces
49
what causes the variation in melting point in period 3?
difference in bond strength and structure
50
what is the melting point like for Na, Mg and Al in period 3? (3)
1) melting point increases from Na to Al due to greater positively charged metal ions
2) more electrons are released as free electrons across the metals
3) electrostatic attraction between electrons and metal cations increases from Na to Al
51
what is the melting point like for Si in period 3? (2)
1) giant covalent structure
2) strong covalent bonds which require a lot of energy to break causes high melting point
52
what is the melting point like for P4, S8 and Cl2 in period 3? (3)
1) simple covalent molecules
2) weak London forces between them
3) low similar boiling points as it does not require a lot of energy overcome intermolecular forces
53
what is the melting point like for Ar in period 3? (2)
1) noble gas with a full outer shell of electrons
2) stable molecule so London forces are weak
54
what valence shell configuration do group 2 metals have and how does this link to how they react? (2)
1) s2
2) group 2 metals lose these 2 electrons in redox reactions to form 2+ ions
55
what happens when group 2 elements are exposed to air? (2)
1) tarnish in air
2) forming a coat of metal oxide
56
what is the trend in reactivity down group 2? (3)
1) increases down the group
2) because ionisation energies decrease down the group
3) due to increased electron shielding and atomic radius which makes outer electrons easier to lose in order to react
57
explain the trend in reactivity in terms of the first and second ionisation energies of group 2 elements down the group? (2)
1) both first and second ionisation energy of group 2 elements decrease down the group
2) due to greater atomic radius and increased shielding which makes outer shell electron easier to lose
58
describe the relative reactivities of group 2 elements Mg--> Ba shown by their redox reactions with water? (2)
1) group 2 metals react with water in a redox reaction to produce a metal hydroxide and hydrogen gas
2) metal hydroxide forms as an alkaline solution
e.g Mg + 2H2O --> Mg(OH)2 + H2
Mg is oxidised from 0 to +2
H is reduced from +2 to 0
59
describe the trend in reactivity of group 2 elements down the group with water (3)
1) increases down the group
2) Ca, Sr and Ba all react with water with increasingly vigorous reactions
3) reactions all form metal hydroxide and hydrogen gas
60
describe the relative reactivities of group 2 elements Mg--> Ba shown by their redox reactions with oxygen
group 2 metals react with oxygen to form metal oxides
e.g
2Mg + O2 --> 2MgO
61
describe the relative reactivities of group 2 elements Mg--> Ba shown by their redox reactions with dilute acids?
give examples with all 3 acids
group 2 metals react with dilute acids to produce hydrogen gas and a salt
e.g
X + 2HCl --> XCl2 + H2
X + H2SO4 --> XSO4 + H2
X + 2HNO3 --> X(NO3)2 + H2
62
what is the solubility of group 2 oxides like? (3)
1) group 2 oxides increase in solubility down the group
2) more hydroxide ions released in solution
3) more alkaline solution down the group
63
describe the action of water on group 2 oxides and the approximate pH of any resulting solutions, including the trend in increasing alkalinity (3)
1) group 2 oxides react with water forming an alkaline solution of a metal hydroxide
e.g
CaO + H2O --> Ca(OH)2
2) hydroxides are only slightly soluble in water- once the solution is saturated any further, hydroxide ions form solid precipitate
3) more alkaline solution down the group as more OH- release as solubility increases
64
what is the outer shell configuration for group 2?
how does this influence the most common type of reaction of group 2 elements? (2)
1) s2 electron configuration
2) the loss of these electrons occurs in redox reactions to form 2+ ions
65
explain the trend in reactivity in terms of the first and second ionisation energies of Group 2 elements down the group (3)
1) Group 2 elements react losing 2 electrons, forming 2+ cations
2) The easier it is to lose electrons (i.e. the lower the first and second ionisation energies), the more reactive the element.
3) So because ionisation energy decreases down the group (due to increase in atomic radius+shielding) reactivity increases down the group.
66
describe the uses of some group 2 compounds as bases, including equations (2)
1) Ca(OH)2 in agriculture to neutralise acid soils- has to be used in moderation
2) Mg(OH)2 and CaCO3 as ‘antacids’ to neutralise stomach acids/ treating indigestion, CaCO3 is an insoluble base so it can be swallowed in indigestion tablets.
equation:
- Mg(OH)2(s) + 2 HCl(aq) → MgCl2(aq) + 2 H2O(l)
- CaCO3(s) + 2 HCl(aq) → CaCl2(aq) + H2O(l) + CO2(g)
- H+(aq) + OH-(aq) -> H2O(l)
67
what are some features of the halogens? (4)
1) most reactive non-metallic group
2) do not occur in elemental form
3) occur as stable halide ions dissolved in sea water/ combined with Na or K as solid deposits
4) all halogens exist as diatomic molecules
68
what is the trend in boiling point for the halogens? (3)
1) increases down the group
2) London forces get stronger due to more electrons
3) so bigger dipoles form bewitch require more energy to overcome
69
describe the outer shell electron figuration of the halogens and how this affects how they react
1) s2p5
2) gain one electron in many redox reactions to form 1- ions
70
what do halogens form in their solid state?
lattices with simple molecular structures
71
what is the appearance of the halogens at room temperature? (4)
1) F2--> pale yellow gas
2) Cl2 --> pale green has
3) B2--> red-brown liquid
4) I2--> shiny black-grey solid
72
what is the most common type of reaction for the halogens? (2)
1) redox
2) each halogen atom is reduced and acts as an oxidising agent (gains an electron from another element which loses it)
73
what is the general half equation for the reduction of halogens in halogen displacement reactions?
X2(aq) + 2e- --> 2X- (aq)
(more reactive one)
74
what is the general half equation for the oxidation of halides in halogen displacement reactions?
2X- (aq) --> X2 (aq) + 2e-
(less reactive one)
75
in aqueous solutions what do Cl2, Br2 and I2 appear as? (3)
1) Cl2 (aq) --> pale green
2) Br2 (aq) --> orange
3) I2 (aq) --> brown
76
why is cyclohexane added to halogens in aqueous solutions?
to make the colour more clear
77
in the cyclohexane layer, what do Cl2, Br2 and I2 appear as? (3)
1) Cl2 (aq) --> pale green
2) Br2 (aq) --> orange
3) I2 (aq) --> violet
78
what do the results of halogen displacement reactions show?
that the reactivity of halogens deceases down the group
79
what happens during a halogen displacement reaction?
a solution of each halogen is added to aqueous solutions of the other halides
(halogen and halide will be different elements)
80
give an example of a halogen displacement reaction and give the ionic and half equations
Cl2 (aq) + 2NaBr (aq) --> 2NaCl (aq) + Br
chlorine is more reactive than bromine
ionic equation:
Cl2 (aq) + 2Br- (aq) --> 2Cl- (aq) + Br2 (aq)
half equations:
Cl2(aq) + 2e- --> 2Cl-
2Br- --> Br2 (aq) + 2e-
81
what is the general rule for writing ionic equations in halogen displacement reactions?
eliminate metal ion- only thing that remains unchanged/ spectator ion
82
explain the trend in reactivity of the halogens from the decreasing ease of forming 1- ions (4)
1) decreases down the group as it is harder to gain an electron/ their oxidising ability decreases
2) due to extra electron shells which leads to a larger atomic radius
3) more shielding
4) weaker attraction to the nucleus so it is harder to gain the electron they need to react
83
explain the trend in reactivity of the halogens illustrated by reactions with other halide ions (2)
1) the reactivity of halogens can be shown by their displacement reactions with other halide ions in solutions
2) a more reactive halogen can displace a less reactive halogen from a halide solution of the less reactive halogen
84
what is a disproportionation reaction?
the same element is both oxidised and reduced in a reaction, can be illustrated by analogous reactions.
85
what is one example of a disproportionation reaction?
reaction of chlorine with water used to purify water
Cl2 (g) + H2O(l) --> HCl (aq) + HClO (aq)
(hypochlorous acid)
Cl is both reduced and oxidised
86
what is another example of a disproportionation reaction?
reaction of chlorine with aqueous sodium hydroxide
Cl2 (g) + 2NaOH (aq) --> NaClO (aq) + NaCl (aq) + H2O (l)
(sodium hypochlorite)
Cl is both oxidised and reduced
87
what is the trend in oxidising power down the halogens? (4)
decreases
1) ability to attract electrons decreases
2) more electron shells so greater atomic radius
3) more shielding
4) weaker electrostatic attraction to attract electron and be reduced
88
what are the pros and cons of chlorine use in water treatment? (3)
pros: kills bacteria
cons: risks with hazards of toxic chlorine gas
possible risks from formation of chlorinated hydrocarbons
89
what is the problem with the test for halide ions? how can you resolve this? (2)
1) colours are hard to tell apart
2) add dilute ammonia solution
90
describe the test for halide ions and the ionic equation
1) add nitric acid
2) add silver nitrate
3) results:
AgCl- --> white precipitate
AgBr- --> cream precipitate
AgI--> yellow precipitate
ionic equation:
Ag+ (aq) + X- --> AgX (s)
91
describe the results of adding dilute ammonia solution after the halide test
Cl --> precipitate dissolves
Br --> no change/ cream precipitate
I --> no change/ yellow precipitate
92
describe the results of adding concentrated ammonia solution after the halide test
Cl --> precipitate dissolves
Br --> precipitate dissolves
I --> no change/ yellow precipitate
93
describe the test for carbonates (4)
1) add a dilute acid (preferably nitric),
2) CO32- will react with H+ (aq) forming CO2 (g)
3) bubbles of CO2 will be produced
4) bubble through limewater which will turn cloudy
94
what is the ionic equation for the carbonate test
2H + (aq) + CO3 2-(aq) --> CO2 (g) + H2O (l)
95
describe the test for sulfates (3)
1) add nitric acid to eliminate any carbonate ions
2) add barium chloride or barium nitrate
3) will result in a white precipitate with Ba2+ (aq) forming CO2 (g)
96
what is the ionic equation for the sulfate test?
Ba 2+ (aq) + SO4 2- (aq) --> BaSO4 (s)
97
what is the general ionic equation for the halide test?
Ag+ (aq) + X- (aq) --> AgX (s)
98
how would you identify an unknown solution? (6)
1) test for carbonates
2) test for sulfates
3) test for halides
4) use nitric acid instead of sulfuric acid/ hydrochloric in the tests to avoid adding Cl- and SO4 2- ions to sample
5) use barium nitrate instead of barium chloride in the test for sulfates to avoid adding Cl- ions
6) This will prevent false results as both BaCO3 and Ag2SO4 are insoluble.
99
describe the test for ammonium ions (NH4+)
`(4)
1) add dilute sodium hydroxide and heat
2) will produce ammonia and water
3) hold damp red litmus paper in the mouth of test tube
4) paper will turn blue (alkali)
100
what is the ionic equation for the test for ammonium ions?
NH4 + (aq) + OH- (aq) --> NH3 (g) + H2O (l)
101
why should you do the test for carbonates before the test for sulfates?
to eliminate any CO32- ions as they will give a positive result for the sulphate test- remove false positives
102
what is enthalpy and its symbol?
measure of heat energy in a chemical system
symbol: H
103
what are some chemical reactions accompanied by?
why? (2)
1) enthalpy changes that are either exothermic (negative energy change) or endothermic (positive energy change)
2) when chemical reactions take place, changes in chemical energy take place and therefore the enthalpy changes
104
describe the features of an exothermic reaction (6)
1) the enthalpy of the reactants is higher than the products
2) Heat energy is given off by the reaction to the surroundings
3) The temperature of the environment increases
4) The energy of the system decreases
5) There is an enthalpy decrease during the reaction so ΔH is negative
6) more energy is released making bonds than breaking bonds
105
what is a chemical system?
a system of atoms, molecules or ions which make up the chemicals
106
what is an enthalpy change and what is its symbol?
the difference between the enthalpies of the reactants and products
symbol: ΔH
107
what is the general equation for enthalpy change?
H(products)-H(reactants)
108
can enthalpy be measured? (2)
1) enthalpy can not be measured directly
2) however enthalpy change can be measured by an experiment
109
what is the conservation of energy? (2)
1) energy cannot be created or destroyed
2) energy can only be transferred between the system and the surrounding
110
describe the features of an endothermic reaction (6)
1) the enthalpy of the reactants is lower than the products
2) heat energy is taken in from the surroundings
3) the temperature of the environment decreases
4) the energy of the system increases
5) There is an enthalpy increase during the reaction so ΔH is positive
6) more energy is taken in breaking bonds than is released when making new bonds
111
what is activation energy?
the minimum amount of energy required to start a reaction
112
what is one reason that some reactions may happen slowly?
only a small proportion of reactant particles have the activation energy
113
what are the standard conditions? (4)
1) standard pressure: 100kPa
2) standard temperature: 298k (25 degrees)
3) standard concentration: 1 mol dm-3
4) standard state: the state of the substance at standard conditions.
114
what is the enthalpy change of reaction?
enthalpy change that accompanies a reaction in the molar quantities shown in a chemical equation, under standard conditions
115
what is the enthalpy change of formation?
enthalpy change that takes place when one mole of a compound is formed from its elements under standard conditions
116
give an example of the enthalpy change of formation and explain
1/2 N2 (g) + 3/2 H2 (g) --> NH3 (g)
the compound needs to remain as one mole so in order to balance it is necessary to balance the reactants in fractions
117
what is the enthalpy change of combustion?
the enthalpy change that takes place when one mole of a substance reacts completely with oxygen under standard conditions
118
give an example of the enthalpy change of combustion and explain
CH4 (g) + 2O2 (g) --> CO2 (g) + 2H2O (l)
substance needs to remain as one mole but everything else can be balanced.
119
what is the enthalpy change of neutralisation?
the enthalpy change when the reaction of an acid and a base form one mole of water under standard conditions
120
give an example of the enthalpy change of neutralisation and explain
HNO3 (aq) + 1/2Ba(OH)2 (aq) --> 1/2 Ba(NO3)2 (aq) + H2O (l)
water needs to remain as one mole so everything else needs to be balanced with fractions
121
what equation do you use to calculate heat energy which you need to calculate enthalpy changes?
q = mcΔT
q- heat energy (J)
m- mass of water being heated (g or cm3)
c- specific heat capacity of water (4.18)
ΔT- temperature change (K/C)
122
what equation do you use to calculate enthalpy change after calculating heat energy?
q/mol. of limiting reactant = Jmol-1
divide by 1000 to get kJmol-1
123
what do you need to measure when carrying out an enthalpy change of combustion experiment? (3)
1) weigh spirit burner before and after for mol of limiting reactant
2) measure temperature of water before and after
3) mass of water
124
how accurate is the experimental value when doing the enthalpy change of combustion experiment? (4)
1) heat is lost to the surroundings (not all heat goes toward heating the water)
2) evaporation of alcohol from the wick (mass is bigger than it should be on the balance)
3) assuming complete combustion has occured
4) may have occurred under non-standard conditions
125
what experiment do you carry out to determine the enthalpy change of reaction? (4)
1) reaction is carried out in a polystyrene cup with a lid
2) thermometer is placed in polysterene cup
3) measure temperature
4) plot a graph time on x axis and extrapolate for temp. change
126
how do you calculate the enthalpy change of reaction? (2)
1) q= mcΔT
2) ΔrH = q/mol of limiting reactant
- if question gives 2 liquids then add the two volumes together
- if question doesn't specify limiting reactant, work it out by calculating the moles of each reactant then writing out the symbol equation
127
how can you improve experimental values in determining the enthalpy change of reaction? (3)
1) take temperature every 30 seconds/ frequent time intervals
2) plot a graph of temperature v time (time on the x axis)
3) extrapolate to correct for heat loss
128
bond enthalpies are always...
endothermic (+ve enthalpy value)
129
what are the limitations of average bond enthalpies? (2)
1) the actual bond enthalpy can vary depending on the chemical environment of the bond
2) an average bond enthalpy is calculated using bond enthalpies from different chemical environments not the actual bond enthalpy of the bond
130
why does Br2 have a lower boiling point than ICl despite having a similar relative molecular mass/ despite iodine being lower down the group? (2)
1) Br2 has induced dipole- dipole interactions whereas ICl has London forces and permanent dipole-dipole interactions
2) requires more energy to break/overcome
131
what is the equation for the second ionisation energy of strontium
Sr+ (g) --> Sr2+ (g) + e-
132
why does universal indicator paper turn red and then white when chlorine comes into contact with it? (2)
1) turns red as HCl is produced which is a strong acid
2) turns white as it has been bleached by HClO
133
what aqueous reagent can you add to distinguish between AgBr and AgI
1) conc. ammonia solution
2) AgBr will dissolve but AgI won't
134
how can you calculate a specified bond enthalpy when given the rest of the bond enthalpies and the total energy change? (3)
1) write equation and draw out each bond (write equation specific to the question e.g enthalpy change of formation etc)
2) calculate bond enthalpies on each side of the reaction
3) write algebraic equation to work out the bond, labelling it as x.
135
what is the purpose of Hess's law?
allows us to calculate enthalpy changes indirectly
136
how can you use Hess's law to calculate an enthalpy change of reaction from enthalpy changes of formation? (6)
1) construct enthalpy cycle by writing equation on the top and elements on the bottom and drawing arrows from elements to the equation
2) use standard enthalpy changes of formation to calculate enthalpy of each side of the reaction
3) draw out enthalpy cycle arrows, 1 going straight across and 2 as a curve
4) flip sign on the left hand side of the equation as the arrows go in opposite direction
5) enthalpy change of formation for elements is always zero
6) calculate enthalpy change of reaction by adding the 2 enthalpies together
137
how can you use Hess's law to calculate an enthalpy change of reaction from enthalpy changes of combustion? (6)
1) construct enthalpy cycle by writing equation on the top and CO2 and H2O (if reaction doesn't involve carbon use products from reaction instead) on the bottom and drawing arrows from equation to the elements
2) use standard enthalpy changes of combustion to calculate enthalpy of each side of the reaction
3) draw out enthalpy cycle arrows, 1 going straight across and 2 as a curve
4) flip sign on the right hand side of the equation as the arrows go in opposite direction
5 ) calculate enthalpy change of reaction by adding the 2 enthalpies together
138
define rate of reaction
measures how fast a reactant is used up or how fast a product is produced
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what is the equation and units for rate of reaction
1) change in amount (conc/vol/mass)/ time (sec)
- usually conc
2) mol dm -3 sec-1
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how do you calculate rate of reaction from a graph if it asks to calculate rate 'over' x seconds? (3)
1) draw 2 points from 0 seconds to x seconds
2) connect with a ruler and into a triangle to do rise/run
3) do change in y/ change in x
units in mol dm-3 s-1
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how do you calculate rate of reaction from a graph if it asks to calculate rate 'at' x seconds? (3)
1) draw a tangent to the point
2) pick two points on the tangent
3) do change in y/change x
units in mol/ dm-3 s-1
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why is rate of reaction lower as a reaction progresses? (3)
1) reactants are being used up
2) less reactant particles available
3) less frequent successful collisions
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why will a reaction eventually stop?
reactant particles eventually run out
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describe collision theory (3)
1) 2 particles must collide to create a reaction, in order for successful collisions to occur:
2) particles must collide with enough energy/activation energy
3) and must collide in the correct orientation
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how does increasing concentration affect rate of reaction? (3)
1) increases rate of reaction
2) more particles in the same volume
3) more frequent successful collisions
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how does increasing pressure affect rate of reaction? (3)
1) increases rate of reaction in reactions which have at least one reactant as a gas
2) same particles in a smaller volume
3) more frequent successful collisions
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how does increasing surface area affect rate of reaction? (3)
1) increases rate of reaction in reactions which have at least one reactant as a solid
2) more particles exposed on the surface
3) more frequent successful collisions
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how can you alter the enthalpy change of neutralisation experiment to prevent heat loss? (2)
1) use a polystyrene cup
2) place a lid on top of the cup
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what material should a calorimeter be made from?
copper
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what are the measurable properties to follow the rate of reaction? (3)
1) gas volume
2) mass/conc of reactants/products
3) colours
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what 2 ways can you monitor rate of reaction in reactions that produce gases? (2)
1) monitoring volume of gas at regular intervals
2) monitoring loss of mass using a balance
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what is a catalyst?
a substance which speeds up a chemical reaction without being used up
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how does a catalyst speed up rate of reaction? (2)
1) a catalyst works by providing an alternative pathway for the reaction which has a lower activation energy
2) this speeds the reaction up because more of the particles have energy greater than the activation energy so more frequent successful collisions
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what is a heterogeneous catalyst? (2)
1) a catalyst in a different state to the reactants
2) reaction happens on the surface of the catalyst by adsorption and desorption
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what is homogeneous catalyst? (2)
1) a catalyst in the same state as the reactants
2) works by forming an intermediate species
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what are the advantages of using a catalyst (4)
1) great economic importance as only a small amount of catalyst can catalyse a lot reactants
2) reduces energy demand from combustion of fossil fuels so resulting reduction in CO2 emissions
3) benefits to the environment of improved sustainability weighed against toxicity of some catalysts.
4) can use lower pressures + temperatures so cheaper/less energy used/ more sustainable
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what are the disadvantages of using a catalyst
most only work for one type of reaction
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what is the Boltzmann distribution?
the spread of molecular energies in gases is shown by the Boltzmann distribution
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explain how the presence of a catalyst increases the rate of reaction (3)
1) activation energy decreases
2) so there is a greater proportion of molecules with energy greater than the activation energy
3) so more frequent successful collisions
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explain how an increase in temperature increases the rate of reaction (3)
1) activation energy stays the same
2) however a greater proportion of the molecules will have energy greater than the activation energy
3) so more frequent successful collisions
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describe and explain how the rate of reaction is affected by a decrease in temperature (3)
1) a decrease in temperature will decrease the rate of reaction
2) because a smaller proportion of the molecules have energy greater than the activation energy
3) less frequent successful collisions
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where does a dynamic equilibrium exist? (3)
1) a dynamic equilibrium exists
in a closed system
2) when the rate of the forward reaction is equal to the rate of the reverse reaction
3) and the concentrations of reactants and products do not change
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use le chataliers principle to explain what will happen to the yield of C if the conc. of A is increased in the reaction
A + B -><- C + D (2)
1) equilibrium moves to the right to reduce the concentration of A
2) so the yield of C increases
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use le chataliers principle to explain what will happen to the yield of C if the conc. of D is increased in the reaction
A+ B -><- C + D (2)
1) equilibrium moves to the left to reduce the concentration of D
2) so the yield of C deceases
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use le chataliers principle to explain what will happen to the yield of C if the pressure is increased in the reaction
2A(g) + B (g) -><- C(g) + D(g) (2)
1) equilibrium moves to the right with the fewer gas molecules to decrease the pressure
2) so the yield of C increases
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use le chataliers principle to explain what will happen to the yield of C if the pressure is decreased in the reaction
2A(g) + B (g) -><- C(g) + D(g) (2)
1) equilibrium moves to the left with the most gas molecules to increase the pressure
2) so the yield of C decreases
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use le chataliers principle to explain what will happen to the yield of C if the temperature of this reaction is increased (forwards reaction is exothermic)
A + B -><- C + D (2)
1) equilibrium moves to the left to try decrease the temperature as the backwards reaction is endothermic
2) so the yield of C decreases
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use le chataliers principle to explain what will happen to the yield of C if the temperature of this reaction is decreased (forwards reaction is exothermic)
A + B -><- C + D (2)
1) equilibrium moves to the right to try increase the temperature as the forwards reaction is exothermic
2) so the yield of C increases
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summarise using le chataliers principle to determine equilibria position in temperature changes (2)
1) if temperature is increased, the equilibria moves in the direction which is endothermic
2) if the temperature is decreased, the equilibria moves in the direction which is exothermic
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what is the effect of a catalyst on equilibrium position (3)
1) a catalyst increases the rate of
of both forward and backwards reactions in an equilibrium by the same amount
2) resulting in an unchanged position of equilibrium
3) so does not change yield/ has no effect
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in the reaction
C2H4 + H2O -><- C2H5OH
energy change: -46kj/mol
compromise:
300 degrees
60 atm
why are these temperatures and pressures used? (7)
1) forward reaction is exothermic
2) decreasing temp favours forward reaction
3) yield of ethanol increases
4) a lower temp would give a slower rate of reaction
5) so increases the rate of reaction without compromising the yield
6) 60atm is safer and less expensive
7) gives maximum yield and faster reaction
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when must a compromise be made in industry? what happens? (3)
1) when a factor increases product yield but decreases rate of reaction
2) so reactant conditions are chosen to give a relatively good product yield and a relatively fast rate of reaction
3) safety and economics are also taken into account when determining conditions used in industrial reactions
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what techniques and procedures can be used to investigate changes to the position of equilibrium for changes in concentration and temperature.
monitoring colour change depending on yield being produced/ used up
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in the haber process why is ammonia removed as soon as its formed? (3)
1) so that concentration of ammonia remains low
2) so equilibrium shifts to the right
3) and ammonia yield increases
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what is equilibrium constant, Kc used for? (2)
1) indicates the position of equilibrium for a reaction
2) used in homogeneous equations
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how do you calculate equilibrium constant, Kc? (2)
1) conc. of product/ conc of reactants
2) any variation in the no. of moles raises the concentration of that substance to a power with the same value as the number of moles
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write the expressions for the equilibrium constant, Kc of this reaction:
H2 (g) + I2(g) -><- 2HI(g)
Kc= [HI(g)]2/ [H2(g)] [I2(g)]
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how can you estimate the position of equilibrium from the magnitude of Kc (3)
1) if Kc > 1, there are more products than reactants at equilibrium so equilibrium position lies right
2) if Kc= 1, there is the same amount of reactant and products at equilibrium
3) if Kc < 1, there are more reactants than products at equilibrium, so equilibrium position lies left
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exam q: catalytic converters on cars are less effective on short journeys compared to long journeys. They only start to become effective after the first 5 miles. Suggest why the catalytic converter is not as effective in the first 5 miles of a car journey (2)
1) the temperature is too low at the start of the journey
2) only a small proportion of reactant molecules have energy greater than the activation energy
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exam q: sodium thiosulfate reacts with hydrochloric acid according to the following equation:
Na2S2O3 (aq) + 2HCl (aq) --> S (s) + SO2 (g) + 2NaCl(aq) + H2O (l)
this reaction can be used to investigate the factors which affect the rate of reaction. The two reactants are mixed in a conical flask and the time taken for the solution to become opaque and obscure a cross beneath the flask is measured.
Why does the solution become opaque?
sulfur is produced which is a white precipitate/solid
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exam Q: the Kc of a reaction is 6.4 x 1^-3, when repeated at a higher temperature, the Kc was 2.1x10^-2, deduce with reasons, whether the forward reaction is exothermic or endothermic (2) b
1) Kc increases so higher temperatures produce more products/ increase in product concentration
2) if equilibrium position shifts right (toward products) at higher temperatures then the forward reaction is endothermic
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exam Q: chlorine reacts with aqueous sodium hydroxide to form bleach. write the equation and state the conditions for this reaction (2)
1) equation: 2NaOH + Cl2 --> NaCl + NaClO + H2O
2) conditions: cold and dilute sodium hydroxide
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exam Q
