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a level chemistry > module 3: periodicity > Flashcards

module 3: periodicity Flashcards

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1
Q

define periodicity

A

a regular, periodic variation of properties of elements with atomic number and position in the periodic table

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2
Q

what is the second ionisation energy

A

the energy required to remove one electron from each ion in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions

X+(g) –> X2+(g) + e-

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3
Q

what are the factors influencing ionisation energies

A
  1. nuclear charge
  2. atomic radius
  3. electron shielding
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4
Q

how does 1st ionisation energy change as you go down group 2

A

it decreases
the number of filled shells increases as you go down a group
so shielding and distance from nucleus increases, making it easier to remove the outermost electron

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5
Q

how does 1st ionisation energy change as you move across a period

A

it increases

atomic radius is reasonably constant
shielding is constant
nuclear charge increases

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6
Q

how does 1st ionisation energy change between the last element of one period and the first of the next

A

rapid decrease

atomic radius increases
shielding increases
nuclear charge increases

nuclear charge is outweighed by the other two factors

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7
Q

when do you see the largest changes in successive ionisation energies

A

the largest jumps occur between changes of subshells

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8
Q

why do giant covalent substances form

A

they form when a substance’s bonding capacity can’t be satisfied in the form of a molecule. The bonds between atoms indefinitely continue, forming a large lattice

happens with carbon+silicon because each atom can form 4 strong covalent bonds; this is why carbon has several allotropes

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9
Q

outline diamond and explain it’s properties

A

each C atom is covalently bonded to 4 others, it has a tetrahedral shape

  1. insoluble in ALL solvents
  2. can’t conduct (no delocalised e-)
  3. good thermal conductor, vibrations travel easily through the stiff lattice
  4. it is extremely hard; used in drills and saws
  5. very high melting point
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10
Q

outline graphite and explain it’s properties

A

each C is covalently bonded to 3 others in a layered hexagonal structure (120˚ bond angle)
- the 4th outer e- of each C is delocalised between the layers
- weak induced dipole-dipole forces hold the layers together

  1. the induced dipole-dipole interactions between the layers easily break, so they can side over each other making it slippery so it is used as a dry lubricant and in pencils
  2. conducts electricity; delocalised e-
  3. insoluble due to strong covalent bonds
  4. very high melting point due to strong covalent bonds
  5. used to make strong, lightweight sports equipment - the layers are quite distanced in comparison to the length of covalent bonds, so it’s less dense
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11
Q

outline graphene and explain it’s properties

A

it is a single graphite layer; it is only one atom thick, 2D

  1. best known electrical conductor - the delocalised electrons are free to move along, above and below the sheet bc of lack of layers
  2. very strong - delocalised e- strengthen covalent bonds between C atoms
  3. a single graphene layer is transparent, flexible and very light
    - possible material for touchscreens
  4. high strength, low mass and good electrical conductivity
    - high speed electronics and aircraft technology
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12
Q

outline metallic bonding and explain properties of metals

A

giant metallic lattice structure
electrons in the outer shell of metal atoms are delocalised and free to move about the metal, leaving a positive metal cation
the cations are electrostatically attracted to delocalised electrons, forming a lattice of closely packed cations in a sea of delocalised electrons

  1. atoms with more delocalised electrons have a higher melting point because bonding is stronger when there are more
  2. the size of the metal ion and the lattice structure can affect melting point, a smaller ionic radius hold delocalised electrons closer to the nucleus
  3. good thermal conductors - delocalised electrons pass kinetic energy to each other
  4. no bonds hold specific ions together so ions can slide past each other when the structure is pulled, so metals are malleable (can be hammered into sheets) and ductile (drawn into wires)
  5. good electrical conductors - delocalised electrons can move and carry a charge
  6. insoluble except in liquid metals - metallic bonds are very strong
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13
Q

what happens in group 7 displacement reactions

A

the more reactive element displaces the less reactive element in solution

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14
Q

what trends are found as you descend group 7

A
  • boiling points increase; increased number of electrons increases induced dipole-dipole interactions between molecules
  • reactivity decreases because atomic radius increases
  • electron shielding increases even tho nuclear charge increases (denis)
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15
Q

how do group 7 elements act in reactions

A

they are the most reactive non metals and act as strong oxidising agents, meaning that they gain electrons
reactivity and oxidising power decreases as you go down the group

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16
Q

what are the colours of halogens in water vs in cyclohexane (an organic solvent)

A

chlorine in water = pale-green
chlorine in cyclohexane = pale-green

bromine in water = orange
bromine in cyclohexane = orange

iodine in water = brown
iodine in cyclohexane = violet

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17
Q

how can you test for a halide ion

A

precipitation reactions
add ethanol and silver nitrate solutions
then add dilute/concentrated ammonia
precipitates form faster as you go down the list

chloride ion = white precipitate, soluble in dilute ammonia

bromide ion = cream precipitate, soluble in concentrated ammonia

iodide ion = yellow precipitate, insoluble even in concentrated ammonia

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18
Q

what is disproportionation

A

when the same element is oxidised and reduced

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19
Q

what disproportionation reactions do we need to know

A

Cl2(aq) + H2O –> HClO(aq) + HCl(aq)
chlorine is added to water to kill bacteria to make water safe to drink, hydrochloric and chloric acid are formed

Cl2(aq) + 2NaOH(aq) –> NaCl(aq) + NaClO(aq) + H2O(l)
bleach is formed when cold, dilute, aqueous sodium hydroxide and chlorine react at room temperature

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20
Q

how do you test for carbonate ions

A
  1. add a dilute acid
  2. test with limewater solution if there are bubbles
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21
Q

how do you test for sulfate ions

A
  1. add a few drops of dilute HCl, then aqueous barium chloride/nitrate to the unknown solution
  2. a white precipitate of barium sulfate will form if sulfate ions are present
    - barium nitrate prevents a false positive if a halide test is later used unlike barium chloride
  • why does this work? most sulfates are soluble but barium sulfate is very insoluble
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22
Q

how do you test for ammonium ions

A
  1. react the solution with warm, aqueous NaOH, if ammonium ions are present, they will react with OH- to form ammonia gas
  2. ammonia can then be tested for using damp red litmus paper which it will turn blue
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23
Q

what sequence of tests do you carry out if there’s only 1 unknown ion and why

A
  1. carbonate - only carbonate ions will produce bubbles with acid so it can then be ruled in/out
  2. sulfate - adding barium nitrate can also form barium carbonate, this is also white and insoluble so by doing it second we will know that it isn’t barium carbonate
  3. halide - silver sulfate and carbonate both also form as precipitates
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24
Q

what sequence of tests for ions do you do if you know there is more than 1 unknown ion

A
  1. carbonate - if it bubbles, continue adding acid until bubbling stops to ensure removal of carbonate ions
  2. sulfate - add an excess of barium nitrate, any sulfate ions will precipitate as barium sulfate which can be filtered out
  3. halide - add silver nitrate to the remaining solution, any formed precipitate must include a halide, adding ammonia can confirm which one
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25
how does reactivity change down group 2
it increases because atomic radius and electron shielding increases, this reduces nuclear attraction and makes it easier to lose an outer electron
26
how do group 2 elements react with water
react more as you move down the group forms metal hydroxide and hydrogen Be - no reaction Mg - ribbon, reacts Very slowly, a few bubbles and floats to surface Ca - steadily reacts Sr - reacts fairly quickly Ba - reacts rapidly
27
how do group 2 elements react with oxygen
they burn in oxygen to form solid white oxides
28
what is an exothermic reaction
a chemical system released heat energy to its surroundings delta H is negative increase in temp if surroundings energy loss from the system is balanced by energy gain in surroundings
29
what is an endothermic reaction
when a chemical system takes in heat energy from it’s surroundings delta H is positive there’s a decrease in temperature of surroundings energy gain from system is balanced by energy gain in surroundings
30
exothermic reaction
31
endothermic reaction
32
what are the conditions for standard enthalpy change
standard pressure - 100kPa standard temp- 298K standard conc. - 1mol/dm^3 standard state - state under these conditions
33
what is the standard enthalpy change of reaction
enthalpy change that accompanies a reaction in the molar quantities shown in a chemical equation under standard conditions with all reactants and products in their standard states
34
standard enthalpy change of formation
enthalpy change when one mole of a compound is formed from its elements under standard conditions with all reactants and products in their standard states all elements have an enthalpy change of formation of 0kJ/mol
35
define standard enthalpy change of combustion
enthalpy change when one mole of a substance reacts completely with oxygen under standard conditions with all reactants and products in their standard states
36
standard enthalpy change of neutralisation
energy change accompanying the reaction of an acid by a base to form 1 mole if H2O (l), under standard conditions with all reactants and products in their standard states the value is always the same for all neutralisation reactions
37
what is average bond enthalpy
the energy required to break one mole of a specified bond type in a gaseous molecule as this energy varies depending on the environment of the bond being broke, this is taken into account in the calculation
38
how does pressure influence rate of reaction
higher pressures compress gases into a smaller volume, thus increasing concentration so that collisions are at a higher frequency
39
how does surface area impact rate of reaction
grinding a solid increases the SA/V ratio, so the frequency of collisions increases as there are more surfaces for collisions to occur
40
how do catalysts affect reaction rate
they increase it without being used up they may react to form an intermediate but are then regenerated they lower activation energy through providing an alternate reaction pathway
41
describe the Boltzmann distribution curve and how it changes with temperature and catalysts
line starts at 0 bc no molecules have 0 energy line never again meets the x axis bc there's no max energy the area under the curve = total no. of particles
42
define homogeneous catalysts
the catalysts and reactants are in the same physical state, usually both aqueous
43
define heterogeneous catalysts
the catalyst and reactants are different physical states, usually where reactants are gases but the catalyst is a solid
44
what are economic and sustainability advantages of catalysts
- using a catalyst can mean lower temps and pressures can be used, saving energy costs - saving energy costs means less us of fossil fuels/electricity which means less CO2 emissions - catalysts can enable reactions to be used which have better atom economies
45
outline Hess’ Law
arrows point down in combustion
46
what equation is used in calculating bond enthalpy
enthalpy of reaction (delta H) = bonds broken - bonds formed
47
how many electrons can fill each of the 1st 4 shells
1st - 2 2nd - 8 3rd - 18 4th - 32
48
are combustion values negative or positive
negative
49
are formation values positive or negative
they can be either
50
what is a dynamic equilibrium
a reversible reaction in which the rates of the forwards and backwards reaction are the same and the concentrations of reactants and products remain constant
51
outline and explain Le Chatelier's principle
The position of the equilibrium of a system changes to minimise the effect of any imposed change in conditions. increase in temp - shifts in favour of endo route to decrease temp and vise versa (if deltaH is negative then the forwards reaction is exothermic) conc - if reactant conc. increases, the equilibrium shifts to the right to rid the extra reactant by making more product and vice versa with increased product conc. pressure - increasing pressure shifts the equilibrium to the side with less moles to reduce the pressure and vice versa
52
what is the effect of a catalyst on a reversible reaction
- increase the rate of reaching equilibrium and the overall rate, but not the amount of product formed - do not shift the equilibrium in either direction - economically viable
53
why is compromise needed with industrial reversible reactions
high temps and pressures use lots of ENERGY low temps and pressures are SLOW
54
how does rate change with the forwards and reverse reactions over time
forward slows down over time because reactants are used up and as more products are formed, the reverse reaction speeds up these rates will eventually be identical; this is dynamic equilibrium
55
how can you calculate Kc
the [] indicates that concentrations are in moldm-3 units should be calculated as such products over reactants
56
how will Kc change with temperature
if the change increases product conc, Kc increases if the change increases reactant conc, Kc decreases
57
what is homogenous equilibria
when the species involved all have the same states/phases
58
what is heterogenous equilibria
when species involved have different states/phases
59
what method can you use in calculating Kc when given you are not given all/the equilibrium values
ICE Initial Change Equilibrium
60
how would you describe the solid giant covalent lattices shown by carbon and silicon
networks of atoms bonded by strong covalent bonds
61
how does oxidising power of halogens change going down the group
it decreases
62
how does reducing power of halide ions change as you move down the group
it increases
63
what are some sources of error in a calorimetry experiment
evaporation of the fuel heat loss incomplete combustion inadequate water stirring
64
what is a thiosulfate ion
S2O3^2-
65
what does a directly proportional graph look like
a straight line graph that goes through the origin
66
how can you calculate Kc of the backwards/forwards reaction when given Kc for the opposite reaction
do 1/Kc works because the two Kc’s are reciprocals of each other
67
evaluate the use of chlorine in water treatment
+ it kills bacteria and is used to treat drinking water and in swimming pools this prevents disease eg. cholera -- can react with organic hydrocarbons to make chlorinated hydrocarbons which have been linked to suppressed immune systems and cancer -- chlorine gas is toxic
68
how could water be purified, without chlorine
physically with filtration or distillation chemically by UV light
69
how do group 2 oxides react with water and how does this change down the group
they react with water to form the metal ion and OH- the solutions get more alkaline as you go down the group because of increasing solubility
70
how can certain group 2 compounds be used and what are the equations
some group 2 oxides, hydroxides and carbonates can be used as bases eg. Ca(OH)2 can neutralise acidic soils in agriculture Ca(OH)2 (s) + 2H+ (aq) → Ca2+ (aq) + 2H2O (l) eg. Mg(OH)2 and CaCO3 act as antacids for stomach indigestion Mg(OH)2 (s) + 2HCl (aq) → MgCl2 (aq) + 2H2O (l)
71
how do you calculate k in a rate equation
sub in and divide UNITS!

a level chemistry (15 decks)

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