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a level chemistry > module 3 section 1 - The periodic table > Flashcards

module 3 section 1 - The periodic table Flashcards

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1
Q

describe how atomic radius vary across a period (4)

A
  • decrease
  • nuclear charge increases
  • no change in shielding
  • greater attraction between nuc and outer shell
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2
Q

describe how electronegativity varies across a period (4)

A
  • increases
  • nuclear charge increases
  • no change in shielding
  • greater attraction between nuc and outer shell
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3
Q

describe how first ionisation energy varies across a period (4)

A
  • increases
  • nuclear charge increases
  • no change in shielding
  • greater attraction between nuc and outer shell
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4
Q

describe the general structure and forces of group 1,2 and 3 elements in period 2 and 3 (4)

A
  • giant metallic lattices
  • lattice structures
  • delocalised e-
  • very strong electrostatic forces OA between cations and sea of delocalised e-
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5
Q

how are giant metallic latices kept in a neatly arranged lattice ? (1)

A
  • positively charged cations repel
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6
Q

describe the structure and forces of group 4 elements in period 2 and 3 (2)

A
  • giant covalent lattices
  • strong electrostatic attraction between bonded electrons and nuclei of bonded atoms
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7
Q

describe the structure and forces of group 5,6,7 and 0 elements in periods 2 and 3 (3)

A
  • smell mlcls
  • finite mlcls
  • weak IDDI between mlcls
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8
Q

state the properties of giant metallic lattices (5)

A
  • malleable
  • ductile
  • high MP + BP
  • insoluble
  • good conductors of electricity
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9
Q

explain how giant metallic lattices are malleable (2)

A
  • layers can slide over eachother w out breaking bond
  • can be bent into shape
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10
Q

explain how giat metallic lattice are ductile (1)

A

layers can be pulles into wire, 1 atom thick

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11
Q

explain why giant metallic lattices have high MP and BP (3)

A
  • alot energy needed to overcome
  • strong electrostatic of attraction]
  • between cations and sea of delocalised e-
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12
Q

explain the solubility of giant metallic lattices (2)

A
  • not soluble
  • interactions between polar mlcls and metallic lattice more likely to lead to reactions
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13
Q

explain the conductivity of giant metallic lattices (3)

A
  • good conductors of electricity as molten and solid
  • delocalised e- move and carry charge
  • throughout whole structure
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14
Q

describe the structure of the element boron (3)

A
  • semi metal ( metalloid )
  • semi conductor
  • similar structure to carbon and silicon
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15
Q

state the properties of giant covalent lattices (3)

A
  • insoluble
  • do not conduct electricity
  • high MP and BP
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16
Q

why do giant covalent lattices not conduct electricity (1)

A
  • no delocalised e- or mobile ions
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17
Q

why do giant covalent lattices have a high MP + BP (3)

A
  • very srong attractive forces
  • between nuclei and shared pair of e-
  • alot energy needed to overcome
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18
Q

state 4 examples of giant covalent lattices

A
  • diamond
  • graphite
  • graphene
  • silicon oxide
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19
Q

state the properties of simple covalent mlcls (2)

A
  • dont conduct electricity
  • low MP and BP
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20
Q

explain why simple covalent mlcls do not conduct electricity (1)

A
  • no deolcalised e- or mobile ions
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21
Q

explain why simple cov mlcls have low MP and BP (2)

A
  • weak IDDI between mlcls
  • dont need much energy to overcome
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22
Q

name 2 multiatomic mlcls and state their formula (2)

A

phosphorous - P4
sulfur - S8

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23
Q

state 2 monoatomic mlcls (2)

A
  • neon
  • argon
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24
Q

how does BP and MP vary along group5,6,7,8 elements ?

A
  • larger mlcl = more e- = higher MP + BP
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25
Q

does sulfur or chlorine have a higher MP ? (3)

A
  • sulfur
  • S8 = 16 X 8 = 128e-‘s
  • Cl2 = 17 X 2 = 34 e-
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26
Q

explain why lithium is an S-block element (1)

A
  • outermost electron is in an S - orbital
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27
Q

suggest why group 1 metals are called alkali metals give examples to justify your answer (3)

A
  • they react with water to form alkaline solutions
  • 2Na(s) + H2O (l) = 2NaOH(aq) + H2 (g)
  • 2K (s) + H2O (l) = 2KOH (aq) + H2 (g)
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28
Q

describe periodicity along a period (3)

A
  • atomic number increases
  • same number of shells
  • atomic radium decreases
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29
Q

describe periodicity down a group (2)

A
  • similar outer shell electron configuration
  • similar chem and physical properties
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30
Q

state what determins the order of elemnts in the periodic table (1)

A

atomic number

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31
Q

describe the trend in atomic radius across a period (4)

A
  • atomic radius decreases
  • as number of protons ( stomic no.) increases
  • nuclear charge increases
  • attraction between nuc + outermost e- increases
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32
Q

identify 3 elements that will have an electron configuration ending with d3 (3)

A
  • vanadium
  • niobium
  • tantalum
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33
Q

justify the position of the d - block in the periodic table (2)

A
  • 3d subshell higher in energy than 4s subshell
  • but lower in energy than 4p subshell
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34
Q

what makes hydrogen not fit group 1 ? (4)

A
  • doesnt conduct heat/ electricity
  • very low BP
  • doesnt form cations readily
  • forms mostly cov bonds
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35
Q

what makes hydrogen fit group 1 ? (1)

A
  • outer shell configuration contains 1 s-orbital electron only
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36
Q

what makes helium fit group 0 ? (5)

A
  • monoatomic
  • very low MP + BP
  • gas at RTP
  • inert
  • full outer shell configuration
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37
Q

what makes helium not fit group 0 ? (1)

A
  • outer shell configuration contains 2 s-orbital electrons only ( no P subshell)
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38
Q

describe and explain the trends in electronegativity for Li - Fr (4)

A
  • decreases
  • atomic radius increases
  • increased sheilding
  • reduced nucelar attraction to outermost e-
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39
Q

describe and explain trends in reactivity for Li - Fr (3)

A
  • increases
  • reduced nuc attraction between nuc and outermost e-
  • form cations more readily
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40
Q

suggest why hydrogen have a higher electronegativity than lithium (3)

A
  • only has 1 e-
  • no e- sheilding
  • nuc charge is stroger
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41
Q

suggest why hydrogen forms covalent bonds (1)

A
  • smaller diff in electroneg between hydrogen and other metals (<1.8)
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42
Q

suggest why hydrogen is a gas at RTP (3)

A
  • hydrogen exists as small mlcls
  • IDDI between mlcls are weak
  • little energy needed to overcome
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43
Q

give a chemical explanation for why potassium is placed immediately after argon in the periodic table (1)

A
  • potassium atoms have 1 more proton than argon
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44
Q

suggest and justify wether oxygen or nitrogen would have a lower first ionisation energy (2)

A
  • oxygen
  • repulsion from paired electrons in 2p subshell make it easier to remove an e- in an ox atom
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45
Q

write the equation for 1st ionisation energy

A

A(g) = A+(g) + e-

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46
Q

write an equation for the first ionisation energy of sodium (1)

A

Na(g) = Na+ (g) + e-

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47
Q

write and equation for the first ionisation energy of oxygen (1)

A

O (g) = O+ (g) + e-

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48
Q

describe the trend in 1st ionisation energy down a group (4)

A
  • decreases
  • increased sheilding
  • decreased nuc attraction between nuc and outermost e-
  • easier to remove an e-
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49
Q

describe the trend in first ionisation energy along a period (5)

A
  • generally increases
  • decreased atomic radius
  • same number of shells
  • increased nuc charge
  • increased nuc attraction between nuc and outermost -
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50
Q

what do large jumps in ionisation energy show (1)

A

evidence for change in shell

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51
Q

what do small jumps in ionisation energy show (1)

A

evidence for change in sub shell

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52
Q

write an equation for 2nd ionisation energy (1)

A

A+(g) = A2+ (g) + e-

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53
Q

write an equation for 3rd ionisation energy (1)

A

A2+ (g) = A3+ (g) + e-

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54
Q

explain why successive ionisation energies increase (4)

A
  • nuc attraction shared between all e- in the shell
  • e- removed but no change to nuc charge
  • slightly greater attraction between nuc and outermost e-
  • more energy needed to overcome
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55
Q

describe the trend in successive ionisation energies (2)

A
  • increase as more e- are removed
  • increase is not constant
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56
Q

what does the increase in successive ionisation energies depend on (1)

A

electron configuration

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57
Q

what 3 factors affect ionisation energy (3)

A
  • atomic radius
  • nuclear charge
  • electrons shielding
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58
Q

explain how atomic radius affects first ionisation energy (3)

A
  • larger atomic radius =
  • weaker attraction between nuc and outermost e-
  • = lower ionisation energy
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59
Q

explain how nuclear charge affects ionisation energy (4)

A
  • increased no. protons
  • greater nuc charge
  • harder to remove an e-
  • higher ionisation energy
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60
Q

explain how electron ehilding affects ionisation energy (4)

A
  • increased shielding
  • reduced attraction between nuc and outermost e-
  • easier to remove an e-
  • lower ionisation energy
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61
Q

compare the first ionisation energies of berylium and boron (4)

A
  • highest energy e- in boron is in a 2p subshell
  • in berylium it is in a 2 s subshell
  • 2p subshell in boron has higher energy than 2s subshell in berylium
  • 2p electron in boron easier to remove than 2 s e-‘s in berylium
  • 1st IE in boron is less than in berylium
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62
Q

compare the 1st IE of nitrogen and oxygen (4)

A
  • in nit + ox highest energy e-‘s are in a 2p subshell
  • in ox the paired e-‘s in one of the 2p orbitals repel
  • easier to remove an e- from oxygen atom than nit atom
  • first IE of ox = less
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63
Q

what are group 2 metals commonly used as ? (1)

A

reducing agents

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64
Q

which group 2 element would be the best reducing agent, explain why (6)

A
  • radium
  • more shells
  • more sheilding
  • wekaer nuc charge
  • weaker attraction between nuc + outermost e-
  • forms cations more readily
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65
Q

describe the trend of reactivity going down group 2 (2)

A
  • reactivity increases
  • as cations form more readily
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66
Q

state trhe key trends which occur going down group 2 (2)

A
  • 1st and 2nd IE decrease
  • reactivity increases
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67
Q

write a general equation for the reaction of group 2 metals with oxygen (2)

A

M(s) + O2 (g) = 2MO (s)

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68
Q

write a balanced symbol equation for the reaction of magnesium with oxygen (2)

A

Mg(s) + O2 (g) = 2 MgO (s)

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69
Q

write a balanced symbol equation for the reaction of berylium with oxygen (2)

A

Be(s) + O2 (g) = 2BeO (s)

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70
Q

write the general formula for the reaction of group 2 metals with water (2)

A

M (s) + 2H2O (l) = M(OH)2 (s) + H2 (g)

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71
Q

what observations can be made during the reaction of a group 2 metal with water ? (2)

A
  • effervescence
  • alkaline solution
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72
Q

describe the reaction of beryllium with water (1)

A

no reaction with water

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73
Q

describe the reaction of magnesium with water (2)

A
  • very slow reaction w water
  • vigorous reaction w steam but produces MgO
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74
Q

what 2 reagents can be reacted with magnesium to produce magnesiuym oxide ? (2)

A
  • oxygen
  • steam
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75
Q

explain with equations the differences in chemistry and observations between the reactions of magnesium with water and steam (6)

A
  • water: Mg (s) + 2H2O (l) = Mg(OH)2 (s) + H2 (g)
  • very slow reaction - slight effervescence produced
  • hydroxide produced - alkaline solution ( around 10)
  • steam: Mg(s) + H2O(g) = MgO (s) + H2 (g)
  • fast reaction - faster effervescence from H2 produced
  • MgO produced - no alkaline solution
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76
Q

write the general equation for the reaction of group 2 metals with dilute acids (2)

A

M (s) + 2HX (aq) = MX2 (aq) + H2 (g)

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77
Q

explain why observations would be different when magnesium reacted with sulfuric acid compared to barium (6)

A
  • Mg (s) + H2SO4 (aq) + MgSO4 (aq) + H2
  • colourless solution of magnesium sulfate formed
  • as magnesium sulfate in soluble
  • Ba (s) + H2SOO4 (aq) = BaSO4 (s) + H2 (g)
  • white precipitate of barium sulfate formed
  • bcs barium sulfate is insoluble
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78
Q

describe solubility of group 2 carbonates as you move down group 2 (1)

A

solubility decreases ( become insoluble)

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79
Q

describe the solubility of group 2 hydroxides in water as you move down group 2 (1)

A

increases

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80
Q

describe how group 2 oxides react with water (4)

A
  • release hydroxide ions
  • forming an alkaline solution
  • higher conc of OH- ions = more alkaline solution
  • so solution gets more alkalne down group
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81
Q

write the ionic equation for a group 2 oxide reacting with water

A

MO (s) + H2O (l) = M2+ (aq) + 2OH- (aq)

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82
Q

explain why solutions produced by group 2 metal oxides reacting with water increase in alkilinity as you move down group 2 (3)

A
  • greater solubility of metal hydroxides
  • more OH- ions released
  • group 2 hydroxides become more soluble down group
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83
Q

describe the use of Mg(OH)2 refering to its PH and solubility (3)

A
  • PH10
  • partially soluble in water
  • used to neutralise stomach acid
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84
Q

which alternative metal carbonate can be used to neutralise stomach acid ? (1)

A

CaCO3

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85
Q

describe what Ca(OH)2 is used for referring to its PH and solubility (3)

A
  • PH 11
  • resonably soluble in water
  • used in agriculture to neutralise acidic soil
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86
Q

describe the trend in solubility of group 2 sulfates down the group (1)

A
  • become less soluble down the group
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87
Q

explain the trnd in PH of each hydroxide solution down the group 2 (3)

A
  • greater solubility of group 2 metal hydroxides
  • greater conc of OH- ions released
  • increased alkalinity = higher PH
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88
Q

suggest a reason for the trend in solubility going down group 2 metal hydroxides (5)

A
  • solubility increases
  • cations increase in size
  • increased electropositivity
  • increased ionic character
  • decrease in lattice enthalpy ( strength between cations + OH anions)
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89
Q

what is thermal decomposition ? (3)

A
  • the breakdown of a compound
  • into 2 or more different substances
  • using heat
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90
Q

write the general equation for thermal decomposition (1)

A

MCO3 (s) = MO(s) + CO2 (g)

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91
Q

describe and explain when the heat energy needed for the thermal decomposition of a group 2 metal carbonate decreases (4)

A
  • greater charge density of metal ion
  • more polarised carbonate ion becomes
  • greater polarisation
  • less heat needed to break C- O bond
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92
Q

describe the trend in charge density as you go down group 2 metal ions (3)

A
  • decreases
  • charge remains 2+
  • atomic radius increases
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93
Q

what will have a higher charge density, berylium or magnesium ? (1)

A

berylium

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94
Q

what does a high charge density mean in terms of thermal stability ? (1)

A

low thermal stability

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95
Q

describe the trend in energy needed for thermal decomposition going down group two metal carbonates (6)

A
  • energy increases
  • as charge density decreases
  • thermal stability increases
  • carbonate ion becomes less polarised
  • less polarisation
  • more enrgy needed to overcome C - O bond
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96
Q

give the formula of magnesium carbonate (1)

A

MgCO3(s)

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97
Q

predict the group 2 nitrate that would require the least energy to thermally decompose

A

Be(NO3)2

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98
Q

explain why magnesium carbonate is less thermally stable that calcium carbonate (4)

A
  • Mg ion has a higher charge density than Ca ion
  • as has less sheilding but same nuc charge
  • able to polarise carbonate ion
  • breaking C - O bond
99
Q

compare the thermal stability of sodium carbonate with magnesium carbonate (2)

A
  • sodium carbonate more thermally stable
  • as sodium ion has lower charge density than Mg ion
100
Q

suggest why solutions need to be acidified before testing for sulfate ions using barium nitrate (3)

A
  • remove any carbonate ions
  • if carbonate ions were prsent they would also form a white precipitate w barium ions
  • identical; in appearence as both insoluble
101
Q

state the type of reaction which takes place when barium reacts with oxygen to form barium oxide (1)

A

redox

102
Q

name the reagent that reacts with barium oxide to form barium hydroxide, state the type of reaction it is and write out the equation (3)

A
  • water
  • BaO (s) + H2O (l) = Ba(OH)2 (aq)
  • combination
103
Q

name the reagent that reacts with magnesium nitrite to form magnesium hydroxide, state the type of reaction it is and write out the equation (3)

A
  • water
  • Mg3N2 (aq) + 6H2O (g) = 3Mg(OH)2 (s) + 2NH3 (g)
  • displacement
104
Q

descrube and explain the trend reactivity of group 2 metals with water (4)

A
  • reactivity increases down group
  • increased e- sheilding
  • no change in nuc charge
  • form cations more readily
105
Q

suggest why the reactions of Ca - Ba with sulfuric acid are slower than the reactions of Be or Mg (2)

A
  • Ca - Ba form insoluble sulfates
  • formation of insoluble sulfate layer on metal surface prevents further reactions with acid
106
Q

what is the physical colour and state of fluorine ?

A

pale yellow gas

107
Q

what is the physical colour and state of chlorine ?

A

green gas

108
Q

what is the physical colour and state of bromine ?

A

red-brown liquid

109
Q

what is the physical colour and state of iodine ?

A

black solid

110
Q

describe the physical structure of halogen molecules (3)

A
  • diatomic mlcls
  • weak IDDI between mlcls
  • simple mlclr structures
111
Q

what type of structures do halogen molecules form ? (1)

A

simple molecular structures

112
Q

describe the trend in boiling points down the halogens (5)

A
  • increase down group
  • mlcl becomes larger
  • more e-
  • stronger IDD forces between mlcls
  • more energy needed to overcome IDD forces
113
Q

what type of molecules do halogens form at RTP ? (1)

A

diatomic molecules

114
Q

state the outer electron configuration of the halogens (1)

A

s2 p5

115
Q

describe the trend in reactivity going down the group 7 halogens (5)

A
  • reactivity decreases down group
  • atomic radius increases
  • greater e- sheilding
  • less nuc attraction between nuc + outermost e-
  • forms anions less readily
116
Q

descfribe what happend during a displacment reaction (1)

A

a more reactive halogen displaces a less reactive halogen

117
Q

write an ionic equation to show chlorine displacing bromide ions

A

Cl2(aq) + 2Br- (aq) = Br2 (aq) + 2Cl- (aq)

118
Q

justify using ionic equations and oxidation numbers a halogen that could be used to reduce aqeuous sodium bromate (V) to bromine. (6)

A
  • I2 (aq) + 2BrO- 3 (aq) = Br2 (aq) + 2IO3 - (aq)
  • I went from 0 to +5 = oxidised
  • Br went from +5 to 0 = reduced
  • iodine will reduce bromate (V) at it is a stronger reducing agent
  • due to greater nuc shielding
  • so less attraction to outer e-
119
Q

state the polarity of organic layers in displacement observations

A

non polar

120
Q

state the polarity of aqueous layers in displacement observations

A

polar

121
Q

dscribe how you are able to identify between polar and non polar solvents in a displacement observation (2)

A
  • polar and non polar solvents are immiscible
  • form seperate phases
122
Q

describe how you are able to distinguish between a halogen and a halide during a displacement observation (2)

A
  • non polar halogen will dissolve in non polar organic layer
  • polar halide will dissolve in polar aqueous layer
123
Q

describe the state and appearence of chlorine in water (2)

A

pale green solution

124
Q

describe the state and appearence of bromine in water (2)

A

orange solution

125
Q

describe the state and appearence of iodine in water (2)

A

brown solution

126
Q

state the obervation reaction and equation which occurs when potassium chloride (aq) reacts with chlorine (aq) (3)

A
  • green solution
  • no reaction
127
Q

state the obervation reaction and equation which occurs when potassium chloride (aq) reacts with bromine (aq) (3)

A
  • yellow/ orange solution
  • no reaction
128
Q

state the obervation reaction and equation which occurs when potassium chloride (aq) reacts with iodine (aq) (3)

A
  • brown solution
  • no reaction
129
Q

state the obervation reaction and equation which occurs when potassium bromide (aq) reacts with chlorine (aq) (3)

A
  • yellow / orange solution
  • cl displaced Br
  • Cl2 (aq) + 2Br- (aq) = Br2 (aq) + 2Cl- (aq)
130
Q

state the obervation reaction and equation which occurs when potassium bromide (aq) reacts with bromine (aq) (3)

A
  • yellow / orange solution
  • no reaction
131
Q

state the obervation reaction and equation which occurs when potassium bromide (aq) reacts with iodine (aq) (3)

A
  • brown solution
  • no reaction
132
Q

state the obervation reaction and equation which occurs when potassium iodide (aq) reacts with chlorine (aq) (3)

A
  • brown solution
  • cl displaces I
  • Cl2 (aq) + 2I- (aq) = I2 (aq) + 2Cl- (aq)
133
Q

state the obervation reaction and equation which occurs when potassium iodide (aq) reacts with chlorine (aq) (3)

A
  • Cl2 (aq) + 2Br- (aq) = Br2 (aq) + 2Cl- (aq)
134
Q

state the obervation reaction and equation which occurs when potassium iodide (aq) reacts with bromine (aq) (3)

A
  • brown solution
  • Br displaces I
  • Br2 (aq) + 2I- (aq) = I2 (aq) + 2Br- (aq)
135
Q

state the obervation reaction and equation which occurs when potassium iodide (aq) reacts with iodine(aq) (3)

A
  • brown solution
  • no reaction
136
Q

what colour is chlorine in non polar solvents such as hexane (1)

A

pale yellow - green

137
Q

what colour is bromine in non polar solvents such as hexane (1)

A

yellow - orange

138
Q

what colour is iodine in non polar solvents such as hexane (1)

A

purple

139
Q

explain how you can use a displacement reaction to identify organic and aqueous layers (4)

A
  • use a seperating funnel
  • shake hexane w dilute aqueous halide solution
  • 2 layers form - aqueous + organic
  • organic on top
140
Q

sodium iodid is shaken with bromine, describe the clours of each layer and explain using an ionic equation what this shows (4)

A
  • organic layer: purple
  • aqueous layer: orange
  • Br2 (aq) + 2I- (aq) = I2 (aq) + 2Br- (aq)
  • iodine displaced by bromine
141
Q

sodium bromide is shaken with chlorine, describe the clours of each layer and explain using an ionic equation what this shows (4)

A
  • organic layer: orange/layer
  • aqueous layer: pale green/colourless
  • Cl2 (aq) + 2Br- (aq) = Br2 (aq) + 2Cl- (aq)
  • bromine displaced by chlorine
142
Q

sodium chloride is shaken with iodine, describe the clours of each layer and explain using an ionic equation what this shows (4)

A
  • organic layer: pale yellow - green
  • aqusous layer : brown
  • no reaction between iodine ad sodium chloride
143
Q

Solutions such as iodine and bromine can appeara similar orange colour depending on concentration explain what can be done to tell them apart (2)

A

• Add cyclohexane to shaken mixture
• non polar halogen dissolves more readily un cyclohexane than water

144
Q

What colour is chlorine in cyclohexane

A

Pale green

145
Q

What colour is bromine in cyclohexane

A

Orange

146
Q

What colour is iodine incyclohexane

A

Violet

147
Q

Which halogen is the strongest oxidising agent?

A

Fluorine

148
Q

What can caloric acid act as

A

A weak bleach

149
Q

How can you demonstrate that caloric acid acts as a weak bleach (3)

A

• Add indicator to solution of chlorine in water
• indicator first turns red
• then colour disappears because of bleaching effect

150
Q

Write the equation for the reaction of chlorine with water, what are the 2 products

A

CL2 (aq) + H2O(l) = HClO (aq) + HCl(aq)

  • chloric acid
  • hydrochloric acid
151
Q

Explain why the reaction of chlorine water is limited (1)

A
  • Because of low solubility of chlorine in water
152
Q

Describe a reaction which produces more chlorate ions , Cl o -)than is produced when chlorine reacts with water (5)a

A

-Chlorine with cold dilute aqueous sodium hydroxide -
- much more chlorine dissolves

CL2(aq) + 2NAOH(aq) = NaClO(aq) + NaCl(aq)+ H2O (l)
- disproportination reaction
- resulting solution contains a large concentration of chlorate sons from sodium chlorate

153
Q

what type of process is ionisation ? (1)

A

endothermic

154
Q

ionisation energy decreases when going down a group, what does this act as evidence for ? (2)

A
  • shells exist
  • supports Bohr model of the atom
155
Q

explain why there is a drop in ionisation energy going from group 2 to group 3 (5)

A
  • change in subshell
  • outer e- in group 3 elements in p subshell but group 2 in s subshell
  • p orbital has a slightly higher energy
  • p orbital found further from nuc
  • p orbital has additional sheilding provided by s electrons

these p and s orbitals are found in the same shell

156
Q

explain why there is a drop in ionisation energy between group 5 and 6 (4)

A
  • P orbital repulsion
  • in group 5 elements e- is removed from sngly occupied orbital
  • in group 6 elements an e- is being removed from orbital containing 2e-
  • repulsion between 2e- in orbital means e- are easier to remove
157
Q

what is an allotrope (3)

A
  • diff forms
  • of same element
  • in same state
158
Q

name 3 allotropes of carbon (3)

A
  • diamond
  • graphite
  • graphene
159
Q

how can carbon atoms is each carbon atom in diamond bonded to ? (1)

A

4

160
Q

state the shape which atoms arrange themselves into in a diamond (1)

A

tetrahedral

161
Q

state the properties of diamond (5)

A
  • v high MP
  • v hard
  • good thermal conductor
  • cant conduct electricity
  • insoluble
162
Q

describe the covalent bonds in diamond (1)

A

lots of strong cov bonds

163
Q

state the properties of graphite (6)

A
  • feels slippery
  • sheets can slide over eachother
  • conduct electricity
  • less dense than diamond
  • very high MP
  • insoluble
164
Q

describe the structure of graphite (4)

A
  • C atoms aranged in sheets
  • each C atom forms 3 cov bonds
  • 4th e- is deolcalised between sheets
  • sheets bonded togtether by weak IDDI
165
Q

explain why graphite is slippery (2)

A
  • weak IDDI between layers easily broken
  • so sheets can slide over eachother
166
Q

explain why graphite can conduct electricity (3)

A
  • delocalised e- between sheets
  • mobile
  • free to move through full structure + carry elec charge
167
Q

explain why graphite is less dense than diamond (1)

A

layers far apart

168
Q

why does graphite have a high MP ? (1)

A

strong cov bonds

169
Q

explain why graphite is insoluble (1)

A

cov bonds in sheets too strong to break

170
Q

state 3 uses of graphite (3)

A
  • dry lubricant
  • in pencils
  • strong, lightweight sports equipment
171
Q

what is graphene ? (1)

A

1 layer of graphite

172
Q

describe the structure of graphene (3)

A
  • sheet of carbon atoms joined in hexagons
  • one atom thick
  • 2D cmpnd
173
Q

state the properties of graphene (3)

A
  • best known electrical conductor
  • extremely strong
  • a single layer is transparent + v light
174
Q

why is graphene very strong ? (1)

A
  • delocalised e- strengthen cov bonds between C atoms
175
Q

state 3 uses of graphene (3)

A
  • high speed electronics
  • aircraft tehnology
  • touchscreens
176
Q

describe the MP and BP of ionically bonded mlcls

A

high

177
Q

describe the MP and BP of simple cov mlcls

A

low

178
Q

describe the MP and BP of giant cov lattices

A

high

179
Q

describe the MP and BP of giant metallic lattices

A

high

180
Q

describe the state of ionically bonded mlcls at RTP

A

solid

181
Q

describe the state of simple cov mlcls at RTP

A

sometimes solid but usually liquid or gas

182
Q

describe the state of giant cov lattices at RTP

A

solid

183
Q

describe the state of giant metallic lattices at RTP

A

solid

184
Q

do ionically bonded mlcls conduct electricity when solid ?

A

no

185
Q

do simple cov mlcls conduct electricity when solid ?

A

no

186
Q

do giant cov lattices conduct electricity when solid ?

A

no

187
Q

do giant metallic lattices conduct electricity when solid ?

A

yes

188
Q

do ionically bonded mlcls conduct electricity when liquid ?

A

yes

189
Q

do simple cov mlcls conduct electricity when liquid ?

A

no

190
Q

do giant cov lattices conduct electricity when liquid ?

A

will generally sublime

191
Q

do giant metallic lattices conduct electricity when liquid ?

A

yes

192
Q

are ionic mlcls soluble in water ?

A

yes

193
Q

are simple cov mlcls soluble in water ?

A

depends how polarised mlcl is

194
Q

are giant cov lattices soluble in water ?

A

no

195
Q

are giant metallic lattices soluble in water ?

A

no

196
Q

what products are formed when a group 2 metal reacts with water ? (2)

A
  • metal hydroxide
  • hydrogen
197
Q

what products are formed when a group 2 metal burns in oxygen ? (1)

A

solid white oxide

198
Q

write the ionic equation for the reaction of calcium oxide with water

A

CaO(s) + H2O(l) = Ca2+ (aq) + 2OH-(aq)

199
Q

describe the reaction which occurs when group 2 oxides react with water (3)

A
  • form metal hydroxides
  • whoch dissolve hydroxide ions
  • making solutions strongly alkaline
200
Q

what are group 2 compounds used for ? (1)

A

neutalise acidity

201
Q

which group 2 compoun d is used in agriculture to neutralise acidic soild ?

A

Ca(OH)2

202
Q

which group 2 comounds are used in indegestion tablets to neutralise stomach acids ? (2)

A
  • Mg(OH)2
  • CaCO3
203
Q

what type of molecules are halogen molecules ?

A

diatomic

204
Q

describe the trend in volatility moving down group 7 (1)

A
  • decreases
205
Q

describe how bleach is made (2)

A
  • mix chlorine gas with cold dilute aquesous sodium hydroxide
  • forming sodium chlorate ( bleach)
206
Q

write the equation for the formation of bleach

A

2NaOH (aq) + Cl2 (g) = NaClO(aq) + NaCl (aq) + H2O (l)

207
Q

explain how chlorine can be used to kill bacteria in water (5)

A
  • mix chlorine w water
  • undergoes disproportiniation
  • form HCL + chloric acid
  • aquesous chloric acid ionises to ake chlorate ions
  • chlroate ions kill bacteria
208
Q

what is the physical state of fluorine ?

A

gas

209
Q

what is the physical state of chlorine ?

A

gas

210
Q

what is the physical state of bromine ?

A

liquid

211
Q

what is the physical state of iodine ?

A

solid

212
Q

what are the benefits to using chlorine for water treatements (5)

A
  • kills disease causing microorganisms
  • some Cl stays in water and prevents reinfection
  • prevents growth of algae
  • eliminates bad taste + smell
  • removes discoloration cause by organic cmpnds
213
Q

what are the risks assosiated with using chlorine to treat water (3)

A
  • Cl gas irritates resp system if breathed in
  • liq Cl causes chem burns
  • water contains many organic cmpnds and chlorine reacts w these cmpnds to form chlorinated hydrocarbons = carcinogenic
214
Q

what are some alternatives to using chlorine to treating water ? (2)

A
  • ozone
  • ultraviolet light
215
Q

state some dissadvantages assosiated with using ozone to treat water (2)

A
  • £££ to produce
  • short Half Life so treatement isnt permanent
216
Q

state some dissadvantages assosiated with using Ultraviolet light to treat water (2)

A
  • innefective in cloudy water
  • water treatement not permanent
217
Q

how does UV light kill microorganisms ? (1)

A

damaging their DNA

218
Q

what is the formula for ozone ?

A

O3

219
Q

in what order should we test for ions ? (3)

A
  • test for carbonates
  • test for sulfates
  • test for halides
220
Q

explain how you can test for carbonates (1)

A
  • add dilute acid
221
Q

what will be present if a carbonate is present after adding dilute acid to your sample ? (1)

A

CO2 will be released

222
Q

write the equation for the reaction which occurs when testing for carbonates

A

CO3^2- (s) + 2H+(aq) = CO2(g) + H2O(l)

223
Q

write the equation for the reaction which takes place when dilute hydrochloric acid is added to calcium carbonate when testing for carbonates

A

CaCO3(s) + HCl(aq) = CO2(g) + H2O(l) + CaCl2(aq)

224
Q

explain how you can test for carbon dioxide (2)

A
  • bubble gas through test tube of limewater
  • CO2 turns limewater cloudy
225
Q

explain how you can test for sulfate ions (4)

A
  • add dilute HCl
  • followed by barium chloride solution
  • white ppt = barium sulfate
  • telling you a sulfate is present
226
Q

write the equation for the reaction which takes place when testing for sulfate ions

A
  • Ba2+ (aq) + SO4^2- (aq) = BaSO4 (s)
227
Q

write the equation for the reaction which takes place when testing for sulfate ions in a solution of sodium sulfate

A

Na2SO4 (aq) + BaCl2 (aq) = BaSO4 (s) + 2 NaCl (aq)

228
Q

explain how you would test for halides (3)

A
  • add nitric acid
  • then silver nitrate solution
  • if chloride, bromide or iodide present a ppt will form
229
Q

describe the ppt which will form if chloride ions are present (1)

A

white ppt

230
Q

describe the ppt which will form if bromide ions are present (1)

A

cream ppt

231
Q

describe the ppt which will form if iodide ions are present (1)

A

yellow ppt

232
Q

why wont silver fluoride give any ppt ? (1)

A

it is soluble

233
Q

how can you use ammonia to tell chloride, bromise and iodide ions apart ? (3)

A
  • AgCl dissolves in dilut NH3 + conc NH3
  • AgBr doesnt dissolve in dilute NH3 but dissolves in conc NH3
  • AgI doesnt dissolve in either dilute or conc NH3
234
Q

explain how you can test for ammonia gas (3)

A
  • ammonia gas = alkaline
  • use damp peice of red litmus paper
  • turns blue
235
Q

what is the formula for ammonia gas ? (1)

A

NH3

236
Q

what is the formula for ammonium ions (1)

A

NH4^+

237
Q

explain how you can test for ammonium ions (3)

A
  • add some sodium hydroxide to substance in a test tube
  • warm mixture
  • if ammonia given off = ammonium ions present
238
Q

write the ionic equation for the reaction which takes place when testing for ammonium ions

A

NH4^+ (aq) + OH- (aq) = NH3 (g) + H2O (l)

239
Q

write the equation for the reaction which takes place when testing for ammonium ions in a sample of ammonium chloride

A

NH4Cl (aq) + NaOH (aq) = NH3 (g) + H2O (l) + NaCl (aq)

240
Q

why does the red litmus paper need to be damp when testing for ammonia ? (2)

A
  • so ammonia gas can dissolve
  • and make colour change
241
Q

why do we add dilute acid first when testing for ions ? (1)

A
  • remove any carbonate ions present
  • bcs barium carbonate is also a white ppt
242
Q

why doe we test for sulfate ions before testing for halide ions (1)

A
  • silver sulfate would form
  • which is a white ppt
243
Q

how can we get around the issue of false positives when testing for carbonate ions, sulfate ions and halide ions ? (2)

A
  • adding dilute acid first to test samples
  • to get rid of any unwanted anions
244
Q

why do we test for halides last ? (1)

A
  • silver carbonate and silver sulfate are both white ppt

a level chemistry (9 decks)

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