Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

chemistry a level > periodic table and energy > Flashcards

periodic table and energy Flashcards

You may prefer our related Brainscape-certified flashcards:
AP® Chemistry flashcards Periodic Table flashcards Chemistry 101 flashcards Organic Chemistry 101 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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

what is ionisation energy?

A

measures how easily an atom loses electrons to form positive ions

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

what is the first ionisation energy equation for sodium?

A

Na (g) –> Na+(g) + e-

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

what is the general rule in successive ionisation energies?

A

no. of electrons = ionisation energies

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

every time an electron is removed… (3)

A

1) there is a greater nuclear attraction
2) atomic radius is slightly smaller
3) successive ionisation energies will increase

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

what do trends in 1st ionisation energy provide evidence for?

A

the existence of shells and sub shells

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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 by the general trend 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.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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 in comparison to the configuration of nitrogen

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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 than expected by the general trend 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 overcome

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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 in comparison to the configuration of phosphorus

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

what 3 things affect first ionisation energy? describe the trend in these across a group and down a period

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)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

why does shielding remain constant across a period?

A

the number of energy levels is the same

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
24
Q

How are elements classified into SPDF block elements?

A

classification is based on the name of orbitals which receive the last electron.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
Q

what can successive ionisation data be used for? (3)

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
26
Q

what is electron shielding?

A

when the inner shells create a barrier that blocks the attractive forces

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
27
Q

what is metallic bonding?

A

strong electrostatic attraction between cations and delocalised electrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
28
Q

what is ionisation energy measured in?

A

kJmol-1

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
29
Q

when do successive ionisation energies occur?

A

when further electrons are removed after first ionisation energy

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
30
Q

why does successive ionisation energy require more energy? (2)

A

1) as electrons are removed, the electrostatic 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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
31
Q

what is the structure of metals? (3)

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
32
Q

what should you remember when drawing a diagram of metallic structure?

A

when asked to draw a specific metal, add the charge to the cation
e.g
Ca would be 2+

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
33
Q

what are the properties of metals? (3)

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
34
Q

which 3 elements form giant covalent structures?

A

Boron, Carbon and Silicon

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
35
Q

what are the properties of giant covalent structures? (3)

A

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.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
36
Q

what is the periodic trend in melting points across period 2? (4)

A

1) increases between Li and Be as Be forms 2+ ions so there is a stronger attraction to electrons within the metallic structure
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 monoatomic ion/group 0

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
37
Q

what are the giant metallic structures in period 2?

A

Li, Be

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
38
Q

what are the giant metallic structures in period 3?

A

Na, Mg, Al

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
39
Q

what are the simple molecular structures in period 2?

A

N2, O2, F2, Ne

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
40
Q

what are the simple molecular structures in period 3?

A

P4, S8, Cl2, Ar

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
41
Q

what is a giant covalent lattice?

A

a solid giant covalent lattice is a network of atoms bonded by many strong covalent bonds

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
42
Q

what are the properties of giant covalent structures? (3)

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
43
Q

what are the allotropes of carbon?

A

diamond, graphite and graphene

44
Q

what are the features of diamond? (3)

A

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

45
Q

what are the features of graphite? (3)

A

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

46
Q

what are the features of graphene? (4)

A

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

47
Q

what causes the variation in melting point in period 2?

A

due to different bond strength and structure

48
Q

what is the melting point like for Li and Be in period 2? (4)

A

1) Li and Be have metallic bonding
2) melting point increases from Li to Be because Be has a greater positive charge of 2+ when forming ions
3) more electrons as free electrons in Be lattice
4) electrostatic forces are greater than that in lithium

49
Q

what is the melting point like for B and C in period 2? (2)

A

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

50
Q

what is the melting point like for N2 O2, F2 and Ne in period 2? (2)

A

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

51
Q

what causes the variation in melting point in period 3?

A

difference in bond strength and structure

52
Q

what is the melting point like for Na, Mg and Al in period 3? (4)

A

1) metals with metallic bonding
2) melting point increases from Na to Al due to greater positively charged ions
3) and more electrons are released as free electrons across the metals
4) electrostatic attraction between electrons and metal cations increases from Na to Al

53
Q

what is the melting point like for Si in period 3? (2)

A

1) giant covalent structure
2) strong covalent bonds which require a lot of energy to break causes high melting point

54
Q

what is the melting point like for P4, S8 and Cl2 in period 3? (3)

A

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

55
Q

what is the melting point like for Ar in period 3? (2)

A

1) noble gas, exists as individual atoms with a full outer shell of electrons
2) stable molecule so London forces are weak

56
Q

what valence shell configuration do group 2 metals have?

A

ns2

57
Q

what happens when group 2 metals react?

A

they lose 2 electrons to form 2+ ions

58
Q

what happens when group 2 elements are exposed to air?

A

tarnish in air, forming a coat of metal oxide

59
Q

what is the trend in atomic radius down group 2?

A

increases due to additional electron shells

60
Q

what is the trend in reactivity down group 2? (2)

A

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 electron easier to lose

61
Q

what is the trend in ionisation energy down group 2? (2)

A

1) decrease down the group
2) due to greater atomic radius and increased shielding which makes outer shell electron easier o lose

62
Q

how do group 2 compounds react with water? (3)
give an example

A

1) group 2 metals react with water in a redox reaction to produce a metal hydroxide and hydrogen
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

63
Q

what is the speed like for reactions of group 2 compounds and water? how can this be fixed? give an example (4)

A

1) slow with liquid water
2) reaction can be much faster with steam as it provides the reaction with extra energy
3) when steam is used the magnesium burns with a bright white flame to form hydrogen and magnesium oxide, a white powder
e.g Mg + H2O –> MgO + H2

64
Q

what is reactivity like with water down the group 2 metals?

A

1) increases
2) Ca, Sr and Ba all react with water with increasingly vigorous reactions
3) reactions all form metal hydroxide and hydrogen gas

65
Q

how do group 2 metals react with chlorine? give an example (4)

A

1) group 2 metals react with chlorine to produce metal chlorides
2) these are all white precipitates
3) as you move down the group, reactions become more vigorous as elements are more reactive
e.g
Mg + Cl2 –> MgCl2

66
Q

how do group 2 metals react with oxygen? give an example (3)

A

1) group 2 metals react with oxygen to form oxides
2) once the reaction has been initiated it is vigorous
e.g
2Mg + O2 –> 2MgO

67
Q

how do Sr and Ba react with oxygen? give an example (2)

A

1) Sr and Ba can react with excess oxygen and heat energy to form metal peroxides
e.g
Sr + O2 –> SrO2

68
Q

how do group 2 metals react with dilute acids? give examples with all 3 acids (4)

A

1) 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

69
Q

what is the solubility of group 2 hydroxides like? (3)

A

1) group 2 hydroxides increase in solubility down the group
2) more hydroxide ions released in solution
3) more alkaline solution down the group

70
Q

how do group 2 oxides react with water? (4)

A

1) form an alkaline solution of the metal hydroxide
e.g
CaO + H2O –> Ca(OH)2
2) hydroxides are only slightly soluble in water
3) once the solution is saturated any further, hydroxide ions form solid precipitate
4) more alkaline solution down the group as more OH- release

71
Q

what is the outer shell configuration for group 2?
how does this influence the most common type of reaction of group 2 elements? (2)

A

1) s2 electron configuration
2) the loss of these electrons occurs in redox reactions to form 2+ ions

72
Q

explain the trend in reactivity in terms of the first and second ionisation energies of Group 2 elements down the group (3)

A

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.

73
Q

what are some features of the halogens? (4)

A

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

74
Q

what is the trend in boiling point for the halogens? (3)

A

1) increases down the group
2) London forces get stronger due to more electrons
3) so bigger dipoles form

75
Q

what do halogens form in their solid state?

A

lattices with simple molecular structures

76
Q

what is the appearance of the halogens at room temperature? (4)

A

1) F2–> pale yellow gas
2) Cl2 –> pale green has
3) B2–> red-brown liquid
4) I2–> shiny black-grey solid

77
Q

what is the most common type of reaction for the halogens? (2)

A

1) redox
2) each halogen atom is reduced and acts as an oxidising agent (gains an electron from another element which loses it)

78
Q

what is the general half equation for the reduction of halogens in halogen displacement reactions?

A

X2(aq) + 2e- –> 2X- (aq)
(more reactive one)

79
Q

what is the general half equation for the oxidation of halides in halogen displacement reactions?

A

2X- (aq) –> X2 (aq) + 2e-
(less reactive one)

80
Q

in aqueous solutions what do Cl2, Br2 and I2 appear as? (3)

A

1) Cl2 (aq) –> pale green
2) Br2 (aq) –> orange
3) I2 (aq) –> brown

81
Q

why is cyclohexane added to halogens in aqueous solutions?

A

to make the colour more clear

82
Q

in the cyclohexane layer, what do Cl2, Br2 and I2 appear as? (3)

A

1) Cl2 (aq) –> pale green
2) Br2 (aq) –> orange
3) I2 (aq) –> violet

83
Q

what do the results of halogen displacement reactions show?

A

that the reactivity of halogens deceases down the group

84
Q

what happens during a halogen displacement reaction?

A

a solution of each halogen is added to aqueous solutions of the other halides
(halogen and halide will be different elements)

85
Q

give an example of a halogen displacement reaction and give the ionic sonf half equations

A

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-

86
Q

what is the general rule for writing ionic equations in halogen displacement reactions?

A

eliminate metal ion- only thing that remains unchanged/ spectator ion

87
Q

what is the trend in reactivity in the halogens? (4)

A

1) decreases down the group as it is harder to gain an electron
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

88
Q

what is a disproportionation reaction?

A

the same element is both oxidised and reduced in a reaction

89
Q

what is one example of a disproportionation reaction?

A

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

90
Q

what is another example of a disproportionation reaction?

A

reaction of chlorine with aqueous sodium hydroxide
Cl2 (g) + NaOH (aq) –> NaClO (aq) + NaCl (aq) + H2O (l)
(sodium hypochlorite)
Cl is both oxidised and reduced

91
Q

what is the trend in oxidising power down the halogens?

A

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

92
Q

what are the pros and cons of chlorine use in water treatment? (3)

A

pros: kills bacteria
cons: risks with hazards of toxic chlorine gas
possible risks from formation of chlorinated hydrocarbons

93
Q

what is the problem with the test for halide ions? how can you resolve this? (2)

A

1) colours are hard to tell apart
2) add dilute ammonia solution

94
Q

describe the test for halide ions

A

1) add nitric acid
2) add silver nitrate
3) results:
AgCl- –> white precipitate
AgBr- –> cream precipitate
AgI–> yellow precipitate

95
Q

describe the results of adding dilute ammonia solution after the halide test

A

Cl –> precipitate dissolves
Br –> no change/ cream precipitate
I –> no change/ yellow precipitate

96
Q

describe the results of adding dilute ammonia solution after the halide test

A

Cl –> precipitate dissolves
Br –> precipitate dissolves
I –> no change/ yellow precipitate

97
Q

describe the test for carbonates (3)

A

1) add a dilute acid (preferably nitric)
2) bubbles of CO2 will be produced
3) bubble through limewater which will turn cloudy

98
Q

what is the ionic equation for the carbonate test

A

2H + (aq) + CO3 2-(aq) –> CO2 (g) + H2O (l)

99
Q

describe the test for sulfates (3)

A

1) add nitric acid to eliminate any carbonate ions
2) add barium chloride or barium nitrate
3) will result in a white precipitate

100
Q

what is the ionic equation for the sulfate test?

A

Ba 2+ (aq) + SO4 2- –> BaSO4 (s)

101
Q

what is the general ionic equation for the halide test?

A

Ag+ (aq) + X- (aq) –> AgX (s)

102
Q

how would you identify an unknown solution? (6)

A

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.

103
Q

describe the test for ammonium ions (NH4+)

A

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)

104
Q

what is the ionic equation for the test for ammonium ions?

A

NH4 + (aq) + OH- (aq) –> NH3 (g) + H2O (l)

105
Q

why should you do the test for carbonates before the test for sulfates?

A

to eliminate any CO32- ions as they will give a positive result for the sulphate test- remove false positives

chemistry a level (6 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2024 Bold Learning Solutions. Terms and Conditions