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Chemistry > Module 3 - The Periodic Table And Periodicity > Flashcards

Module 3 - The Periodic Table And Periodicity Flashcards

(115 cards)

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1
Q

What is the magnitude of an ionisation energy influenced by??

A

Nuclear charge

Atomic radius

Shielding

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2
Q

Effect of nuclear charge

A

The greater the nuclear charge the greater the attraction on the outer electrons

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3
Q

Effect of atomic radius

A

The further the electrons is from the nucleus, the weaker the attraction

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4
Q

Effect of shielding

A

Full inner shells of electrons exert a repelling effect on outer electrons. This reduces the attraction between the nucleus and outer electrons

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5
Q

Trends in first ionisation energy

Decrease down a group

A

Nuclear charge increases

Atomic radius increases so outer electron is further from the nucleus

Shielding increases as the number of full inner shells increases

Increase in distance and shielding outweighs the increased nuclear charge

Nuclear attraction on outer electron decreases

First ionisation energy decreases

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6
Q

Trends in first ionisation energy

General increase across a period

A

Nuclear charge increases

Atomic radius decreases so outer electron is closer to the nucleus

Shielding stays the same as electrons are added to the same shell

Nuclear attraction on outer shell increases

First ionisation energy increases

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7
Q

Decrease between groups 2 and 3

A

2p orbitals have slightly more than 2s

S electrons provide slightly greater shielding of the p electron

Less energy needed to remove outer electron despite increased nuclear charge

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8
Q

Decrease between groups 5 and 6

A

Spin pairing occurs in 2p subshell of oxygen

Paired electrons in a 2p orbital of oxygen repel each other

Less energy needed to remove outer electron despite increased nuclear charge

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9
Q

What is the atomic orbital

A

Region around the nucleus that can hold up to 2 electrons with opposite spin

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10
Q

General trend in the first ionisation energies of the first 11 elements

A

Nuclear charge increases
Atomic radius decreases

Outer electron is closer to the nucleus, shielding stays the same as electrons are added to the same shell

So ionisation energy increases across a period

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11
Q

Group number

A

Number of outer shell electrons

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12
Q

Period number

A

Number of electron shells

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13
Q

Periodicity

A

The repeating trend in physical and chemical properties across periods of the period table

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14
Q

Structure of a metallic lattice

A

Cations are in fixed positions giving the shape of the metal

Delocalised electrons are mobile and can move throughout the structure

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15
Q

What is a metallic bond??

A

A strong electrostatic attraction between a lattice of positive ions and a sea of delocalised electrons

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16
Q

Electrical conductivity in giant metallic lattices

A

Delocalised electrons can move and carry charge

Electrons will move towards the positive terminal, electrons will be supplied by the negative terminal

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17
Q

Melting and boiling points in giant metallic lattices

A

High due to large amount of energy needed to overcome the strong electrostatic attractions between positive ions and delocalised electrons

Metallic bond strength depends on number of outer shell electrons to be delocalised and charge density of the metal ion

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18
Q

Solubility in giant metallic lattices

A

All insoluble

Some will react with water

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19
Q

Structure of giant covalent lattices

A

Atoms held together in a network of strong covalent bonds forming a giant covalent lattice

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20
Q

Electrical conductivity of giant covalent lattices

A

Non-conductors, no delocalised electrons all outer shell electrons are used in bonding

Exception - graphite - delocalised electrons are present which can move

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21
Q

Melting points and boiling points of a giant covalent lattice

A

High due to the large amount of energy needed to break the many strong covalent bonds

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22
Q

Solubility of giant covalent lattices

A

All insoluble

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23
Q

What is a homologous series??

A

A series of organic compounds having the same functional group but with each successive member differing

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24
Q

What is the general formula??

A

The simplest algebraic formula for a member of a homologous series

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25
What is the structural formula??
The minimal detail that shows the arrangement of atoms in a molecule
26
What is the displayed formula??
The relative positioning of atoms and the bonds between them
27
What is the skeletal formula??
The simplified organic formula showing only a carbon skeleton and associated functional groups
28
Functional group
A group of atoms responsible for the characteristic reactions of a compound
29
Formula for an alkyl group
CnH2n+1
30
What does aliphatic mean??
A compound containing carbon and hydrogen Joined in straight chains, branched chains or non-aromatic rings
31
What does alicyclic mean??
An aliphatic compound arranged in non-aromatic rings with or without side chains
32
What does aromatic mean??
A compound containing a benzene ring
33
What does saturated mean??
Single carbon carbon bonds only
34
What does unsaturated mean??
The presence of multiple carbon-carbon bonds
35
What is a structural isomer??
Compounds with the same molecular formula but different structural formula
36
Why does ionisation energy increase across a period??
Increasing nuclear charge No change in shielding Atomic radius decreases Nuclear attraction increases More energy needed to remove outer electron
37
Why does ionisation energy decrease down a group??
Atomic radius increases More full inner shells so more shielding These factors outway the increased nuclear charge Nuclear attraction decreases Less energy is needed to remove the outer electron
38
Why do successive ionisation energies increase??
Same nuclear charge Fewer electrons Electrons pulled closer to the nucleus Nuclear attraction increases More energy needed to remove the next electron
39
What is atomic orbital??
Area around the nucleus that can hold 2 electrons with opposite spin
40
What is an orbital??
A region around the nucleus that can hold up to 2 electrons with orbital spin
41
What is a shell??
Region around the nucleus which contains orbitals with the same principle quantum number
42
What is a subshell??
A group of orbitals with the same principle quantum number
43
How to detect carbonate ions are present
Add dilute nitric acid
44
Indication that carbonate ions are present when dilute nitric acid is added
Effervescence
45
How to prove gas produced is carbon dioxide??
Bubble gas through limewater which will turn cloudy
46
How to detect sulphate ions
Add nitric acid and barium nitrate solution
47
What will occur if sulphate ions are present when nitric acid and barium nitrate solution is added??
A white precipitate of barium sulphate shows the presence of sulphate ions
48
Equation of the qualitative analysis of carbonate ions
2H+ + CO32- ~~~> CO2 + H20
49
Equation of the qualitative analysis of sulphate ions
Ba2+ + SO42- ~~~> BaSO4
50
How to detect for the presence of a halide ion
Add nitric acid and then silver nitrate solution to an aqueous solution of the sample
51
After adding nitric acid and silver nitrate solution to a sample, what will indicate a halide ion is present
A white, cream or yellow precipitate will mean a halide ion is present
52
Equation of the qualitative analysis of halide ions
Ag+ + X- ~~~> AgX
53
How to determine WHICH halide ion is present
Add ammonia
54
If a Cl- halide ion is present what will be the colour of the precipitate when ammonia is added
White
55
If a Br- halide ion is present what will be the colour of the precipitate when ammonia is added
Cream
56
If a I- halide ion is present what will be the colour of the precipitate when ammonia is added
Yellow
57
What is the solubility of a Cl- ion in NH3
Soluble in dilute NH3
58
What is the solubility of a Br- ion in NH3
Soluble in concentrated NH3
59
What is the solubility of an I- ion in NH3
Insoluble in concentrated NH3
60
Test for carbonates
Add nitric acid
61
Test for sulphates
Add barium nitrate
62
Test for halides
Add silver nitrate
63
Which ion should be tested for first and why??
By testing for carbonate ions first, subsequent tests are only needed if the unknown is not a carbonate
64
When does the sulphate test not work??
When carbonate ions are present If carbonate ions are present, a false positive could be produced
65
Qualitative analysis of ammonium ions
Add sodium hydroxide solution
66
What happens when sodium hydroxide solution is added to a solution containing ammonium ions??
Ammonia gas is produced No effervescence as the ions are very soluble
67
Equation for the qualitative analysis of ammonium ions
NH4+ + OH- ~~~> NH3 + H2O
68
How to test for ammonia gas
Warm the mixture to release ammonia gas Test the gas with damp RED litmus, the alkaline gas will turn the litmus BLUE
69
When is limestone used in the blast furnace??
In the extraction of iron
70
When limestone is heated what is formed??
CaO (quicklime) and CO2
71
What happens when limestone is firstly heated then water is added??
Slaked lime is formed Ca(OH)2
72
What happens when limestone is heated then excess water is added??
Limewater is formed Ca(OH)2
73
What happens when CO2 is added to limewater??
Calcium carbonate is formed Limewater turns milky because a precipitate is formed
74
What happens when excess CO2 is added to limewater
Ca(HCO3) is formed Present in hard water
75
Why is limestone heated with clay in a kiln
To make cement
76
When cement is mixed with sand, water and crushed rock what is formed??
Concrete
77
Why is limestone used in lakes and soils
To neutralise acidity
78
What is oxidation
Loss of electrons Gain of oxygen
79
What is reduction??
Gain of electrons Loss of electrons
80
Physical properties of group 2 elements
* High melting and boiling point (higher than group 1 but lower than transition elements) * atomic radii increases upon descent down a group * increased shielding outweighs increased nuclei charge * ionisation energy decreases down the group
81
Redox reactions of group 2 elements
All group 2 elements react readily with o2, h2o and dilute acids
82
What is the trend in ionisation energies down group 2
Sum of the 1st and 2nd ionisation energies decrease down the group
83
Reactions of group 2 elements with oxygen
All burn in oxygen Reactivity increases down a group
84
Observation of magnesium reacting with oxygen
Silver solid burns with a bright white light to form a white solid
85
Equation of magnesium reacting with oxygen
Mg (s) + O2 (s) ~~~> 2MgO (s)
86
Reaction of group 2 elements with water
Reactivity increases down a group
87
Reaction of beryllium with water
No reaction
88
Reaction of magnesium with water
Reacts with steam Silver solid reacts slowly with effervescence to produce a cloudy solution Resulting solutions are alkaline Alkalinity increases down a group as the hydroxide becomes more soluble
89
Reaction of ca/sr/Ba with water
Reacts with cold water
90
Group 2 reactions with dilute acids
All react with dilute acids Reactivity increases down the group
91
Uses of group 2 compounds
Limestone Quarried for building materials Products used to neutralise acidic soil/emissions
92
Uses of Mg(OH)2 and CaCO3
Antacids Neutralise excess stomach acid and prevent acid indigestion
93
Uses of BaSO4
Barium meal Insoluble so not absorbed into the blood stream X-rays cannot pass through so enables X-rays of the digestive tract to be taken
94
Physical properties of the halogens
Low mp and bp (increases down the group due to increasing strength of induced dipole dipole interactions) Atomic radio increases down the group Solubility - non polar molecules dissolve well in non polar solvents
95
Appearance of cl2 at rtp
Pale green gas
96
Appearance of Br2 at rtp
Orange/brown liquid
97
Appearance of I2 at rtp
Silver/grey solid
98
Colour of cl2 in a non polar solvent
Pale green
99
Colour of br2 in a non polar solvent
Orange
100
Colour of i2 in a non- polar solvent
Purple/pink
101
Colour of cl2 in water
Very pale green
102
Colour of br2 in water
Orange
103
Colour of i2 in water
Yellow/brown
104
Oxidising strength in halogens
Decreases down a group
105
Displacement reactions in halogens
An element higher up in the group will displace an ion lower down from its compound
106
Reaction of chlorine with water
Chlorine disproportionates
107
Equation to show the disproportionation of Cl2
Cl2 + H2O ~>~
108
What is disproportionation
The same element is simultaneously oxidised and reduced
109
Disproportionation of chlorine with an alkaline
2NaOH + Cl2 ~~~> NaCl + NaOCl + H2O
110
Conditions of a reaction between chlorine and an alkali
Cold (below room temp) Using dilute NaOH
111
Ionic equation EXAMPLE with silver nitrate
X- + Ag+ ~~~> AgX
112
What is used to distinguish between Cl and Br/I
Use Silver nitrate
113
Uses of chlorine
Water treatment Chlorine bleach Bromine manufacture Production of Organochlorine compounds such as PVC, solvents and Pesticides
114
Pros of cl2 in water treatment
Kills bacteria Prevents spread of disease through water supply
115
Cons of cl2 in water treatment
Toxic so careful handling required May react with hydrocarbons in water forming harmful substances

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