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x A-level Chem > Module 3: Section 1 > Flashcards

Module 3: Section 1 Flashcards

1
Q

How are elements arranged in the periodic table?

A

Arranged by proton number

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2
Q

Define 1st Ionisation Energy

A

Energy needed to remove 1 mol of electrons from 1 mol of gaseous atoms.

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3
Q

Is Ionistaion energy process endothermic or exothermic?

Explain why.

A

Endothermic process as you have to put energy in to ionise an atom/molecule.

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4
Q

What 3 factors affect ionistaion energy?

Describe how these factors change

A
  1. Nuclear charge. Increases by 1 across a period. The more protons in an atom, the more positively charged it is and the stronger the attraction with the electrons.
  2. Atomic radius. Increases as you go down a group, so reduces nuclear attraction to outer electrons.
  3. Shielding. Increases down a group, reducing the attraction between the nucleus and electrons.
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5
Q

Why is there a drop in Ionisation Energy between groups 2 and 3?

A

Group 2 = s block
Group 3 = p block.
P orbital has slightly more energy than an s orbital in the same shell, so the electron is found further from the nucleus. P orbital also has additional shielding, provided by the s electrons.
These factors override the increase in nuclear charge.

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6
Q

Why is there a drop in Ionisation Energy between groups 5&6?

A

In group 5 the electron is being removed from a singly occupied orbital.
In group 6, the electron is being removed from an orbital containing 2 electrons.
The repulsion between the 2 electrons in the orbital means that the electrons are easier to remove from shared orbitals.

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7
Q

What structure do Diamond, Graphite and graphene all have?

A

They are all giant covalent lattices.

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8
Q

Diamond, graphite and graphene are all allotropes of carbon. what does that mean?

A

They are different forms of the same element in the same state.

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9
Q

What element are Diamond, Graphite and graphene all made up of?

A

Carbon

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10
Q

How can carbon atoms form giant covalent lattices such as Diamond, Graphite and graphene.

A

Each carbon atom can form 4 strong covalent bonds.

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11
Q

Describe the shape of Diamond.

A

In diamond, each carbon atom is covalently bonded to 4 other carbon atoms and are arranged in a tetrahedral shape.

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12
Q

Describe the properties of diamond, due to the strong covalent bonds.
(4)

A
  1. Very high mpt.
  2. Extremely hard.
  3. Good heat conductor - vibrations can travel through the stiff structure easily.
  4. Can’t conduct electricity - electrons are held in place.
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13
Q

Describe the shape of graphite.

A

Carbon atoms are arranged in hexagons, covalently bonded with 3 bonds each.
The 4th electron of each carbon atom is delocalised between the layers.
The sheets are bonded by weak induced dipole-dipole forces.

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14
Q

Describe the properties of graphite

4

A
  1. Weak forces between layers, so sheets can slide over eachother - graphite feels slippery.
  2. Conducts electricity - electrons are free to move and carry the charge.
  3. Very high mpt due to strong covalent bonds.
  4. Insoluble - covalent bonds are too strong to break.
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15
Q

Describe the structure of graphene.

A

Graphene is a sheet of carbon, just one atom thick. Each carbon atom has 3 covalent bonds and 1 delocalised electron.

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16
Q

Describe the properties of graphene.

(3)`

A
  1. Best known electrical conductor - delocalised electrons are free to move and without the layers they can move quickly above and below the sheet.
  2. Extremely strong - delocalised electrons strengthen the covalent bonds.
  3. A single layer of graphene is transparent and very lightweight.
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17
Q

Describe the structure and bonding in metals.

A

Electrons in the outermost shell of the atom are delocalised. This leaves a positively charged metal cation (Mg2+, Na+). These metal cations are electrostatically attracted to the delocalised negative electrons.
This is metallic bonding.
There is a lattice of metal cations and a “sea” of delocalised electrons.

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18
Q

Give 3 properties of metals.

A

Good heat conductors - electrons can pass kinetic energy to eachother.
Good electrical conductors - electrons are free to move and carry the charge.
Metals are insoluble, due to the strength of the metallic bonds.

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19
Q

How does the bonding in metals change as there are more delocalised electrons?

A

More delocalised electrons = the stronger the attraction and metallic bonding.
Results in a higher mpt.

20
Q

Describe the bonding in simple molecular structures, and give examples.

A 
Covalent bonds between the atoms are very strong.
Intermolecular forces (induced dipole-dipole forces) between molecules are very weak, giving low mpts and bpts.
More atoms = stronger induced dipole-dipole forces.
Examples = O2, CO2, I2
21
Q

Explain the trend in reactivity down group 2.

A

Down the group Ionistaion Energy decreases as there is an increased atomic radius and shielding effect.
It is easier to remove elctrons, making the element more reactive, as it can form positive 2+ ions (cations)

22
Q

What happens when group 2 elements react with water?

A

They produce a hydroxide and hydrogen.

M + 2H2O —> M(OH)2 + H2

23
Q

What happens when group 2 elements react with oxygen?

What do you see?

A

They form oxides (solid white in colour)

2M + O2 —> 2MO

24
Q

What happens when group 2 elements react with dilute acid?

A

They produce a salt and hydrogen.

M + 2HCl —> MCl2 + H2

25
Q

Why are the solution of the oxides of group 2 metals and water strongly alkaline?

A

Oxides of group 2 metals react readily with water to form metal hydroxides, which dissolve. The hydroxide ions (OH-) make the solution highly alkaline (pH 12-13)

CaO + H2O —> Ca 2+ + 2OH-

26
Q

Give 2 uses of group 2 compounds.

A
  1. Lime - Ca(OH)2 is used in agriculture to neutralise acidic soils
  2. Mg(OH)2 and CaCO3 are used in indigestion tablets.
27
Q

Give the ionic equation of neutralisation.

A

H+ + OH- —> H2O

28
Q

What is an alkali?

A

An alkali is a base that is soluble in water.

29
Q

Give the colour of the following halogens:

  1. Fluorine (F2)
  2. Chlorine (Cl2)
  3. Bromine (Br2)
  4. Iodine (I2)
A
  1. Fluorine = Yellow
  2. Chlorine = Green
  3. Bromine = Red/Brown
  4. Iodine = Grey
30
Q

Describe the structure, mpt and bpt of the halogens.

A

Halogens exist as diatomic molecules (2 atoms joined by a single covalent bond)
Mpt & Bpt increase down the group, due to the increasing strength of induced dipole-dipole forces.

31
Q

Why do halogens get less reactive down the group?

A

They gain electrons to form 1- ions.
They reduce and oxidise another substance, making them oxidising agents.
Down the group,atomic radius and shielding increases. This makes it harder for larger atoms to attract the electron needed to form an ion.

32
Q

When will a halogen displace a halide?

A

A halogen will displace a halide from solution if the halide is below it in the periodic table.
E.g. Cl2 will displace Br- and I-
Br2 will displace I-
I2 will not displace F-, Cl- or Br-

33
Q

Describe the test for halides including;

  1. method
  2. ppt colours
  3. additional test to be sure
A
  1. Add nitric acid to remove ions that may interfere with the test.
  2. Add silver nitrate solution (AgNO3)
  3. A precipitate of silver nitrate is formed.
    Ag+ + X- —> AgX
Cl- = white ppt
Br- = cream ppt
I- = yellow ppt.

To be extra sure, you can add ammonia.

Cl- = dissolves in dilute NH3
Br- = dissolves in conc NH3
I- = insoluble in conc NH3
34
Q

Define disproportionation.

A

When the same species simultaneously undergoes oxidation and reduction.

35
Q

Give the equation for making bleach.

A

2NaOH + Cl2 —> NaClO + NaCl + H2O

NaClO = sodium chlorate (I) solution - bleach

36
Q

Give 3 uses of bleach.

A

Water treatment
Bleach paper & textiles
Clean toilets.

37
Q

Give the disproportionation equation of chlorine in water

A

Cl2 + H2O —> HCl + HClO

38
Q

Give 3 benefits of adding chlorine to water.

A

Kills disease-causing microorganisms (bacteria)
Prevents reinfection down the supply
Prevents growth of algae, eliminating bad tastes and smells.

39
Q

Give 2 risks of adding chlorine to water.

A

Accidents involving chlorine could be fatal, as chlorine gas is harmful if breathed in and liquid chlorine causes severe chemical burns.
Chlorine reacts with organic compounds to form chlorinated hydrocarbons which can be carcinogenic (cancer-causing).

40
Q

Give 2 alternatives to adding chlorine to water, and describe their benefits and negatives.

A

Ozone (O3) : Stromg oxidising agent, so good at killing microorganisms. Expensive to produce & shprt half-life in water, so treatment is temporary.
Ultraviolet Light : kills microorganisms by damaging their DNA. Ineffective in cloudy water.

41
Q

Describe the test for carbonates.

A
  1. Add dilute acid (HCl)
  2. If carbonates are present, CO2 will be released and turn limewater cloudy.

CO3^2- + 2H+ —> CO2 + H2O
carbonate + acid

42
Q

Describe the test for sulfates (SO4^2-)

A
  1. Add dilute HCl
  2. Add barium chloride solution.
  3. If you get a white ppt, it is the barium sulfate, which tells you your substance is a sulfate.
43
Q

Describe the test for ammonium ions (NH4+)

A
  1. Add sodium hydroxide
  2. Warm the mixture.
  3. If ammonium ions are present, a damp piece of red litmus paper will turn blue.
44
Q

When testing for ammonium ions, why must the litmus paper be damp?

A

It must be damp so ammonia gas can dissolve and make the colour change.

45
Q

In metals, where do the delocalised electrons come from?

A

The electrons in the outer most shell of a metal atom are delocalised.
Leaves a positively charged metal cation ion (Ca2+)

46
Q

What is the formula and systematic name for bleach?

A

NaClO = Sodium Chlorate (I) solution

x A-level Chem (13 decks)

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