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Chemistry OCR AS Level > The Periodic Table > Flashcards

The Periodic Table Flashcards

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1
Q

How are elements in the periodic table arranged?

A

Elements in the periodic table are arranged in order of increasing atomic number.

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2
Q

What is a period?

A
  • The name of each horizontal row of elements in the periodic table.
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3
Q

What is periodicity?

A
  • The elements in a period show trends (gradual changes) in properties across the period.
  • SO…..Periodicity is this repeating pattern of trends (both physical and chemical) across each period.
  • e.g. the trend across period 1 is the same as the trend across period 2, 3, 4 etc.
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4
Q

Examples of trends?

A

1st ionisation energy (Physical property)
Electronegativity (Chemical property)
Boiling point and Melting Pont (Physical property)

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5
Q

What is a group?

A

The name of each vertical column of elements in the periodic table.

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6
Q

Why do elements in a group have similar properties?

A

The elements in a group show similar chemical properties because 1) they have same number of electrons in their outershells and these occupy 2) the same type of orbitals (similar electron configurations).

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7
Q

Why are trends similar between different periods and similar between the different groups?

A

The repeating pattern of similarity is caused by the underlying repeating pattern of electron configuration.

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8
Q

How can trends down a group affect periodic trends:

A

Moving across a period element changes from a metal to a non-metal.
Moving down the periods, this change (metal to non-metal or vice versa) moves further to the right.

e.g. at top of group 14 is carbon, a non-metal, whereas at the bottom, is tin and lead, which is a metal.

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9
Q

There are three different types of elements on the periodic table, what are they and how are they different?

A

Metals, non-metals and metalloids/semi-metals.

Metalloids have properties between those of a metal and a non-metal. Examples: boron, silicon, germanium.

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10
Q

As you know, the larger the value of n, the further the shell is from the nucleus and the higher the energy level.
Within a shell, the sub-shell energies increase in the order of? State the exception.

A

s sub-shell < p sub-shell < d sub-shell < f sub-shell

However, the empty 4s orbitals, are at a lower energy level than the empty 3d orbitals. This means 4s fills before 3d.

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11
Q

4s fills before 3d, what about when these orbitals are emptied?

A

4s fills before 3d because the empty 4s orbitals are at a slightly lower energy level than the empty 3d orbitals. However, when there are electrons in the 3d orbitals, the 4s orbitals are now at a higher energy level than the 3d orbitals so when ionization occurs electrons are lost from the 4s orbital before they’re lost from the 3d orbitals.

SO 4s fills first and is emptied first.

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12
Q

The periodic table is divided into sub-shells which…

Draw the periodic table and label all the sub-shells.

A

…corresponds to the position of the atom’s outermost electrons.

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13
Q

How to use the periodic table to determine electron configuration. - use oxygen as an example.

A

Period of oxygen: 2
Orbital section of the period: p-orbital block
Position of element in the sub-shell block: 4th element in 2p block.
This means its 4 outermost electrons of oxygen are found in the 2p sub-shell, written as 2p⁴
So electron configuration of oxygen is 1s²2s²2p⁴

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14
Q

How can electron configuration be shortened?

Write the shortened electron configuration Li, Na, P and O

A 
By basing the inner shell electron configuration on the noble gas that comes before the element in the periodic table. 
Li: 1s²2s¹ or [He]2s¹
Na: 1s²2s²2p⁶3s¹ or [Ne]3s¹
P: 1s²2s²2p⁶3s²3p³ or [Ne]3s²3p³
O: 1s²2s²2p⁴ or [He]2s²2p⁴
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15
Q

Define ionisation?

A

Ionisation is when atoms lose or gain electrons to form positive or negative ions.

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16
Q

What is plasma?

A

Plasma refers to the mixture of positive and negative ions.

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17
Q

What is ionisation energy?

A

The energy needed to remove electrons and form positive ions.

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18
Q

Defien first ionisation energy (of an element)

A

The first ionisation energy is the energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions.

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19
Q

What is the general equation for first ionisation?

A

X(g) → X⁺(g) + e⁻
e.g. Na(g) → Na⁺(g) + e⁻
NOTE THAT THE ATOM AND IONS IN THE EQUATION ARE GASEOUS.

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20
Q

What state is the first (as well as second and third etc.) ionisation energy calculated in?

A

In the gaseous state. Both the atoms and ions must be gaseous.

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21
Q

What is nuclear attraction?

A

Negative electrons are held in their shells by their attraction to the positive nucleus - this is known as nuclear attraction.

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22
Q

How is this nuclear attraction broken? I.e. how are electrons removed?

A

By ionisation, where energy is supplied to the electron to overcome this nuclear attraction (no longer bound to the atom), forming a positive ion.

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23
Q

Which electrons are removed from the atom first and why?

A

Electrons in the outer shell are removed first as they are the furthest away from the nucleus so experience the least nuclear attraction and thus require the least ionisation energy.

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24
Q

So in conclusion, what is ionisation energy (energy required to remove an electron) dependant on?

A

The ionisation energy is dependant on the nuclear attraction of the electron. The greater the nuclear attraction, the greater the ionisation energy.

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25
Q

What is the nuclear attraction of an electron dependant on?

This is also what the ionisation energy is ultimately dependant on.

A

1) Atomic radius/distance from the nucleus
2) Nuclear charge
3) Electron shielding/screening

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26
Q

How does the atomic radius affect the nuclear attraction and why?

A

The larger the atomic radius, the smaller the nuclear attraction experienced by the outermost electrons. This is because the positive charge of the nucleus is further away from the outermost electrons.

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27
Q

How does the nuclear charge affect the nuclear attraction and why?

A

The greater the nuclear charge, i.e. the greater the number of protons, the greater the nuclear attraction experienced by the outershell electrons.

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28
Q

How does electron shielding affect the nuclear attraction.

A

Inner shells of electrons repel the outershell electrons because they are all negatively charged.
This repelling effect is called electron shielding/screening.
The more inner shells of electrons there are, the greater the electron shielding and thus the smaller the nuclear attraction experienced by the outershell electrons.

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29
Q

What is the successive ionisation energies?

A

Successive ionisation energies are a measure of the amount of energy required to remove each electron in turn.

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30
Q

What is second ionisation energy (of an element)?

A

The second ionisation energy is the energy required to remove one electron from each ion in one mole of gaseous +1 ions to form one mole of gaseous 2+ ions.

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31
Q

What is the general equation for second ionisation?

A

X⁺(g) → X²⁺(g) + e⁻
e.g. Na⁺(g) → Na²⁺(g) + e⁻
NOTE THAT THE ATOM AND IONS IN THE EQUATION ARE GASEOUS.
ALSO NOTE THAT THE CHARGE ON THE ION PRODUCED WILL TELL YOU WHICH SUCCESSIVE IONISATION HAS OCCURRED.

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32
Q

What is the general equation for third ionisation?

A

X²⁺(g) → X³⁺(g) + e⁻
e.g. Na²⁺(g) → Na³⁺(g) + e⁻
NOTE THAT THE ATOM AND IONS IN THE EQUATION ARE GASEOUS.

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33
Q

What is the trend in successive ionisation energy?

A

Each successive ionisation energy is higher than the one before.

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34
Q

Why is each successive energy higher than the one before?

A
  • This is because as each electron is removed, there is less repulsion between the remaining electrons (in the shell), so as a result, the remaining electrons will be drawn in slightly closer to the nucleus. A 2+ ion will have small radius than a 1+ ion.
  • The positive nuclear charge will outweigh the negative charge every time an electron is removed. Proton to electron ratio is greater in a 2+ ion than in a 1+ ion.
  • So each time an electron is removed (successively), the nuclear attraction increases. More energy is needed to remove each successive electron.
  • When the electrons in the outer shell have all been removed, electrons from the next shell are removed if more ionisation occurs. This takes more energy than an electron being removed from the same shell, because of the closer proximity of the nucleus (i.e much smaller ionic radius).
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35
Q

When do we use ‘ionic radius’ instead of ‘atomic radius’?

A

When we talk about the first ionisation energy the distance of the electron from the nucleus is effectively the atomic radius, however after the first electron is removed, the distance between the nucleus and the next electron has decreased. It is no longer the same as the atomic radius and is referred to as the “ionic radius” or “distance from the nucleus”

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36
Q

What is the trend in first ionisation energy across the period?

Same trend for second and third ionisation energy

A
  • As you go across each period:
  • Nuclear charge ↑ (number of protons increases) so greater attraction for outershell electrons
  • Atomic/ionic radius ↓ - Electrons are being added to the same energy level = electrons have greater attraction for positive nucleus = so are drawn inwards slightly.
  • Electron shielding ↔ does not change, because the number of inner electron shells remains the same across the period.
  • Therefore, nuclear attraction ↑
  • So going across the period, more energy required to overcome this attraction so first ionisation energy ↑
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37
Q

Where does first ionisation energy drop across the period?

A
  • There is a decrease from Group 2 to 13, and group 15 to 16.
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38
Q

What is the reason for the drop in first ionisation energy from group 2 to group 13?

A
  • The outermost electron in group 13 elements is in a p-orbital whereas the outermost electron in group 2 elements is in an s-orbital.
  • P-orbitals are at a slightly higher energy level than s-orbitals, so are further away from the nucleus. Thus the outermost electron in group 13 experiences less nuclear attraction, so less energy is required to remove it and a lower first ionisation energy is observed
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39
Q

What is the reason for the drop in first ionisation energy from group 15 to group 16?

A
  • Unlike in group 13 to 15, where each of the occupied p-orbitals contain a single electron, but in group 16, the outermost electron is spin-paired in the px orbital (the first orbital of the p orbitals).
  • Electrons that are spin-paired experience some repulsion between them and therefore the outermost electron in group 16 is requires less energy to remove, and a slightly lower ionisation energy is observed.
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40
Q

What is the trend in first ionisation energy down the group?

CAN THE TREND ACROSS THE PERIOD AND DOWN THE GROUP IN FIRST IONISATION ENERGY BE APPLIED TO SECOND IONISATION ENERGY?

A

As you go down the group:

  • Atomic/ionic radius ↑ - Number of electron shells increases, so the distance of the outershell electrons from the nucleus increases.
  • Electron shielding ↑ - due to more inner shells.
  • Nuclear charge also ↑ (number of protons increases) but this is not the determining factor in this case.
  • Nuclear attraction ↓
  • So as you go down the group, less energy is required to overcome this attraction so first ionisation energy ↓
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41
Q

When moving doen the group, nuclear charge increases however, nuclear attraction does not. Why not?

A

The effect of increasing nuclear charge is outweighed by the effect of the increasing distance from the nucleus (atomic radius) and electron shielding.

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42
Q

Define a metallic bond.

A

A metallic bond is the electrostatic attraction between the fixed positive metal ions (cations) and the delocalised electrons that are free to move throughout the giant metallic lattice.

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43
Q

Structure of a metal?

A

3D Giant metallic lattice held together by metallic bonds which act in all directions and A metallic bond is the electrostatic force of attraction between the fixed positive metal ions (cations) and the delocalised electrons that are free to move throughout the giant metallic lattice.
These delocalised electrons are shared between all the metal ions.
Over the whole structure the charges must balance.

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44
Q

Where do the delocalised electrons in the giant metallic lattice come from?

A

The formation of positive metal ions, is due to the outer-shell electrons becoming delocalised. These delocalised electrons are shared between all the metal ions.

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45
Q

Once an electron becomes delocalised can you tell which particular positive ion it originated from?

A

No

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46
Q

Properties of metals due to their giant metallic lattices?

A

High melting and boiling points
Good electrical conductivity
Malleable and Ductile
Insoluble

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47
Q

Why do metals have high metling and boiling points?

A
  • A large amount of energy is required to 1) overcome the strong and large number of metallic bonds in the giant metallic lattice and to 2) dislodge the metal ions from their rigid fixed positions in the lattice.
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48
Q

What else affects the melting point and boiling point of giant metallic lattices?

A

The number of delocalised electrons per ion affects the melting and boiling points. The more delocalised electrons there are per ion, the stronger metallic bonds which result in higher melting and boiling poitns. For example, Mg²⁺ has 2 delocalised electrons per ion so has a higher melting and boiling point than Na⁺ which has only one delocalised electron per ion.

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49
Q

Why do metals conduct electricity.

A

Metals can conduct electricity because the delocalised electrons are free to move throughout the giant metallic lattice, carrying charge.
Metals can conduct electricity as both a solid and a liquid.

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50
Q

Define ductile

Define Malleable

A

Ductile means it can be stretched/drawn out (into wires)

Malleable means it can be hammered into different shapes.

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51
Q

Why are metals malleable and ductile.

A

No metallic bonds hold specific ions together, and because the positive ions are in layers, they can slide over each other easily, when a force is applied.

52
Q

One difference between covalent bonds and metallic bonds

A

Covalent bonds involve localised electrons, metallic bonds involve delocalised electrons.

53
Q

How do the following molecules exist?

Nitrogen, Oxygen, Fluorine, Neon, Phosphorus, Sulfur, Chlorine, Argon

A 
N₂
O₂
F₂
Ne
P₄
S₈
Cl₂
Ar
54
Q

Trend in melting points across period 2 and 3

A

From G1-14, melting point steadily increase, as the elements have giant structures - giant metallic lattices held together by strong metallic bonds and giant covalent lattices held together by strong covalent bonds which requires large amounts of energy to break. If giant metallic, for each successive group, nuclear charge ↑ no of electrons in the outer shell ↑ - so more delocalised electrons per ion, so stronger metallic bonds which results higher mp as more energy is required to break these bonds. If giant covalent, each successive group has more electrons with which to form covalent bonds, so stronger covalent bonds, thus higher mp as more energy is required to break these bonds
Between G14-15, there is a sharp decrease in melting point. This is because the elements in group 15 have simple molecular lattices, whereas the elements in group 14 have giant covalent lattices. SML have weak intermolecular forces, London forces, between them that require less energy to break than the covalent bonds in group 14 elements.
Between G15-18, melting points remain quite low as all these elements have SML

55
Q

Which elements in period 2 and 3 form metallic lattices?
Which elements in period 2 and 3 form covalent lattices?
Which elements in period 2 and 3 form simple molecular lattices?

A

Li, Be, Na, Mg, Al
B, C, Si
N₂, O₂, F₂, Ne, P₄, S₈, Cl₂, Ar

56
Q

What are teh Group 2 elements?

A

The alkaline earth metals.

57
Q

When group 2 elements react (to form ions) what happens?

A

They lose two electrons to obtain a noble gas configuration.

58
Q

Electronic configuration of Group 2 elements?

A

They have their 2 highest energy electrons in an s sub-shell and these are the two electrons in their outer shell.

59
Q

Physical properties of Group 2 metals?

A 
Reasonably high melting and boiling points.
Light metals with low densities
Form colourless (white) compounds.
60
Q

Trend in reactivity moving down group 2?

A

As you go down the group:

  • Atomic radius increases as the number of electron shells increases so the distance between the outer-shell electrons and the nucleus increases
  • Electron shielding increases as the number of inner electron shells increases.
  • Nuclear charge also increases but its not the determining factor here.
  • Therefore, moving down the group the nuclear attraction decreases so less energy required to remove the two electrons so first and second ionisation energy decreases, thus reactivity increases.
61
Q

What type of agent are group 2 metals?

A

Reducing agent. During a reaction, group 2 metals are oxidised from a state of 0 to +2, in the process of reducing another element.

62
Q

How do you write the ionic equations of group 2 metals in a reaction?

A

M → M²⁺ + 2e⁻

0 → +2 (oxidised as oxidation number increases)

63
Q

Describe the observations of the reaction between group 2 metals and oxygen?

A
  • Reacts vigorously with oxygen - gets more vigorous as you go down the group.
  • They burn in oxygen to form solid white ionic oxides only with the general formula MO
64
Q

What type of reaction is the reaction between group 2 metals and oxygen? Explain using the reaction betwwen calcium and oxygen.

A

A redox reaction which is a reaction where both reduction and oxidation take place:
2Ca(s) + O₂(g) → 2CaO(s)
0 → +2 (Ca oxidised as ox no. increases)
0 → -2 (O reduced as ox no. decreases)

65
Q

General equation for the reaction of group 2 metals with oxygen?

A

2M(s) + O₂(g) → 2MO(s)

METAL OXIDES ATE INSOLUBLE SO ARE IN SOLID STATE NOT AQUEOUS.

66
Q

Describe the observations of the reaction between group 2 metals and water?

A
  • Reacts with all group 2 metals except beryllium.
  • Mg reacts very slowly, but reactions gets more vigorous as you go down the group
  • Fizzing is observed as hydrogen gas is released
  • Colour change also seen when universal indicator is added.
  • This is because the metal hydroxides that form, alongside hydrogen gas, dissolve in water to reelase OH⁻ ions, this makes the solution strongly alkaline with a PH of about 10-12.
  • The metal hydroxides formed have the general formula M(OH)₂.
67
Q

What type of reaction is the reaction between group 2 metals andwater? Explain using the reaction betwwen calcium and water.

A

A redox reaction which is a reaction where both reduction and oxidation take place:
Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g)
0 → +2 (Ca oxidised as ox no. increases)
+1 → 0 (H reduced as ox no. decreases)

68
Q

General equation for the reaction of group 2 metals with water?

A

M(s) + 2H₂O(l) → M(OH)₂(aq) + H₂(g)

69
Q

Describe the observations of the reaction between group 2 metals and dilute acids?

A
  • More vigorous as you go down the group
  • Fizzing is observed as hyrogen gas is released, alongside the formation of a salt.
  • Different acids produce different salts.
70
Q

What type of reaction is the reaction between group 2 metals and dilute acid? Explain using the reaction betwwen calcium and dilute acid (HCl).

A

A redox reaction which is a reaction where both reduction and oxidation take place:
Ca(s) + 2HCl(aq) → CaCl₂(aq) + H₂(g)
0 → +2 (Ca oxidised as ox no. increases)
+1 → 0 (H reduced as ox no. decreases)

71
Q

General equation for the reaction of group 2 metals with

1) HCl
2) H₂SO₄
3) HNO₃

A

1) M(s) + 2HCl(aq) → MCl₂(aq) + H₂(g)
2) M(s) + H₂SO₄(aq) → MSO₄(aq) + H₂(g)
3) M(s) + 2HNO₃(aq) → M(NO₃)₂(aq) + H₂(g)

72
Q

How else can a group 2 metal hydroxide be formed apart from directly adding the group 2 metal to water?
Give the general equation.

A

Group 2 metal oxides readily react with water to form metal hydroxides, which dissolve in water to release OH⁻ ions forming a strongly alkaline solution with a PH between 10 to 12.
MO(s) + H₂O(l) → M(OH)₂(aq)

73
Q

State the trend in Solubility of Group 2 metal hydroxides?

A

As you go down group 2, the solubility of its metal hydroxides increases.

74
Q

What does a more soluble metal hydroxide mean?

A

The more soluble a metal hydroxide, the more OH⁻ it will release when dissolved, so a more alkaline solution with a higher PH will be produced.
So, increased solubility = increased alkalinity

75
Q

Describe the trend in Solubility of Group 2 metal hydroxides?

  • Compare Ba(OH)₂ with Mg(OH)₂
  • Also mention beryllium oxide.
A
  • Beryllium is at the top of Group 2; beryllium oxide does not dissolve in water.
  • Magnesium forms Mg(OH)₂, which is only slightly soluble in water, so the resulting solution is dilute with a comparitively low OH⁻ concentration. Produces a less alkaline solution.
  • Ba(OH)₂, is much more soluble in water than Mg(OH)₂, so has a higher OH⁻ concentration, so the resulting solution is more alkaline than a solution of Mg(OH)₂.
76
Q

What happens when carbonates, oxides, or hydroxides react with with dilute acid?
Write out the reaction between calcium hydroxide and hydrochloric acid. Also write out the reaction between Magnesium oxide and hydrochloric acid.

A

They will form salt and water (and sometimes CO₂).
MgO(s) + 2HCl(aq) → MgCl₂(aq) + H₂O(l)
Ca(OH)₂(s) + 2HCl(aq) → CaCl₂(aq) + 2H₂O(l)

Note how hydroxides are solid when it is a reactant, but when produced it is always aqueous. Oxides are always solid.

77
Q

Uses of Group 2 compounds (carbonates, hydroxides, and oxides)?

A

Neutralising acidic soils
Indigestion remedies
Building and construction uses

78
Q

How are Group 2 compounds used in neutralising acidic soils?

A

If soil is too acidic, adding Ca(OH)₂ will neutralise it, as it is an alkali. However adding too much of it will make the soil alkaline.

79
Q

How are Group 2 compounds used in indigestion remedies? Give equation of any reactions taking place.

A

Indigestion is the build up of too much stomach acid (HCl acid). Antacids such as Mg(OH)₂ (milk of magnesia) or CaCO₃ are used to neutralise the excess acid, producing salt and water: Mg(OH)₂(s) + 2HCl(aq) → MgCl₂(aq) + 2H₂O(l)

80
Q

How are Group 2 compounds used in building and construction?

A

CaCO₃ is an extremely useful building material - present in limestone and marble. Used in the manufacture of steel and glass also

81
Q

Drawback of using Group 2 carbonates as building materials?

A

They react readily with acids. Most rainwater has an acidic PH, which leads to gradual erosion of limestone/marble objects such as buildings and statues.

82
Q

Name all halogens going down group 7.

A

Fluorine, Chlorine, Bromine, Iodine, Astatine

83
Q

Properties of halogens?

A

Low melting and boiling points

Exist as diatomic molecules

84
Q

Electron configuration of halogens?

A

They have their 5 highest-energy electrons in an p sub-shell out of a total 7 electrons in their outer shell. The other 2 electrons are in an s sub-shell.

85
Q

What state are each of the halogens at room temperature?

A 
Fluorine - Gas
Chlorine - Gas
Bromine - Liquid
Iodine - Solid
Astatine - Solid
86
Q

Describe the trend in boiling points down Group 7 Halogens?

A

As you go down the group, boiling point increases.
This is because each successive element has an extra shell of electron (more electrons in the atom). As a result, stronger London forces occur between the molecules, so more energy is required to break these intermolecular forces.

87
Q

State and explain the reactivity and electronegativity of halogens?

A

Halogens are highly reactice and highly electronegative. This is because they are strong oxidising agents (attracting electrons of other atoms, oxidising them).
During the formation of ionic compounds, halogen atoms gain an electron to have a 1- charge, obtaining a noble gas configuration

88
Q

Describe the trend in reactivity and oxidising power down the halogens.

A

As you go down the group:

  • Atomic radius ↑ - Number of electron shells increases, so the distance of the outershell electrons from the nucleus increases.
  • Electron shielding ↑ - because the number of inner electron shells also increases
  • Nuclear charge also ↑ (number of protons increases) but this is not the determining factor in this case.
  • Nuclear attraction ↓
  • Thus moving down the group, the ability of the halogens to gain an electron i.e by oxidising another element ↓ so reactivity decreases ↓.
89
Q

When does a redox reaction occur between halogens?

A

Redox reactions occur between aqueous solutions of halide ions (e.g. Cl-(aq), Br-(aq), I-(aq)) and aqueous solutions of halogens.

90
Q

When does a displacement reaction occur between a aqueous solution of halide ions and an aqueous solutions of halogens? (the redox reaction) State the rule

A

A more reactive halogen will oxidise and displace a halide of a less reactive halogen - this is a displacement reaction

91
Q

How can we indicate that a redox reaction(displacement reaction) has occured between two halogens?

A

Halogens form different coloured solutions in different solvents, so its colour changes that indicates whether a redox reaction has occured.

92
Q

What colour is Cl₂, Br₂, I₂ in water?

A

Cl₂ - Pale green
Br₂ - Orange
I₂ - Brown

93
Q

What colour is Cl₂, Br₂, I₂ in cyclohexane?

A

Cl₂ - Pale green
Br₂ - Orange
I₂ - Violet

94
Q

The colours of Br₂ and I₂ in water are very similar (orange and brown respectively). How can these be better distinguished?

A

Dissolve both solutions in cyclohexane, Br₂ will remain orange, whereas I₂ will be violet

95
Q

Which halogens does chlorine oxidise? And right equations for this as well as the colour of the final solution in water and cyclohexane.

A

Cl₂ oxidises both Br- and I- ions (say bromide ions and iodide ions not bromine and iodine)
Cl₂(aq) + 2Br-(aq) → 2Cl-(aq) + Br₂(aq) Final solution is orange in both water and cyclohexane.
Cl₂(aq) + 2I-(aq) → 2Cl-(aq) + I₂(aq) Final solution is brown in water but violet in cyclohexane

WRITE THE OXIDATION NUMBER OF THESE^^

96
Q

Which halogens does bromine oxidise? And right equations for this as well as the colour of the final solution in water and cyclohexane.

A

Bromine oxidises iodide ions.

Br₂(aq) + 2I-(aq) → 2Br-(aq) + I₂(aq) Final solution is brown in water but violet in cyclohexane

97
Q

What is disproportionation?

A

Disproportionation is a reaction in which the same element is both reduced and oxidised.

98
Q

Which reactions does dispropotionation take place in and their applications?

A
  • Reaction of chlorine with water - used in water purification
  • Reaction of chlorine with COLD DILUTE aqueous sodium hydroxide (used in bleach formation)

*COLD DILUTE IS VERY IMPORTANT TO MENTION

99
Q

Why is the reaction of chlorine with water used in water purification?

A

Chlorine kills bacteria, making water safe to drink.

100
Q

Why might the reaction of chlorine with water not be used in water purification?

A

This is because chlorine may react with organic compounds in the water to form chlorinated hydrocarbons which can cause cancer.

101
Q

What is produced when chlorine reacts with water? Write an equation.

A

Chlorine reacts with water forming a mixture of hydrochloric acid, HCl, and chloric(I) acid, HClO.
Cl₂(aq) + H₂O(l) → HClO(aq) + HCl(aq)

102
Q

Using oxidation numbers show that the reaction of chlorine with water is a disproportionation reaction.

A

Cl₂(aq) + H₂O(l) → HClO(aq) + HCl(aq)
0 (Cl₂) → +1 (Cl in HClO) - ox no. ↑ so Cl oxidised
0 (Cl₂) → -1 (Cl in HCl) - ox no. ↓ so Cl reduced

103
Q

Why is the reaction of chlorine with cold dilute aqueous sodium hydroxide used in bleach formation?

A

Chlorine is only slightly soluble in water and has a mild bleaching action (HClO is a mild bleach)
Household bleach is formed when cold dilute aqueous sodium hydroxide and chlorine react at room temperature (NaClO is the bleach)

104
Q

What is produce when chlorine reacts with cold dilute aqueous sodium hydroxide? Write an equation

A

Cl₂(aq) + 2NaOH(aq) → NaCl(aq) + NaClO(aq) + H₂O(l)

105
Q

Using oxidation numbers show that the reaction of chlorine with cold dilute aqueous sodium hydroxide is a disproportionation reaction.

A

Cl₂(aq) + 2NaOH(aq) → NaCl(aq) + NaClO(aq) + H₂O(l)
0 (Cl₂) → +1 (Cl in NaClO) - ox no. ↑ so Cl oxidised
0 (Cl₂) → -1 (Cl in NaCl) - ox no. ↓ so Cl reduced

106
Q

What type of tests do chemists use to identify part of a substance (usually an ion)?

A

Qualitative tests that are usually carried out on a small scale.

107
Q

An examples of a quantitative test?

A

A precipitation reaction. This is a reaction usually between two aqueous solutions to produce an insoluble solid.

108
Q

Why are these identification tests carried out on a small scale?

A

There are usually only small quantities available, so we do not want to waste too much of the substance on too many identification steps.

109
Q

Which anions do you need to be able to identfiy?

A

CO₃²⁻ (Carbonate ions)
SO₄²⁻ (Sulfate ions)
Cl⁻, Br⁻, I⁻ (Halide ions)

110
Q

Method used to identify carbonate ions? State the positive test observations.

A

1) Add dilute strong acid (e.g. HCL) to the suspected carbonate.
2) Collect any gas formed + bubble/pass it through clear limewater

Positive test observations are:
Fizzing occurs/colourless gas produced and it turns the limewater cloudy.

111
Q

What is limewater?

A

A dilute aqueous solution of calcium hydroxide.

112
Q

Why does the positive test prove for the carbonate ion. Include any equations.

A

Cabonate compounds contain the carbonate ion, CO₃²⁻ , which react with acids to produce carbon dioxide, which in turn turns limewater cloudy.

CO₃²⁻(aq) + 2H⁺(aq) → H₂O (l) + CO₂(g)

113
Q

Method used to identify sulfate ions? State the positive test observation.

A

1) Add dilute strong HCl and barium chloride to the suspected sulfate
Positive test observations are:
A white precipitate of barium sulfate is produced.

114
Q

Why does the positive test prove for the sulfate ion? Include any equations

A

Sulfate compounds contain the sulfate ion, SO₄²⁻, which react with barium ions to form the precipitate barium sulfate, BaSO₄:
Ba⁺(aq) + SO₄²⁻(aq) → BaSO₄(s)

115
Q

The test for sulfate ions is a precipitation reaction. What is a precipitate?

A

An insoluble solid salt

116
Q

Method used to identify halide ions?

A

1) Dissolve suspected halide in water, then add nitric acid
2) Next, add an aqueous solution of silver nitrate
3) Note the colour of any precipitate formed.
4) If the colour is hard to distinguish, add aqueous ammonia (first dilute (D), then concentrated (C)) to the formed precipitates.
5) Note the solubility of the precipitate formed in aqueous ammonia.

*IT IS THE PRECIPITATE THAT YOU ARE DISSOLVING!!!

117
Q

Positive test observations for the halide test? Include equations FOR THE FORMATION OF THE PRECIPITATE.

A
  • Silver chloride: White precipitate, soluble in D + C NH₃(aq)
    Ag⁺(aq) + Cl⁻(aq) → AgCl(s)
  • Silver bromide: Cream precipitate, soluble in C NH₃(aq) only
    Ag⁺(aq) + Br⁻(aq) → AgBr(s)
  • Silver iodide: Yellow precipitate, soluble in D NH₃(aq) only
    Ag⁺(aq) + I⁻(aq) → AgI(s)
118
Q

Why doesn’t silver fluoride produce a precipitate in the first place?

A

It is because it is soluble.

119
Q

What are the observations of solubility?

A

If the halide precipitate dissolves, the precipitate disappears and the solution becomes colourless.

120
Q

Which cations do you need to be able to identigy?

A

NH₄⁺ (the ammonium ion)

121
Q

Method to identify the ammonium ion? State positive test observations.

A

1) Add sodium hydroxide solution to the suspected ammonium compounds and warm very gently
2) Test any gas evolved with red litmus paper

Positive test observations:
Ammonia gas will turn red litmus paper blue
Ammonia gas has a distinctive smell

122
Q

When testing for ammonium ions why should we take care?

A

Ammonia gas is hazardous.

123
Q

Why does the positive test prove for the ammonium ion? Include any equations.

A

The ammonium ion reacts with hydroxide ions to produce ammonia and water.
NH₄⁺(aq) + OH⁻(aq) → NH₃(aq) + H₂O(l)
The ammonia is heated so it becomes a gas that is tested with litmus paper.

124
Q

When testing unknown substances what order should the identification tests take place?

A

1) Carbonate test
2) Sulfate test
3) Halide test

125
Q

Why should identification tests be done in this specific order?

A
  • This is because, if you did the sulfate test first and added barium ions to the sample substance, both BaCO₃ and BaSO₄ form a white precipitate, so if you had not already ruled out carbonate ions you would not be able to know if the precipitate was barium carbonate or barium sulfate.
  • If you did the halide test before the sulfate test, and added silver nitrate to the sample substance, both silver halide and silver sulfate, Ag₂SO₄, form a precipitate so if you had not already ruled out sulfate ions you wouldnt know if the precipitate was silver sulfate or silver halide.
126
Q

Why is nitric acid added to the suspected halide solutions first?

A

The acid, HNO₃, is added to the suspected halide first as it reacts with any carbonate ions present, removing them and thus stopping a false posiitve result for chloride ions from occuring as carbonate ions also react with silver nitrate solution to produce a white precipitate.

127
Q

Why can you not use HCl?

A

HCl will introduce Cl- ions to the sample which may have not contained chloride ions previously, giving an false positve result for silver chloride.

Chemistry OCR AS Level (9 decks)

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