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Question 1

Properties of the alkali (group 1) metals

Answer 1

The alkali metals are the elements in Group 1 of the periodic table- they are lithium, sodium, potassium, rubidium, caesium and francium.
The alkali metals all have one electron in their outer shell. This gives them similar chemical properties. They’re all very reactive and they react in similar ways.

Question 2

Physical Properties of the alkali (group 1) metals

Answer 2

• Hardness- soft and can be cut by knife
• Low melting and boiling point
• softness increases as you go down the group
• They’re all silvery solids that have to be stored in oil and handled with forceps (they can cause chemical burns on the skin).

Question 3

Reactivity of the alkali metals

Answer 3

The more readily a metal loses its outer electrons, the more reactive it is. Group 1 metals readily lose their single outer electron to form a 1+ ion with a stable electronic structure. This makes Group 1 metals very reactive.

Question 4

Why does the reactivity increase as you go down group one?

Answer 4

As you go down group one the atoms get bigger so the outer electron will get further away from the nucleus. This means that the attraction from the nucleus becomes less, so the electron becomes more easily lost and so the metals become more reactive towards the bottom of group 1.
⬇ group 1
atom gets bigger
outer electron = further away from nucleus
attraction nucleus = less : electron = more easily lost : metals = more reactive ⬇ group 1

Question 5

alkali metals Reaction with water

Answer 5

The alkali metals react with cold water to form a metal hydroxide and hydrogen gas. The general equation for this reaction is:
alkali metal + water → metal hydroxide + hydrogen

Question 6

Metal hydroxide solutions

Answer 6

The hydroxides that are formed when the alkali metals react with water will dissolve in water to give alkaline solutions. This is where the name ‘alkali metals’ comes from. The reaction of sodium with water is illustrated in Figure 3. The reaction is similar for the other alkali metals but becoming more vigorous as you go down the group.

Question 7

Description of Group One reaction with water

Answer 7

The reactivity of Group 1 metals with water (and dilute acid) increases down the group because the outer electron is lost more easily in the reaction. This results in the reaction becoming more violent:
Lithium will move around the surface of the liquid, fizzing furiously.
Sodium and potassium do the same, but also melt in the heat of the reaction. Potassium gets hot enough to ignite the hydrogen gas being produced with a coloured flame.

Question 8

Properties of the halogens – group 7

Answer 8

Group 7 is made up of fluorine, chlorine, bromine, iodine and astatine (see Figure 1). These elements are known as the halogens. They all have 7 electrons in their outer shell, so they all have similar chemical properties.
The halogens exist as diatomic molecules (e.g., Cl2 Br2). Sharing one pair of electrons in a covalent bond gives both atoms a full outer shell.

Question 9

Physical properties of the halogens – group 7

Answer 9

The halogens also have similar physical properties. As you go down Group 7, melting points and boiling points of the halogens increase. This means that at room temperature:
* Chlorine (Cl2) is a fairly reactive, poisonous, pale green-yellow gas.
* Bromine (Br2) is a poisonous, red-brown liquid, which gives off an orange vapour at room temperature.
* iodine (I2) is a poisonous, dark grey crystalline solid which gives off a purple vapour when heated.

Question 10

Test for chlorine

Answer 10

Chlorine (Cl) bleaches damp litmus paper, turning it from blue to white (it may turn red for a moment first though - that’s because a solution of chlorine is acidic.)

damp blue litmus paper -> red -> bleached white

Question 11

Reactivity and electronic structure in Group 7

Answer 11

As you go down Group 7, the reactivity of the halogens decreases. A halogen atom only needs to gain one electron to form a 1-ion (known as a halide ion) with a stable electronic structure (see Figure 3).

Question 12

Why does the reactivity increase as you go up group 7?

Answer 12

As you go down group 7 the atoms get bigger so the outer electron will get further away from the nucleus. This means that the attraction from the nucleus becomes less, so the electron becomes harder to gain and so the metals become more reactive towards the top of group 7.

Question 13

Reaction with metals

Answer 13

The halogens will react vigorously with some metals to form salts called ‘metal halides’. Halogens higher up in Group 7 are more reactive because they can attract the outer electron of the metal more easily. For example:
* Chlorine reacts with sodium to form sodium chloride - Cl2 + 2Na → 2NaCl
* Bromine reacts with potassium to form potassium bromide - Br2 + 2K → 2KBr

Question 14

Reaction with hydrogen

Answer 14

Halogens can also react with hydrogen to form hydrogen halides. For example:
* Chlorine reacts with hydrogen to form hydrogen chloride - CI2 + H2→ 2HCI
* iodine reacts with hydrogen to form hydrogen iodide – I2 + H2→ 2HI
Hydrogen halides are soluble, and they can dissolve in water to form acidic solutions. For example, Hydrogen chloride is a gas. It dissolves in water to form an aqueous solution of hydrochloric acid.

Question 15

Displacement reactions

Answer 15

A more reactive halogen can displace (kick out) a less reactive halogen from an aqueous solution of its salt. For example:
Chlorine is more reactive than bromine, so chlorine will displace bromine from an aqueous solution of its salt (a bromide). chlorine + potassium bromide → bromine + potassium chloride
A less reactive halogen will not displace a more reactive halogen from the aqueous solution of the more reactive halogen’s salt. For example:
If you mixed bromine with sodium chloride, nothing would happen there wouldn’t be any reaction. This is because chlorine is more reactive than bromine, so the bromine cannot displace the chlorine from the chloride salt.

Question 16

Displacement reactions - Redox reactions

Answer 16

Like all displacement reactions (see p.153), the halogen displacement reactions are redox reactions. In redox reactions, both reduction and oxidation occur. The halogens gain electrons (reduction) whilst the halide ions lose electrons (oxidation). For example:
Chlorine is more reactive than bromine (it’s higher up Group 7). If you add chlorine water (an aqueous solution of Cl,) to potassium bromide solution, the chlorine will displace the bromine from the salt solution.
The chlorine is reduced to chloride ions, so the salt solution becomes potassium chloride:
chlorine + potassium bromide → bromine + potassium chloride
An ionic equation shows that the bromide ions are oxidised to bromine:
Cl₂ + 2Br- → Br₂ + 2Cl-

Question 17

Displacement reactions - method

Answer 17

You can use displacement reactions to show the reactivity trend of halogens.
1. Measure out a small amount of a halide salt solution in a test tube.
2. Add a few drops of a halogen solution to it and shake the tube gently.
3. If you see a colour change, then a reaction has happened the halogen has displaced the halide ions from the salt. If no reaction happens, there won’t be a colour change the halogen is less reactive than the halide and so can’t displace it.
4. Repeat the process using different combinations of halide salt and halogen.

Question 18

Displacement reactions – displaying reactivity trends

Answer 18

The table below shows what should happen when you mix different combinations of chlorine, bromine and iodine water with solutions of the salts potassium chloride, potassium bromide and potassium iodide. All the starting solutions of the halide salts are initially colourless.
* Chlorine displaces both bromine and iodine from salt solutions.
* Bromine can’t displace chlorine, but it does displace iodine.
* iodine can’t displace chlorine or bromine.

Question 19

Properties of the Group 0 elements

Answer 19

Group 0 elements are called the noble gases and include the elements helium, neon and argon (plus a few others).
They all have a full outer shell of electrons. For most of the noble gases, this means there are eight outer electrons. Helium, however, only has electrons in the first shell, which only needs two electrons to be filled. As their outer shell is energetically stable they don’t need to give up or gain electrons to become more stable. This means they are more or less inert - they don’t react with much at all. As they are inert they’re non-flammable (they won’t set on fire).
At room temperature they all exist as colourless monatomic gases single atoms not bonded to anything else.

Question 20

Uses of the noble gases - inertness

Answer 20

The noble gases are very unreactive, so are often used to provide an inert atmosphere. For example:
* Argon provides an inert atmosphere in filament lamps (light bulbs). It’s non-flammable, so it stops the very hot filament from burning away.
* Flash photography uses the same principle- argon, krypton and xenon - stop the flash filament from burning up in the high temperature flashes.
* Argon and helium can also be used to protect metals that are being welded. The inert atmosphere stops the hot metal reacting with oxygen.

Question 21

Uses of the noble gases - Density

Answer 21

Group 0 gases have low densities. Helium and neon are less dense than air. For example:
Helium is used in airships and party balloons. Since helium has a lower density than air, it makes balloons float. It is also non-flammable which makes it safer to use than hydrogen gas (which used to be used for the same purpose but is dangerously flammable).

Question 22

Predicting properties of elements in Group O

Answer 22

As you go down Group 0, the density of the noble gases increases. The melting and boiling points of the noble gases also increase as you go down the group. As with other groups, trends in physical properties within Group 0 mean that you can use information about some elements to predict the properties of other elements in the group. Here are some examples:

The densities of helium and argon are 0.2 kg m³ and 1.8 kg m-3 respectively. Neon comes between helium and argon in the group, so you can predict that its density will be roughly halfway between their densities:
(0.2+1.8) ÷ 2 = 2.0 ÷ 2 = 1.0
Neon should have a density of about 1.0 kg m-3.
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The boiling points of neon, argon and krypton are -246 °C, -186 °C and -153 °C respectively. Using this information, you can calculate the average gap between boiling points and therefore estimate the boiling point of xenon.
Gap between neon and argon: (-186) (-246) = 60 °C Gap between argon and krypton: (-153)-(-186) = 33 °C Average gap: (60+33) + 2 = 46.5 °C
You know that boiling points increase as you go down Group 0. So, xenon will have a higher boiling point than neon, argon and krypton. You can predict xenon’s boiling point by adding 46.5 to the boiling point of krypton.
Estimated boiling point of xenon: (-153) + 46.5=-106.5 °C