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Flashcards in Topic 2 Deck (39):
1

Understand the terms group and period

Groups:
A group is a column (goes down) in the periodic table. All elements in a group have similar chemical properties (because they have the same number of electrons on their outer shell), however the properties of each elements (such as reactivity) often gradually change as you go down a group (as the atomic number increases).

Periods:
A period is a row (goes right) in the periodic table. Properties of elements in the same period are not similar and quite often change.

2

Recall the positions of metals and non-metals in the Periodic Table

*Red line under B, Si, As, Te, At that separates metals and non-metals*

Left + Metals and Right = Non metals

The group of elements between group 1 and group 3 are called the transition metals. They don't tend to follow normal patterns and don't have a group number.

3

Explain the classification of elements as metals or non-metals on the basis of their electrical conductivity and the acid-base character of their oxides

- Metals are conductors that form metal oxides that are alkaline
- (Most) non-metals are ones that don't conduct and form non-metal oxides that are acidic

4

Understand why elements in the same group of the Periodic Table have similar chemical properties

Elements in the same column have the same number of electrons in their outer shells. This means they react and bond similarly. This is because they need to lose or gain the same number of electrons to become stable and have full outer shells

5

Understand that the noble gases (Group 0) are a family of inert gases and explain their lack of reactivity in terms of their electronic configurations.

Nobel gases (group 0, the last column in the Periodic Table) are inert, meaning they don't reach with many things at all. This is because the noble gases have a full outer shell, so they have no need to gain or lose electrons

6

Describe the reactions of these elements with water and understand that the reactions provide a basis for their recognition as a family of elements

*Group 1 Metals*

Lithium (Li), Sodium (Na) and Potassium (K) all react vigorously with water. The rest further down the group is not safe to experiment with.

Lithium:
- Floats
- Fizzes
- Solid eventually disappears
- Produces some heat but doesn't melt

Sodium:
- Floats
- Heat given off, causing it to melt into a ball
- Gradually disappears
- Sometimes a white trail of sodium hydroxide can be seen, but it usually ends up dissolved in the water
- Moves around (due to hydrogen given off)
- If it gets stuck on the side of the container, it will burns with an orange flame

Potassium:
- Burns with a lilac flame
- Reacts similarly to sodium but quicker
- Gives off heat, causing it to melt into a ball

7

Describe the relative reactivities of the elements in Group 1

As you go down the group, the metals because more reactive

8

Explain the relative reactivities of the elements in Group 1 in terms of distance between the outer electrons and the nucleus.

Group 1 elements only have 1 electron in their outer shell. Electrons are held to the atom by the forces of attraction between the nucleus (which has protons, so a positive charge) and the electrons themselves (which have a negative charge). If there are more shells, the last electron (which will be lost if the metal reacts with anything) is further away from the nucleus, so the forces of attraction between the nucleus and the electron are weaker. The more shells, the further away, and it's more reactive.

As you go down the group, they become more reactive because there is an increase in the number of shells.

9

Recall the colours and physical states of the elements at room temperature

Fluorine (F) - Gas, Yellow
Chlorine (Cl) - Gas, Yellow-green
Bromine (Br) - Liquid, Brownish red
Iodine (I) - Solid, Purple
Astatine (At) - Solid, Black

10

Make predictions about the properties of other halogens in this group (Group 7)

As you go down the group:
- Reactivity decreases
- Colour darkens
- Melting/boiling point increases
- Particles become closer together (i.e. fluorine is a gas but astatine is a solid)

11

Understand the difference between hydrogen chloride gas and hydrochloric acid

Hydrogen chloride gas is HCl(g)
When dissolved in water, it becomes HCl(aq) because the ions become detached: they dissociate, leaving the separate H+ and Cl- ions. H+ is acidic, which is what makes it an acid

Dissociate - Disconnect or Separate

12

Explain, in terms of dissociation, why hydrogen chloride is acidic in water but not in methylbenzene

If HCl is dissolved in water, it dissociates (Splits up) into H+ and Cl- ions. This solution (hydrochloric acid) is acidic because it contains H+ ions.

However, if HCl is dissolved in methylbenzene, it doesn't dissociate, therefore no H+ ions are present so the solution is not acidic.

13

Describe the relative reactivities of the elements in Group 7

Group 7 elements become less reactive as you go down the group

At the top, the positive charge of the proton in the nucleus is close to the surface (as there are few shells) this makes it easy for them to pull in the one electron they need to become stable, meaning they are very reactive.

Lower down where there are more shells, the pull of the proton is further from the surface making it less easy to pull in another electron.

14

Describe experiments to demonstrate that a more reactive halogen will displace a less reactive halogen from a solution of one of its salts

A displacement reaction is basically a reaction where the more reactive element displaces (pushes out) a less reactive element from a compound.

Chlorine is more reactive than iodine (as it is higher up in group 7). Therefore, if you add chlorine water to potassium iodide solution, the chloride will react with the potassium to form potassium chloride (basically, it displaces the iodine). The iodine is displaced from the salt (potassium iodide) and just kind of gets left in the water solution (this turns the solution brown),

15

Understand these displacement reactions as redox reactions

When a more reactive halogen displaces a less reactive halogen, it's called a redox reaction. This is because an element has gained something and the other has lost something.

OIL RIG:
Oxidisation Is Loss (of electrons)
Reduction Is Gain (of electrons

16

Recall the gases present in air and their approximate percentage by volume

The air is made up of approximately:
- 21% oxygen
- 78% nitrogen
- 1% other gases (including 0.9% argon and 0.037% carbon dioxide

17

Explain how experiments involving the reactions of elements such as copper, iron and phosphorus with air can be used to investigate the percentage by volume of oxygen in air

Copper, iron and phosphorus all react with air.
If you know the volume of air that you have, then react it with one of these, then re-measure the volume of air, What has been lost is all oxygen that reacted.

Copper (google for image):
- When copper is heated, it reacts with oxygen in the air to make copper (II) oxide
- This reaction uses up oxygen
- If you heat an excess of copper in a tube and pass it over two syringes, you can use the markers on the syringes to work out how much oxygen has been used up.
*New X (after reaction is cooled and cooled)/Starting X = X% must be oxygen

Iron (google for image):
Iron will react with oxygen to produce rust, this means iron will remove oxygen from the air.
- Soak some iron wool in acetic acid (this acid will catalyse the reaction)
- Push the iron wool at the bottom of a test tube and invert the tube into a beaker of water
- Mark the level of the water in the tube at the beginning
- Leave the experiment for a set amount of time (eg
1 hour)
- Mark the level of water in the tube at the end of the experiment

In conclusion, over time, the level of the water will rise in the test tube. This is because the iron reacts with the oxygen in the air, making iron oxide (The water rises as it takes the place the oxygen took up).

To work out the percentage of air that is oxygen, mark the level of the water in the tube at the beginning and end of the experiment, then fill up the tube to each mark and pour the contents into a measuring cylinder to find out the volume of air at the start and end. Using the difference between the start and end volumes, work out the % that has been used to (should be approx 20%).

Phosphorus:
You can do a similar experiment with white phosphorus. White phosphorus smoulders in air to produce phosphorus oxide. Use the same calculation method as with iron.

18

Describe the laboratory preparation of oxygen from hydrogen peroxide, using manganese(IV) oxide as a catalyst

Hydrogen peroxide will decompose to form oxygen and water. However, this process is very slow so manganese (IV) oxide is added to speed up the decomposition (acts as a catalyst). The oxygen produced can be collected in two ways...

1. Over water
- Connect a delivery tube to bubble the gas into a n upside-down measuring cylinder in a beaker/bowl full of water.

2. In a gas syringe
- You can use a gas syringe to collect pretty much any gas. Just connect it to the flash that the hydrogen peroxide is decomposing in.

NOTE: The equation for this reaction is 2(H2O2) (aq) -> 2(H2O) (l) + O2 (g)

19

Describe the reactions of magnesium, carbon and sulfur with oxygen in air, and the acid-base character of the oxides produced

When anything is burnt, it reacts with oxygen in the air to form oxides (which can have either acidic or basic character).

Magnesium:
When magnesium burns in air, it produces a bright white flame and a white powder is formed (this is magnesium oxide). Magnesium oxide is slightly alkali when dissolved in water.

The equation for the reaction (burning Mg in air) is 2Mg(s) + O2(g) -> 2MgO(s)

Carbon:
Carbon will only burn in air if it is very strongly heated. It burns with a yellowy-orangey flame and produces carbon dioxide (as a gas). Carbon dioxide is slightly acidic when dissolved in water.

The equation for the reaction (burning C in air) is C(s) + O2(g) -> CO2(g)

Sulfur:
Sulphur burns (in air) with a pale blue flame and produces sulfur oxide which is acidic when dissolved in water.

The equation for the reaction (burning S in air) is S(s) + O2(g) -> SO2(g)

20

Describe the laboratory preparation of carbon dioxide from calcium carbonate and dilute hydrochloric acid

Dilute HCl will react with calcium carbonate (marble chips) to produce calcium chloride, water and carbon dioxide.

Method:
- Put marble chips at the bottom of a flash
- Fill the flash with hydrochloric acid, doesn't have to be full, just covering the marble chips
- Immediately attach a bung with a delivery tube into an upturned test tube in water
- The carbon dioxide will be collected in this tube

Equations:
Hydrochloric acid + Calcium carbonate -> Calcium chloride + Water + Carbon dioxide

2HCl(aq) + CaCO3(s) -> CaCl2(Aq) + H2O(l) + CO2(g)

21

Describe the formation of carbon dioxide from the thermal decomposition of metal carbonates such as copper(II) carbonate

The thermal decomposition (heating) of metal carbonates will produce CO2 as in thermal decomposition, the substance being heated will break down into simpler substances.

Method:
- Put some copper (II) carbonate (its a green powder) into a test tube and insert a bung with a deliver tube at the top
- Camp the test tube at a 90 degree angle and insert the delivery tube into another test tube (that is positioned vertically)
- Heat the copper (II) carbonate with a bunsen burner

NOTE: Because CO2 is denser than air, the downward delivery method can be used.

Equations:
Copper (II) carbonate -> Copper oxide + Carbon dioxide
CuCO3(g) -> CuO(s) + CO2(g)

22

Describe the properties of carbon dioxide, limited to its solubility and density

Carbon dioxide is more dense than air (which is why it can be collected using the downward delivery method). It is also soluble in water at high pressure.

23

Explain the use of carbon dioxide in carbonating drinks and in fire extinguishers, in terms of its solubility and density

Carbon dioxide makes soft drinks fizzy. As CO2 is slightly soluble in water and dissolves in drinks when under high pressure, a slightly acidic solutions forms due to the formation of carbonic acid. It eventually goes flat because one the bottle is opened, the CO2 is escaping.

CO2 is more dense than air, this makes it perfect for fire extinguishers. The CO2 will sink onto flames and 'suffocate' them (stop the oxygen getting to the flames). As fire needs oxygen to burn, the fire will go out as no oxygen can get to it.

NOTE: CO2 fire extinguishers are only used when water extinguishers aren't safe. For example, when putting out an electrical fire.

24

Understand that carbon dioxide is a greenhouse gas and may contribute to climate change.

Carbon dioxide is a green house gas. This means it traps heat in the planet and stops the sun's rays being reflected away from the earth. Excess levels of carbon dioxide contribute to climate change, causing the earth to warm up at unnatural rates.

25

Describe the reactions of dilute hydrochloric and dilute sulfuric acids with magnesium, aluminium, zinc and iron

Acid + Metal -> Salt + Hydrogen

Magnesium:
- Reacts vigorously with cold dilute acids
- Produces lots of bubbles

Aluminium:
- Little reaction with cold dilute acids (as it has a protective aluminium oxide layer)
- Reacts vigorously with warm dilute acids and produces a lot of bubbles

Zinc:
- React slowly with dilute acids but more strongly if you heat them up

Iron:
- Same reactions as zinc

NOTE: The more reactive the metal, the faster the reaction. For example, very creative metals (such as sodium) will react explosively, as the reaction is very quick. The speed of reaction is given off by the amount of bubbles that are given off.

26

Describe the combustion of hydrogen

If hydrogen is burnt in air, water is produced (initially as water vapour, as it is hot, but later condenses to water if cooled)

2(H2) + O2 -> 2(H20)

27

Describe the use of anhydrous copper(II) sulfate in the chemical test for water

Anhydrous copper (II) sulphate will turn from white to blue if water is present.

To test for water, all you need to do is add anhydrous copper (II) sulphate, which is a white powder, to the substance you are testing. If the anhydrous copper (II) sulphate turns from white to blue, water is present.

This is a reversible reaction, meaning if you heat the blue powder, it will turn white again. This is because the water has 'left' (it has evaporated)

NOTE: Anhydrous means without water, hydrates means with. Therefore, if water is present, anhydrous copper (II) sulphate will turn to hydrates copper (II) sulphate.

*This does not show pure water, just that the substance contains water molecules*

28

Describe a physical test to show whether water is pure.

If the substance you are testing boils at 100 degrees or freezes at 0 degrees, it is pure water.

29

Understand that metals can be arranged in a reactivity series based on the reactions of the metals and their compounds: potassium, sodium, lithium, calcium, magnesium, aluminium, zinc, iron, copper, silver and gold

Reactivity Series (Top are most reactive and bottom are least reactive):
- Potassium
- Sodium
- Lithium
- Calcium
- Magnesium
- Aluminium
- ZInb
- Iron
- *Lead*
- Coppper
- Silver
- Gold
- *Platinum*

30

Describe how reactions with water and dilute acids can be used to deduce the following order of reactivity: potassium, sodium, lithium, calcium, magnesium, zinc, iron and copper

Metals high up in the series (potassium, sodium, lithium and calcium) react very vigorously with water.

Metals in the middle (magnesium, zinc and iron) react with steam but don't react with cold water.

Copper won't react with either steam or water.

The more reactive the element is with water and dilute acid, the further up the series the element is positioned.

31

Deduce the position of a metal within the reactivity series using displacement reactions between metals and their oxides, and between metals and their salts in aqueous solutions

Any metal higher in the reactivity series will displace one lower down from its oxide. For example...

To find out whether magnesium or copper is more reactive (higher in the series), just add magnesium to copper (II) oxide. The magnesium will displace copper, meaning that magnesium is more reactive than the copper. Alternatively, you could add copper to magnesium oxide, no reaction would take place. This is because copper is less reactive than magnesium (lower in the series) as it will not displace the magnesium.

This must mean magnesium is above copper in the reactivity series.

It's the same thing with metals and a solution of their salt (the more reactive metal will displace a less reactive metal). For example...

To find out whether zinc or copper is more reactive, add zinc to a solution of copper (II) sulphate. The zinc will displace the copper (As it is more reactive), meaning it is positioned higher than copper in the reactivity series.

NOTE: In this particular reaction, the blue colour of the copper (II) sulphate solution fades as colourless zinc sulphate solution is formed.

32

Understand oxidation and reduction as the addition and removal of oxygen respectively

Oxidisation is when something gains oxygen.
Reduction is when something loses oxygen.
(In terms of oxygen gain/loss)

OIL RIG:
Oxidisation is Loss (of electrons)
Reduction is Gain (of electrons)

33

Understand the terms redox, oxidising agent, reducing agent

Redox reaction - A reaction in which both reduction and oxidation are taking place.

Oxidising agent - A substance in that is capable of oxidising another substance

Reducing agent - A substance in that is capable of reducing another substance.

34

Describe the conditions under which iron rusts

Rusting ONLY occurs when iron is in contact with both water and oxygen (from the air). The chemical reaction that is taking place is oxidation of iron to form iron (III) oxide (oxidation reaction), water then bonds to the iron (III) oxide and forms hydrated iron (III) oxide - this is rust.

Iron + Oxygen -> Iron (III) Oxide
Iron (III) Oxide + Water -> Hydrated Iron (III) Oxide

35

Describe how the rusting of iron may be prevented by grease, oil, paint, plastic and galvanising

These all create a 'barrier around the iron, stopping water and/or oxygen from reaching it.

Paint/plastic can also be decorative and can be used on big or small structures (its versatile) and can be decorative

Grease/oil can only be used on moving parts

Galvanise is a protective layer of zinc

36

Understand the sacrificial protection of iron in terms of the reactivity series.

The sacrificial method involves placing a more reactive metal (such as zinc) with the iron. Water and oxygen then react with the sacrificial metal rather than the iron (as its more reactive than the iron).

Zinc is often used as it is more reactive than iron, so zinc will be oxidised instead of iron. A coating of zinc could be sprayed onto the iron object (galvanising). Another method is to bolt big blocks of zinc onto the iron.

37

Describe tests for the cations:
i Li+, Na+, K+, Ca2+ using flame tests
ii NH4+, using sodium hydroxide solution and identifying the ammonia evolved
iii Cu2+, Fe2+ and Fe3+, using sodium hydroxide solution

Flame Tests:
To do a flame test, dip a platinum wire loop in dilute HCl then hold it in a flame, take it out once the flame burns without a colour (this means the platinum is clean). Dip the loop in the sample your testing and put it back in the flame, the colour of the flame will tell you what metal ion is in the substance.

Li+ - Burns with a crimson red flame
Na+ - Burns with a yellow-orange flame
K+ - Burns with a lilac flame
Ca2+ - Burns with a brick red flame

Sodium Hydroxide Solution and Identify the Ammonia Evolved (NH4+):
You can check for ammonia gas using damp red litmus paper (will turns from red to blue if ammonia is present).

To test for ammonium ions in a unknown substance, add a few drops of sodium hydroxide solution to a solution of the unknown substance (in a test tube). Hold a piece of litmus paper near the top of the test tube, if ammonia is being given off, ammonium ions are present in the unknown substance.

Sodium Hydroxide:
Metal hydroxides are insoluble and precipitate out of solution when formed.

For this test, add a few drops of sodium hydroxide solution to a solution with what your testing. This will form an insoluble hydroxide. Some of these hydroxides have characteristics colours, this can be used to tell which metal ions are present in the unknown solution.

Cu2+ - Blue
Fe2+ - Sluggish Green
Fe3+ - Reddish Brown

38

Describe tests for the anions:
i Cl-, Br- and I-, using dilute nitric acid and silver nitrate solution
ii SO4^2-, using dilute hydrochloric acid and barium chloride solution
iii CO3^2-, using dilute hydrochloric acid and identifying the carbon dioxide evolved

i) Using diluted nitric acid & silver nitrate solution (depending on the precipitate)
- Chloride Ions (Cl-) + Nitric Acid + Silver Nitrate -> White Precipitate (Silver Bromide)
- Bromide Ions + Nitric Acid + Silver Nitrate -> Cream Precipitate (Silver Bromide)
- Iodide Ions + Nitric Acid + Silver Nitrate -> Yellow Precipitate (Silver Iodide)

ii) Using dilute hydrochloric acid & barium chloride soltuion
- Sulphate Ions + Hydrochloric Acid + Barium Chloride -> White Precipitate (Barium Sulphate)
-SO4(2-) + HCl + BaCl2(2+) -> BaSO4

iii) Using dilute hydrochloric acid (& identifying if CO2 is produced)
- Carbonate + Hydrochloric Acid -> Salt + water + carbon Dioxide
Test for CO2:
When bubbled through lime water, will turn it cloudy

39

Describe tests for the gases:
i hydrogen
ii oxygen
iii carbon dioxide
iv ammonia
v chlorine

Hydrogen:
Squeaky pop test - put a lit splint into a test tube of gas. If it pops then hydrogen is in the tube

Oxygen:
Oxygen will relight glowing splint

Carbon Dioxide:
If bubbled through lime-water and it changes cloudy, then carbon dioxide is present.

Ammonia:
Damp red litmus paper will turn blue if ammonia is present

Chlorine:
Turns damp blue litmus paper red, then white (as it bleaches)