Hydrogen and water

Note: I'm adding Labels at the end of posts so that when you click on them, it will show ALL the posts in that section related to it. This will make this blog easier to use, as I don't post in a specific order, rather I answer or post according to what people needed. Hope this helps and please feedback. :) 

Section 2: Chemistry of the Elements; part e) Hydrogen and water

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

Metals above hydrogen in the reactivity series will react with acids to form a salt (e.g. magnesium sulfate or zinc chloride) and hydrogen. The metals are 'displacing' hydrogen. The higher the metal in the series, the more violent the reaction. (This is why if you put copper in acids, you won't see a reaction, as it is below hydrogen in the reactivity series. However, it does react with concentrated nitric acid but we're not concerned with that now.)

Metal + dilute hydrochloric acid à metal chloride + hydrogen

Metal + dilute sulfuric acid à metal sulfate + hydrogen

Magnesium

Magnesium reacts vigorously with cold dilute acids, and the mixture becomes very warm as heat is produced. There is rapid fizzing (effervescence) and a colourless gas is evolved, which pops with a lighted splint (the test for hydrogen). The magnesium gradually disappears and a colourless solution of magnesium sulfate or chloride is formed. 

--The reactions between magnesium and hydrochloric acid or sulfuric acid are similar because it is reacting with the hydrogen ions. All acids in solutions have hydrogen ions. Although hydrochloric acid has chloride ions, and sulfuric acid has sulfate ions, these are spectator ions. They do not participate in the reaction and are unchanged by it. 

You can rewrite the equations as ionic equations. In the case of hydrochloric acid:

Mg (s) + 2H⁺ (aq) + 2Cl^- (aq) à Mg²⁺ (aq) + 2Cl^- (aq) + H₂ (g)

You can see that the chloride ions weren’t changed by the reaction. It is a spectator ion, so we leave it out of the ionic equation. Leaving out the spectator ions produces the ionic equation:

Mg(s) + 2H⁺ (aq) à Mg²⁺ (aq) + H₂ (g)

Repeating this with sulfuric acid:

Mg (s) + 2H⁺ (aq) + SO₄^2- (aq) à Mg²⁺ (aq) + SO₄^2- (aq) + H₂ (g)

Again, leaving out the spectator ion which is the sulfate ion in this case.

Mg(s) + 2H⁺ (aq) à Mg²⁺ (aq) + H₂ (g)

So the reactions look the same because they are the same. All acids in solution contain hydrogen ions. That means that magnesium will react with any simple dilute acid in the same way. 

Aluminium

Aluminium is slow to start reacting, but after warming it reacts very vigorously. There is a very thin, but very strong, layer of aluminium oxide on the surface of the aluminium, which stops the acid from getting to it. On heating, the acid removes this layer, and the aluminium can show its true reactivity. With dilute hydrochloric acid:

2Al (s) + 6HCl (aq) à 2AlCl₃ (aq) + 3H₂ (g)

Zinc and Iron

Zinc and iron react slowly in the cold, but more rapidly on heating. Their reactions are less vigorous than that of aluminium, and iron less than zinc of course, as it is below zinc in the reactivity series. Zinc forms zinc chloride or sulfate and hydrogen. The iron forms iron (II) sulfate or iron (II) chloride and hydrogen. For example:

Zn (s) + H₂SO₄ (aq) à ZnSO₄ (aq) + H₂ (g)

Fe (s) + 2HCl (aq) à FeCl₂ (aq) + H₂ (g)

2.27 describe the combustion of hydrogen

Hydrogen reacts violently with oxygen in the presence of a flame to give water. It could explode if there was a lot of hydrogen. But a lighted splint placed at the mouth of a test tube of hydrogen will just give a squeaky pop as the hydrogen reacts with oxygen in the air. The lighted splint and a squeaky pop heard is the test for hydrogen. 

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

Anhydrous copper (II) sulfate is white, anhydrous being without water, it is dry (an--without, hydrous--related to water). Whereas hydrated copper (II) sulfate crystals are bright blue, the water is what gives it the colour, and is part of the structure. To show that the water is part of the structure, there is a '.' [dot] in the formula: 

CuSO₄·5H₂O 

^You see the dot in the middle? That shows the water is part of the copper sulfate crystal structure. 

So that is a chemical test for water, just add it to anhydrous copper (II) sulfate and watch it turn blue!

Adding water to anhydrous copper sulphate

2.29 describe a physical test to show whether water is pure

Heat the water and use a thermometer to check if it boils at exactly 100°C. Pure water boils at exactly 100°C. Or you can cool it until it freezes, it should freeze at exactly 0°C. My teacher said it's safer to state both, as pure water is the only substance that has these specific boiling and freezing points, whereas another substance might boil/freeze at either temperature. 

Hope this helped!

Group 1 elements-lithium, sodium and potassium

Update: I've included the trends, physical and chemical properties of Group 1 elements as requested by someone. Some information taken from Edexcel Chem textbooks. :) 

Section 2: Chemistry of the elements

part b) Group 1 elements-lithium, sodium and potassium

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

2.7 recall the relative reactivities of the elements in Group 1

Alkali metal	Hydroxide solution produced	Gas produced	Rate of gas produced
Lithium	Lithium hydroxide	Hydrogen	Fairly vigorous
Sodium	Sodium hydroxide	Hydrogen	Vigorous
Potassium	Potassium hydroxide	Hydrogen	Very vigorous
Rubidium	Rubidium hydroxide	Hydrogen	Explosive
Caesium	Caesium hydroxide	Hydrogen	Extremely explosive

As you can see, the reactions get more violent as you go down the group, telling you that the metals are more reactive down the group. 

Describe what happens when sodium is added to water:

When sodium is added to water, it reacts very quickly and vigorously. The reaction is exothermic and the heat produced melts the sodium. The molten sodium darts around the water surface and a yellow flame is seen. You may see a bit of fizzing/bubbling (effervescence) as hydrogen is evolved. 

Remember MM-FF. Melts, moves, floats, fizzes 
Li, Na and K are all less dense than water, hence they float on water. 

Write word equations to show the reactions of lithium, sodium and potassium with water:

Lithium + water à lithium hydroxide + hydrogen

Sodium + water à sodium hydroxide + hydrogen

Potassium + water à potassium hydroxide + hydrogen

So you see they all form hydroxides, and since all group 1 metal compounds are soluble it dissolves to form an alkali. (All alkalis are just soluble bases, but not all bases are alkalis, for example metal oxides are bases but not all of them are soluble to form alkalis.. e.g. Copper (II) oxide.)

If they ask you about what colour the solution will turn if universal indicator is added, it will turn blue or purple. This is because the solution is alkaline, and if they ask you to state what ion causes this, it’s the OH^- ion (hydroxide ion). They are called alkali metals for a reason!

2.8 explain the relative reactivities of the elements in Group 1 in terms of distance between the outer electrons and the nucleus. (single science)

As you go down the group the metals become more and more reactive. This is because their atoms get bigger, so the outer shell electrons are further away from the nucleus. So the electrostatic force between the nucleus and the outer shell electron (OSE) is weaker, hence it is easier to lose the OSE (alkali metals have only 1 OSE). The atoms want to lose the OSE to form full outer shells so they are more stable and unreactive after that. 

Also, you can think of how as the atoms get bigger, there are more electron shells in between the OSE and the nucleus. You can think of them as 'shields', so the force of the nucleus on the OSE is weaker the bigger the atom. 

The following is extra information someone has requested. 

Physical Properties

	Melting Point (°C)	Boiling Point (°C)	Density (g/cm3)
Lithium - Li	181	1342	0.53
Sodium - Na	98	883	0.97
Potassium - K	63	760	0.86
Rubidium - Rb	39	686	1.53
Caesium - Cs	29	669	1.88

• You will notice that the melting and boiling points of the elements are very low for metals, and decrease as you go down the group.
• Their densities tend to increase, but potassium has a lower density than sodium, so the densities don't increase that tidily. Lithium, sodium and potassium are all less dense than water, hence will float on it. 
• The metals are very soft and can be easily cut with a knife. They get softer as you go down the Group. They are shiny and silver when freshly cut, but tarnish within seconds on exposure to air. 

Storage and handling

All these metals are extremely reactive, and get more reactive as you go down the Group. They all react quickly with air to form oxides, and react between rapidly and violently with water to form strongly alkaline solutions of the metal hydroxides.

To prevent them from reacting with oxygen/water vapour in the air, lithium, sodium and potassium are stored under oil. Rubidium and caesium are so reactive that they have to be stored in sealed glass tubes to stop any possibility of oxygen getting to them. 

Great care must be taken not to touch any of these metals with bare fingers. There could be enough sweat on your skin to give a reaction producing lots of heat and a very corrosive metal hydroxide. 

Compounds of the alkali metals

All Group 1 metal ions are colourless. That means that their compounds will be colourless or white unless they are combined with a coloured negative ion. For instance,  Potassium dichromate (VI) is orange, because the dichromate (VI) ion is orange. Group 1 compounds are typical ionic solids and are mostly soluble in water.

Summary of the main features of the Group 1 elements:

Group 1 elements:

• are metals
• are soft with melting points and densities very low for metals
• have to be stored out of contact with air or water
• react rapidly with air to form coatings of the metal oxide
• react with water to produce an alkaline solution of the metal hydroxide and hydrogen gas
• increase in reactivity as you go down the Group
• form compounds in which the metal has a 1+ ion
• have mainly white compounds which dissolve to produce colourless solutions

Reactivity Series

Section 2: f) Reactivity series

Seems boring at first but worth watching, and the teacher's awesome, you'll see. (this links to spec 2.30)

sodium + water in a 40 gallon trash can (y) 

love this one, 'it's coming for you..'

2.30 recall 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

Mnemonic	Element	Symbol	Reactivity
Please	Potassium	K	As you can see these metals (excluding carbon) are above hydrogen in the reactivity series so they react with acids and displace hydrogen gas. Metal + acid à metal salt + hydrogen
Send	Sodium	Na
Little	Lithium	Li
Charles	Calcium	Ca
McClean	Magnesium	Mg
A	Aluminium	Al
Common	Carbon	C
Zebra	Zinc	Zn
If	Iron	Fe
The	Tin	Sn
Lame	Lead	Pb
Horse	Hydrogen	H	H+ ions are responsible for acidic properties.
Can’t	Copper	Cu	These elements are below hydrogen so they do not react with acids. (Acids contain H+ ions) Exception: Copper reacts with concentrated nitric acid, the nitrate ions oxidize copper. But that’s not really important)
Munch	Mercury	Hg
Some	Silver	Ag
Grass	Gold	Au
Properly	Platinum	Pt

2.31 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

Reactions with cold water: 

Basically, you will see that the higher up the metal in the reactivity series, the more vigorous the reaction. For example, the reaction of some alkali metals and water:

Alkali metal	Hydroxide solution produced	Gas produced	Rate of gas produced
Potassium	Potassium hydroxide	Hydrogen	Very vigorous
Sodium	Sodium hydroxide	Hydrogen	Vigorous
Lithium	Lithium hydroxide	Hydrogen	Fairly vigorous

They love asking about sodium. 

"Describe what happens when sodium is added to water" and stuff.. 

Well when sodium is added to water, it reacts very quickly and vigorously. It's an exothermic reaction and the heat produced causes the sodium to melt. The molten sodium darts around the water surface and a yellow flame is seen. You may see a bit of fizzing/bubbling (effervescence) as hydrogen is evolved. 

Remember MM-FF. Melts, moves, floats, fizzes

With Calcium, it reacts gently with cold water you may see some bubbles and calcium hydroxide is formed, or better known to some as limewater. So you will see a white precipitate forming as hydroxides are actually insoluble unless it's an alkali metal hydroxide. (All alkali metal salts and hydroxides are soluble.) 

Magnesium reacts slowly in water but reacts vigorously with steam. The reason why magnesium doesn't really react with cold water is that it becomes coated with magnesium hydroxide, which is insoluble, so it prevents water coming into contact with the magnesium. Magnesium also burns with a bright white flame and white magnesium oxide ash is formed. 

Zinc and iron don't react with cold water, but they react with steam to form oxides. Neither metal burns like magnesium. 

Copper doesn't react with water or steam as it is below hydrogen in the reactivity series. 

Reactions with acids: 

I'd say potassium, sodium, lithium and calcium are probably too reactive to react with dilute acids and would be quite dangerous. Too reactive to add safety to acids. They're already pretty violent with water. Basically the reaction would be exothermic and a lot of heat is produced, hence the hydrogen evolved could ignite and catch fire. 

Metal + acid à metal salt + hydrogen

You can tell the metal's position in the reactivity series by seeing how many bubbles are formed, or how fast. 

You can also: add a small piece of the metal to some cold water. If there is any rapid reaction, then the metal must be above magnesium in the reactivity series. If there isn't any reaction, add a small amount of the metal to some dilute hydrochloric acid or dilute sulphuric acid. If there isn't any reaction in the cold, warm/heat it carefully. 

If there's still no reaction, the metal is probably below hydrogen in the reactivity series. If there is a reaction, then it is somewhere between magnesium and hydrogen. 

2.32 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 compound. So for example a reaction with magnesium and copper (II) oxide will result in the magnesium displacing (pushing out) the copper from its oxide, so the magnesium basically replaces it. 

Magnesium + copper (II) oxide à magnesium oxide + copper

Mg (s) + CuO (s) à MgO (s) + Cu (s)

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, the reaction between zinc and copper (II) sulphate solution:

The copper is displaced by the more reactive zinc. The blue colour of the copper (II) sulphate solution fades as colourless zinc sulphate solution is formed. 

Zinc + copper sulphate à zinc sulphate + copper

Zn (s) + CuSO₄ (aq) à ZnSO₄ (aq) + Cu (s) 

2.33 understand oxidation and reduction as the addition and removal of oxygen respectively

Oxidation could mean the addition of oxygen, and reduction could mean the removal of it, but also remember that it can be about electrons: 

OILRIG = Oxidation is Loss, Reduction is Gain

So if something has lost electrons, it is oxidised.

Likewise if something has gained electrons, it has been reduced. 

2.34 understand the terms redox, oxidising agent and reducing agent

A redox reaction is a reaction in which both reduction and oxidation are occurring. They always go together. 

An oxidising agent is a substance that causes another substance to be oxidised. So it causes something else to lose electrons, and gains these electrons itself. So the oxidising agent itself is reduced. *This confuses people!! Remember that oxidising agent doesn't get oxidised, don't let the name fool you.

An example of good oxidising agents are the halogens. Especially fluorine, which is super reactive. They only need to gain one electron to get a full outer shell so they easily oxidise other elements, such as the alkali metals which only need to lose one electron too. A common example is sodium chloride-NaCl--your common table salt. 

To oxidise something can also involve oxygen, where oxygen is added to a substance. (See previous specification point)

A reducing agent is a substance that reduces something else. So it causes the substance to gain electrons, by losing electrons itself. So the reducing agent is said to be oxidised. It can also be taken as the reducing agent takes away oxygen from the other substance, such as: 

Magnesium + copper (II) oxide à magnesium oxide + copper

So here the magnesium is the reducing agent, whilst the copper (II) oxide is the oxidising agent. 

2.35 recall the conditions under which iron rusts

Iron rusts in the presence of oxygen and water. Rusting is accelerated in the presence of electrolytes such as salt. 

Note: Many metals corrode, but it is only the corrosion of iron that is referred to as rusting. 

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

Obviously to prevent rusting you need to keep oxygen and water away from the iron. You can do this by painting it, or coating it in oil/grease, or covering it with plastic. But once the coating is broken, the iron will rust. 

Coating the iron with a metal below it in the reactivity series (such as tin) is just a barrier method. Once the layer of tin on the iron is scratched, a tin can, for example, will rust very quickly. This is because the iron is more reactive than the tin and the tin won't prevent it. 

2.37 understand the sacrificial protection of iron in terms of the reactivity series.

Galvanised iron is iron that is coated with a layer of zinc. It serves as a barrier to air and water. But unlike tin, if it is scratched, the iron still doesn't rust. This is because zinc is more reactive than iron, and so corrodes instead of the iron. So the zinc is 'sacrificed' for the iron.

Galvanising is the term used only when iron/steel is coated with a protective layer of zinc, with other metals, it is sacrificial protection. 

Oxygen and Oxides

Note: This has just been updated as of 5/08, I expanded and changed bits on Spec 2.25, right at the bottom of this post. :) 


2.16 recall the gases present in air and their approximate percentage by volume

This is for unpolluted, dry air. Normally air contains a little water vapour too!

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

An experiment was carried out to find the percentage of air that is oxygen.  100 cm³ of air was passed from side to side over copper that was being heated with a Bunsen.  All the oxygen in the air will react with the copper.  No air could get in or out of the system while it was passed to and fro between the syringes.  As it was passed to and fro, the volume of air went down.  It was passed until the volume stopped decreasing, and a few minutes later the volume of remaining air was recorded.  There was 79 cm³ left. This shows that 21cm3 of the original 100cm3 of air was oxygen, because it was the oxygen that reacted with the copper to form black copper oxide. During this experiment, you should see the copper go black as it forms copper (II) oxide.

copper + oxygen à copper (II) oxide

2Cu (s) + O₂ (g) à 2CuO (s)

http://www.gcsescience.com/w10-air-oxygen.htm

2.18 describe the laboratory preparation of oxygen from hydrogen peroxide

Oxygen is most easily made in the lab from hydrogen peroxide solution using manganese (IV) oxide as a catalyst. The reaction is known as the catalytic decomposition (splitting up using a catalyst) of hydrogen peroxide. 

2H₂O₂ (aq) à 2H₂O(l) + O₂(g)

The oxygen produced is collected in an inverted glass cylinder by the downward displacement of water in a trough.

2.19 describe the reactions with oxygen in air of magnesium, carbon and sulphur, and the acid-base character of the oxides produced

When a substance burns in air, it reacts with oxygen gas and is said to be oxidised. 

2Mg (s) + O₂ (g) à 2MgO (s)

Magnesium burns with a bright white flame to give a white, powdery ash of magnesium oxide. The product is basic.

C (s) + O₂ (g) à CO₂ (g)

Carbon burns with a yellow flame, to give colourless carbon dioxide. The product is acidic.

S (s) + O₂ (g) à SO₂ (g)

Sulphur burns with a bright blue flame to give colourless sulphur dioxide, the product is acidic. Remember, sulphur dioxide gas is poisonous, and forms acid rain, so it's acidic! 

Any metal forms basic oxides, any non-metal forms acidic oxides. 

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

The marble chips are basically calcium carbonate. Remember acids react with carbonates to give a salt, carbon dioxide and water. 

Acid + metal carbonate à salt + carbon dioxide + water

 CaCO₃ (s) + 2HCl (aq) à CaCl₂ (aq) + CO₂ (g) + H₂O (l)

It displaces the air as it is denser. 

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

Most carbonates decompose (split up) when heated to from the metal oxide and carbon dioxide. 

In the case of copper (II) carbonate, this is a green powder which decomposes to give copper (II) oxide, which is a black powder and carbon dioxide when heated. 

CuCO₃ (s) à CuO (s) + CO₂ (g)

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

Carbon dioxide is a colourless, odourless gas, denser than air and slightly soluble in water. 

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

It is used in carbonate (fizzy) drinks because it dissolves in water under pressure and when you open the bottle/can, the pressure falls and the gas bubbles out of the solution.

It is used in fire extinguishers because it is a dense gas, denser than oxygen and can displace it, basically pushing it out of the way so that no more oxygen reaches the fire. Thereby putting it out. 

2.24 recall the reactions of carbon dioxide and sulphur dioxide with water to produce acidic solutions

Rain is naturally slightly acidic because of dissolved carbon dioxide. Sulphur dioxide makes it even more acidic.

2.25 recall that sulphur dioxide and nitrogen oxides are pollutant gases which contribute to acid rain, and describe the problems caused by acid rain

Acid rain is formed when acidic air pollutants such as sulphur dioxide and nitrogen dioxide dissolve in rainwater. The sulphur dioxide and nitrogen oxides mainly come from power stations and factories burning fossil fuels, or from motor vehicles. The acid rain produces many problems. 


Sulphur dioxide dissolves in water in the atmosphere to form sulphurous acid (H₂SO₃). In the presence of oxygen in the air, this acid is slowly oxidised to sulphuric acid (H₂SO₄).

Oxides of nitrogen also contribute to acid rain. In the presence of oxygen and water, nitrogen dioxide is converted to nitric acid.

nitrogen dioxide + water + oxygen à acid rain

4NO₂ (g) + 2H₂O (l) + O₂ (g) à 4HNO₃ (aq)

The pH value of unpolluted rainwater is usually slightly below 7. This is because carbon dioxide in the air dissolves in rainwater to form carbonic acid, which is a weak acid.

CO₂ (g) + H₂O (l) à H₂CO₃ (aq)

However, acid rain is much more acidic than rain that only contains carbonic acid. Acid rain has a pH value of 4 or less. Which is slightly less acidic than vinegar, which is at around pH 3. 


So, what are the effects of acid rain?

• Acid rain reacts with metals and with carbonates in marble and limestone (calcium carbonate). When this happens, metal bridges and stone buildings are damaged, even statues if they're made of limestone. 
• Acid rain can reduce the pH of natural water bodies from between 6.5 and 8.5 to below 4, which will kill fish and other aquatic life. The water is then too acidic to support life.
• Acid rain also leaches important nutrients from the soil and destroys plants. Without these nutrients, plant growth is stunted. In some cases, acid rain dissolves aluminium hydroxide (Al(OH)₃) in the soil to produce Al³⁺  ions, which are toxic to plants. Forests throughout most of Central and Eastern Europe have been destroyed in this manner by acid rain. The plants are literally 'sick' and dying. 

acid rain eroding a statue