Tuesday, 12 June 2012

Ionic Compounds

Note: Updated as of 2/1/13, and this blog is no longer being updated because I'm done with IGCSEs and am doing IB, which is seriously hectic. I apologise that this blog is not totally complete but it's all I have and I'm sorry. 

f) Ionic compounds

1.27 describe the formation of ions by the gain or loss of electrons

Ions are atoms or molecules with an electric charge due to the gain or loss of electrons. If electrons are lost, the ion has a positive charge. Metals tend to do this, so they form cations (positive ions), so normally elements from group 1-3 will form cations.

If electrons are gained, the ion has a negative charge. Non-metals tend to do this, and they form anions (A-Negative-ION - ANION). So elements from group 5-7 will form anions. Group 0/8 are the noble gases and are inert + unreactive, so they do not form ions.

1.28 understand oxidation as the loss of electrons and reduction as the gain of electrons

OILRIG -  Oxidation Is Lost, Reduction Is Gain

1.29 recall the charges of common ions in this specification

Positive ions/Cations
Negative ions/Anions
Name of ion
Name of ion
Copper (I)
Copper (II)
Iron (II)
Lead (II)
Nickel (II)
Iron (III)

1.30 deduce the charge of an ion from the electronic configuration of the atom from which the ion is formed

So if the electronic configuration is 2.8.1, you can see that the atom has one outer shell electron only. And so it only needs to lost that to have a full outer shell. So its ion would have a positive 1 (1+) charge, as the electronic configuration would be 2.8. Basically, if there are less outer shell electrons to lose to have a full outer shell, then the charge will be positive. (Here, it is easier to lose than gain, because the ion would have to gain SEVEN electrons to have a full outer shell!)

Another example, if the electronic configuration is 2.8.7, then the atom only needs to gain 1 outer shell electron to have a full outer shell. So the ion formed would have the electronic configuration of 2.8.8, and so the charge would be negative 1 (you’re gaining one electron, which has a negative charge (1-) ). Thus if it is easier to gain electrons, (it is here, rather than losing 7 outer shell electrons), then the charge will be negative.

Here, the Sodium (Na) has lost one electron. It doesn't have equal numbers of protons and electrons anymore; it has one less electron than protons (or you can think of it as one more proton than electrons), so it has a 1+ charge. 

The chlorine, on the other hand, has gained one electron. So it has one more electron than proton, thus it has a negative 1 charge (1-). The formula for sodium chloride is NaCl. 

The magnesium atom loses 2 electrons to an oxygen atom, and they both have full outer shells now. It is common that ions form noble gas structures like this to become more stable and unreactive like the group 0/8 elements. 

The magnesium oxide is held together by very strong attractions between the ions. The ionic bonding is stronger here than in sodium chloride as this time you have 2+ ions attracting 2- ions. The greater the charge, the greater the attraction.

The formula for magnesium oxide is MgO.

1.31 explain, using dot and cross diagrams, the formation of ionic compounds by electron transfer, limited to combinations of elements from Groups 1, 2, 3 and 5, 6, 7

This is a common example, it doesn't matter which element has dots/crosses for their electrons,  the important thing is to make it clear that the electrons are transferred from one atom to another to make 2 ions. The one losing an electron here is sodium, so it becomes a positive ion called a cation, and the chlorine atom gains an electron and becomes a chloride ion (an anion).

Diagram of bonding in magnesium oxide. A magnesium ion (2,8)2+ gives two electrons to an oxide ion (2,8)2-. Both ions have full highest energy levels
Magnesium oxide: MgO
Even though only one magnesium ion and one oxide ion is shown, the actual equation is:
2Mg + Oà 2MgO
(remember that oxygen is diatomic)

Diagram of bonding in calcium chloride. A calcium ion (2,8,8)2+ gives one electron to a chloride ion (2,8,8)- and another electron to another chloride ion (2,8,8)-. All three ions have full highest energy levels
Calcium chloride: CaCl2
However, here you have to show 2 chloride ions because calcium loses 2 electrons, and 2 chlorine atoms gain an electron each to form 2 chloride ions.

1.32 understand ionic bonding as a strong electrostatic attraction between oppositely charged ions

1.33 understand that ionic compounds have high melting and boiling points because of strong electrostatic forces between oppositely charged ions
The strong electrostatic forces between oppositely charged ions are ionic bonds as mentioned in the previous spec point. And these require a lot of energy to break, hence high melting/boiling points. 

Wednesday, 6 June 2012

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 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) + SO42- (aq) à Mg2+ (aq) + SO42- (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 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: 
^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!