Les cours

Membres

• Liste membres
• Nos Dossiers

Before going on more challenging and more interesting things, we will look into the periodic tables available to you in the IB. If they still appear to you to be boring, outstandingly unnecessary chunks of obscure information, I urge you to read on...

Dimitri Mendeleev, in his research, noticed patterns in the properties and atomic weights of the elements. To extend his comprehension on the subject, he made a card for each of the 63 elements known at his time. The cards contained the name of the element, its symbol, atomic weight and its chemical and physical properties. Then, he arranged the cards... a bit like he used to do in his favorite card game : patience ! The elements were ordered in rows by an increase in atomic mass and in columns by their properties. With his method, he corrected mistakes in the atomic masses of certain elements (for example Berillyum which atomic mass was believed to be 14 instead of 9) and was able to predict the existence of 10 other elements and their properties, 7 of which have been discovered while the 3 remaining elements don't exist.

The very basics

Elements are arranged in the periodic table by increasing atomic number. Each column forms a group containing the elements with same number of valence electrons (number of electrons present on the last shell containing electrons). Each row is a period and the elements of one row have all the same number of shells.

In a group, the elements have roughly the same chemical and physical properties which means that across a period, the properties change but also that they change in sensibly the same way whatever the period you consider. It means the same patterns repeat themselves (for chemical and physical properties) on each period. This is called the periodicity.

Here is an example of a periodic table (for electro-negativities — I will tell you all about it in this course). Click on the image to enlarge it.

Metals are good conductors of heat and electricity. They are shiny and malleable... and you will know more about what they are and why they have such properties in the following course ! It's good having a little bit of suspense. Non-metals are not good conductors of heat and electricity. Metalloids are elements which possess both properties of metals and properties of non-metals. Inside that category, Silicon is a special case : it is a semi-conductor. Semiconductors are, as their name says, not conductors, but not completely isolated. Under certain conditions (a small flow of electrons, for example), they can show properties of conductors. This flexibility makes them very used in electronics. Be careful, not all metalloid are semi-conductors !

NOTE : Semi-conductors are used basically everywhere in modern technology : from computers to cell-phones, they are essential in building electronic circuits.

The physical properties

The atomic radius and the ionic radius of the elements are the first properties (and the easiest ) we will be looking at.

The atomic radius is defined as half the distance between the nuclei of two bonded atoms of the same element. Here, you need only logic to answer any question the IB might ask you. Elements in a group have the same number of valence electrons but a different number of shells. As a result, as you go down a group (for example going from Li to Rb), elements gain shells. Because electrons are on the next, further shell form one element to the next, the atomic radius increases down a group. However, along a period, the electrons are all on the same shell, only the number of protons and of electrons increase. Because the electrons stay at the same distance and the number of protons increase, the electrons are more held by the protons in the element at the end of the period than in the one at the start : the atomic radius decreases across a period (for example going from Li to F).

The ionic radius is a bit more tricky. You have to think first about wether the atom will gain or lose electrons when it becomes an ion. If it gains electrons, it becomes negative and is called an anion. If it loses electrons, it becomes positive and is called a cation. Okay, but how do YOU know (not your teacher, you ) whether it will become a cation or an anion? You can think of atoms as being very very lazy beings, they will only do what costs them less trouble and less work. I will give you two examples. For the Beryllium atom which has two outer (or valence) electrons, which is the most convenient : losing 2 electrons or attracting 6 electrons ? Obviously, losing two electrons is simpler than finding 6 electrons. For the Oxygen atom, which possesses 6 outer electron, which is the most convenient : losing 6 electrons or gaining 2 electrons ? Gaining two electrons will be easier ! Why should it "makes to move" 6 electrons when it can avoid it and move only 2 electrons ?

Now, you should consider what happens in terms of size to anions and cations compared to the atomic size of their neutral atom. In an anion, by losing all the valence electrons, you lose a shell, therefore their ionic radius will be smaller. In cation, you gain electrons but no proton, the more electrons you gain, the harder it is for the protons to hold them... therefore explaining that ionic radius of cations are larger than the atomic radius of their neutral atom (parent atom would look better on your exam paper ).

Along a group, the ionic radius will behave the same way as the atomic radius. In fact, the element down the group still has one more shell than the element just below it because BOTH have lost one shell in the anion's case, and their number of shells just didn't change in the cation's case ! Along a period, things are a bit trickier. In a period, cations are bigger than anions. However, the ionic radius decreases as we "go down" the anions and then as we go down the cations. All the anions of a period have the same number of electrons, we can say they are isoelectronic. The same remark can be made about the cations of the same period.

For example : Sodium has the same number of electrons than Magnesium, but one proton less... Sodium is larger than Magnesium. Similarly, Phosphorus ion has the same number of electrons than Sulfur's ion but one less proton, and so the radius of the phosporus ion is bigger than the one of the Sulfur ion. The radius of Phosphorus is also bigger than the radius of sodium. 

As mentioned in the summary, I have decided to move the melting point section of this courses to the Ionic and metallic bonding section. Why? Because otherwise I would be talking about forces you have still no idea about and it would not be very productive. If you however feel the absolute need to read it now, it is in the following course: Ionic and metallic bonding.

Fortunately (or unfortunately), atomic and ionic radius are not the only things you can understand at this point.

We have already talked about ionisation energy, but how does it varies along a period? Ionization energy is the energy needed to remove one electron from an atom in the gaseous state. Down a group, the outer electrons go further and further away from the nucleus. Therefore, the electrosatic attraction (Don't forget, it's just the "holding" !) of the protons on the outer electrons becomes weaker and less and less energy is required to remove an outer electron. The ionisation energy decreases down a group. On a similar logic, the electrostatic attraction on the outer electrons increase down a period (the number of protons in the nuclei increases), and, as a result, the ionisation energy increase down a period.

Finally, we are coming to the end of our chapter on physical properties. Electro-negativity is a measure of the attraction an atom has for a shared pair of electrons when it is covalently bonded to another atom. Along a group, the smaller the radius of the atom, the higher the electro-negativity (because a smaller radius means a greater electrostatic force). The electro-negativity diminishes across a group. Across a period, we go from elements tending to become cations to elements tending to become anions (wanting to gain electrons), so the electro-negativity will increase.

NOTE : covalent bonding occurs between two elements having a difference in electro-negativity values of less than 1.8 . Fluorine, Nitrogen and Oxygen are the more electronegative elements of the periodic table.

The chemical properties

Elements in the same group have the same chemical properties. We will, for now, consider only 2 groups : the alkali metals (group 1) and the halogens (group 7).

NOTE : The noble gases (group 8) rarely react with other elements as they already have a full shell.

The alkali metals

As you obviously know from the preceding chapters, all the elements of this group are isoelectonic (same number of electrons... But you remembered this, right ?). They all have one valence electron and they react by losing this electron. Because they have only one electron to lose, they are very reactive metals... In fact, they are so reactive that they must be kept in liquid paraffin to prevent them with reacting with the air !

The reaction of the elements of group 1 with water become more vigorous as we go down the group... the outer electrons being further and further and being therefore more readily lost.

©Author : http://openlearn.open.ac.uk/ made available by Creative Commons attribution

The flames are due to hydrogen gas which is a highly flammable. But what is the common equation for this reaction ? Let's do it for Lithium :

2 Li(s) + 2H₂O `|->` 2Li⁺ (aq) + 2OH^- (aq) + H₂ (g)

For all the other elements, just replace the symbol of Lithium by the symbol of that element, and you are done  !

The halogens

The halogens possess 7 valence electrons  and want to gain one electron to have a complete outer shell of electrons. When they do so, they become ions and are called the halide ions. The last shell is further and further away from the nucleus as we go down the group, there are more and more electrons repelling the possible extra electron (negative charges repeal each other) and the electrostatic attraction of the protons become lower with the increasing distance... So the reactivity of the halogens decreases as we go down group 7.

NOTE: Knowing the electron arrangement and knowing how many electrons each shell need to be full, you can deduce how many electrons the atom would need to gain. If you are unsure about the number of electrons in each shell, please go to the electron arrangement section in the Atomic Structure course. 

It means any halide ion can lose an electron and give it to an halogen above it in the group (except fluorine - first element of the group and therefore the most reactive). Here is an example of how you would write such a reaction :

Cl₂ (aq) + 2Br^-(aq) `->` 2Cl^-(aq) + Br₂ (aq)

The equation will be similar for any pair of reacting halogen/halide ion. The halide  ion in the reactants corresponds the less reactive halogen while the halogen in the reactants in the more reactive of the two.

If you want to know which halide ion is present in a solution, simply add a silver nitrate solution in your mixture. Silver halides (each having a specific colour) will be formed. X is any of the halide ions.` `

Ag⁺ (aq) + X^- (aq) `|->` AgX (s)

Note : under ligth, the silver halides will transform into a silver metals (Ag + `1/2` X₂). This principle is used in photography with 3 of the silver halides: AgCl, AgBr, AgI. For example, when light comes onto the photographic paper, some of the AgCl crystals inside it are changed into Ag and Cl₂... Therefore forming a dark image on the film. 

Reaction between Group 1 and Group 7

One element of group 1 can react with one element of Group 7 to form an ionic salt. The equation of the reaction can be written in the following way :

2Na (s) + Cl₂ (g) `|->` 2 Na⁺Cl^- (s)

You can write the product as NaCl (s), I was in fact just showing you the electrons gained and lost by each element at the end of the reaction. For any pair of halogen/alkali metal, you just have to replace the Na by the alkali metal you are interested in and the Cl by the halogen. As you can see once you got the rule, it's fairly simple .

The more reactive the two elements, the more vigorous the reaction.

NOTE : The salt you use in your food is none other than sodium chlorine (NaCl) !

The oxides at the SL level

Oxides are the association of elements with the oxygen element. For example, both SO₃ and Na₂O are oxides. Looking specifically at period 3, we see that :

• sodium oxide and magnesium oxide are basic and react to form hydroxides (ex: Na(OH))
• Aluminium oxide is amphoteric and can therefore act as a base or an acid.
• Sulphur trioxide and phosphorus pentoxide (P₄O₁₀) react with water to form their respective acids (namely : sulfuric acid and phosphoric acid).

I know all of this about acids and bases doesn't do much sense for you now... But in the IB curriculum, they wanted it to be put into the periodic table section to emphasize on the fact that for all periods metal oxides tend to be basic and non-metal oxides tend to be acidic (I hope you see the periodicity here !). However, I urge you NOT to learn this by heart but to first read up to the chapter on acids and bases and then come back to this section, understand it and learn it.

Finally after you explored all the properties of elements in the periodic table, I feel you are ready to finally handle "real" chemistry... So please jump ahead to the following section : ionic and metallic bonding !