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Introductory Chemistry

The Periodic Table and Atomic Structure

Flame colours of Group 1 and 2 elements

One way to study the arrangement of electrons in atoms is to disturb the electrons and then observe what happens as they go back to their original arrangement. A good way to do this is to heat compounds in a Bunsen burner flame and study the characteristic colours produced. These colours are caused by the electrons losing energy they have gained from the flame.

Electrons absorb energy and become excited (they jump from a lower energy orbital to a higher energy orbital). The electrons are not stable in their new quantum shell and so release energy so they can return to their ground state. This energy is released in the form of visible light at certain frequencies, giving each element a different colour, which is why we can identify them using the flame test.

All compounds of a particular element give the same flame colour, but the chlorides are the best to use because they vapourise relatively easily in a Bunsen burner flame. You need to know the procedure of the flame test:

Safety Notes!

Concentrated HCl is corrosive. Some barium salts are toxic; others harmful if inhaled or swallowed.
  1. Clean a platinum or nichrome wire in concentrated hydrochloric acid, HCl (on a watch glass) and a hot Bunsen flame. Continue this until the wire produces no colour in the flame.
  2. Pour the impure acid away and take a fresh portion. Dip your clean wire into the acid and then into the compound.
  3. Hold the wire in the flame and observe the flame colour through a diffraction grating or direct vision spectroscope and look for the coloured lines that make up the spectrum.

You also need to know the colours emitted by the elements:

Element | Colour Emitted
Sodium Yellow
Potassium Lilac
Rubidium Lilac
Caesium Pink
Beryllium No colour
Magnesium No colour
Calcium Brick Red
Strontium Bright Red
Barium Apple Green

Beryllium and magnesium have no flame colour because the frequency of electromagnetic wave they emit is beyond the visible spectrum for humans - i.e. ultraviolet and beyond.

The ionisation energy of an element is the energy required to remove one mole of electrons from one mole of atoms of the element in the gaseous state to form gaseous ions, and is given the symbol Em. The values of ionisation energies are given in kilojoules per mole, kJmol -1.

When the flame colour is viewed through a direct vision spectroscope, a number of different coloured lines may appear. This is because the electrons, which have been excited, need to get down to a more stable orbital, but in order to do this, they may have to go down to the orbital in a number of steps i.e. go to a lower energy orbital first, then the stable one.

There will be more lines in the emission spectrum of strontium than sodium because strontium has more electrons than sodium etc.

The Interpretation of Patterns in the Ionisation Energies

Quantum Shells and Sub-Orbitals

As we know from GCSE, electrons arrange themselves in groups of 2,8,8,18 etc. The reason for this is because of sub-orbitals, which we shall now explore.

There are four types of sub-orbitals (also called sub-levels) in an atom. These are called s, p, d, and f. The s sub-orbital can hold 2 electrons, the p sub-orbital can hold 6 electrons, and the d sub-orbital can hold 10 electrons. (At AS level, we do not need to know about the f sub-orbital). The amount that are filled depends on the number of electrons the element has, and this obviously increases as we go across the periodic table.

The d sub-orbital starts in the fourth quantum shell, but starts at the number 3. (See below)

Note

s orbitals are circular, p orbitals are in a firgure-of-eight.

We can use the periodic table to read off the electronic structure of the elements. (See Fig. 1)

The periodic table showing where the sub-orbitals begin Fig. 1 The periodic table showing where the different sub-orbitals types begin.

The way we write the electron configuration is like this:

e.g. Chlorine (Cl - 17 electrons): 1s2, 2s2, 2p6, 3s2, 3p6

As you can see, these numbers are just read off the periodic table because you know where the sub-orbital groups are. The first number is the quantum shell number, the letter is the type of sub-orbital, and the superscript number is the number across, when numbering from the beginning of the sub-orbital. More examples are:

Sub-Orbitals and Electron Spin

An electron spin is an intrinsic property of electrons. Electron spin is denoted by either an up or a down arrow, ↿ or ⇂. This denotes a pair of electrons with opposite spin. Depending on which sub-orbitals are filled, we either pair up these arrows or leave them as a lone arrow. Lets take some examples:

From this, we can see that a p sub-orbital fills up single arrows until it is half full (3). Then the next electron pairs up with the previous to make a new lone pair. In fact:

It is this reason why there is a peak at half-filled sub-shells as well as full sub-shells and full shells, for the first ionisation energy of the elements. (See Fig. 2)

Patterns in the First Ionisation Energies for the First Twenty Elements

The first ionisation energies of the first twenty elements are shown in the graph in Fig. 2.

Periodic patterns in the first ionisation energies of the elements Fig 2. Periodic patterns in the first ionisation energies of the elements.

What reason is there for the peaks in the ionisation energies of the elements?

Therefore, the order of stability is:

We can therefore conclude that the order of stability which exists among members of a group of elements are the result of their similar electron arrangement.

Electron Affinity

The energy change which takes place when each of the atoms in a mole of gaseous atoms accquires an electron to become a slightly-charged negative ion is known as the electron affinity, Eaff.

Cl (g) + e- → Cl- (g) Eaff = −349 kJmol-1

Note that this is for a mole of atoms (Cl), and not molecules (Cl2)

Properties of Ionic Compounds

The Electrical Conductivity of Solutions

We looked at the electrical properties of solutions using an AC circuit with a low voltage power pack, a torch bulb and carbon rods to test the solutions (concentration 0.1M). AC is used so that electrolysis does not happen (which it would using DC), so we do not decompose it.

Safety!

Copper(II) sulphate solution is harfmul. Sodium hydroxide is irritant at 0.1M

These all conducted electricity, but it depends on their solubility in water.

Migration of Ions using Copper(II) Sulphate
Interpretation of the Experiments

Ionic compounds are made up of ions and they all have opposite charges. They are also free to move when in solution, so conduct electricity. The same is true when the compound is molten.

Water is a polar solvent (has a charge). Because the charge is not very strong, it is noted with delta, δ, either positive or negative charge. Ionic compounds dissolve in polar solvents.

Fig. 3 Water is a polar solvent.

Solution in water; Oxygen is attracted to Na+, Hyrdrogen is attracted to Cl-

The slight charges on water molecules attract the ions of the ionic solid. Eventually the ionic solid is pulled apart and dissolves.

Cyclohexane is a hydrocarbon molecule and is covalent. There are no charges on the molecule to attract the ions of the ionic solid. The solid will not dissolve. If the bonds between a compound are too strong, the water charges may not be strong enough to pull it apart.

Cations Anions
Positive Negative
Usually made by metals (H+ is exception) Usually made by non-metals
Atoms lose electrons Atoms gain electrons
Ion is smaller than atom because:
  • One less quantum shell is filled
  • Remaining electrons are held more tightly
    • proton:electron ratio has increased
 Ion is larger than atom because:
  • Full quantum shell
  • More electrons so proton:electron ratio has decreased
    • electrons are less tightly held
Group | Outer electron | Gain/Lose | Ions
1 1 Lose 1 1+
2 2 Lose 2 2+
3 3 Lose 3 3+
4 4 None -
5 5 Gain 3 3-
6 6 Gain 2 2-
7 7 Gain 1 1-
8 8 None -

Within a crystal, the ions are held in place by strong forces of attraction between oppositely charged ions, and by the repulsions of similarly charged ions. These forces are non-directional, unlike the forces in a covalent bond, which are directional. So the ions are arranged in giant crystal lattices in which attraction and repulsive forces are balanced.

Electron density maps are used to show the orbitals around an atom.

Electron Arrangements in Atoms

Most compounds formed between metallic and non-metallic elements are ionic, which metals forming cations and non-metals forming anions. Dot and cross diagrams can be used to show this (see Fig. 3).

Fig 3. Dot and Cross diagram to show the formation of lithium oxide.

Because the compounds formed will have full outer shells, they will form a Noble gas structure.

Metallic Bonds

Metallic bonding is one of the three main types of bonding. Significant properties of metals are their high melting points, their high electrical conductivities and their high thermal conductivities. A simple model of bonding in a solid metal consists of positive metal ions surrounded by a sea of mobile electrons (see Fig. 4).

A metallic bond model Fig 4. Model of metallic structure of aluminium and the sea of mobile electrons.

Notice the pattern of electrons and protons; they can not just be drawn randomly.

The sea of electrons bond the metal ions tightly into the lattice and generally results in a higher melting point. Since the strong bonding forces are still present in the liquid, many metals have a wide temperature range over which they remain liquid. The mobility of the electrons provides a means of conducting the electricity and heat.