The arrangement of the elements

The periodic table of the elements is a method of showing the chemical elements in a table with the elements arranged in order of increasing atomic number. Most of the work that was done to arrive at the periodic table that we know can be attributed to a Russian chemist named Dmitri Mendeleev. Mendeleev designed the table in 1869 in such a way that recurring ("periodic") trends or patterns in the properties of the elements could be shown. Using the trends he observed, he left gaps for those elements that he thought were “missing”. He also predicted the properties that he thought the missing elements would have when they were discovered. Many of these elements were indeed discovered and Mendeleev's predictions were proved to be correct.

To show the recurring properties that he had observed, Mendeleev began new rows in his table so that elements with similar properties were in the same vertical columns, called groups. Each row was referred to as a period. Figure 1 shows a simplified version of the periodic table. The full periodic table is reproduced at the front of this book. You can view an online periodic table at periodic table.

**Figure 1:** A simplified diagram showing part of the periodic table. Metals are given in gray, metalloids in light blue and non-metals in turquoise.

Definitions and important concepts

Before we can talk about the trends in the periodic table, we first need to define some terms that are used:

Atomic radius

The atomic radius is a measure of the size of an atom.
Ionisation energy

The first ionisation energy is the energy needed to remove one electron from an atom in the gas phase. The ionisation energy is different for each element. We can also define second, third, fourth, etc. ionisation energies. These are the energies needed to remove the second, third, or fourth electron respectively.
Electron affinity

Electron affinity can be thought of as how much an element wants electrons.
Electronegativity

Electronegativity is the tendency of atoms to attract electrons. The electronegativity of the elements starts from about 0.7 (Francium (Fr)) and goes up to 4 (Fluorine (F))
A group is a vertical column in the periodic table and is considered to be the most important way of classifying the elements. If you look at a periodic table, you will see the groups numbered at the top of each column. The groups are numbered from left to right starting with 1 and ending with 18. This is the convention that we will use in this book. On some periodic tables you may see that the groups are numbered from left to right as follows: 1, 2, then an open space which contains the transition elements, followed by groups 3 to 8. Another way to label the groups is using Roman numerals.
A period is a horizontal row in the periodic table of the elements. The periods are labelled from top to bottom, starting with 1 and ending with 7.

For each element on the periodic table we can give its period number and its group number. For example, B is in period 2 and group 13. We can also determine the electronic structure of an element from its position on the periodic table. In Chapter 4 you worked out the electronic configuration of various elements. Using the periodic table we can easily give the electronic configurations of any element. To see how this works look at the following:

**Figure 2**

We also note that the period number gives the energy level that is being filled. For example, phosphorus (P) is in the third period and group 15. Looking at the figure above, we see that the p-orbital is being filled. Also the third energy level is being filled. So its electron configuration is: $[Ne] 3 s^{2} 3 p^{3}$ . (Phosphorus is in the third group in the p-block, so it must have 3 electrons in the p shell.)

Periods in the periodic table

The following diagram illustrates some of the key trends in the periods:

**Figure 3:** Trends on the periodic table.

Ionization energy

Table 4 summarises the patterns or trends in the properties of the elements in period 3. Similar trends are observed in the other periods of the periodic table. The chlorides are compounds with chlorine and the oxides are compounds with oxygen.

Table 1: Summary of the trends in period 3

Element	$_{11}^{23} Na$	$_{12}^{24} Mg$	$_{13}^{27} Al$	$_{14}^{28} Si$	$_{15}^{31} P$	$_{16}^{32} S$	$_{17}^{35} Cl$
Chlorides	$NaCl$	${MgCl}_{2}$	${AlCl}_{3}$	${SiCl}_{4}$	${PCl}_{5}$ or ${PCl}_{3}$	$S_{2} {Cl}_{2}$	no chlorides
Oxides	${Na}_{2} O$	$MgO$	${Al}_{2} O_{3}$	${SiO}_{2}$	$P_{4} O_{6}$ or $P_{4} O_{10}$	${SO}_{3}$ or ${SO}_{4}$	${Cl}_{2} O_{7}$ or ${Cl}_{2} O$
Valence electrons	$3 s^{1}$	$3 s^{2}$	$3 s^{2} 3 p^{1}$	$3 s^{2} 3 p^{2}$	$3 s^{2} 3 p^{3}$	$3 s^{2} 3 p^{4}$	$3 s^{2} 3 p^{5}$
Atomic radius	Decreases across a period.
First Ionization energy	The general trend is an increase across the period.
Electro-negativity	Increases across the period.
Melting and boiling point	Increases to silicon and then decreases to argon.
Electrical conductivity	Increases from sodium to aluminium. Silicon is a semi-conductor. The rest are insulators.

Note that we have left argon ( $_{18}^{40} Ar$ ) out. Argon is a noble gas with electron configuration: $[Ne] 3 s^{2} 3 p^{6}$ . Argon does not form any compounds with oxygen or chlorine.

Exercise 1: Periods on the periodic table

Problem 1:

Use Table 4 and Figure 3 to help you produce a similar table for the elements in period 2.

Answer 1:

Table 1
Element	$_{3}^{7} Li$	$_{4}^{9} Be$	$_{5}^{11} B$	$_{6}^{12} C$
Chlorides	$LiCl$	${BeCl}_{2}$	${BCl}_{3}$	${CCl}_{4}$
Oxides	${Li}_{2} O$	$BeO$	$B_{2} O_{3}$	${CO}_{2}$ or $CO$
Valence electrons	$2 s^{1}$	$2 s^{2}$	$2 s^{2} 2 p^{1}$	$2 s^{2} 2 p^{2}$
Atomic radius	Decreases across the period.
First Ionization energy	Increases across the period.
Electro-negativity	Increases across the period.
Melting and boiling point	Increases to carbon and then decreases to neon.
Electrical conductivity	Increases to boron and then decreases. Boron is a semi-conductor. Lithium and beryllium are conductors. The rest are insulators.

Table 2
Element	$_{7}^{14} N$	$_{8}^{16} O$	$_{9}^{19} F$	$_{10}^{20} Ne$
Chlorides	${NCl}_{3}$	no compounds, but oxygen does combine with chlorine in compounds called chlorine oxides	no compounds	no compounds
Oxides	${NO}_{2}$ or $NO$	No compounds. Oxygen combines with itself to form $O_{2}$ .	no oxides, but fluorine does combine with oxygen in compounds called oxygen fluorides.	no compounds
Valence electrons	$2 s^{2} 2 p^{3}$	$2 s^{2} 2 p^{4}$	$2 s^{2} 2 p^{5}$	$2 s^{2} 2 p^{6}$
Atomic radius	Decreases across the period.
First Ionization energy	Increases across the period.
Electro-negativity	Increases across the period.
Melting and boiling point	Increases to carbon and then decreases to neon.
Electrical conductivity	Increases to boron and then decreases. Boron is a semi-conductor. Lithium and beryllium are conductors. The rest are insulators.

Problem 2:

Refer to the data table below which gives the ionisation energy (in kJ·mol⁻¹) and atomic number (Z) for a number of elements in the periodic table:

Table 3
Z	Name of element	Ionization energy	Z	Name of element	Ionization energy
1		1310	10		2072
2		2360	11		494
3		517	12		734
4		895	13		575
5		797	14		783
6		1087	15		1051
7		1397	16		994
8		1307	17		1250
9		1673	18		1540

Fill in the names of the elements.
Draw a line graph to show the relationship between atomic number (on the x-axis) and ionisation energy (y-axis).
Describe any trends that you observe.
Explain why:
1. the ionisation energy for $Z = 2$ is higher than for $Z = 1$
2. the ionisation energy for $Z = 3$ is lower than for $Z = 2$
3. the ionisation energy increases between $Z = 5$ and $Z = 7$

Answer 2:

Z	Name of element	Ionisation energy	Z	Name of element	Ionisation energy
1	Hydrogen	1310	10	Neon	2072
2	Helium	2360	11	Sodium	494
3	Lithium	517	12	Magnesium	734
4	Beryllium	895	13	Aluminium	575
5	Boron	797	14	Silicon	783
6	Carbon	1087	15	Phosphorus	1051
7	Nitrogen	1397	16	Sulphur	994
8	Oxygen	1307	17	Chlorine	1250
9	Fluorine	1673	18	Argon	1540

3)The ionisation energy increases from Z=3 to Z=10. The ionisation energy also increases from Z=11 to Z=18. Between Z=2 and Z=3 there is a sharp decrease in ionisation energy. This decrease is also seen between Z=10 and Z=11. The ionisation energy increases between Z=1 and Z=2.

4) Ionisation energy depends on the electron configuration of an atom. If the outer shell is full the ionisation energy will be high and it will be hard to lose electrons. (Ionisation energy is the energy needed for an atom to lose an electron.)

a) Since the atom with Z=2 has a full outer shell it is harder to remove electrons and so the ionisation energy will be higher. Z=1 has only 1 electron in its outermost shell.

b) The atom with Z=3 has 1 electron in its outermost shell. The atom with Z=2 has a full outer shell. This means that the ionisation energy for Z=3 will be lower than for Z=2 since it will be easier to remove the electron from the atom with Z=3.

c) As you increase the atomic number the number of electrons in the outermost shell increases. As the number of electrons in the outermost shell increase the energy needed to remove one of those electrons also increases. In other words the atom holds onto its electrons more tightly and the ionisation energy increases.

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Everything

Science

The arrangement of the elements

Definitions and important concepts

Periods in the periodic table

Exercise 1: Periods on the periodic table