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The Arrangement Of The Elements

Chapter 5: The periodic table

5.1 The arrangement of the elements (ESABM)

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 5.1 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.

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Figure 5.1: 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 (ESABN)

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, \(\text{B}\) 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:

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We also note that the period number gives the energy level that is being filled. For example, phosphorus (\(\text{P}\)) 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: \([\text{Ne}]3\text{s}^{2}3\text{p}^{3}\). (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 (ESABO)

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

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Figure 5.2: Trends on the periodic table.

Table 5.1 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.

Element

\(_{11}^{23}\text{Na}\)

\(_{12}^{24}\text{Mg}\)

\(_{13}^{27}\text{Al}\)

\(_{14}^{28}\text{Si}\)

\(_{15}^{31}\text{P}\)

\(_{16}^{32}\text{S}\)

\(_{17}^{35}\text{Cl}\)

Chlorides

\(\text{NaCl}\)

\(\text{MgCl}_{2}\)

\(\text{AlCl}_{3}\)

\(\text{SiCl}_{4}\)

\(\text{PCl}_{5}\) or \(\text{PCl}_{3}\)

\(\text{S}_{2}\text{Cl}_{2}\)

no chlorides

Oxides

\(\text{Na}_{2}\text{O}\)

\(\text{MgO}\)

\(\text{Al}_{2}\text{O}_{3}\)

\(\text{SiO}_{2}\)

\(\text{P}_{4}\text{O}_{6}\) or \(\text{P}_{4}\text{O}_{10}\)

\(\text{SO}_{3}\) or \(\text{SO}_{4}\)

\(\text{Cl}_{2}\text{O}_{7}\) or \(\text{Cl}_{2}\text{O}\)

Valence electrons

\(3\text{s}^{1}\)

\(3\text{s}^{2}\)

\(3\text{s}^{2}3\text{p}^{1}\)

\(3\text{s}^{2}3\text{p}^{2}\)

\(3\text{s}^{2}3\text{p}^{3}\)

\(3\text{s}^{2}3\text{p}^{4}\)

\(3\text{s}^{2}3\text{p}^{5}\)

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.

Table 5.1: Summary of the trends in period 3

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

Periods on the periodic table

Exercise 5.1

Use Table 5.1 and Figure 5.2 to help you produce a similar table for the elements in period 2.

Element

\(_{3}^{7}\text{Li}\)

\(_{4}^{9}\text{Be}\)

\(_{5}^{11}\text{B}\)

\(_{6}^{12}\text{C}\)

Chlorides

\(\text{LiCl}\)

\(\text{BeCl}_{2}\)

\(\text{BCl}_{3}\)

\(\text{CCl}_{4}\)

Oxides

\(\text{Li}_{2}\text{O}\)

\(\text{BeO}\)

\(\text{B}_{2}\text{O}_{3}\)

\(\text{CO}_{2}\) or \(\text{CO}\)

Valence electrons

\(2\text{s}^{1}\)

\(2\text{s}^{2}\)

\(2\text{s}^{2}2\text{p}^{1}\)

\(2\text{s}^{2}2\text{p}^{2}\)

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.

Element

\(_{7}^{14}\text{N}\)

\(_{8}^{16}\text{O}\)

\(_{9}^{19}\text{F}\)

\(_{10}^{20}\text{Ne}\)

Chlorides

\(\text{NCl}_{3}\)

no compounds, but oxygen does combine with chlorine in compounds called chlorine oxides

no compounds

no compounds

Oxides

\(\text{NO}_{2}\) or \(\text{NO}\)

No compounds. Oxygen combines with itself to form \(\text{O}_{2}\).

no oxides, but fluorine does combine with oxygen in compounds called oxygen fluorides.

no compounds

Valence electrons

\(2\text{s}^{2}2\text{p}^{3}\)

\(2\text{s}^{2}2\text{p}^{4}\)

\(2\text{s}^{2}2\text{p}^{5}\)

\(2\text{s}^{2}2\text{p}^{6}\)

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.

Refer to the data table below which gives the ionisation energy (in \(\text{kJ·mol$^{-1}$}\)) and atomic number (Z) for a number of elements in the periodic table:

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

  1. Fill in the names of the elements.

  2. Draw a line graph to show the relationship between atomic number (on the x-axis) and ionisation energy (y-axis).

  3. Describe any trends that you observe.

  4. 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 \(Z = 7\)

Solution not yet available.