# 9-1 Chemistry/The periodic table

## The periodic table

The elements in the periodic table are arranged in order of their atomic, proton, number. Each column, known as a group, contains elements with similar properties as they all have the same number of electrons in their outer shell.

### Metals and non-metals

Metals are elements which react to form positive ions whilst non-metals react to form negative ions. The majority of elements are metals and are found on the left side of the periodic table whilst the few non-metals are found at the top-right.

### Development of the periodic table

A German scientist called Johann Döbereiner put forward his law of triads in 1817. Each of Döbereiner's triads was a group of three elements. The appearance and reactions of the elements in a triad were similar to each other. See the table below:

Alkali metals
symbol A (atomic mass)
Li 7
Na 23
K 39

Döbereiner discovered that the relative atomic mass of the middle element in each triad was close to the average of the relative atomic masses of the other two elements. However, his law worked only for few elements, and hence soon fell out of favour. [1]

#### Newlands' Law of Octaves

An English scientist called John Newlands put forward his law of octaves in 1864. He arranged all the elements known at the time into a table in order of relative atomic mass. When he did this, he found that each element was similar to the element eight places further on. For example, starting at Li, Be is the second element, B is the third and Na is the eighth element. Newlands' table showed a repeating or periodic pattern of properties, but it had problems. For example, he put iron in the same group as oxygen and sulphur, which are two non-metals. As a result, his table was not accepted by other scientists. [2]

#### Mendeleev's periodic table

In 1869, just five years after John Newlands put forward his law of octaves, a Russian chemist called Dmitri Mendeleev published a periodic table. Mendeleev also arranged the elements known at the time in order of relative atomic mass, but he also realized that that the physical and chemical properties of elements were related to their atomic mass in a 'periodic' way, and arranged them so that groups of elements with similar properties fell into vertical columns in his table. [3]

Mendleev's table from 1871

In the table above, you can see that some elements have been marked with a dash (-). Those elements were not discovered at that time, so he marked (predicted) them as eka-<a compound with similar specifications>. For example, gallium and eka-aluminimum have similar properties, as you can see in the table below:

 Eka-Aluminum (Ea) Gallium (Ga) Atomic mass 68 amu 69.9 amu Melting point Low 30.15°C Density 5.9 g/cm 3 5.94 g/cm 3 Formula of oxide Ea 2 O 3 Ga 2 O 3

The element gallium was discovered four years after the publication of Mendeleev’s table, and its properties matched up remarkably well with eka-aluminum, fitting into the table exactly where he had predicted. This was also the case with the element that followed gallium, which was named eventually named germanium.

Mendeleev’s periodic table gained wide acceptance with the scientific community and earned him credit as the discoverer of the periodic law. Element number 101, synthesized in 1955, is named mendelevium after the founder of the periodic table.[4]

## Group 0: Noble gases

The elements on the far right of the periodic table are known as the noble gases, Group 0. They are very unreactive as they have a full outer shell, giving them a naturally stable arrangement of electrons (they do not have to react to form a stable arrangement). So we say they are inert - extremely unreactive and do not take part in chemical reactions and they exist as single atoms (monatomic).

The boiling points of the noble gases increase with increasing relative atomic mass, so the further down in the group you go, the higher the boiling point. This is as a result of greater intermolecular forces as atoms get bigger resulting in more energy being needed to break the bonds and change state. However, compared to metals they all have very low boiling and melting points and usually exist as gases at room temperature.

## Group 1: Alkali Metals

The elements on the far left of the periodic table are known as the Alkali Metals or Group 1 (since they all have 1 electron in their outer shell). In order to form a stable arrangement, they need to lose 1 electron, resulting in them becoming a positive ion (1+). All alkali metals are soft (as they can be easily cut with a knife) and have relatively low melting and boiling points.

This shows sodium hydroxide, the product when sodium reacts with water.

Alkali metals get there name from their reaction with water:

${\displaystyle {\ce {2Na + 2H2O -> 2NaOH + H2}}}$

${\displaystyle {\text{Sodium + Water}}\rightarrow {\text{Sodium hydroxide + hydrogen}}}$

Alkali metals react with water to produce an alkali metal hydroxide and hydrogen. The hydroxide is an alkali, with a pH greater than 7, which is able to dissolve in water, turning a universal indicator blue or purple.

The reactivity of alkali metals increases going down the group since there is a higher shielding effect created by more filled electron shells, meaning that electrons can be lost more easily. This means that if francium were to react with water, there would be an explosion.

${\displaystyle {\ce {4Na + O2 -> 2Na2O}}}$

${\displaystyle {\text{Sodium + Oxygen}}\rightarrow {\text{Sodium oxide}}}$

Alkali metals react with oxygen to produce an alkali metal oxide. This explains why when you cut an alkali metal the shiny surface quickly dulls as an oxide layer forms, having reacted with oxygen. It also explains why alkali metals burn vigorously when you place them in a jar filled with oxygen. The oxide will form as a white smoke.

${\displaystyle {\ce {2Na + Cl2 -> 2NaCl}}}$

${\displaystyle {\text{Sodium + Chlorine}}\rightarrow {\text{Sodium chloride}}}$

Group 1 metals react with chlorine to produce (chlor)ides, a compound containing chlorine and another element.

## Group 7: Halogens

Halogens are the Group 7 elements in the periodic table and have 7 electrons in their outer shell. The halogens are non-metals and consist of diatomic (a pair of halogen atoms bonded together) molecules which are gases at room temperature. As you go down Group 7, the halogens become less reactive, have higher melting and boiling points. This is the exact opposite of alkali metals which get more reactive as you go down.

Halogens react with metals to produce salts which are held together by ionic bonds. The halogen becomes a halide ion with 1- charge. when they bond with metals. In order to write the salt formed, all you need to remember is that the net charge of the new compound formed must be 0. For example, sodium (an alkali metal) reacting with chlorine (a halogen) makes sodium chloride as seen above. Another example would be:

${\displaystyle {\ce {Fe + Br2 -> Fe^{3+}Br^{-1}_{3}}}}$

${\displaystyle {\text{Iron + Bromine}}\rightarrow {\text{Iron (III) bromide}}}$

As we see, iron (a transition metal) reacts with bromine (a halogen) to make iron bromide. Remember that halogens are diatomic.

Halogens react covalently with non-metals to form compounds which all have a simple molecular structure:

${\displaystyle {\ce {H2 + Cl2 -> 2HCl}}}$

${\displaystyle {\text{Hydrogen + Chlorine}}\rightarrow {\text{Hydrogen chloride}}}$

### Displacement reactions of Halogens

A more reactive halogen can displace a less reactive halogen from an aqueous solution of its salts. In simpler terms, the halogen further up the group will displace a halogen further down the group because reactivity decreases going down the group.

${\displaystyle {\ce {Cl2 + KI -> KCl + I2}}}$

${\displaystyle {\text{Chlorine + Pottasium iodide}}\rightarrow {\text{Pottasium chloride}}+{\text{iodine}}}$

Notice how diatomic halogens lose their pairs when they react ionically. This is because they can obtain a full outershell by covalent bonding with the metal.