The early periodic table
With the discovery of over 50 elements by the 1860s, scientists began to try to sort the elements into a logical sequence by identifying patterns in their chemical properties. The work of John Newlands and Dmitri Mendeleev in developing early periodic tables, ultimately led to the development of the modern periodic table.
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 a pattern among the early elements. The pattern showed that each element was similar to the element eight places ahead of it. For example, starting at Li (lithium), Be (beryllium) is the second element, B (boron) is the third and Na (sodium) is the eighth element. He then put the similar elements into vertical columns, known as groups.
Newlands' table showed a repeating or periodic pattern of properties, but this pattern eventually broke down. By ordering strictly according to atomic mass, Newlands was forced to put some elements into groups which did not match their chemical properties. For example, he put iron (Fe), which is a metal, in the same group as oxygen (0) and sulfur (S), which are two non-metals.
As a result, his table was not accepted by other scientists.
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 did some other things that made his table much more successful. He realised 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.
Sometimes this method of arranging elements meant there were gaps in his horizontal rows or 'periods'. But instead of seeing this as a problem, Mendeleev thought it simply meant that the elements which belonged in the gaps had not yet been discovered.He was also able to work out the atomic mass of the missing elements, and so predict their properties. And when they were discovered, Mendeleev turned out to be right. For example, he predicted the properties of an undiscovered element that should fit below aluminium in his table. When this element, called gallium, was discovered in 1875, its properties were found to be close to Mendeleev's predictions. Two other predicted elements were later discovered, lending further credit to Mendeleev's table.
Evaluation of Newlands and Mendeleev
- Ordered elements by atomic weight
- Included only the elements known at the time
- Maintained a strict order of atomic weights
- Every eighth element had similar properties (Newlands’ Law Of Octaves)
- Was criticised by other scientists for grouping some elements with others when they were obviously very different to each other
- Ordered elements by atomic weight
- Left gaps for elements he predicted would be discovered later
- Swapped the order of some elements if that fitted their properties better
- Elements in groups had similar properties
- Was seen as a curiosity to begin with, but then as a useful tool when the predicted elements were discovered later
The modern periodic table
Dmitri Mendeleev’s early periodic table was further refined in the early 20th century in light of the discovery of protons, neutrons and electrons. This allowed elements to be placed in appropriate groups according to atomic numbers instead of atomic masses, which produced the periodic table we use today.
Dmitri Mendeleev put the elements in order of their relative atomic mass, and this gave him some problems. For example, iodine has a lower relative atomic mass than tellurium, so it should come before tellurium in Mendeleev's table.In order to get iodine in the same group as other elements with similar properties - such as fluorine, chlorine and bromine - he had to put it after tellurium, which broke his own rules.However, the discovery of protons, neutrons and electrons in the early 20th century allowed Mendeleev’s table to be refined into the modern periodic table. It involved an important modification – the use of atomic number to order the elements. An element’s atomic number (also called proton number) is the number of protons in its atoms.
Using atomic number instead of atomic mass as the organising principle was first proposed by the British chemist Henry Moseley in 1913. It explained why Mendeleev needed to change the order of some of the elements in his table. For example, tellurium has a higher atomic mass than iodine, but iodine has a higher atomic number than tellurium. So, even though he didn't know why, Mendeleev was (as it turned out) right to place iodine after tellurium.
The arrangement of the moderd periodic table
Columns in the table - groups.
The elements in a vertical column are in the same group. The main groups are labelled groups 1-7, with the noble gases in group 0. All elements in a group have similar chemical properties.The elements in a group all have the same number of electrons in their highest occupied energy level (also referred to as the outer shell). This is why they have similar chemical properties.An element’s group number is the same as the number of electrons in its highest occupied energy level (outer shell). For example, all the metals in Group 2 have 2 electrons in their highest occupied energy level (outer shell).
Rows in the table - periods
Elements in a horizontal row are in the same period. The periods are numbered from top to bottom.The period number is the same as the number of occupied energy levels (shells). For example, magnesium is in period 3 – its atoms have three occupied energy levels. Calcium is in period 4 – its atoms have four occupied energy levels.
Trends within the periodic table
Elements within different groups within the periodic table have different physical and chemical properties. This determines the kinds of reactions these elements have. Different groups also show different trends, in terms of reactivity, as you move down a group. This can also determine how violently a reaction occurs - or whether it happens at all.
Goup one elements; The elements in group 1 are called the alkali metals. They belong to the left-hand column in the periodic table. They are very reactive and must be stored in oil to avoid contact with air or water.The alkali metals are soft, reactive metals. They react vigorously with water and become more reactive as you go down the group.
The alkali metals have the following properties in common:
Physical and chemical trends in Group 1
Physical and chemical trends in Group 1
Melting and boiling points
The alkali metals all have low melting points and boiling points compared to other metals. The melting points and boiling points decrease as you go down the group.
As you go down the group, the metals become more reactive. Lithium (at the top) is the least reactive, while francium (which is at the bottom) is the most reactive.
You will probably see lithium, sodium and potassium at school, but rubidium and caesium are considered to be too reactive to use in the classroom. Francium is radioactive and very rare - there are only a few grams of it in the whole of the Earth's crust at any time.
The elements in the centre of the periodic table - between groups 2 and 3 - are called the transition elements. They are all metals. They include most of the commonly-used metals, such as iron, copper, silver and gold.
- Melting points; High (except mercury, which is liquid at room temperature)
- Reactivity; Low (do not react so vigorously with water or oxygen)
- Strength; Strong and hard
- Density; High
- Compounds; Coloured
hydrogen + nitrogen ammonia
3H2(g) + N2(g) 2NH3(g)
Many transition elements form ions with different charges. This means that iron oxide can exist in two forms, iron(II) oxide, FeO, and iron(III) oxide, Fe2O3.
Group 7 elements
The elements in Group 7 are called the halogens. They belong to the column second from the right in the periodic table. The halogens are all toxic, but this can be a useful property. Chlorine is used to sterilise drinking water and water in swimming pools. Iodine is used in antiseptics to treat wounds.
The halogens have the following properties in common:
Physical and Chemical trends in Group 7
Melting and boiling point; The halogens have low melting points and low boiling points. You can see from the graph that fluorine, at the top of Group 7, has the lowest melting point and lowest boiling point in the group. The melting points and boiling points then increase as you go down the group.
Colour; The halogens become darker as you go down the group. Fluorine is very pale yellow, chlorine is yellow-green and bromine is red-brown. Iodine crystals are shiny purple-black but easily turn into a dark purple vapour when they are warmed up.
Reactivity; The halogens become less reactive as you more down the group. Fluorine (at the top of the group) is the most reactive, while astatine (at the bottom) is the least reactive.
Displacement reactions in the halogens
Halogens react with metals to form ionic compounds, which dissolve in water. The reacting of the halogens also decreases as you move down the group.These two principles can be used to explain displacement reactions. In these reactions, a more reactive halogen can displace a less reactive halogen from an aqueous solution of its salt.For example, chlorine is more reactive than bromine, so it can displace bromine from bromide compounds:
chlorine + sodium bromide → sodium chloride + bromine
Cl2(g) + 2NaBr(aq) → 2NaCl(aq) + Br2(aq)
If you test different combinations of the halogens and their salts you can work out a reactivity series for the halogens. The most reactive halogen displaces all the other halogens from solutions of their salts, while the least reactive halogen is always displaced. It works just the same whether you use a sodium salt or a potassium salt.Test your understanding using this animation in which chlorine, bromine and iodine are added to various halogen salts. Note carefully the products which are present in the test tube after each reaction.