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The Periodic Table

Chemistry

Chemistry

THE PERIODIC TABLE

Dmitri Mendeleev is credited as being the Father of the modern periodic table. In 1869 he arranged the 50 or so known elements in order of atomic number, Z, putting elements with similar properties in the same vertical group, and leaving gaps for unknown elements, yet to be discovered.

When the elements were later discovered, they were found to have the properties predicted by Mendeleev’s table.

Groups and Periods

Elements in the same period have the same number of shells, but the number of electrons occupying the last shell increase from left to right i.e. from one to eight.

 

The number of shells increases down a group. However, the number of electrons in the last shell of each element is the same. Elements in a given group in the periodic table share many similar chemical and physical properties. The Periodic Table came about through attempts by people to group elements according to their chemical properties.

 

The modern periodic table is very useful for giving a summary of the atomic structure of all the elements.

Some of the Groups have Names and some have Numbers.

 

Chemical Families

Patterns and Properties

A Group is a vertical column of chemically and physically similar elements. The alkali metals are in group 1 on the left of the periodic table.

The elements in this group are Hydrogen (H), Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Cesium (Cs) and Francium (Fr). They have all only one electron in their outermost shells.

Since the atomic number, hence number of shells increases down the group, the atomic radius increases down the group.

 

Why atomic or ionic radius increases down the group:

From one element to the next, an extra shell of electrons is added. This increases the electron ‘bulk’ and the outer electrons are increasingly less strongly held. The radii of the adjacent Group 2 atom is smaller than Group 1 atom on the same period, because the nuclear charge has increased by one unit (L to R ), but is attracting electrons in the same shell.

 

Similarly the radii of Group 2 M2+ ion is smaller than the adjacent Group 1 M+ ion on the same period, because the nuclear charge has increased by one unit (L to R ), but is attracting the same number of electrons in the same shells. The alkali metals are all highly reactive, losing their one outer electron to form a 1+ ion with non-metals. They give up 1 electron easily as losing 1 is easier than gaining 7 to complete the octet.

 

They all have the common properties of metals, being silvery-grey in colour, and good conductors of heat and electricity. They are unusually soft, and can easily be cut with a knife. When freshly cut, they rapidly tarnish by reaction with oxygen to form an oxide layer, which is why they are stored under oil. The first three members, lithium, sodium and potassium, are unique in being the only metals which are less dense than water (they float!).

 

Ionization energy (I.E.) decreases down the group

This is the energy required to remove one mole of electrons from the outermost shell of an atom to form a positively charged ion.

M(g) M+(g) + e-

This process can be repeated again to give the second ionization energy.

 

This is more difficult than the first ionization energy because we are removing a negative electron from a positive ion.

M+(g) M2+(g) + e-

And again…

M2+(g) M3+(g) + e-

 

It is possible to continue in this way until all of the electrons on an atom have been removed. As you go down the group from one element down to the next, the atomic radius gets bigger due to an extra filled electron shell. The outer electrons are further and further from the nucleus and are also shielded by the extra full electron shell of negative charge.

 

Therefore the outer electrons are less and less strongly held by the positive nucleus and so less and less energy is needed to remove them. Successive ionization energies always increase e.g. … 3rd > 2nd > 1st, because the same nuclear charge is attracting fewer electrons and on average closer to the nucleus. BUT note the 2nd IE for Group 1, and the 3rd IE for Group 2, show a particularly significant increase in IE compared to the previous ionization energy or energies. This is due to removing an electron from an electronically highly stable full inner shell and puts an upper limit on the chemically stable oxidation state.

 

Why reactivity increases down the group

When an alkali metal atom reacts, it loses an electron to form a singly positively charged ion e.g. Na Na+ + e- (in terms of electrons 2.8.1 2.8 and so forming a stable ion with a noble gas electron arrangement).

As you go down the group from one element down to the next the atomic radius gets bigger due to an extra filled electron shell.

 

The outer electron is further and further from the nucleus and is also shielded by the extra full electron shell of negative charge.

Therefore the outer electron is less and less strongly held by the positive nucleus.

This combination of factors means the outer electron is more easily lost, the M+ ion more easily formed, and so the element is more reactive as you go down the group.

The reactivity argument mainly comes down to increasingly lower ionization energy down the group.

 

See also:

ISOTOPES

THE STRUCTURE OF THE ATOM AND THE PERIODIC TABLE

USES OF HYDROGEN GAS

PROPERTIES OF HYDROGEN GAS

HYDROGEN

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