Chemistry

PERIODS

Characteristics of periods

The first period starts with hydrogen (H) and ends with helium (He). It has just two elements H (Z=1) and He (Z = 2).

H has one electron in the first-shell. He has 2 electrons in the first-shell.

As we have seen in the chapter on the structure of atoms, the first-shell can hold only 2 electrons.

Thus the first period is complete.

It has to be borne in mind that the place of hydrogen is unique in the periodic table.

It has been placed above the alkali elements starting with Li in group 1A.

This is because H has valency 1 just as the other alkali elements.

But the properties of hydrogen otherwise are very different from the other group 1A alkali elements Li, Na, K, Cs, etc.

Now let us see the next periods: periods 2 and 3.

The second period starts with Li (Z=3), where the first-shell is filled and the next shell is starting to fill. After Li the next element is beryllium (Be, Z=4).

Its first-shell is complete and it has 2 electrons in the second shell. The maximum number of electrons held in the second shell is 8.

So the period has 8 elements, in which each element’s second shell is getting filled.

The last element in the period is neon (Ne, Z=10).

Neon’s both first and the second shell are completely filled.

A similarly situation occurs for the third period.

Here the next shell after second shell or the third shell is getting filled.

The maximum number of electrons in the third shell is 8.

Thus across the period, starting with element sodium (Na, Z=11) the third-shell has 1 electron; and the period ends with argon (Ar, Z= 18) which has 2 electrons in the first, 8 electrons in the second shell and 8 electrons in the third shell.

Trends across Period 3

Now let us look at some of the chemical and physical properties in a particular period. What we will learn from one period, will hold true for all the other periods.

Consider the third period.

The figure below shows how the electronic configuration is changing as we go from left to right in the period.

The number of valence electrons is increasing in an integral fashion.

The change in the valency is according to the tendency to give up or borrow electrons.

Thus elements in the same period have consecutive atomic numbers and different valencies.

If we see the atomic radii across the period, we will notice that the size decreases across the period.

Now let us consider the metallic character of the elements in the third period. Figure below shows the same.

We have proper metals in the first and the second places: sodium (Na) and magnesium (Mg) are alkali and alkaline-earth metals.

They give up the electrons in the last shell very easily.

They are shiny in colour and conduct electricity.

After Mg comes aluminum (Al).

Al has 3 electrons in its outermost shell and behaves like a metal.

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The next element is silicon (Si).

It has 4 electrons in its outermost shell.

It thus needs to borrow four electrons or give up all its four electrons to form a stable shell.

Si does not do any of these; instead it binds tetrahedral most of the time.

Thus Si behaves neither like a metal nor like a non-metal.

Hence it is called as a metalloid. After Si, come three elements: phosphorus (P), sulphur (S) and chlorine (Cl).

All the three are non-metals.

Thus while moving from left to right in the period, the metallicity decreases.

Also the chemical reactivity first decreases and then increases.

As discussed before, the chemical reactivity depends on how easily the outermost orbit gives off or borrows electrons to make a stable orbit.

The two extremes of the third period, namely Na and Cl are very reactive.

But Na is very electro-positive in nature, whereas Cl is very electro-negative in nature.

If we look at the nature of the oxides formed by the elements in period three, we see that sodium oxide is basic in nature.

The next oxide, namely magnesium oxide is also basic in nature.

At the other extreme, chlorine oxide, sulphur oxides and phosphorus oxides are acidic in nature.

The mid-elements like Al, Si have their oxides behave in both acidic and basic manner, depending on the oxidation conditions.

Such oxides are said to be amphoteric in nature.

First Ionization Energy across Period 3

First ionization energy generally increases going across Period 3.

However, it needs more detailed consideration than the trend in Group 2 because:

  • The first ionization energy drops between magnesium and aluminium before increasing again.
  • The first ionization energy drops between phosphorus and sulphur before increasing again.

Explanation of this trend

General increase across the period

The first ionization energy is the enthalpy change when one mole of gaseous atoms forms one mole of gaseous ions with a single positive charge.

It is an endothermic process, i.e. is positive.

A general equation for this enthalpy change is:

Going across Period 3:

  • There are more protons in each nucleus so the nuclear charge in each element increases.
  • Therefore the force of attraction between the nucleus and outer electron is increased, and
  • There is a negligible increase in shielding because each successive electron enters the same energy level.
  • So more energy is needed to remove the outer electron.

Trend in atomic radius of Period 3 elements

Atomic radius decreases going across Period 3.

Explanation of this trend

Going across Period 3:

  • The number of protons in the nucleus increases so …
  • The nuclear charge increases …
  • There are more electrons, but the increase in shielding is negligible because each extra electron enters the same principal energy level …
  • Therefore the force of attraction between the nucleus and the electrons increases …
  • So the atomic radius decreases.

As the number of electrons in each atom increases going across Period 3, you might expect the atomic radius to increase.

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This does not happen, because the number of protons also increases and there is relatively little extra shielding from electrons in the same principal energy level.

Trend in electrical conductivity

Electrical conductivity increases going across Period 3 from sodium to aluminium, then decreases to silicon.

The remaining elements have negligible conductivity.

Explanation of this trend

For an element to conduct electricity, it must contain electrons that are free to move.

In general, metals are good conductors of electricity and non-metals are poor conductors of electricity.

Sodium, magnesium and aluminium

Sodium, magnesium and aluminium are all metals.

They have metallic bonding, in which positive metal ions are attracted to delocalized electrons.

The delocalized electrons are free to move and carry charge. Going from sodium to aluminium:

  • The number of delocalized electrons increases…
  • There are more electrons which can move and carry charge…
  • So the electrical conductivity increases.

Silicon

Silicon is a metalloid (an element with some of the properties of metals and some of the properties of non-metals).

Silicon has giant covalent bonding. It has a giant lattice structure similar to that of diamond, in which each silicon atom is covalently-bonded to four other silicon atoms in a tetrahedral arrangement.

This extends in three dimensions to form a giant molecule or macromolecule.

Silicon is called a semiconductor because:

  • The four outer electrons in each atom are held strongly in covalent bonds …
  • Few electrons have enough energy at room temperature to enter the higher energy levels so there are few delocalized electrons and silicon is a poor conductor but …
  • At higher temperatures more electrons are promoted to the higher energy levels …
  • So there are more delocalized electrons to move and carry charge.

Non-metals

The remaining elements in Period 3 do not conduct electricity:

  • In phosphorus, sulphur and chlorine, the outer electrons are not free to move and carry charge because they are held strongly in covalent bonds…
  • In argon (which exists as single atoms) the outer electrons are not free to move and carry charge because they are held strongly in a stable third energy level.

Trends in melting and boiling points

The trends in melting points and boiling points going across Period 3 are not straightforward, and need more detailed consideration than the trends in Group 2:

  • Melting points generally increase going from sodium to silicon, then decrease going to argon (with a “bump” at sulphur).
  • Boiling points generally increase going from sodium to aluminium, then decrease to argon (again with a “bump” at sulphur).

Explanation of the trends

Melting

When a substance melts, some of the attractive forces holding the particles together are broken or loosened so that the particles can move freely around each other but are still close together.

The stronger these forces are, the more energy is needed to overcome them and the higher the melting temperature.

Boiling

When a substance boils, most of the remaining attractive forces are broken so the particles can move freely and far apart.

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The stronger the attractive forces are, the more energy is needed to overcome them and the higher the boiling temperature.

Sodium, magnesium and aluminium

Sodium, magnesium and aluminium are all metals.

They have metallic bonding, in which positive metal ions are attracted to delocalized electrons.

Going from sodium to aluminium:

  • The charge on the metal ions increases from +1 to +3 (with magnesium at +2) …
  • The number of delocalized electrons increases …
  • So the strength of the metallic bonding increases and …
  • The melting points and boiling points increase.

Silicon

Silicon is a metalloid (an element with some of the properties of metals and some of the properties of non-metals).

Silicon has giant covalent bonding.

It has a giant lattice structure similar to that of diamond, in which each silicon atom is covalently-bonded to four other silicon atoms in a tetrahedral arrangement.

This extends in three dimensions to form a giant molecule or macromolecule.

Silicon has a very high melting point and boiling point because:

  • All the silicon atoms are held together by strong covalent bonds …
  • Which need a very large amount of energy to be broken.

PHOSPHORUS, SULPHUR, CHLORINE AND ARGON

These are all non-metals, and they exist as small, separate molecules.

Phosphorus, sulphur and chlorine exist as simple molecules, with strong covalent bonds between their atoms.

Argon exists as separate atoms (it is monatomic).

Their melting and boiling points are very low because when these four substances melt or boil, it is the van der Waal’s forces between the molecules which are broken.

These bonds are very weak bonds so little energy is needed to overcome them.

Sulphur has a higher melting point and boiling point than the other three because:

  • Phosphorus exists as P4molecules
  • Sulphur exists as S8molecules
  • Chlorine exists as Clmolecules
  • Argon exists individual Ar atoms
  • The strength of the van der Waal’s forces decreases as the size of the molecule decreases
  • So the melting points and boiling points decrease in the order S8> P4> Cl2> Ar

Summary of the characteristics of elements in a period:

  • The atomic numbers are consecutive.
  • The number of valence electrons in the elements increases incrementally from left to right.
  • The elements of the same period have different valencies.
  • The atomic radii decrease while going from left to right in a period.
  • Metallic character reduces while going from left to right in a period.
  • Chemical reactivity is highest at the two extremes and is the lowest in the centre.

The reactivity on the left extreme is most electro-positive whereas on the extreme right it is most electro-negative.

  • Oxides formed of elements on the left are basic and of elements on the right are acidic in nature.

Oxides of elements in the centre are amphoteric.

See also:

THE NOBLE GASES – GROUP 0

THE HALOGENS – GROUP 7

TREND IN ATOMIC RADIUS

THE PERIODIC TABLE

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