Trends of the Periodic Table

 Trends In The Periodic Table

Trends in Atomic Radii

The atomic radius (covalent radius) of an atom is defined as half the distance between the nuclei of two atoms of the same element that are joined together by single covalent bond.

This distance is known as bond length and maybe measured by methods known as 'X-ray diffraction' and 'electron diffraction'.

No values of atomic radii for noble gases as they dont form covalent bonds.

If you examine values of atomic radii a definite pattern emerges 
> values of atomic radius decrease along any one period
> values of atomic radius increase down any one group

Atomic Radius increases down groups:

i) New Shells (of Electrons): Outermost electrons are further away from the nucleus.

ii) Screening Effect of Inner Electrons: Inner shells of electrons help to shield the outer electrons from the positive charges of the nucleus

Atomic Radius decreases across periods:

i)Increasing effective nuclear charge: No. of protons increasing from left to right, therefore a greater attraction force of outer electrons will draw shell closer to nucleus, as the number of protons increase.

ii) No increase in the screening effect: Extra electrons is added to same outer shell left to right across a period Therefore no screening effect as no new shells formed, hence the atoms get smaller.

Trends in Electronegativity

The values of electronegativity decrease down the group

i) Increasing atomic radius: outermost electrons becoming further away from the attractive force of nucleus, hence smaller attraction between nucleus and shared pair of electrons ( electronegatvity decreases)

ii)Screening effect of inner electrons: the atomic radius decreases across a period therefore outermost electrons are being drawn closer to attractive force of nucleus (electronegativity decreases)

The values of electronegativity increase across periods:

i) Increasing effective nuclear charge: no. of protons increasing across period leading to a greater force of attraction between the outermost electrons + nucleus, hence the electrons involved in bonding are being more strongly attracted to the nucleus. (electronegativty increases

ii) Decrease atomic radius: the atomic radius decrease across a period therefore outermost electrons are being drawn closer attractie force of nucleus (electronegativity increases)

Trends in Ionisation Energy

The values of ionisation energy decrease down the group

i) Increasing atomic radius: outermost electrons becoming further away from the attractive force of nucleus, making it easier to remove an electron.

ii)Screening effect of inner electrons: going down a group, the full inner shells of electrons shield the outermost electrons from the attractive force of the nucleus making it easier to remove an electron.

The values of ionisation energy increase across periods:

i) Increasing effective nuclear charge: no. of protons increasing across period leading to a greater force of attraction between the outermost electrons + nucleus, this requires more energy to remove an electron(energy increase)

ii) Decrease atomic radius: going across a period, the atomic radius decreases hence outermost electrons are being drawn closer to attractive forces of nucleus, requiring more energy to remove electron(i.e energy increases)

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The first ionisation energy of an atom is the minimum energy required to completely remove the most loosely bound electron from a neutral gaseous atom in its ground state

Na(g) →Na⁺ (g)+e^-

The second ionisation energy of an element refers to the removal of a second electron from the positive ion formed when the first electron is removed.

Na⁺(g) →Na²⁺ (g)+e^-

Ionisation energy values are gotten using the emission spectra of atoms.

Alkali metals are the elements with lowest first ionisation energy ⇒ tend to lose electrons readily, hence very reactive

Noble gases are the element with the highest first ionisation energy⇒ these gases have very little tendency to lose electrons, hence very unreactive

EXCEPTIONS TO GENERAL TREND ACROSS PERIOD

Beryllium+Nitrogen have higher value than expected. These irregularities in first ionisation energy trends can be explained by the fact that any sublevel that is completely filled or exactly half filled have extra stability⇒hence requiring more energy than expected to break these stable arrangements by removing an electron

Further evidence for existance of energy levels

Using Potassium K (1s22s22p63s2 3p64s2 ) as example;

If you removes 1 electron from n=4 energy level, it will have the lowest ionisation energy value as it is the easiset to remove(first ionisation energy). The second electron will be more difficult to remove as being taken from K⁺ ion.

 The completey filled 3p6 sublevel has extra stability, so more energy is required to move electrons. When removing eight more electrons from the n=3 sublevel, values steadily increase as electrons removed.

Then a large jump in ionisation energy is seen as we move onto the n=2 sublevel. Finally another jump is seen as we move on to the n=1 energy level. 

-There is a steady increase in ionisation values across sublevels

-Large increase in ionisation energy whenever electron is removed from a new energy level due to filled sublevels having extra stability.