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There is also a big increase in ionization energy for removal of an electron from an inner shell (lower n value). As the effective charge increases, or as the distance of the electron from the nucleus decreases, the greater the attraction between the nucleus and the electron.
Valency of an element is defined as the number of hydrogen atoms or the number of chlorine atoms or double the number of oxygen atoms that combine with one atom of the element. Atomic volume of an element is the volume in cm3, occupied by one gram atom of the element in the solid state and hence it is also called gram atomic volume. Atomic radius is the most probable distance from the centre of the nucleus to the outermost shell of electrons.
Variation of atomic radius in a period: As we move from left to right in a period, the atomic radii goes on decreasing.
It is defined as the amount of energy required to remove one valence electron from an isolated neutral gaseous atom resulting in the formation of a monovalent positive ion.
The relative tendency or power of an element in a molecule to attract the shared pair of electrons towards itself is known as its electronegativity. Each Period has lower ionisation energies than the previous one (see Period 2 Li to Ne compared to Period 3 Na to Ar) - this is because another shell has been added, making the electrons further from the nucleus and also shielding them more. The general trend within a period is up - this is because there is no additional shielding, the nuclear charge is increasing and (don't forget to mention) atomic size decreases across a period meaning that the electrons are getting closer to the nucleus. There is a drop from Be to B (and Mg to Al) - in both cases because the electron in Be (and Mg) came from an s orbital and in B (and Al) it came from a higher energy p orbital. There is a drop from N to O (and P to S) this is because the electronic structure of N ends with 2p3 and for O it is 2p4 - the fourth electron in the p subshell is forced to spin-pair.
You won't be asked about ionisation energies beyond element 18 because the trends within the D Block is less easily explained. A polar covalent bond is one in which one atom has a greater attraction for the electrons than the other atom. A quantity termed 'electronegativity' is used to determine whether a given bond will be nonpolar covalent, polar covalent, or ionic. T HE P ERIODIC T ABLE We can use the electron configuration to split up the valence electrons into sub-levels. E FFECTIVE N UCLEAR C HARGE In any atom the nucleus exerts an attractive force on the electrons. T RENDS IN I ONIC R ADII The positive ions are smaller than the negative ions, as the former have only two occupied electron shells and the latter have three. I ONIZATION E NERGY The first ionization energy of an element is the energy required to remove one mole of electrons from one mole of gaseous atoms.
T RENDS IN I ONIZATION E NERGIES There are some small exceptions to the increasing trend across a period: Ionization energy for a p sub-shell is lower than for an s sub-shell. T RENDS IN I ONIZATION E NERGIES ? Down a group ionization energy decreases as the outer electron is further from the pull of the nucleus.

E LECTRONEGATIVITY Electronegativity is the ability of an atom to attract electrons in a covalent bond. Melting points of alkali metals (group 1) decrease down the group as the metallic bonds weaken – valence electrons are further from the nucleus so the attraction between the delocalized electrons and the positive ions decreases.
M ELTING P OINTS Melting points will increase with stronger bonding and intermolecular forces. M ELTING POINTS ACROSS A PERIOD Across a period the bonding changes from metallic (strong) to giant covalent (very strong) to van der Waals forces between molecules (weak). C HEMICAL P ROPERTIES Chemical properties of an element are largely dependent on the number of electrons in the outer shell. T HE A LKALI M ETALS, GROUP 1 Reactivity increases down the group as the valence electron is further from the attraction of the nucleus and ionization energy decreases. Its valence electrons have a higher principle quantum number, and are therefore further away from the nucleus, than C or B. These elements are called transition elements as they show transitional properties between S and P-block elements.These elements contain partially filled d-orbitals and hence they are called d-block elements.
For example, the atoms of all the elements of group 1 (alkali metals) have one electron in their valence shell while those of group 2 (alkaline earth metals) have two electrons in their outermost shell. Elements having the valency of 1,2,3 and 4 are called monovalent, divalent, trivalent and tetravalent respectively.The valency of representative elements is usually equal to the number of electrons in the outermost orbit or group number 10. This means that 349 KJ of energy is released when one mole of chlorine atoms change into Cl- ions. This explains the big difference between the ionic radii of the Si 4+ and Si 4- ions and the discontinuity in the middle of the table. This is because p orbitals are slightly higher in energy than s orbitals (in the same period). Melting points of halogens (group 7) increase down a group as van der Waals forces increase with molar mass. This means that groups tend to have similar chemical properties - they react in a similar way. Thus, alkali metals have ns1 configuration while alkaline earth metals have ns2 configuration. This means that 496 KJ of energy would remove an electron from each Na atom in one mole of Na atoms. Explanation for the first four trends should be given in terms of the balance between the attraction of the nucleus for the electrons and the repulsion between electrons. The effective charge increases with the nuclear charge as there is no change in the number of inner electrons. The decrease in ionic radius is due to the increase in nuclear charge with atomic number across the period. Number of protons increases across period 3 so effective nuclear charge increases and ionization energy increase with it.

There is also a decrease from the 5th element to the sixth as the p sub-shells start to be doubly filled. Electronegativity increases from left to right across a period owing to the increase in nuclear charge, resulting in an increased attraction between the nucleus and the bond electrons. The melting points increase and then decrease accordingly with the changes in strength of bonding. In the long form of the periodic table, the elements are arranged in order of increasing atomic number.
The effective nuclear charge experienced by an atom’s outer electrons increases with the group number of the element. The increased attraction between the nucleus and the electrons pulls the outer shell closer to the nucleus.
It is easier to remove the 6 th electron as it is repelled by its partner whereas the 5 th electron is not paired so it takes more energy to remove it. The bonding changes from metallic (Na, Mg and Al) to giant covalent (Si) to weak van der Waals forces between molecules (P 4, S 8 and Cl 2 ) and single atoms (Ar). They have complete outer shells of electrons so have the highest ionization energies for each period.
In the IB data booklet, the representative groups in the Periodic Table are numbered from1 to 7 and the last column is labeled as “0”. Across period attraction between the nucleus and the outer electrons increases as the nuclear charge increases so electrons are pulled in more and atomic and ionic radii decrease.
Therefore, as we move across a period, each element acquires a new electronic configuration as shown below in the case of second period elements. Compounds of xenon, krypton and argon have been made but it requires special conditions to create these. It may be noted that in the case of noble gases, we measure van der Waal's radii instead of covalent radii since these elements do not form covalent bonds. There is increased electron-electron repulsion in the outer shell so they move farther apart and increase the radius of the outer shell.
Variation in a PeriodAs we move across a period, the number of valence electrons increases. For example, sodium has one valence electron (3s1), Magnesium has two valence electrons (3s2) while aluminum has three valence electrons (3s23p1). As a consequence, the valency of the elements changes in a period as we move from the left towards the right.

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