Periodicity T3

Periodicity: the regular repetition of chemical and/or physical properties of elements arising from patterns in their electron configuration.

• Column are groups and rows and periods

(Elements ending electron configuration letter depends on the group they are in)

• Metals are on the left hand of the periodic table 

• Non-metals (NM) are on the right hand of the periodic table 

• Metalloids are what have both metallic and NM properties

1. Effective Nuclear Charge: How strong the nucleus holds on to the valence electron, how strong the nucleus is attracting the electrons.   

• Decreases down the group

• Stays the same across the period left to right

EX: Lithium (no.3) and francium (no.87) both in group 1. Hence have 1 valence electron.  

Li would have High electron nuclear charge as there are less electrons shielding the valence electron. 

• Valence electrons are added to the same energy level left to right

EX: all elements from Li to Fr will have the same energy level

2. Atomic Radius: Distance between the nucleus the outermost electron

• Increase down a group

We add electrons as we go to higher energy levels, increasing the atomic radius down the group. 

• Decreases from left to right across the period 

Due to electrons being added to the same energy level, it pulls the electrons closer to itself, decreasing the atomic radius across a period from left to right. 

3. Ionic Radius: Distance between nucleus of an ion and its outermost shell

• Metals lose e- to become cations (+ ions)

• NM gain e- to become anions (- ions)

• Cations lose more electrons than protons, making electrostatic force between the nucleus greater, hence ions are smaller than its parent atom. 

• Anions have more electrons than protons, making less repulsion causing the radius to spread out more, hence are always larger than its parent atom. 

• Positive ions decrease from 1-14 bc were losing more valence electrons causing a smaller ionic radius, for negative ions there is a spike at 14 and then continues to decrease. 

4. Ionization Energy: Energy it takes to lose one mole of electrons from one mole of gaseous atoms in their ground state. 

• Metals have low ionization energy as they easily give away electrons

• NM have high IE as they gain electrons so it takes a lot more energy to take an electron.

• Increase across a period from left to right (Going from Metals to NM)

• Decreases down a group due to increased  shielding, being farther away from the nucleus

5. Electron Affinity: Energy change when an electron is added to an isolated atom in gaseous state. (How much atoms like electrons)

• Increase period from left to right

• Mainly due to NM wanting to gain electrons

• Decreases down a group as they is more shielding

6. Electronegativity: Measure of attraction that an atom has for a pair of electrons covalently bonded to another atom. (How strongly atoms will attract electrons)

• Increase period from left to right

• Mainly due to NM wanting to gain electrons

• Decreases down a group as they is more shielding

• Most electronegative element is F (Fluorine)

• Least electronegative element is Cs (Cesium)

7. Melting points 

Group 1

• Melting point Decreases going down the group as atomic radius is larger

• Larger radius means atoms are farther apart from each other,making it easier to break resulting in lower boiling point

Group 17-Halogens

• They form molecules formed by covalent bonds (Have intermolecular forces)

• The dispersion forces increase as radius increases

• Hence melting point increases down the group

Group 18

• 8 VE-

• Stable due to octet rule

• Non reactive 

• Colorless and odourless

Group 17

• 7 VE- (Very reactive as they only need 1 electron)

• Reactive (Often react with grp1 metals to form halides like Nacl, NaBr, etc)

• Diatomic (EX: F2, CL2, I2)

• Colored