A Levels - Periodic Table Guide

The periodic table is one of the most powerful tools in chemistry, organising the elements in a way that reveals patterns in their properties and behaviors. For A-Level H2 Chemistry students, a deep understanding of the periodic table is essential, as it underpins much of the subject’s advanced concepts, from atomic structure to chemical bonding and reactivity trends.

In this article, let’s delve deeper into the trends and variations of atomic, physical, and chemical properties.

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Trends and Variations in Electronic Configuration

The modern periodic table is arranged in order of increasing atomic number (the number of protons in an atom’s nucleus). The elements are organized into rows called periods and columns called groups.

Across a period, elements have

Same number of outermost quantum shell
Similar core electrons (e.g., Na: 1s22s22p63s1, Al: 1s22s22p63s23p1)
Periodic changes in physical and chemical properties

Down a group, elements have

Same number of valence electrons (e.g., Li: 1s22s1, Na: 1s22s22p63s1)
Similar physical and chemical properties

Several key periodic trends emerge from the arrangement of elements in the periodic table, and these trends are critical for understanding the chemical behavior of elements.

Trends and Variations in Atomic and Physical Properties

There are a few factors that commonly affect the atomic and physical properties of elements, such as atomic and ionic radius, electronegativity, and ionisation energy.

Number of filled quantum shells
Nuclear charge, Z

Positive charge of nucleus due to positively charged protons, which provide attraction on negatively charged electrons

Shielding effect by inner shells of electrons

Valence electrons shielded from attraction of nucleus by inner shells of electrons, thus full attractive force is not experienced

The net attractive force of the nucleus on the valence electrons, also termed as effective nuclear charge, Zeff, depends on both the nuclear charge and the screening effect.

Zeff = Z – screening effect

Let’s see how we can apply these factors to explain atomic and physical properties of elements across the periodic table.

Atomic Radius

Down the group	Across the period
↑ number of protons, ↑ nuclear charge (Z)	↑ number of protons, ↑ Z
↑ number of inner shell of electrons, ↑ shielding effect	Number of inner shell of electrons is constant, shielding effect relatively constant
↑ screening effect outweighs ↑ Z as valence electrons are further away from the nucleus	↑ effective Z
↓ effective Z	↑ effective Z
↓ strength of electrostatic attraction between nucleus and valence electrons	↑ strength of electrostatic attraction between nucleus and valence electrons
↑ atomic radius	↓ atomic radius

Electronegativity

Electronegativity is the relative tendency of an atom to attract electrons in a covalent bond.

Down the group	Across the period
↑ number of protons, ↑ Z	↑ number of protons, ↑ Z
↑ number of inner shell of electrons, ↑ shielding effect	Number of inner shell of electrons is constant, shielding effect relatively constant
↑ screening effect outweighs ↑ Z as valence electrons are further away from the nucleus	↑ effective Z
↓ effective Z	↑ effective Z
↓ strength of electrostatic attraction between nucleus and valence electrons	↑ strength of electrostatic attraction between nucleus and valence electrons
↑ atomic radius	↓ atomic radius
↓ tendency to attract a bonding pair of electrons	↑ tendency to attract a bonding pair of electrons
↓ electronegativity	↑ electronegativity

Ionic Radius

Down the group	Across the period
↑ number of protons, ↑ Z	↑ number of protons, ↑ Z
↑ number of inner shell of electrons, ↑ shielding effect	Number of inner shell of electrons is constant, shielding effect relatively constant
↑ screening effect outweighs ↑ Z as valence electrons are further away from the nucleus	↑ effective Z
↓ effective Z	↑ effective Z
↑ atomic radius	↓ atomic radius
↓ strength of electrostatic attraction between nucleus and valence electrons	↑ strength of electrostatic attraction between nucleus and valence electrons
↓ energy required to remove valence electrons	↑ energy required to remove valence electrons
↓ 1st IE	↑ 1st IE
	Take note of the following abnormalities: From Group 2 to 13 (e.g., Be to B – 1s22s2 to 1s22s22p1) Less energy required to remove a 2p electron than a 2s electron as the 2p electron is further from nucleus From Group 15 to 16 (e.g., N to O – 1s22s22p3 to 1s22s22p4) Less energy required to remove an electron from paired electron than unpaired electron due to electrostatic repulsion between paired electrons

Cation	Anion
Atoms lose some/all valence electrons to form cations	Electrons added to outermost quantum shell of atoms to form anions
Cations have 1 less quantum shell of electrons than neutral atom	Increase electron-electon repulsion in outermost quantum shell of electrons
Ionic radii of cations SMALLER than atomic radii of corresponding atoms	Ionic radii of anions LARGER than atomic radii of corresponding atoms

Isoelectronic Series (e.g., Na+ to Si4+, P3- to Cl-)

↑ number of protons, ↑ Z

Number of inner shell of electrons is constant, shielding effect relatively constant

↑ effective Z

↑ strength of electrostatic attraction between nucleus and valence electrons

↓ ionic radius

First Ionisation Energy, IE

First Ionization Energy is defined as the amount of energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous singly charged positive ions.

X(g) → X+(g) + e−

Down the group	Across the period
↑ number of protons, ↑ Z	↑ number of protons, ↑ Z
↑ number of inner shell of electrons, ↑ shielding effect	Number of inner shell of electrons is constant, shielding effect relatively constant
↑ screening effect outweighs ↑ Z as valence electrons are further away from the nucleus	↑ effective Z
↓ effective Z	↑ effective Z
↑ atomic radius	↓ atomic radius
↓ strength of electrostatic attraction between nucleus and valence electrons	↑ strength of electrostatic attraction between nucleus and valence electrons
↓ energy required to remove valence electrons	↑ energy required to remove valence electrons
↓ 1st IE	↑ 1st IE
	Take note of the following abnormalities: From Group 2 to 13 (e.g., Be to B – 1s22s2 to 1s22s22p1) Less energy required to remove a 2p electron than a 2s electron as the 2p electron is further from nucleus From Group 15 to 16 (e.g., N to O – 1s22s22p3 to 1s22s22p4) Less energy required to remove an electron from paired electron than unpaired electron due to electrostatic repulsion between paired electrons

Melting and Boiling Point

Electrical Conductivity

Volatility

Volatility is the tendency to vapourise. It depends on the boiling point of substance.

Down Group 17

Halogens have simple covalent structure
(exist as discrete molecules held together by weak instantaneous dipole-induced dipole interactions)

↑ number of electrons

↑ strength of instantaneous dipole-induced dipole interactions

↑ energy required to overcome interactions

↑ BP

↓ volatility

Trends and Variations in Chemical Properties

Apart from atomic and physical properties, there are also trends and variations in the chemical properties of the elements across the Periodic Table. Particularly, we are interested in the Period 3, Group 2, and Group 17 elements.

Chemical Properties of Chlorides of Period 3 Elements

Chloride	NaCl	MgCl2	AlCl3	SiCl4	PCl5
Chemical bonding	Ionic	Ionic (with covalent character)	Covalent (with ionic character)	Covalent
Structure	Giant ionic		Simple covalent
MP/BP	High		Low
Oxidation Number	+1	+2	+3	+4	+5

Chloride	Reaction with Water		pH
NaCl	Hydration occurs through ion-dipole interactions to form neutral solution Does not hydrolyse in water (no reaction with water) due to low charge density of Na+	Hydration: NaCl(s) → Na+(aq) + Cl–(aq)	7
MgCl2	Hydration occurs through ion-dipole interactions Undergoes slight hydrolysis in water to form slightly acidic solution due to higher charge density of Mg2+	Hydration: MgCl2(s) + 6H2O(i) → [Mg(H2O)6]2+(aq) + 2Cl–(aq) Hydrolysis: [Mg(H2O)6]2+(aq) + H2O(l) [Mg(H2O)5(OH)]+(aq) + H3O+(aq)	6.5
AlCl3	Hydration occurs through ion-dipole interactions Undergoes substantial hydrolysis in water to form acidic solution due to higher charge density of Al3+ → able to polarise water molecules → weaken O-H bonds in water molecules → produce H3O+	Hydration: AlCl3(s) + 6H2O(l) → [Al(H2O)6]3+(aq) + 3Cl–(aq) Hydrolysis: [Al(H2O)6]3+(aq) + H2O(l) [Al(H2O)5(OH)]2+(aq) + H3O+(aq)	3
SiCl4	Hydration does not occur Undergoes complete hydrolysis in water to form very acidic solution	Hydrolysis: SiCl4(l) + 2H2O(l) → SiO2(s) + 4HCl(aq)	2
PCl5	Hydation does not occur Undergoes complete hydrolysis in water to form very acidic solution	Hydrolysis: PCl5(l) + 4H2O(l) → H3PO4(s) + 5HCl(aq)	2

Chemical Properties of Oxides of Period 3 Elements

Oxide	Na2O	MgO	Al2O3	SiO2	P4O10	SO3
Chemical bonding	Ionic		Ionic (with covalent character)	Covalent
Structure	Giant ionic			Giant covalent	Simple covalent
MP/BP	High				Low
Type of oxides	Basic		Amphoteric	Acidic

Oxide	Reaction with Water		pH
Na2O	Readily dissolves in water to give alkaline solution	Na2O(s) + H2O(l) → 2NaOH(aq)	13
MgO	Dissolves slightly in water to give a weak alkaline solution	MgO(s) + H2O(l) Mg(OH)2(aq)	9
Al2O3	Insoluble in water due to high lattice energy	–	–
SiO2	Insoluble in water due to high lattice energy	–	–
P4O10	Readily dissolves in water to give an acidic solution	P4O10(s) + 6H2O(l) → 4H3PO4(aq)	2
SO3	Readily dissolves in water to give an acidic solution	SO3(g) + H2O(l) → H2SO4(aq)	2

Oxide	Reaction with Acid/Base
Na2O	React with acid to form salt and water Does not react with alkali	Na2O(s) + 2HCl(aq) → 2NaCl(aq) + H2O(l)
MgO	React with acid to form salt and water Does not react with alkali	MgO(s) + 2HCl(aq) → MgCl2(aq) + H2O(l)
Al2O3	React with BOTH acid and alkali to form salt and water	Al2O3(s) + 6HCl(aq) → 2AlCl3(aq) + 3H2O(l) Al2O3(s) + 2NaOH(aq) + 3H2O(l) → 2NaAl(OH)4(aq)
SiO2	React with concentrated alkali to form salt and water	SiO2(s) + 2NaOH(aq) → Na2SiO3(aq) + H2O(l)
P4O10	React with alkali to form salt and water	P4O10(s) + 12NaOH(aq) → 4Na3PO4(aq) + 6H2O(l)
SO3	React with alkali to form salt and water	SO3(g) + 2NaOH(aq) → Na2SO4(aq) + H2O(l)

Chemical Properties of Group 2 Elements

Chemical Property	Trend	Explanation
Reducing Power	Increases down the group	↑ number of protons, ↑ Z ↑ number of inner shell of electrons, ↑ shielding effect ↑ screening effect outweighs ↑ Z as valence electrons are further away from the nucleus ↓ effective Z ↑ atomic radius ↓ strength of electrostatic attraction between nucleus and valence electrons ↓ energy required to remove valence electrons ↓ 1st IE ↑ tendency of metal to be oxidised EӨ(M2+/M) more negative ↑ reducing power
Thermal Stability of Group 2 Compounds	Increases down the group	↑ catonic radii of M2+ ↓ charge density (charge to size ratio) ↓ ability to polarise electron cloud of large anion (e.g., CO32-) Bonds weakened to smaller extend ↑ thermal stability

Chemical Properties of Group 17 Elements

Chemical Property

Trend

Explanation

Oxidising Power

Decreases down the group

- ↑ number of protons, ↑ Z
- ↑ number of inner shell of electrons, ↑ shielding effect
- ↑ screening effect outweighs ↑ Z as valence electrons are further away from the nucleus
- ↓ effective Z
- ↓ strength of electrostatic attraction between nucleus and valence electrons
- ↑ atomic radius
↓ tendency to attract a bonding pair of electrons

↓ electronegativity
↓ ease of gaining electrons to form halide ions during oxidation
EӨ less positive
↓ oxidising power

Thermal Stability of Hydrogen halides

Decreases down the group

- ↑ atomic radius
- ↑ diffusion of valence orbitals
- ↓ effectiveness of orbital overlap between smaller H and larger X atoms
- ↓ energy required to break H-X bond
↓ thermal stability

Final Comments

‘Periodic Table’ is a heavy topic in ‘A’ Levels H2 Chemistry, incorporating content from across different topics. However, it is definitely not difficult once you are able to understand the main group and period patterns associated with the periodic table. Do remember to practice with examples and regularly revisit periodic trends to reinforce your understanding. Best of luck, and may your enthusiasm for chemistry continue to inspire you in all your future pursuits!

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Prepared by: Tan Wee Leng

This article is written based on the 9729 Chemistry GCE Advanced Level H2 Syllabus. https://www.seab.gov.sg/docs/default-source/national-examinations/syllabus/alevel/2024syllabus/9729_y24_sy.pdf

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