Unit 1: Periodic Table and Periodic Properties — Long Questions

Creation of the periodic table
The reasons for the creation of periodic table are:
• Symbol of chemistry: The periodic table of elements is considered as the symbol of chemistry.
• Systematic categorization: The systematic categorization and arrangement of elements in periodic table provides easy way to study the wide subject of chemistry.
• Periodicity: The main reason for the creation of periodic table is periodicity of elements and their compounds. Periodicity means that properties of arranged elements are repeated after regular intervals.
• Turning point: One of the most important turning point in the history of science is the creation of periodic table.
• Tabular form: All the 118 elements are arranged in tabular form in the current periodic table based on their atomic number, electronic configuration, and recurrent chemical characteristics.
• Framework: The periodic table offers a framework for researching the periodic behaviours of elements is a significant accomplishment.

Historical background of Periodic table
(i) Dobereiner's Triads
In 1829, Dobereiner grouped the elements into triads (a group of three elements) with similar properties, noticing that the atomic weight of the middle element was roughly the average of the other two.
Examples: Lithium, sodium and potassium (⁷Li, ²³Na, ³⁹K) are included in the same triad.
(ii) Newland's law of Octaves
• English chemist John Newland, in 1864, first time observed periodicity in the 62 known elements.
• He studied that the properties of every eighth element were similar when arranged by the increasing order of their atomic masses.
• He classified the elements into groups so that every eight element resembled the first element in properties.
(iii) Lother Meyer's curves
Lother Meyer developed his famous curves by plotting a graph between the atomic weights and atomic volumes of elements in 1869. Theses curves also showed periodicity.
(iv) Mendeleev's Periodic table
Russian chemist Dmitri Mendeleev considered the father of the Periodic Table. He arranged 63 elements in his periodic table named Mendeleev's Periodic table in 1869. Some features of his table are:
Mendeleev's periodic law: The properties of the elements are periodic functions of their atomic masses.
• Groups: The vertical columns of elements in the periodic table are called groups. He arranged elements in eight groups in his periodic table.
• Periods: The horizontal rows of elements in the periodic table are called periods. He arranged elements in seven periods.
• Gaps: The success of his table was hidden in leaving gaps for undiscovered elements and predicting their atomic masses and properties which proved accurate when these elements were practically found.
(v) Modern Periodic table
In 1913, Moseley determined the exact atomic numbers of known elements using X-ray emission, resolving flaws and discrepancies in Mendeleev's table by arranging the elements by atomic numbers instead of atomic masses.
The Modern Periodic law: The modern periodic law states that "the properties of elements are periodic functions of their atomic numbers".

Main features of the Modern Periodic table
The classification of elements in the modern periodic table helps in the easier understanding of their properties.
Atomic number: All the 118 elements are arranged in groups and periods in modern period the table in ascending order of their respective atomic numbers.
Groups and Periods: There are seven horizontal rows called periods and eighteen vertical columns called groups. (In older versions of the table, there were 18 vertical groups were divided into two types of groups: Eight A-Groups and Ten B-Groups).
Properties: In the periodic table, elements within the same group exhibit similar chemical properties because they have the same number of valence electrons. However, they show a gradual change in physical properties from top to bottom in a group.
Valance electrons: Elements in a period show a gradual change in properties moving from left to right in periods, due to variation in number of valence electrons.
Representative and Transition elements: Main group elements (group 1, 2, 13 – 18) and transition elements (group 3 – 12).
Metals, Non-metals and Metalloids: Most of the elements are metals in periodic table, some of them are non-metals, only few of them are metalloids.
Block and Families: Other than groups and periods in the periodic there are different ways of grouping the elements into various blocks, families and categories just to enhance understanding.

Classification of elements
Elements can be broadly classified as metals, non-metals and metalloids.
(i) Metals
Definition: Metals are elements which tend to lose electrons to form positive ions.
Location: Elements to the left of the stair – step line in periodic table are considered as metals.
Examples: Iron, Copper, Gold and Silver etc. are metal elements in periodic table.
(ii) Non-Metals
Definition: Non-metals are elements which tend to gain electrons to form negative ions.
Location: Elements to the far right side of the stair – step line in the periodic table are non-metals. The exception is hydrogen, the first element of the periodic table.
Examples: Chlorine, Sulfur and Phosphorous, Carbon and Hydrogen etc. are non-metal elements in periodic table.
(iii) Metalloids
Definition: The metalloids exhibit some properties of metals and some of non-metals.
Location: The metalloids separate the metals and non-metals on a periodic table. Mostly periodic tables have a "stair-step line" on the table identifying the element groups. The line begins at boron (B) and extends down to polonium (Po) including Si, Ge, As, Sb and Te.
Examples: Si, Ge, As and Sb etc. are metalloid elements in the periodic table.

BLOCKS IN PERIODIC TABLE
Blocks in Periodic table
Definition: Elements in the periodic table can be classified based on the subshells containing their valence electrons.
Example: The blocks of periodic table are given in the following:
s-Block
• The valence electrons of elements in the first two groups (1 and 2) are in the "s" subshells, so these elements are placed in the s-block.
• It is present at extreme left side of the periodic table.
p-Block
• The elements in groups 13 to 18 having their valence electrons in their "p" subshells so they are included in p-block.
• It is located at the right side of the periodic table.
d-Block
• The elements in groups 3 to 12 having "d" subshells in the process of filling so they are included in d-block, and they are called transition elements.
• Transition means in between two things, as these elements are located in between s and p blocks so they are given the name transition elements.
f-Block
• The elements in the two series at the bottom of the table (known as Lanthanides and Actinides) are categorized as f-block elements.
• Their "f" subshells are in the process of filling.

Families in Periodic table
Definition: An element family is a set of elements sharing common properties are placed in the same group of periodic table.
There are six famous families of elements in the periodic table:
• Alkali metals (Li, Na, K, Rb, Cs, Fr)
• Alkaline earth metals (Be, Mg, Ca, Sr, Ba: Ra)
• Transition metals (Sc, Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn etc.)
• Chalcogens (O, S, Se, Te, Po, Lv)
• Halogens (F, Cl, Br, I, At, Ts)
• Noble gases (He, Ne, Ar, Kr, Xe, Rn, Og)
(i) Alkali Metals
Definition: Elements in the group 1 of the periodic table are known as alkali metals because they produce alkalis when they react with water.
Example: Sodium and potassium are notable examples of these elements.
Characteristics
• Alkali metals are characterized by one valence electron.
• Alkali metals show low densities, relatively low melting points, and low ionization energies.
• These are the most reactive metals.
• They are soft metals.
(ii) Alkaline Earth Metals
Definition: Group 2 elements are metals primarily found in the earth and form alkalis, hence they are referred to as alkaline earth metals.
Example: Calcium and magnesium are notable the examples of alkaline earth metals.
Characteristics
• These elements have two electrons in their valence shell, making them divalent.
• They are metallic solids that are harder and denser than alkali metals.
• Easily oxidized with high thermal and electrical conductivities.
• They are soft metals but harder than alkali metals.
(iii) Transition Elements
Definition: The transition metals make up the largest family of elements in the middle of periodic table.
Example: They include four series of d-block elements (groups 3 to 12), as well as the lanthanides and actinides (f-block elements) found in the two rows at the bottom of periodic table.
Characteristics
• They exhibit high thermal and electrical conductivities, high melting points, high density, and variable oxidation states.
• They mostly form coloured compounds.
• They form complex compounds.
(iv) Chalcogens
Definition: The group 16 elements are called chalcogens because most ores of copper (Greek Chalko= ores, gen = forming) are oxides or sulphides.
Example: In this group, oxygen & sulphur are non-metals, Se, Te, Po are metalloids and Livermorium is a metal.
Characteristics
• They have six valance electrons.
• They mostly show – 2 oxidation state.
(v) Halogens
Definition: Elements in group 17, known as halogens are non-metallic. The term "halogen" means "salt-former" because these elements easily react with alkali metals and alkaline earth metals to form stable halide salts.
Example: Fluorine, chlorine, bromine and iodine etc. are halogens.
Characteristics
• Halogens are highly reactive non-metals with high electron affinities.
• Halogens can easily accept one electron to complete their outermost shell.
(vi) Noble gases
Definition: The noble gases are a group of unreactive elements present at the extreme right of the periodic table in Group 18.
Example: Helium, argon, krypton etc. are noble gases.
Characteristics
• Stable electronic configuration (complete outermost shell).
• They are almost entirely unreactive under normal conditions and rarely form compounds with other elements.
• These elements are mono-atomic in nature.

Electronic configuration
Definition: The arrangement of electrons in sub-shells or orbitals around the nucleus in an atom is called electronic configuration.
Periodic arrangement: The periodic arrangement of elements in the periodic table provides information about the physical properties, such as their physical state and atomic radii as well as their electronic structure and chemical reactivity.

Electronic configuration helps in the following ways in arrangement of periodic table.
• Period Number: The period number indicates the principal quantum number (n), representing the number of electronic shells surrounding the nucleus.
Example: An element X in the 3rd period have three shells, with its valence electrons located in the 3rd shell.
• Sub-shell and block: The specific subshell where the valence electrons are found, depends on the element's block (azimuthal quantum number). If an element X in the 3rd period is in the s-block, its valence electrons are in the 3s subshell.
• Group Number: The group number indicates the number of valence electrons.
Example of Mg: An element X in the 3rd period and group 2 has two valence electrons in its outermost shell. Thus, the element Z in the 3rd period and group 2 (s-block) has two valence electrons in the 3s subshell, which means that X would be magnesium (Mg).
Example of Al: Another example to relate period number and group number with electronic configuration and position of element in period table.
X belongs to group 13 and period 3
The element X belongs to group 13 of periodic table so it has 3 valence electrons; and it is found in period 3 so it has three shells around its nucleus.
It means that the 3 valence electron are in the 3rd shell which means element is Al.
The configuration will be: 1s² 2s² 2p⁶ 3s² 3p¹

1st shell 2nd shell 3rd shell
• Chemical Properties: Understanding the periodic arrangement of elements provides an explanation of an element's electronic configuration, which is essential for understanding its chemical properties and behavior.

Atomic radius
Definition: The atomic radius is a measure of the size of an atom, it is half of the distance between two identical atoms bonded together.
• The atomic radius can vary depending on the type of bond (covalent, Metallic & van der Waal's force) or the state of the atom.
• The radius can be different in a covalent bond compared to an ionic bond.
Example: The atomic radius of sodium atom is 186pm.
Unit: The atomic radius is typically measured in picometers (pm) or Angstroms ( Å ).
Factors affecting atomic radius
The factors affecting the atomic radius are
(i) Atomic number
(ii) Effective nuclear charge
(iii) Shielding effect of inner electrons.
Periodic trends in atomic radius:
Variation in periods
Atomic radius decreases across a period (from left to right) in the periodic table.
Reason: Increasing nuclear charge in period from left to right, which pulls the electronic cloud closer, results in decrease in atomic radius.
Variation in group
Atomic radius increases down a group (from top to bottom) in the periodic table.
Reason: The additional electronic shells are added moving down the group, so more shielding makes the atom larger despite the increase in nuclear charge (which is outweighed), as a result atomic radius increases.

IONIC RADIUS
Ionic radius
Definition: It is defined as the distance form the nucleus of an ion to the outermost electron shell.

Unit It is measured in picometers (pm) or angstroms ( Å ).
Example: The cationic radius of K⁺¹ is 133 pm.
Cationic radius: When an atom loses one or more electrons to become a positive ion, it generally becomes smaller than the neutral atom.
Reason: The loss of electrons reduces electronic repulsion and allows the remaining electrons to be pulled closer to the nucleus, so cationic radius is generally smaller than the parent atom.
Anionic radius: When an atom gains one or more electrons, it forms an anion, it generally becomes larger than the neutral atom.
Reason: The addition of electrons increases electronic repulsion, as a result the nuclear pull on electrons decreases and the electron cloud expands, so the anionic radius is larger than its parent atom.
Periodic trends in ionic radius:
Variation in period
When we move across a period from left to right, the ionic radius of cations and anions decreases.
Reason: Increasing nuclear charge which pulls the electrons closer for cations. The anionic radius also decreases across a period because the increasing nuclear charge also pulls the electrons closer to the nucleus.
Variation in group
Both cations and anions increase in size as we move down a group.
Reason: The principal quantum number (n) increases, leading to an increase in the number of electron shells. Consequently, the distance between the nucleus and the outermost electrons becomes larger, outweighing the effect of increased nuclear charge. The additional electron shells make the ions larger.

Definition of 1st Ionization Energy: The energy needed to remove one electron from each atom in one mole of atoms of the element in the gaseous state to form one mole of gaseous 1+ ions is known as 1st ionization energy (ΔHi1).
Example: 1st Ionization Energy
Na₍ₐ₎ ─→Na⁺₍ₐ₎ + e⁻ ΔHi1 = 496 kJ mol⁻¹
Ca₍ₐ₎ ─→Ca⁺₍ₐ₎ + e⁻ ΔHi1 = 590 kJ mol⁻¹
Definition of 2nd Ionization Energy: If a second electron is removed from each ion in a mole of gaseous 1+ ions, it is called the 2nd ionization energy (ΔHi2).
Example: 2nd Ionization Energy
Ca⁺₍ₐ₎ ─→Ca²⁺₍ₐ₎ + e⁻ ΔHi2 = 1150 kJ mol⁻¹
Definition of 3rd Ionization Energy: If third each electron is removed from each ion in a mole of gaseous 2+ ions, it is called the 3rd ionization energy (ΔHi3).
Example: 3rd ionization energy
Ca²⁺₍ₐ₎ ─→Ca³⁺₍ₐ₎ + e⁻ ΔHi3 = 4940 kJ mol⁻¹
An element can have several ionization energies; the exact number corresponds to its atomic number.
Factors affecting the ionization energy
The magnitude of the ionization energy of an element depends upon the following factors;
(i) Nuclear charge
Greater the effective nuclear charge, greater is the electrostatic force of attraction, more difficult is the removal of an electron from the atom. So, ionization energy increases with an increase in the effective nuclear charge.
(ii) Size of the atom or ion
In bigger atoms force of attraction between the nucleus and the outermost electrons is weaker. Therefore, the ionization energy decreases as the size of the atom increases and vice-versa.
(iii) Electronic arrangement
It is observed half-filled and completely-filled orbitals are found to be more stable. Therefore, the ionization energy is higher when an electron is to be removed from a fully-filled or half-filled shells.
• Noble gases have highest ionization energies in their respective periods. It is due to highly stable fully-filled shells (ns²npⁿ).
• Oxygen has lower ionization energy than nitrogen due to the half filled sub-shell of nitrogen.
The electronic configurations of oxygen and nitrogen are:
₇N = 1s² 2s² 2p₁ᵞ 2p₁ᵞ 2pᶻ₁ (more stable) ΔHi1 = 1403 kJmol⁻¹
₈O = 1s² 2s² 2pₓ² 2pᵞ¹ 2pᶻ¹ (less stable) ΔHi1 = 1365 kJmol⁻¹
(iv) Shielding effect
Greater the shielding effect, easier it is to remove the valence electrons from an atom. Larger the number of inner electrons, greater is the shielding effect, therefore, lower is the ionization energy.
(v) Spin-Pair Repulsion
When electrons are spin-paired in the same orbital, the repulsion between them can lead to a slightly lower ionization energy compared to removing unpaired electron.
Reason: The paired electrons experience increased repulsion, making it slightly easier to remove one of the paired electrons.
Manganese (Mn) has two spin-paired electrons in its 4s orbital. The ionization energy to remove one of these paired electrons is relatively lower due to increased repulsion between the electron pairs. In contrast chromium (Cr) has one unpaired electron in its 4s orbitals. Removing one of these unpaired electrons requires more energy due to the absence of spin-pairing repulsion.
₂₄Cr = [Ar] 3d⁵4s¹ (more stable)
₂₅Mn = [Ar] 3d⁵4s² (less stable)
Periodic trends in ionization energy
Variation in groups
Going down in a group ionization energy decreases.
Reason: The nuclear charge increases but as the size of the atom and the number of electrons causing the shielding effect also increase therefore ionization energy decreases from top to bottom.
Group 1: In Group 1, the ionization energies decrease in the following order: Li> Na> K>Rb>Cs.
Example: The 6s valence electron of Cs is farther from the nucleus and thus easier to remove compared to the 5s valence electron of Rb.
Variation in period
When we move from left to right across a period the ionization energy generally increases.
Reason: Number of shells remains unchanged while the effective nuclear charge increases, making it more difficult to remove an electron. Although the number of electrons also increases across a period, the shielding effect within the same shell is same so not considered. So, the ionization energy increases.
Graph
• The trend of ionization energies of period (1-3) is shown in Figure.
• The figure also reveals that noble gases have the highest values of ionization energy due to complete outermost shell in them, the removal of electron is extremely difficult.
• Alkali metals have lowest values of ionization energy.
Abnormality: Group number 13, 16 elements show abnormality in general trend. Actual trend of ionization energy of second period elements is Li < B < Be < C < O < N < F < Ne.

Definition of 1st Electron Affinity: The first electron affinity, (ΔH°ea1), is the enthalpy change involved when 1 mole of electrons is added to 1 mole of gaseous atoms to form 1 mole of gaseous uni-negative ions under standard conditions.

Example of Cl (Chlorine)
Electron affinity of chlorine atom.
Cl(g) + e− ――→ Cl−(g) ΔH°ea1 = −348.8kJmol−1

This is amount of energy released when 6.02 ×10²³ atoms of chlorine in the gaseous state are converted into Cl−(g) ions. Since, energy is released, so first electron affinity carries negative sign.

Definition of 2nd Electron Affinity: The second electron affinity, ΔH°ea2, is the amount of energy required to add electrons to 1 mole of uni-negative gaseous ions to form 1 mole of gaseous 2- ions under standard conditions.

Example of oxygen When first electron is added to a neutral oxygen atom, 141kJ mol−1 energy is released.
O(g) + e− ――→ O−(g) ΔH°ea1 = −141kJmol−1

But 844 kJmol−1 of energy is absorbed on adding second electron to a uni-negative (O−) ion.
O(g) + e− ――→ O−(g) ΔH°ea2 = +844kJmol−1

The net enthalpy changes for the formation of the oxide ion (O²−) can be calculated by adding the first and second electron affinities
ΔH°ea1 + ΔH°ea2 = (−141) + (844) = +701kJmol−1

Factors affecting electron affinity

Important factors affecting the magnitude of electron affinity values of elements are as follows:

(i) Size of atom
For small sized atoms the attraction of the nucleus for the incoming electron is stronger. Thus, smaller is the size of the atom, greater is its electron affinity.

(ii) Nuclear charge
Greater the magnitude of nuclear charge of an element stronger is the attraction of its nucleus for the incoming electron. Thus, with the increase in the magnitude of nuclear charge, electron affinity also increases.

(iii) Electronic configuration of atom

Group 15 The electron affinity is low when the electron is added to a half filled sub-shell than that for partially filled one. Electron affinity values of 'N' and 'P' group-15 (V-A), atoms are very low. This is because of the presence of half-filled 'np' orbitals in their valence shell (N = 2s² 2p³, P = 3s² 3p³). These half-filled p-subshells, being very stable, have very little tendency to accept any extra electron to be added to them.

Group 18 Noble gases group-18 (VIII-A) have stable ns² np⁶ configuration and hence the atoms of these gases, do not accept any extra electron, This is evident from their positive 1st electron affinities.

Periodic trends in electron affinity

Variation in groups
Electron affinity generally decreases down the group.

Reason As the atomic size increases down the group, the larger electron cloud causes the incoming electron to experience less attraction from the nucleus.

This trend is observed in the halogens (At<I<Br<F<Cl).

Variation in periods
Electron Affinity generally increases in period from left to right.

Reason This is firstly due to increase in the nuclear charge, which attracts additional electrons more strongly and secondly due to decreasing atomic radius.

Abnormality Group number 2, 15 and 18 of periodic table show abnormality in general trend moving left to right. Actual order of electron affinity of 2nd period elements.
Ne<Be<N<Li<B<C<O<F

Definition Electronegativity is the power of an atom to attract shared pair of electrons toward itself in a molecule.

Example The electronegativity of F is 4.0 on Pauling scale.

Unit It has no unit.

Scale Linus Pauling, an American chemist, developed a scale of dimensionless electronegativity values, which range from just below one for alkali metals to a maximum of four for fluorine. Higher electronegativity values signify a stronger attraction for electrons compared to lower values.

Factors affecting Electronegativity

(i) Atomic size
A larger atomic size will result in a lower value of electronegativity. This is because electrons being far away from the nucleus will experience a weaker force of attraction.

Example The electronegativities of halogens in group 17 are in the order:
F > Cl > Br > I

(ii) Effective nuclear Charge
A higher value of the effective nuclear charge will result in a greater value of electronegativity, because an increase in nuclear charge causes greater attraction to the bonded electrons. This is why the electronegativity in a period increases from left to right. The electronegativity of Li in period 2 is 1.0 and F has a value of 4.0.

Periodic Trends in Electronegativity

Variation in periods
When we move from left to right along the period, the electronegativity increases,

Reason This is due to increasing nuclear charge and decrease in atomic size.

Variation in groups
In the groups, it decreases from top to bottom in periodic table.

Reason This is due to the increase in size by the addition of shells and increasing shielding effect.

Example The halogen group, the electronegativity value decreases from fluorine (4.0) to iodine (2.5) as shown in a part of the periodic table in figure.

• Normally metals being on the left side of the periodic table, possess lower electronegativity.
• Non-metals have higher values of electronegativity than those of metals.
• Metals are electropositive and non-metals are electronegative, respectively.

Figure provides a summary of all the variation trends in various physical properties of elements in the periodic table.

Definition

The metallic character of elements is typically their tendency to lose electrons

Explanation

• The elements on the left side of the periodic table have a greater tendency to lose their outermost electrons to achieve noble gas configuration.
• Elements on the right side of the table tend to gain electrons.
• Elements on the left side of the periodic table are metals that form positive ions.
• Elements on the right side, particularly in the right corner are nonmetals that form negative ions.
• Metallic character of an element largely depends on its valence shell electronic configuration.

Periodic Trends in metallic character

Variation in periods The metallic character of the elements decreases.

Reason The increase in nuclear charge pulls the electron could closer to the nucleus, making it more difficult for the atom to lose electrons and thereby decreasing metallic character (ease of losing electron) makes the element more reactive.

Example Cesium is far more reactive and electropositive than sodium or lithium.

Variation in groups Metallic character increases as one moves down in a group of the periodic table.

Reason The increases in atomic size and the shielding effect, which reduce the nuclear attraction on the valence electrons.

The increase in metallic character (ease of losing electron) makes the element more reactive.

Reactions of Na and Mg

(i) With water

Reaction of Sodium Sodium is more reactive than magnesium towards water.
Na reacts vigorously with water to form sodium hydroxide and hydrogen gas.
2Na(s) + 2H2O(l) ――→ 2NaOH(aq) + H2(g)

Reaction of Magnesium Mg reacts more slowly in forming magnesium hydroxide and hydrogen gas. However, magnesium reacts with steam more vigorously to make magnesium oxide and hydrogen gas.
Mg(s) + 2H2O(l) ――→ Mg(OH)2(aq) + H2(g)
Mg(s) + H2O(g) ――→ MgO(s) + H2(g)

(ii) With Oxygen

Reaction of Sodium Sodium burns in oxygen with a golden yellow flame to produce a white solid mixture of sodium oxide and sodium peroxide. Sodium is kept under kerosene oil to prevent its reaction with air. It reacts vigorously with oxygen in open air to form peroxide.
2Na(s) + O2(g) ――→ Na2O2(s)

Under special conditions like limited O2 or high temperature, sodium oxide is formed.
4Na(s) + O2(g) ――→ 2Na2O(s)

Reaction of Magnesium Magnesium burns in oxygen with an intense white flame to give white solid magnesium oxide. Its reaction with oxygen is slow in comparison with sodium.
2Mg(s) + O2(g) ――→ 2MgO(s)

(iii) With Chlorine: Chlorine reacts with both metals to give soluble salts. It reacts exothermically with sodium, golden yellow flame is seen and white solid, sodium chloride is formed. Magnesium also reacts with chlorine to give white solid, magnesium chloride.
2Na(s) + Cl2(g) ――→ 2NaCl(s)
Mg(s) + Cl2(g) ――→ MgCl2(s)

Oxides

Definition Oxides are binary compounds formed by the reaction of oxygen with other elements.

Ionic and covalent nature of oxides of period 3

• Oxides of group 1, 2 & 3 (Na2O) have more ionic character.
• These oxides exist as giant ionic lattices with strong electrostatic forces between oppositely charged ions.
• Oxides of group 4, 5, 6 & 7 (SO2) are more covalent.
• These oxides exist as covalent molecules with weak intermolecular forces.
• This transition is a result of the increasing electronegativity and decreasing ionic character.

Classification of Oxides

The classification of oxides is done into neutral, amphoteric and basic or acidic based on their characteristics.

(i) Basic oxides

Definition A basic oxide is an oxide that when combined with water gives off an alkali.

Example Na2O, CaO, BaO.

Characteristics Metals react with oxygen to give basic oxides. These oxides are usually ionic in nature. Group 1 and 2 form basic oxides when react with oxygen. Group 2 hydroxides solubility increases down the group so alkalinity also increases down the group.
Na2O(s) + H2O(l) ――→ 2NaOH(aq)
CaO(s) + H2O(l) ――→ Ca(OH)2(aq)

(ii) Amphoteric oxides

Amphoteric oxides are oxides that can react with both acids and bases. This means they have the ability to behave as either an acid or a base, depending on the conditions. Aluminium oxide (Al2O3) is insoluble in water but reacts with hydrochloric acid to form aluminium chloride and water, and with sodium hydroxide to form sodium aluminate and water.
Aℓ2O3(s) + 6HCℓ(aq) ――→ 2Aℓ(Cℓ)3(g) + 3H2O(t)
Aℓ2O3(s) + 2NaOH(aq) ――→ 2NaAℓO2(s) + H2O(t)

(iii) Acidic oxides

An acidic oxide is an oxide that when combined with water gives off an acid. Non-metals react with oxygen to form acidic oxides which are held together by covalent bonds. Silicon dioxide is acidic oxide as it can react with bases.
Examples of acidic oxides in period 3 are P2O3, P2O5, P2O5, SO3, SO2
P2O5(s) + H2O(t) ――→ H3PO3(aq)
SO2(g) + H2O(l) ――→ H2SO3(aq)

Reactions of these oxides with base are given below:
SiO2(s) + 2NaOH(aq) ――→ Na2SiO3(aq) + H2O(t)
P2O5(s) + 6NaOH(aq) ――→ 2Na3PO4(aq) + 3H2O(t)
SO2(g) + 2NaOH(aq) ――→ Na2SO4(aq) + H2O(l)

Chlorides

Definition Chlorine forms binary compounds with other elements known as chlorides.

Ionic and covalent of chlorides

Chlorides of group 1, 2 and 3 (NaCl) are predominately ionic. Chlorides of elements of group 4, 5, 6 and 7 (PCl5) are covalent. The covalent character in chlorides increases due to decrease in electronegativity between the halogen and the other atom.

Classification of chlorides on the bases of pH

These chlorides show characteristic behaviour when we add them into water, resulting in solutions that can be acidic or neutral.

(i) Neutral chlorides

Neutral chlorides are salts, when dissolved in water, produce a neutral solution with a pH close to 7. At the start of period 3 chloride, sodium and magnesium do not react with water. The solutions formed contain the positive metal ions and negative chloride ions surrounded by water molecules. These ions are now known as hydrated ions and this process is known as hydration.
The following reaction occurs:
NaCℓ(s) ――→ Na⁺(aq) + Cℓ−(aq) (pH = 7)
MgCℓ2(s) ――→ Mg²⁺(aq) + 2Cℓ−(aq) (pH = 6.5)

Group 1 and group 2 chlorides are also neutral with few exceptions.

(ii) Acidic chlorides

If we move in period 3, from aluminium to sulphur all chlorides react with water to make acidic solution with pH less than 7 this process is called hydrolysis. When AlCl3 is added to water, aluminium and chloride ions in solution. Al³⁺ ion is hydrated and cause a water molecule to lose an H⁺ ion, this process is hydrolysis. This turns the solution acidic.
The following reaction occurs:
AℓCℓ3(g) + 3H2O(t) ――→ Aℓ(OH)3(g) + 3HCℓ(aq)

Other examples of acidic chlorides are given below.
SiCℓ4(l) + 2H2O(t) ――→ SiO2(s) + 4HCℓ(aq)
PCℓ3(g) + 3H2O(t) ――→ H3PO3(s) + 3HCℓ(aq)

Oxidation number

Definition The oxidation number of an atom is the formal charge on that atom in a molecule or ion.

• The oxidation number is also referred to as the oxidation state.
• In ionic compounds the oxidation number of an atom is defined as the charge which appears on the, ions.

The oxidation numbers in oxides and chlorides of the third period.

The oxidation number of an element of 3rd Period in its oxide or chloride corresponds to the oxide or chloride the number of electrons used for bonding and is always positive because oxygen and chlorine are more electronegative than any of these elements.

The oxidation number matches the group number, reflecting the total number of valence electrons.

Examples Following are some examples of oxidation state of oxide and chlorides of third period:

• In the oxides of third period, the oxidation number increases from +1 in Na to + 6 in S.
• In chlorides of third period, the oxidation number increases from +1 in Na to + 5 in P.
• Phosphorus and sulfur exhibit several oxidation numbers because they can expand their octet by exciting electrons into empty 3d orbitals.
• In SO2, sulfur has an oxidation number of + 4 because only four electrons are used for bonding.
• In SO3, sulfur has an oxidation number of + 6 because all six electrons are used for bonding.

Consider table 1.2 for oxidation states of various elements of the periodic table.

Table 1.2 Oxidation Numbers in Oxides and Chlorides of 3rd Period elements

Oxide | Oxidation Number | Chloride | Oxidation Number
Na in Na2O | +1 | Na in NaCl | +1
Mg in MgO | +2 | Mg in MgCl2 | +2
Al in Al2O3 | +3 | Al in AlCl3 | +3
Si in SiO2 | +4 | Si in SiCl4 | +4
P in P4O10/ P in P4O6 | +5/+3 | P in PCl5 | +5
S in SO3 | +6 | P in PCl3 | +3
S in SO2 | +4 | S in SCl2 | +2