WBCHSE Class 11 Chemistry Notes For Periodic Properties Of Elements and Their Functions

Classification of Elements and Periodicity of Properties

Man has been fascinated by the chemical properties of the natural substances in his environment from almost the dawn of civilization. Initially, his curiosity was fuelled entirely by his need to improve his living conditions. He learned to harden clay by baking it at a high-temperature thousands of years ago.

• Then, quite by accident, he discovered metals, which changed his life completely and led to the study of alchemy. Alchemy was not really a scientific study of metals.
• It was based on this belief in the existence of a magical substance (Philosopher’s Stone) that could transform base metals into gold.

However, modern chemistry owes a lot to alchemists, who in their search for the Philosopher’s Stone, tested all the known substances and increased man’s knowledge of the properties of these substances.

Read and Learn More WBCHSE For Class11 Basic Chemistry Notes

• Early scientists, who inherited the knowledge passed on by the alchemists, could afford to study the properties of the few known elements individually. But as knowledge grew, and more and more elements came to be known, this approach proved to be cumbersome.
• Scientists started looking for a way to group together similar elements so that it would be possible to study their properties and the properties of their compounds more systematically.
• After years of effort and several attempts at classifying all the known elements, they succeeded in coming up with an arrangement in which similar elements, or elements which have similar properties, are grouped together.
• This arrangement, called the periodic table, has systematized the study of elements and their compounds. It has provided a framework for organizing the vast amount of information available on the chemical behavior of elements.

History Of The Periodic Table

All the attempts made to classify elements, initially, were based on their atomic weights. This was because Dalton had proposed, earlier, that different elements have different types of atoms, characterized by their atomic weight.

Dobereiner’s triads: J W Dobereiner, a German chemist, classified similar elements into groups of three, or triads, such that the atomic weight of the central element was approximately the arithmetic mean of the atomic weights of the other two. He also pointed out that the properties of the middle element were in between those of the other two members.

These groups were called Dobereiner’s triads, for example, lithium (7), sodium (23) and potassium (39), calcium (40), strontium (88) and barium (137), and chlorine (35), bromine (80) and iodine (127). The major drawback of this classification was that it could be applied to only a limited number of elements.

Newlands’ law of octaves: An English chemist, John Newlands, came up with this arrangement in 1865, based on the observation that when the lighter elements are arranged in order of increasing atomic weight, every element is similar to the element eight placed from it in the list.

Hus was called Newlnnds’ law of octaves because it was like a musical scale in which the first note is the same as the eighth. Thus, the properties of lithium are similar to those of the element eight places from it, i.e., sodium.

This arrangement too was discarded because it could not be applied to elements with atomic weights greater than 40 u. Also, the discovery of noble gases meant that the eighth element was no longer similar to the first. However, for his work Newlands was awarded the Davy Medal in 1887 by the Royal Society, London.

Lothar Meyer’s arrangement: Julius Lothar Meyer (1830-1895) was a German chemist who studied the atomic volumes, melting points, boiling points, and other physical properties of various elements.

• He found that if a graph is plotted between the atomic volumes and atomic weights of the various elements, similar elements occupy similar positions on the curve.
• The most strongly electropositive alkali metals (Li, Na, K, Rb, Cs) occupy the peaks on the curve, the less electropositive metals (Be, Mg, Ca, Sr, Ba) occupy descending positions on the curve and the most electronegative elements (F, Cl, Br, I) occupy ascending positions on the curve.

On the basis of his observations, Meyer proposed in 1869 that the physical properties of the elements are a periodic function of their atomic weights.

Mendeleev’s periodic law: Around the same time, DI Mendeleev (1834-1907), a Russian chemist, was busy trying to arrange the elements in some order. He found that if they are arranged in increasing order of atomic weights, their chemical properties vary in a regular pattern. He proposed that the chemical properties of the elements are a periodic function of their atomic weights.

• When Mendeleev came to know of Lothar Meyer’s conclusion, he combined his own proposal with Meyer’s, and the Jaw came to be known as Lothar Meyer-Mendeleev law or Mendeleev’s periodic law.
• According to Mendeleev: The properties of the elements, as well as the formulae and properties of their compounds depend on the atomic weights of the elements. This means that when the elements are arranged in increasing order of atomic weight, – elements with similar properties occur at regular intervals in the list.

Mendeleev’s Periodic Table

Mendeleev’s arrangement of elements in increasing order of atomic weights is called Mendeleev’s periodic table. He grouped together similar elements and was insightful and courageous enough to leave blank spaces (on his table) for elements he believed would be discovered later.

The elements predicted by him were discovered later, and remarkably enough, were found to possess the properties he had predicted they would have. For example, gallium and germanium, called eka-aluminum and eka-silicon by Mendeleev, were discovered after he proposed his table and are very similar to aluminum and silicon respectively, as he had foreseen.

Mendeleev’s periodic table had eight vertical columns (groups). Later, when the inert gases were discovered, a new group called the zero group was added to the table.

Vertical Columns There are nine vertical columns in Mendeleev’s periodic table called groups. These groups are designated 1, 2, 3, 4, 5, 6, 7, 8 and zero. With the exception of the 8th and zero groups, every group is divided into two subgroups A and B. Group 8 contains nine elements in three sets of three each, and group zero comprises the inert gases.

Horizontal rows The seven horizontal rows are called periods. The first period contains only two elements, the second and third periods contain eight elements each and are called short periods. The fourth, fifth, and sixth periods contain 18 elements each, while the seventh is incomplete and contains only three elements.

Merits of Mendeleev’s periodic table: Mendeleev’s periodic table was a giant step forward in the study of chemistry.

• Systematic study His classification of the elements made the study of their properties simpler and more systematic. Placing elements which have similar chemical properties in the same group is helpful in that if one knows the properties of one element of a group, one can predict those of the rest of the members of the group.
• Prediction of new elements While arranging the 56 elements known then, Mendeleev left blank spaces for elements which had not been discovered at that time. He predicted the properties of these unknown elements on the basis of their positions in his table. He was proved right later, when these elements were discovered.
• Determination of atomic weights Mendeleev’s periodic table has helped correct errors made in the calculation of tire atomic weights of some elements. For example, beryllium was initially assigned an atomic weight of 13.5 on the basis of its equivalent weight (4.5) and valency, which had been wrongly calculated as 3.

This would have placed it between C and N in the periodic table. But there was no such space in Mendeleev’s table. Besides, its properties showed that its correct position should be between lithium (7) and boron (11), so beryllium was assigned a valency of 2 (atomic weight = 2 x 4.5 = 9) and placed where it belonged.

Drawbacks of Mendeleev’s periodic table: So great was Mendeleev’s contribution to the systematic study of elements that he is usually given credit for the periodic table as we know it today. Nonetheless, the table had the following drawbacks.

• Position of isotopes Isotopes are atoms of the same elements that have different atomic weights. As Mendeleev’s. classification is based on atomic weight, isotopes would have to be placed in different positions. For example, the isotopes of hydrogen, \({ }_1^1 \mathrm{H},{ }_1^2 \mathrm{H},{ }_1^3 \mathrm{H}\) would occupy different positions.
• Position of hydrogen Though hydrogen is placed in Group IA (alkali metals), it resembles the elements of both Group IA and Group 7A (halogens). Thus, the position of hydrogen in the periodic table is controversial.
• Anomalous positioning of elements Some elements with higher atomic weight are placed before those with lower atomic weight. For example, argon (atomic weight = 40) is placed before potassium (atomic weight = 39).

Also, some dissimilar elements are grouped together, while some similar elements are placed in different groups. For example, alkali metals, such as Li, Na, and K (LA) are grouped together with the coinage metals, Cu, Ag, and Au (IB) though their properties are different. On the other hand, some chemically similar elements like Cu (LA) and Hg (2B) are placed in different groups.

Elements of Group 8 The nine elements of Group 8 have not been placed in a proper order. They have been arranged in three triads without any justification.

Modern Periodic Law

Mendeleev’s arrangement of elements, though rational and systematic, was empirical. Several decades later, after the electron had been discovered and the modem theory of atomic structure had been developed, H G J Moseley studied the spectral lines emitted by heavy elements in the X-ray region of the electromagnetic spectrum.

• He found that the X-rays radiated by each element have a characteristic frequency that differs according to a regular pattern. The frequencies were different not because of the change in atomic weight but due to the charge on the nucleus (atomic number).
• Thus, he proposed that the elements should be classified on the basis of atomic number (Z) and not atomic weight. He said that the properties of an element depend on its atomic number, rather than its periodic functions of their atomic numbers.
• This means that if the elements are arranged in order of increasing atomic number, elements that have similar properties will appear at regular intervals. This is known as periodicity.

Cause of periodicity: It is not difficult to understand periodicity if one considers atomic structures and what happens to an atom during a chemical reaction. Neither the nucleus, nor the electrons of the inner shells of an atom participate in a chemical reaction.

• Only the valence electrons participate in chemical reactions, so it is these electrons that govern the chemical properties of atoms. In that case, atoms which have the same kind of arrangement of electrons in their outermost shells should have similar chemical properties.
• Now consider the electronic configurations of any group of elements (for example, the alkali metals or Group 1 in the periodic table).

• They have the same number of valence-shell electrons (Group 1 f elements have one electron in their valence shell). This explains why they have similar properties.
• The atomic number of an element is equal to the number of protons it contains, which is again equal to the number of electrons it has. And an atom’s chemical behaviour is governed by the number of electrons it has.
• That is why it makes more sense to arrange the elements in order of increasing atomic number, rather than in order of increasing atomic weight.

Long Form Of Periodic Table

Several periodic tables have been proposed after it was realised that the atomic number, and not the atomic weight V of elements, should be the basis for the arrangement of elements. The most commonly used one is called the long form of the periodic table. It has eighteen columns and seven rows.

The rows are called periods and the columns contain elements with similar valence-shell configuration. These columns are called groups and elements belonging to the same group (obviously with similar properties) constitute a family, for example, the halogen family (Group 17).

In this form of the periodic table, 14 elements each of the sixth and seventh periods (i.e., lanthanides and actinides) are placed separately at the bottom.

Electronic configurations of the elements and the periodic table: As a result of elucidation of the structure of the atom, it is now recognised that the periodic law is essentially the consequence of the periodic variation in electronic configurations.

The electronic configuration of an element indeed determines the physical and chemical properties of the element and its compounds. The electronic configuration of elements can be best studied in terms of variation in periods and groups of the periodic table.

Periods: You already know that an electron in an atom is characterised by a set of four quantum numbers. The principal quantum number (M) denotes the main energy level or shell.

Each period begins with the filling up of a new energy shell and all the elements of a period have the same number of electron shells, or the same principal quantum number (n) of the outermost shell.

In fact, the number of a period is the same as the principal quantum number of the valence shell of the elements it contains. The number of elements in a period is equal to the number of electrons required to fill the orbitals of that shell.

First period This corresponds to the filling up of the first energy shell (n = 1). This shell has only one orbital, the Is orbital, which can accommodate only two electrons, so the first period has only ttvo elements.

Second period This is associated with filling up of the second energy shell (n = 2). This shell has one 2s and three
2p orbitals, which can accommodate eight electrons. The second period, thus, has eight elements.

Third period This corresponds to the filling up of the third shell (n = 3). This shell has one 3s, three 3p and five 3d orbitals, but the 3d orbitals have higher energy levels than the 4s orbitals. Consequently, the 3d orbitals are filled after the 4s orbitals and the third period involves the filling up of only four orbitals. Thus this period contains only eight elements, from sodium to argon.

Fourth period This corresponds to the filling up of the fourth energy level (n = 4). It starts with the filling up of the 4s orbital after which the five 3d orbitals are filled and then the three 4p orbitals. The 4d and 4f orbitals can be filled only after the 5s orbital. Therefore, this period involves the filling up of nine orbitals and contains eighteen elements, from potassium to krypton.

Fifth period The fifth period is associated with the filling up of the fifth energy level (n = 5). After the 5s orbital the five 4d and three 5p orbitals are filled. Neither the 4f, nor the 5d orbitals can be filled until the 6s orbital is filled. Thus, this period too has only eighteen elements, from rubidium to xenon.

Sixth period This corresponds to the filling up of the sixth energy level (n = 6). Starting with the 6s orbital, it involves the filling up of sixteen orbitals (6s, 4f, 5d and 6p), so it contains thirty-tivo elements, from caesium to radon. The fourteen elements from cerium to lutetium correspond to the filling up of the seven 4f orbitals. These are separated from the main table and placed below it to save space. They constitute the first f-transition series and are called lanthanides.

Seventh period This corresponds to the filling up of the seventh energy level (n = 7). Like the sixth period this should also have contained thirty-two elements, corresponding to the filling up of the 7s, 5f, 6d and 7p orbitals, but it is still incomplete and contains only twenty-five elements.

It includes most of the man-made radioactive elements. The fourteen elements corresponding to the filling of the 5f orbitals are called actinides and constitute the second f-transition series. These too have been placed below the main table to save space and to allow similar elements to come under the same group.

The first three periods are called short periods and the next three are called long periods.

Division of elements into blocks: The long form of the periodic table has four blocks—s, p, d and f. Each block is named after the atomic orbital which receives the last electron during the filling up of orbitals in order of their increasing energies, s Block elements in which the last electron enters the s orbital of their respective valence shells are called s-block elements.

This subshell has only one orbital which can accommodate only two electrons, so there are only two groups of elements in this block—Groups 1 and 2. The general electronic configuration of the outermost shell of these elements is ns^1-2 where n represents the value of the principal quantum number of the valence shell.

Group 1 comprises alkali metals, while Group 2 contains the alkaline earth metals. These elements have the following general characteristics.

1. They are soft metals.
2. They have low ionisation enthalpies and a re highly electropositive.
3. They are very reactive and form ionic compounds.
4. The metallic character and reactivity of the elements increase down the group.
5. They exhibit oxidation states of +1 and +2 as they have only one or two electrons in their valence shell.
6. They are good reducing agents because they lose electrons easily.

p block This block includes elements in which the last electron enters the p subshell of the valence shell, helium (1s²) being the exception. The general electronic configuration of the outermost shell of these elements is ns²np⁶.

This block contains the groups 13 to 18. Group 18 comprises noble gases with completely filled orbitals (ns²np⁶). Thus, each period ends in a noble gas with a closed-shell configuration. The other two important groups are chalcogens (Group 16) and halogens (Group 17).

The p-block elements have the following general characteristics.

1. Most of them are nonmetals and the character increases as we move from left to right across a period.
2. Most of them are highly electronegative.
3. They exhibit variable oxidation states.
4. They form ionic and covalent compounds.
5. Most of them form acidic oxides.

The elements of the s and p blocks together are called representative elements.

• Two exceptions are observed in this classification. The electronic configuration of helium is 1s² so that it should belong to the s block, but it has been placed in the p block with Group 18 elements (noble gases). Its position is justified because it has a completely filled valence shell and shows the characteristic properties of a noble gas. The other exception is hydrogen. The electronic configuration of hydrogen is 1s¹, hence it can be placed in Group 1 (s block).
• However, like halogens it is one electron short of noble-gas configuration. This controversy regarding position of hydrogen in the periodic table has been discussed in detail. It is placed at the top of the alkali metals in Period 1 of the periodic table.

d block Elements in which the last electron enters the d subshell of the penultimate (second last) energy level are called d-block elements. These include elements of groups 3 to 12.

Their general valence shell configuration is (n – 1) d^1-10 ns^1-2 where n represents the outermost energy level. This block contains three rows of ten elements each. The fourth row is incomplete.

Tire rows which are complete are called the first, second and third transition series. They correspond to the filling of the 3d, 4d and 5d orbitals respectively. Members of this block are called transition elements because they represent a change in character from reactive metals on one side to nonmetals on the other.

These elements have the following general characteristics.

1. They are hard metals with high melting points.
2. They exhibit variable valencies and oxidation states.
3. They form ionic as well as covalent compounds.
4. They form coloured complexes.
5. Most exhibit paramagnetism.
6. Most possess catalytic properties.

f block This block contains elements in which the last electron enters the f from the outermost) shell. The general electronic configuration of f-block elements is \((n-2) \mathrm{f}^{1-14}(n-1) \mathrm{d}^{0-1} n \mathrm{~s}^2\), where n is the principal quantum number of the valence shell.

This block comprises two series of elements placed below the main body of the periodic table. Elements of the first series are called lanthanides, while elements of the second series are called actinides. Together they are called the inner transition elements.

These elements have the following general characteristics.

1. They show variable oxidation states.
2. They are metals with high densities and high melting points.
3. They form coloured compounds.
4. Most elements of the actinide series are radioactive.

Elements up to uranium (Z = 92) arc found in nature, except technetium (Z = 43) and promethium (Z = 61), which arc produced from the disintegration of radioactive elements. Elements beyond uranium are called synthetic or transuranic as they are produced synthetically.

Many of these have been made only in nanogram quantities or even less by nuclear reactions and their chemistry is not fully known.

• In addition to the classification of elements into s, p, d and f blocks another classification is done, based on their properties. The elements can be broadly classified into metals and nonmetals.
• Elements on the left-hand side of the periodic table are metals. More than seventy-five per cent of all known elements are metals. Elements at the top right-hand side of the periodic table are nonmetals.
• Metals are generally solids at room temperature (except mercury) with high melting and boiling points. They are malleable and ductile, and good conductors of heat and electricity. Nonmetals are generally solids or gases at room temperature with low melting and boiling points. They are brittle and poor conductors of heat and electricity. The change from metallic to nonmetallic character is gradual. Some elements like germanium, arsenic, silicon and antimony show characteristics of both metals and nonmetals. These elements are called semimetals or metalloids.

Predicting period, group number and block of an element: If you have to predict the group, period and block of an element, proceed step by step.

1. Write the electronic configuration of the element according to the Aufbau principle. Do not change the configuration by considering the relative stability of half-filled and completely filled orbitals. Do not rearrange in order of increasing values of quantum number.
2. The period of the element is the same as the value of the principal quantum number of the valence shell. For example, consider aluminium, whose atomic number is 13. Its electronic configuration is 1s² 2s² 2p⁶ 3s² Sp¹. The value of the principal quantum number of its valence shell is 3, so it belongs to the third period.
3. The block of an element is predicted on the basis of the subshell which receives the last electron. Consider copper, whose electronic configuration should be 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁹.
4. However, due to a slight difference in the energies of 4s and 3d subshells, and stability of the completely filled and half-filled orbitals, the outer electronic configuration of Cu is 3d¹⁰ 4s¹. The last electron is received by the d subshell, so it belongs to the d block.
5. The group number is predicted on the basis of the number of electrons in the outermost or the penultimate shell.

• For an s-block element, the group number is the same as the number of electrons in the valence shell. For example, the electronic configuration of lithium is 1s² 2s¹. It belongs to Group 1 as there is one electron in its valence shell.
• For d-block elements, the group number is obtained by adding 1 to the number of electrons in the d subshell of the penultimate shell. For example, the electronic configuration of copper is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d¹⁰. Its group number is (10 +1) = 11.
• For p-block elements, the group number is obtained by adding 12 to the number of electrons in the p subshell of the valence shell. For example, the electronic configuration of gallium is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p¹. It belongs to Group 13.

Nomenclature of the elements with atomic number more than 100: Enrico Fermi and his co-workers in 1934 made attempts to prepare the elements beyond uranium, The new elements were prepared by the bombardment of uranium with slow neutrons. As of now, elements with numbers up to 112 and 114 have been discovered.

• Earlier the naming of new elements was done y discoverer(s) traditionally. In recent years, disputes have arisen over the original discoveries of some of the elements of atomic number 104 and above, when two scientists from two different countries have common criteria for discovering the same element.
• To avoid such discrepancies IUPAC has recommended that until the c a m or a newly discovered element is proved and its official name announced, a nomenclature is to be followed of name these new elements.
• The nomenclature of new elements and elements which are yet to be discovered is based on the Latin words for the atomic number of the elements. The names are derived using the numerical roots for 0 and numbers 1 to 9.
• The roots are combined together in the sequence of digits which make up the atomic number and (ium) is added at the end. The notation for IUPAC nomenclature of elements is shown in Table 3.7. For example, the IUPAC name and symbol for the element with atomic number 122 may be written from the table. The roots for 1 and 2 are un and bi respectively. Hence the name is un (1) + bi (2) + bi (2) + ium or unbibium and the symbol assigned is Ubb.

Example: Write the IUPAC names and symbols for the elements with atomic numbers 121 and 140. Also, identify the group where these elements would be placed (once discovered) by giving their electronic configuration.
Solution:

The IUPAC names for the elements with atomic numbers 121 and 140 are Unbiunium (Ubu) and Unquadnilium (Uqn) respectively.

• The electronic configuration of elements which are yet to be discovered can be written on the basis of the periodicity in filling of the valence shell orbitals of elements in the periodic table.
• Thus, the electronic configuration of the element with atomic number 121 is [Uuo]7d¹ 8s². Therefore, it will be placed in Group 3.
• The electronic configuration of the element with atomic number 140 is [Uuo]6f¹⁴ 7d⁶ 8s². Therefore, it will be placed in Group 8.

 

Periodic Trends In Chemical Properties

We have learnt about the periodicity shown by atomic properties like atomic size, ionisation enthalpy, electron gain enthalpy and electronegativity. The discussion involving trends in chemical properties can be quite elaborate.

Here we shall restrict ourselves to the periodicity of valency and anomalous properties of second-period elements only. We have already discussed the periodicity of valency earlier in the chapter.

Example: Refer to the periodic table and predict the formulae of compounds which might be formed by the following pairs k, of elements:

1. Carbon and chlorine, and
2. Boron and oxygen.

Solution:

1. Carbon is a Group 14 element with a valency of 4, while chlorine belongs to Group 17 with a valency of 1. Therefore, the formula of the compound formed will be CCl₄.
2. Boron is a Group 13 element with a valency of 3, while oxygen is a member of Group 16 and possesses a valency of 2. Hence, the formula of the compound formed would be B₂O₃.

Anomalous properties of second-period elements: The representative elements of the second period—lithium, beryllium, boron, etc., show anomalous characteristics, as compared to the other members in the respective groups.

• This is attributed to the high ionisation enthalpy, high electronegativity, small size and large charge-to-radius ratio of the first members of the said groups.
• Further these elements showing anomalous characteristics have a close chemical similarity to their diagonal neighbours in the next group of the third period. This is a relationship within the periodic table and is referred to as diagonal relationship. You will study about the anomalous characteristics and similarity in properties of the s block elements due to the diagonal relationship.
• The p-block elements of the second period have a tendency to enter into 7t bonding (or p_π-p_π multiple bonding) with themselves, as in C=C, C≡C and N=N, O=O. They also form π bonds with the other second-period elements. Such bonds exist in oxides of nitrogen, (N=O), cyanide and carboxyl groups (C=N, C=O).
• However, the subsequent members of the same group do not form σ bonds with other elements. For example, in silicon dioxide, there are only a bonds.

Periodic Trends And Chemical Reactivity

By now you are quite aware of the periodic trends in certain fundamental properties of elements. The ionisation enthalpy decreases down a group and increases across a period. The size (atomic and ionic radii) increases from J top to bottom in a group, and decreases from left to right in a period.

• The electronegativity and electron affinity both increase from left to right in a period and decrease from top to bottom in a group. All these properties are related to electronic configuration. But how does this affect the chemical reactivity of elements in the periodic table?
• The elements at the extreme left of the periodic table are very reactive due to their low ionisation enthalpies. The elements at the extreme right are also very reactive due to their high electronegativities or electron affinities.
• The elements in the middle of the periodic table are relatively unreactive. Thus, transition elements are only moderately reactive since their ionisation enthalpies are relatively large and electronegativities or electron affinities are relatively small.
• Tire elements at the extreme left of the periodic table are very reactive due to their tendency to lose an electron and form cations. Similarly, the elements at the extreme right of the periodic table readily accept an electron to form anions, and are, thus, very reactive. The metallic character decreases and nonmetallic character increases on moving from left to right across the periodic table.

The chemical reactivity of an element can be best shown by its reaction with oxygen. Elements at the two extremes of a period form oxides easily. The elements on extreme left form basic oxides (like Na₂O) while elements in the extreme right form acidic oxides (like Cl₂O₇).

The elements in the centre of the periodic table form amphoteric oxides (like Al₂O₃, and As₂O₃) or neutral oxides (like CO, NO, N₂O). Amphoteric oxides behave as acidic oxides with bases and as basic oxides with acids. However, neutral oxides have no acidic or basic properties.

Example: Show that Na₂O is a basic oxide and Cl₂O₇ is an acidic oxide.
Solution:

On reaction with water, Na₂O forms a strong base (NaOH) and so it is a basic oxide.

⇒ \(\mathrm{Na}_2 \mathrm{O}+\mathrm{H}_2 \mathrm{O} \longrightarrow 2 \mathrm{NaOH}\)

On reaction with water, Cl₂O₇ forms a strong acid (HClO₇) so it is an acidic oxide.

⇒ \(\mathrm{Cl}_2 \mathrm{O}_7+7 \mathrm{H}_2 \mathrm{O} \longrightarrow 2 \mathrm{HClO}_7\)

The acidic or basic nature of the oxides can be easily tested by the action of litmus paper on their aqueous solutions.

Example: Predict the formulae of the stable compounds that ivottld be formed by the combination of the following pair of elements.

1. Lithium and oxygen
2. Magnesium and nitrogen
3. Aluminium and iodine
4. Silicon and oxygen
5. Phosphorus and fluorine
6. Element with atomic number 71 (Lu) and fluorine

Solution:

1. Li₂O
2. Mg₃N₂
3. All₃
4. SiO₂
5. PF₃ or PF₅
6. LuF₃