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CHAPTER 1 PERIODIC TABLE AND PERIODIC PROPERTIES
1.1 HISTORICAL BACKGROUND
It is interesting to note that just three centuries ago, less than a dozen elements were known to humanity. By 1700 A.D., only 12 elements—Gold, Silver, Copper, Iron, Lead, Tin, Mercury, Phosphorus, Sulfur, Carbon, Zinc, and Arsenic—were recognized. The remaining elements existed but were either undiscovered or found in combined forms. There was insufficient knowledge at that time to distinguish between an element and a compound, and substances like water and air were present but not understood to be composed of gases. The first element discovered in history was phosphorus, isolated by Hennig Brandt in Hamburg, Germany, in 1669. Over time, more elements were discovered, prompting scientists to organize them systematically. Up to the end of 18th century, Antoine Lavoisier attempted to classify known elements as metals and nonmetals. In 1829, Johann Wolfgang Döbereiner grouped the elements into triads (a group of three) with similar properties, noting that the atomic weight of the middle element was roughly the average of the other two. Examples of such triads include lithium, sodium, and potassium, and chlorine, bromine, and iodine. English chemist John Newlands, in 1864, first time observed periodicity in the 62 known elements, noticing that the properties of every eighth element were similar when arranged by the increasing order of their atomic masses. Later this observation was termed as periodic law. He classified the elements into eight groups with repeating properties, similar to a musical scale (Sa, Re, Ga, Ma, Pa, Da, Ne, Sa). However, his table was not well-received by some contemporaries and the Chemical Society. In the same year, Lothar Meyer of Germany noted that arranging elements by atomic weight resulted in groups with similar chemical and physical properties at periodic intervals. He presented his findings in a graph with atomic weights as ordinates and atomic volumes as abscissae, showing that electropositive elements appeared at the peaks of the curve in order of atomic weights. In 1869, Russian chemist Dmitri Mendeleev, considered the father of the Periodic Table, arranged 63 elements into eight vertical columns by increasing atomic mass, aligning elements with similar properties into vertical groups. The success of his table was hidden in leaving gaps for undiscovered elements and predicting their atomic mass and properties, which proved accurate when these elements were practically found. 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 mass. This significant breakthrough led Moseley to modify the Periodic Law to state that the properties of elements are periodic functions of their atomic numbers.
1.2 MODERN PERIODIC TABLE – FEATURES AND SIGNIFICANCE
Mendeleev’s initial table had numerous discrepancies and flaws. However, most of these defects have been resolved, and many new discoveries have been seamlessly accommodated. The modern periodic table we use today is based on Mendeleev’s concept but differs in that the elements are arranged in increasing order of atomic numbers, not atomic mass. The classification of elements in the modern periodic table helps in the easier understanding of their properties. Following are some of the main features of the modern periodic table: Presently, 118 elements are grouped in the table in ascending order of their respective atomic numbers. There are seven horizontal rows called periods and eighteen
1.3 METALS, NON-METALS AND METALLOIDS:
Elements can be broadly classified as metals, nonmetals and metalloids. Metals are elements which tend to lose electrons to form positive ions. Examples are iron, copper, gold and silver. On the other hand, non-metals are elements which tend to gain electrons to form negative ions. The examples are chlorine, sulfur and phosphorous. The metalloids separate the metals and nonmetals on a periodic table. The metalloids exhibit some properties of metals and some of non-metals. 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. Elements to the left of the line are considered metals. Elements just to the right of the line exhibit properties of both metals and nonmetals and are termed as metalloids or semimetals. Elements to the far right of the periodic table are nonmetals. The exception is hydrogen, the first element on the periodic table.
1.4 BLOCKS IN PERIODIC TABLE:
Elements in the periodic table can be classified based on the subshells containing their valence electrons. For instance, the valence electrons of elements in the first two groups are in the “s” subshells, placing these elements in the s-block. Similarly, transition elements belong to the d-block, and the elements in the two series at the bottom of the table (known as Lanthanides and Actinides) are categorized as f-block elements. The remaining elements in groups 13 to 18, including the inert gases in the last group, belong to the p-block. Knowing the block to which an element belongs provides valuable information about its characteristics, chemical reactivity, oxidation states and other properties such as electronegativity and ionization energy, electron filling, etc..
1.5 FAMILIES IN PERIODIC TABLE:
Elements may be categorized according to element families. An element family is a set of elements sharing common properties. There are five famous families of elements in the periodic table:
Alkali metals (Li, Na, K, Rb, Cs, Fr)
Alkaline earth metals (Be, Mg, Ca, Ba, R.)
Transition metals (Sc, Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn)
Chalcogens (O, S, Se, Te, Po) Halogens (F, Cl, Br, I, At, Ts)
Noble gases (He, Ne, Ar, Kr, Xe, Rn, Og)
i) Alkali Metals:
Elements in the first group of the periodic table are known as alkali metals because they produce alkalis when they react with water. Sodium and potassium are notable examples of these elements. Although hydrogen is not classified as an alkali metal due to its lack of typical group properties, it does share some properties with them. Alkali metals are characterized by having one valence electron, being soft metallic solids with a shiny and lustrous appearance, exhibiting high thermal and electrical conductivity, possessing low densities, relatively low melting points, and low ionization energies.
ii) Alkaline Earth Metals:
Group 2 elements are metals primarily found in the earth and form alkalis; hence they are referred to as alkaline earth metals. Examples include calcium and magnesium. These elements have two electrons in their valence shell, making them divalent. Their characteristic properties include being metallic solids that are harder and denser than alkali metals. They are shiny and lustrous, easily oxidized, with high thermal and electrical conductivities
iii) Transition Elements:
The transition metals make up the largest family of elements in the middle of periodic table. They include four series of d-block elements, as well as the lanthanides and actinides (f-block elements) found in the two rows below. Transition elements are characterized by their shiny, lustrous appearance, hardness, and dense metallic structure. They exhibit high thermal and electrical conductivities, high melting points, high density, and variable oxidation states.
iv) Chalcogens:
The group 16 elements are called Chalcogens because most ores of copper (Greek chalkos) are oxides or sulfides, and such ores contain traces of selenium and tellurium. All chalcogens found in nature can form ores with other elements. The only exception is livermorium, which is synthetically created. The chalcogens are composed of nonmetals, metalloids, and metals, so each element will display slightly different characteristics because of this. The first three elements in the group are needed for human function, while the last three elements are toxic.
v) Halogens:
Elements in group 17, known as halogens, are nonmetallic and belong to the p-block. The term “halogen” means “salt-former” because these elements easily react with alkali metals and alkaline earth metals to form stable halide salts. Examples are fluorine, chlorine, bromine, iodine. Due to having seven electrons in their outermost shell, halogens are highly reactive nonmetals with high electron affinities. Halogens can easily accept one electron to complete their outermost shell. Their melting points and boiling points increase as the relative molecular masses, Mr, increases. The physical state of these elements changes down the group: fluorine and chlorine are gases at room temperature, while bromine is a volatile liquid, and iodine is a solid which sublimes.
vi) Noble Gases:
The noble gases are a group of unreactive elements present at the extreme right of the periodic table in Group 18. Examples include helium and argon. Due to their stable electron configuration (complete outermost shell), they are almost entirely unreactive under normal conditions and rarely form compounds with other elements. These elements are monoatomic in nature. They are odorless, colorless, non-flammable, and exhibit minimal chemical reactivity.
1.6 PERIODIC ARRANGEMENT AND ELECTRONIC CONFIGURATION
Understanding the periodic arrangement of elements in the periodic table offers valuable insight into their physical properties, such as their physical state and atomic radii, as well as their electronic structure and chemical reactivity. The period number indicates the principal quantum number (n), representing the number of electron shells surrounding the nucleus. For example, an element Z in the 3rd period has three electron shells, with its valence electrons located in the 3rd shell. The specific subshell where the valence electrons are found, depends on the element’s block (azimuthal quantum number).
If an element Z in the 3rd period is in the s-block, its valence electrons are in the 3s subshell. Additionally, the group number indicates the number of valence electrons; for instance, an element Z 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 Z would be magnesium (Mg). Here is another example to relate period number and group number with electronic configuration and position of element in period table. X belongs to group 3 and period 3 In above example, 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. The configuration will be: 1s2 2s2 2p6 3s2 3p1 1st 2nd 3rd shell shell Shell 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.
1.7 PERIODIC LAW AND TRENDS OF CHANGE IN PROPERTIES
The Modern Periodic Law states: “The physical and chemical properties of elements are periodic functions of their atomic numbers.” The atomic number corresponds to the number of protons in an atom. This law is the “cornerstone” of the periodic table, indicating that elements with similar properties appear at regular intervals. For instance, when elements are arranged by increasing atomic numbers, sodium, potassium, and cesium exhibit many physical and chemical characteristics similar to lithium, as they are all placed in the same group of the table. However, due to the gradual increase in the number of protons in the nucleus and the addition of new electron shells, the physical and chemical properties of elements vary systematically within a group and a period. Some of the properties that change significantly from left to right across a period or from top to bottom within a group include atomic radius, ionization energy, electron affinity, electronegativity, and metallic character. In this section, we will explore the trends in these properties along a period as well as within a group.
1.7.1 Variation in Metallic Character
The metallic character of elements is typically their tendency to lose electrons. We find that elements on the left side of the periodic table have a greater tendency to lose their outermost electrons to achieve a complete octet. In contrast, elements on the right side of the table tend to gain electrons to complete their outermost shell. Therefore, elements on the left side of the periodic table are metals that form positive ions, while elements on the right side, particularly in the right corner, are nonmetals that form negative ions. Hence one can conclude that the metallic character of an element largely depends on its valence shell electronic configuration. Across the period (from left to right), we find that the outermost shell remains the same, but the number of protons (positive charge) in the nucleus increases. As we move from left to right, the increased positive charge in the nucleus exerts a stronger attraction on the outermost electrons, reducing the tendency of the atoms to lose these electrons. Consequently, the metallic character of the elements decreases. In other words, the increase in nuclear charge pulls the electron cloud closer to the nucleus, making it more difficult for the atom to lose electrons and thereby decreasing its metallic character. Thus, metallic character decreases across a period from left to right.
Contrarily metallic character increases as one moves down in a group of the periodic table. This is due to the fact that the electrons become easier to lose as the atomic size (radius) increases. As the atomic size increases, the electrostatic force of attraction decreases between the nucleus and the electrons, causing the electrons to be held more loosely. The increase in metallic character (ease of losing electron) makes the element more reactive. Hence Cesium is far more reactive and electropositive than sodium or lithium.
1.8 REACTIONS OF Na AND Mg WITH WATER, OXYGEN AND CHLORINE:
1.8.1 With water
Sodium is more reactive than magnesium towards water. Na reacts vigorously with water to form sodium hydroxide and hydrogen while Mg reacts more slowly in forming magnesium hydroxide and hydrogen. However magnesium reacts with steam more vigorously to make magnesium oxide and hydrogen gas
2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
Mg(s) + 2H2O(l) → Mg(OH)2(aq) + H2(g)
Mg(s) + 2H2O(g) → MgO(s) + H2(g)
1.8.2 With Oxygen
Sodium burns in oxygen with an 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) → 2Na2O2(s)
Under special conditions like limited O2 or high temperature, sodium oxide is formed.
4Na(s) + O2(g) → 2Na2O(s)
Magnesium burns in oxygen with an intense white flame to give white solid magnesium oxide.
2Mg(s) + O2(g) → 2MgO(g)
1.8.3 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)
1.9 VARIATION IN ATOMIC RADIUS:
The atomic radius is a measure of the size of an atom. For a neutral single atom, it can be defined as the distance from the nucleus of an atom to the outermost electron shell. Because atoms don’t have well-defined boundaries, the atomic radius is often defined in terms of the distance between the nuclei of two identical atoms that are bonded together through a chemical bond. The atomic radius can vary depending on the type of bond or the state of the atom. For example, the radius can be different in a covalent bond compared to an ionic bond. The atomic radius is typically measured in picometers (pm) or Angstroms (Å). Periodic trends in atomic radius: The factors affecting the atomic radius are: atomic number, effective nuclear charge and shielding effect of inner electrons. Generally, atomic radius decreases across a period (from left to right) in the periodic table due to increasing nuclear charge, which pulls the electron cloud closer. Conversely, atomic radius increases down a group (from top to bottom) because additional electron shells are added, so more shielding makes the atom larger despite the increase in nuclear charge (which is outweighed).
1.10 VARIATION IN IONIC RADIUS
The ionic radius is a measure of the size of an ion in a crystal lattice. It’s typically defined as the distance from the nucleus of an ion to the outermost electron shell, measured in picometers (pm) or angstroms (Å). The concept is useful for understanding how ions interact in compounds and for predicting the properties of substances. When an atom loses one or more electrons to become a positive ion, it generally becomes smaller than the neutral atom. This is because the loss of electrons reduces electronic repulsion and allows the remaining electrons to be pulled closer to the nucleus. Contrarily when an atom gains one or more electrons to become an anion, it generally becomes larger than the neutral atom. This is because the addition of electrons increases electronic repulsion, as a result the nuclear pull on electrons decreases and the electron cloud expands. As you move across a period from left to right, the ionic radius of cations decreases due to the increasing nuclear charge which pulls the electrons closer. For anions, the ionic radius also decreases across a period because the increasing nuclear charge also pulls the electrons closer to the nucleus, even though anions are typically larger than the neutral atoms from which they are formed. On the other hand, both cations and anions increase in size as we move down a group. This is because 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. Understanding these trends helps in predicting the behaviour of ions in various chemical contexts and their role in forming compounds.
1.11 VARIATION IN 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).”
Na(g) → Na+(g) + e– ΔHi1 = 494 kJ mol-1
Ca(g) → Ca+(g) + e– ΔHi1 = 590 kJ mol-1
If a second electron is removed from each ion in a mole of gaseous 1+ ions, we call it the 2nd ionization energy, ΔHi2. Again, using calcium as an example: Ca+(g) → Ca2+(g) + e– ΔHi2 = 1150 kJ mol-1. Removal of a third electron from each ion in a mole of gaseous 2+ ions is corresponding to the 3rd ionization energy. Again, using calcium as an example: 3rd ionisation energy: Ca2+(g) → Ca3+(g) + e– ΔHi3 = 4940 kJ mol-1
An element can have several ionization energies; the exact number corresponds to its atomic number.
1.11.1 Factors affecting the ionization energy:
The magnitude of the ionisation energy of an element depends upon the following factors:
i) Nuclear charge:
The amount of the attractive forces between the nucleus and the electrons depends upon the charge in the nucleus. Greater the effective nuclear charge, greater is the electrostatic force of attraction, more difficult is the removal of an electron from the atom. For this reason, ionisation energy increases with an increase in the effective nuclear charge.
ii) Size of the atom or ion:
The ionisation energy depends upon the coulomb force of attraction between the nucleus and the electrons. This force decreases when the distance between the nucleus and electrons increases. 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 ionisation energy is higher when an electron is to be removed from a fullyfilled or half-filled-orbital.
(a) Noble gases have highest ionisation energies in their respective periods. It is due to highly stable fully-filled orbital (ns2 np6).
(b) Oxygen has lower ionisation energy than nitrogen. The electronic configuration of oxygen and nitrogen are
7N = 1s2 2s2 2px1 2py1 2pz1
8O = 1s2 2s2 2px2 2py1 2pz1
ΔHi = 1403 kJ mol-1
ΔHi = 1365 kJ mol-1
Although, nitrogen has one unit less positive charge in its nucleus than oxygen, but due to the extra-stability of the half-filled orbitals of nitrogen it is difficult to remove an electron from N atom.
(4) Shielding effect:
In multi-electron atoms, the inner electrons shield the outer electrons from the nucleus. The outermost electrons experience lesser attractive force of the nucleus. Greater the shielding, easier it is to remove the valence electrons from an atom. Larger the number of inner electrons, greater is the screening effect, therefore, lower is the ionization energy.
(5) Penetration effect of the electrons:
In multi-electron atoms, the probability of finding an electron near the nucleus follows the order,
s-electrons > p-electrons > d-electrons > f-electrons
So, the s-electrons are more penetrating than p-electrons towards the nucleus and p-electrons are more penetrating than d-electrons and so on. Thus, for the same principle quantum number or the same shell, the ionisation energy of a s-electron is higher than that of a p-electron. Similarly, the ionisation energy of a p-electron is higher than that of a d-electron and so on.
1.11.2 Periodic trends in ionization energy:
Going down in a group, the nuclear charge increases but as the size of the atom and the number of electrons causing the shielding effect also increases therefore ionization energy decreases from top to bottom.
In Group I, the ionization energies decrease in the following order:
Li > Na > K > Rb > Cs
This trend is due to the progressive increase in the size of the valence electron cloud as the principal quantum number “n” increases. For 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. A similar trend is observed in halogens, where fluorine has the highest ionization energy and astatine has the lowest (At < I < Br < Cl < F).
As you move from left to right across a period, the outer electron shell remains unchanged while the effective nuclear charge increases, making it more difficult to remove an electron. Consequently, the ionization energy increases. Although the number of electrons also increases across a period, the shielding effect within the same shell is same so not considered.
The trend of ionization energies of period (1-3) is shown in Fig 1.6. The figure also reveals that Noble gases have the highest values of ionization energy because due to complete outermost shell in them, the removal of electron is extremely difficult, whereas alkali metals have lowest values of ionization energy.
1.12 ELECTRON AFFINITY (ΔHeao)
The first electron affinity, (ΔHea1o), 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. Electron affinity of chlorine atom,
Cl(g) + e– → Cl–(g) ΔHea1o = −348.8 kJ mol-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. The second electron affinity, ΔHea2o 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. For example, when first electron is added to a neutral oxygen atom, 141 kJ mol-1 energy is released.
O(g) + e– → O–(g) ΔHea1o = −141 kJ mol-1
But 798 kJ mol-1 of energy is absorbed on adding second electron to a uni-negative (O–) ion.
O–(g) + e– → O2-(g) ΔHea2o = +798 kJ mol-1
The net enthalpy change for the formation of the oxide ion (O2-) can be calculated by adding the first and second electron affinities
O(g) + 2e– → O2-(g)
ΔHea1o + ΔHea2o = (−141) + (+798) = +657 kJ mol-1
Electron affinities are not easy to measure, so they are not available for all the elements, and some of the values that are available were determined theoretically or were calculated rather than measured. The reason being that the electron is being added against the electrostatic repulsion. This repulsion exists between the negative charge on the electron being added and the negative ion
1.12.1 Factors affecting electron affinity:
Important factors affecting the magnitude of electron affinity values of elements are as follows:
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.
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.
Electronic configuration of atom:
The electron affinity is low when the electron is added to a half filled sub-shell than that for partially filled one. This can be explained by considering the following examples. 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. Noble gases group-18 (VIII-A) have stable ns2 np6 configuration and hence the atoms of these gases, either do not accept any extra electron, or have a very little tendency to accept the electron. This is evident from their positive 1st electron affinities.
1.12.2 Periodic trends in electron affinity:
As the atomic size increases down the group, the larger electron cloud causes the incoming electron to experience less attraction from the nucleus. Consequently, electron affinity generally decreases down the group. This trend is observed in the halogens (At < I < Br < F < Cl). However, within this group, fluorine has a lower electron affinity than chlorine as shown in the Table 1.1. This is due to fluorine's smaller size compared to chlorine which makes it harder to accept another electron.
Generally, electron affinities become more negative as we move from left to across a period. This is firstly due to increase in the nuclear charge, which attracts additional electrons more strongly and secondly due to decreasing atomic radius. Thus the electron affinity values go on increasing when we move from alkali metals to halogens in periods 2, 3, 4, 5 and 6. Though the variation is less regular in a period as compare to the trend from top to bottom within a group. Understanding electron affinities, along with ionization energies, helps in predicting that which atoms are more likely to lose electrons and which are more likely to accept electrons. There are anomalies to these general trend and are being discussed in chapter no 2.
1.12.3 Anomalous behavior:
The sizes of O and F atoms are small as compared to the sizes of S and Cl atoms in their respective groups. The addition of an extra electron to O and F atoms produces high electron density round them. You can say that the repulsion between the electrons already present in the relatively compact 2p orbitals of the valence-shell of these atoms and the extra electron being added to these atoms to get O and F ions increase. Due to the increased electron-electron repulsion, O and F atoms show lesser tendency to attract an electron towards them.
1.13 VARIATION IN ELECTRONEGATIVITY
Electronegativity is a measure of an atom’s attraction for bonding electrons in a molecule relative to other atoms. 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.
1.13.1 Factors Affecting Electronegativity
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. For example, the electronegativities of halogens in group 17 are in the order
F > Cl > Br > I
It is due to the increase in the atomic size from F to I in the group
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. The effective nuclear charge, in turn, depends upon the atomic number, size, and the number of intervening shells in an atom. 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.
Shielding Effect:
The shielding effect is the reduction in the attractive force between the nucleus and an electron in a multi-electron atom due to the inner shell electrons. Electronegativity value decreases as the shielding effect increases and vice versa. The higher number of intervening electrons with the addition of shells will result in an increase in the shielding effect. The electronegativity of alkali metals decrease from top to bottom. This is due to the enhanced shielding effect, which is caused by the addition of shells when we move down the group. Therefore, F is the most electronegative atom among the halogens and I is expected to be the least electronegative.
1.13.2 Periodic Trends in Electronegativity
When we move from left to right along the period, the effective nuclear charge increases and the atomic size decreases. Due to this reason, the value of electronegativity increases across a period in the periodic table. Additionally, the elements on the right side of the periodic table (like fluorine, oxygen) have nearly full valence shells (with one or two vacancies) and are more eager to attract additional electrons to complete their octet. That is why the elements on the left side have low electronegativities and are often referred to as electronpositive elements and those on right side with high electronegativity values are termed as electronegative elements. Table 1 presents the electronegativity values of the elements in period 2 and 3. Chlorine has the highest value of electronegativity (3.0) in period 3, while sodium is the least electronegative element with a value of about 0.9.
However, down the group in the periodic table, the atomic size increases by the addition of shells. The nuclear charge also increases but the effect of the increase in nuclear charge is overcome by the successive addition of shells. Therefore, the effective nuclear charge remains almost the same. Hence, the value of electronegativity decreases from top to bottom. For example, in 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 Fig 1.8.
Normally, metals being on the left side of the periodic table, possess lower electronegativity values than those of non-metals. Hence, metals are electronpositive and non-metals are electronegative, relatively.
Figure 1.9 provides a summary of all the variation trends in various physical properties of elements in the periodic table.
TRENDS IN BONDING IN OXIDES AND CHLORIDES OF PERIOD 3
Oxides of group 1, 2 & 3 (e.g., 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 (e.g., SO2) are more covalent. These oxides exist as covalent molecules with weak intermolecular forces between molecules. This transition is a result of the increasing electronegativity and decreasing ionic character.
Similar to oxides, chlorides of group 1, 2 and 3 (e.g., NaCl) are predominately ionic. Chlorides of elements from group 4, 5, 6 and 7 (e.g., PCl5) are covalent. The covalent character in chlorides increases due to decrease in difference of electronegativity between the halogen and the other atom the higher electronegativity of the central atom.
CLASSIFICATION OF OXIDES AND CHLORIDES:
Oxides:
Oxides are binary compounds formed by the reaction of oxygen with other elements. Oxygen is highly reactive in nature hence it reacts with metals and non-metals to form oxides. The classification of oxides is done into neutral, amphoteric and basic or acidic based on their acid-base characteristics. Along the period three nature of oxides will be explored.
Basic Oxides
A basic oxide is an oxide that when combined with water gives off an alkali. 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. These compounds readily react with water but few exceptions are there. Examples are: Na2O, CaO, BaO. 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)
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. In chemical reactions, they can neutralize both acidic and basic substances. Aluminum 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) → 2AlCl3(s) + 3H2O(l)
Al2O3(s) + 2NaOH(aq) → Na2Al2O4(s) + H2O(l)
This dual reactivity is a characteristic feature of amphoteric oxides, distinguishing them from other oxides that are typically either basic or acidic.
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. These compounds can also be called acid anhydrides. Acidic oxides usually have a low melting and boiling point except for oxides like SiO2 which is insoluble in water and have high melting points as it forms giant molecules. Silicon dioxide is acidic oxide as it can react with bases. Examples of acidic oxides in period 3 are: P2O3, P2O5, SO3, SO2.
P2O3(s) + H2O(l) → H3PO3(aq)
P2O5(s) + H2O(l) → H3PO4(aq)
SO2(g) + H2O(l) → H2SO3(aq)
SO3(g) + H2O(l) → H2SO4(aq)
Reactions of theses oxides with bases are given below:
SiO2(s) + 2NaOH(aq) → Na2SiO3(aq) + H2O(l)
P2O3(s) + 6NaOH(aq) → 2Na3PO3(aq) + 3H2O(l)
SO2(g) + 2NaOH(aq) → Na2SO3(aq) + H2O(l)
1.16 CHLORIDES
Chlorine is a highly reactive nonmetal that forms stable compounds known as chlorides through chemical reactions with various elements. One common example is table salt, an ionic compound consisting of sodium and chloride. These chlorides show characteristic behavior when we add them into water, resulting in solutions that can be acidic or neutral.
1.16.1 Neutral Chlorides
Neutral chlorides are salts that, 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 polar water molecules are attracted to the oppositely charged ions, dissolving these chlorides by breaking down their giant ionic lattices. 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. When sodium chloride (NaCl) is dissolved in water, it undergoes dissociation into its constituent ions.
NaCl(S) → Na+(aq) + Cl–(aq) (pH=7)
MgCl2(S) → Mg+2(aq) + 2Cl–(aq) (pH=6.5)
Group 1 chlorides are all neutral chlorides and group 2 chorides are also neutral with few exceptions.
1.16.2 Acidic Chlorides:
If we move in period 3, from aluminum to sulphur all chlorides react with water to make acidic solution with pH less than 7 this process is called hydrolysis. Aluminium chloride exist as dimer Al2Cl6 which is covalently bonded. Once we add water, dimer breaks and aluminium and chloride ions in solution. Al3+ ion is hydrated and causes a water molecule to lose an H+ ion, this process is hydrolysis. This turns the solution acidic. The following reactions shall occur:
AlCl3(s) + 3H2O(l) → Al(OH)3(s) + 3HCl(aq)
Other examples of acidic chlorides are given below.
SiCl4(l) + 2H2O(l) → SiO2(s) + 4HCl(aq)
PCl3(l) + 3H2O(l) → H3PO3(s) + 3 HCl(aq)
1.17 VARIATION IN OXIDATION NUMBER IN OXIDES AND CHLORIDES
The oxidation number of an atom is the apparent 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 while forming ionic bonds with other heteroatoms. Since an atom can have multiple valence electrons and form multiple bonds, all of them will be assumed to be ionic and assigned an oxidation state equal to the number of electrons involved in the bonding. So, oxidation number or state is a hypothetical case of the assumption of atoms forming an ionic bond. Let’s examine 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 number of electrons used for bonding and is always positive because oxygen and chlorine are more electronegative than any of these elements. The maximum oxidation number matches the group number, reflecting the total number of valence electrons. Consider the following table (Table 1.2) for oxidation states of various elements of the periodic table. In the oxides, the maximum oxidation number increases from +1 in Na to +6 in S. In chlorides, the maximum 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. For instance, in SO2, sulfur has an oxidation number of +4 because only four electrons are used for bonding, while in SO3, sulfur has an oxidation number of +6 because all six electrons are used for bonding.
For the reference see Periodic table here