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Periodic Trends in Properties of Elements – Trends in Properties of Elements and the Factors Affecting the Trends

The majority of us like both playing and watching the outdoor sport of cricket. We might as well have noted that a new over begins after every six deliveries. Depending on the game's format i.e. Twenty-20 Internationals, Test Cricket, or One-Day Internationals, this keeps happening until the allotted number of overs have been used.

But why are we talking about cricket right now?

Does it relate in any way to the concept being addressed here?

It is related to this concept page, yes. In cricket, a new over begins after every legitimate six deliveries. Similarly, the Modern Periodic Table has some chemical and physical characteristics of elements that repeat after certain intervals. This repeating of properties is called the periodic trends in properties of the elements in the periodic table.

On this concept page, we will learn more about the characteristics of the elements, including their trends across the periodic table in terms of factors like ionisation energy, electron affinity, atomic and ionic radii, electronegativity, etc.

TABLE OF CONTENTS

  • The Modern Periodic Law
  • Atomic Radius
  • Ionic Radius
  • Ionisation Energy
  • Metallic Character
  • Electron Gain Enthalpy
  • Electron Affinity
  • Electronegativity
  • Valency and Valence Electrons
  • Chemical Reactivity
  • Facts Based on the Period Trends and Periodic Tables
  • Practice Problems
  • Frequently Asked Questions – FAQ

The Modern Periodic Law

The periodic properties of elements are based on the periodic law. According to Mendeleev, the physical and chemical properties of elements are a periodic function of their atomic weights. Later it was disproved and the modern periodic law was formulated based on Moseley’s experiment.

According to the modern periodic law, “the physical and chemical properties of elements are a periodic function of their atomic numbers”.

The modern periodic table is based on the modern periodic law. The modern periodic table consists of vertical columns called groups and horizontal rows called periods. The elements are arranged based on the increasing atomic numbers.

Every element in the same group has the same outer electronic configuration of electrons, and every element in the same period has the same number of electron shells.

Depending on which orbital the final electron enters, the periodic table can be classified into four different blocks namely, s-block, p-block, d-block, and f-block.

Periodic trends are those which explain certain properties of elements that are present in the periodic table. The periodic trends are atomic radii, ionic radii, ionisation energy, electron gain enthalpy, electronegativity, metallic character, etc.

Atomic Radius

Atomic radius is defined as the size of an atom, and it is determined by measuring the distance between the centre of the nucleus and the outermost electron in an atom.

Atomic radius decreases as we move from left to right across a period in the periodic table and increases as we move from top to bottom down a group.

Covalent radius (for non-metals) and metallic radius (for metals) can be determined by the concept of atomic radius. It also depends on the number of protons and the attraction between the electrons and the protons. Depending on whether the atom is a metal or non-metal, there are three different types of atomic radii. They are

  1. Metallic Radius
  2. Covalent Radius
  3. Van der Waals’ Radius

Metallic Radius: It is half the internuclear distance that separates the metal cores in a metallic crystal.

r=12AB

Covalent Radius: It is one-half the distance between the nuclei of two covalently bonded atoms of the same element in a molecule.

r=12AB

van der Waals Radius: It is one-half of the distance between the nuclei of two identical non-bonded isolated atoms.

van der Waals Radius>Metallic Radius>Covalent Radius

Factors Affecting Atomic Radii

  1. Number of Shells
  2. Effective Nuclear Charge (ZEff)
  3. Screening Effect

Number of Shells: As the number of shells increases, so do the atomic radii.

Example:

Effective Nuclear Charge: As the effective nuclear charge in an atom increases, the atomic radius decreases.

Example:

Screening Effect: More the screening effect of the electrons in the completely filled orbitals on the outer electrons, the more will be the atomic radius.

Example:

Ionic Radius

Ions are formed either by gaining or losing electrons of an atom. Ionic radius is the distance between the centre of the nucleus and the outermost electron in an ion.

Ionic radius decreases as we move from left to right across a period in the periodic table and increases as we move from top to bottom down a group. Generally, anions have a higher ionic radius than cations.

For the same element.

Ionisation Energy

Ionisation energy is the minimum energy required to remove an electron from an isolated gaseous atom in its ground state.

Ionisation energy increases on going from left to right across a period and decreases on moving down a group. The nature of the chemical bonds and the geometry of the molecule depends on this ionisation energy.

Ionisation energy is the quantitative measure of the tendency of an element to lose an electron.

The smaller the ionization energy, easier it is for a neutral atom to change itself into a positive ion.

Factors Affecting Ionisation Energy

  1. Size of the Atom
  2. Effective Nuclear Charge (ZEff)
  3. Screening Effect
  4. Penetration Effect
  5. Electronic Configuration

Size of the Atom: As the size of the atom increases, ionisation energy decreases. This is because as size increases, the distance between the outermost electron and the nucleus increases. Therefore less energy is required to remove the electron from the atom.

Effective Nuclear Charge (ZEff): As the effective nuclear charge increases, the force of attraction between the nucleus and the outermost electron increases. Therefore, more energy is required to remove an electron from the atom.

Screening Effect: More the screening effect in an atom, the less the ionisation energy. This is because, as the inner electrons screen the outermost electrons from the nucleus, the force of attraction between the nucleus and the outermost electrons decreases. Therefore, less energy is required to remove an electron from the atom.

Penetration Effect: As the penetration of electrons in different orbitals increases, ionisation energy increases. The order of energy required to remove electrons from the orbitals in the same shell iss>p>d>f. s orbital is closer to the nucleus and is therefore more penetrated towards the nucleus. Therefore it is easier to remove electrons from p,d and f orbitals as compared to s orbital.

Electronic Configuration: According to Hund’s rule, the stability of half-filled and fully-filled degenerate orbitals is extremely high. Therefore, the removal of electrons from a half-filled or fully-filled orbital required more energy.

The order of IE for fully-filled, half-filled and partially filled orbitals is as follows.

First Ionisation Energy: It is the energy required to remove the outermost electron from a gaseous neutral atom in the ground state.

Example:

In the above example, the first ionisation energy of beryllium is greater than that of boron. This is because the outermost electron in the case of beryllium is present in the s-orbital which is more penetrated towards the nucleus. In the case of boron, the outermost electron is present in the p-orbital which is less penetrated than the s-orbital. Therefore, removing an electron from a s-orbital required more energy than removing an electron from a p-orbital.

Similarly, the first ionisation energy of nitrogen is greater than that of oxygen. This is because the 2p-orbital in nitrogen is half-filled, which gives it more stability than the partially filled 2p orbital in oxygen.

In the 13th group, the first ionisation energy of thallium is greater than that of gallium. This is due to the poor shielding of outer electrons by the inner d-electrons in thallium and the lanthanoid contraction in gallium.

Second Ionisation Energy: The energy required to remove an electron from a mono-positive isolated gaseous ion.

Third ionisation Energy: The energy required to remove the third most loosely bound electron.

For an isolated atom/ion, the removal of an electron from a cation is more difficult than the removal of an electron from a neutral atom.

Therefore,

Generally, the ionisation energy of non-metals is greater than the ionisation energy of metals. Also, noble gases have the highest ionisation energy value in a given period. Caesium has the lowest ionisation energy and is therefore used in photoelectric cells.

Metallic character

Metallic character is attributed to the ability to lose electrons in a chemical reaction. The elements which are present on the left side of the periodic table have a higher metallic character. It decreases from left to right because of the addition of electrons and increases from top to bottom because of the removal of electrons. Ionisation energy is inversely proportional to the metallic character of an element. This is because as the metallic character increases, the atom can easily lose an electron and therefore less energy is required to remove an electron from the atom.

Electron Gain Enthalpy

The change in enthalpy when an electron is added to a neutral gaseous atom to convert it into a gaseous anion is called electron gain enthalpy.

Electron gain enthalpy is positive when energy is absorbed, and negative when energy is released.

When the electron gain enthalpy is positive, energy is absorbed on the addition of an electron, i.e. ΔegH is positive. In this case, the addition of an electron makes the atom unstable.

When the electron gain enthalpy is negative, energy is released on the addition of an electron, i.e. ΔegH is negative. In this case, the addition of an electron makes the atom stable.

Electron gain enthalpy increases from left to right and decreases from top to bottom of the periodic table.

Factors Affecting Electron Gain Enthalpy:

  1. Atomic Size
  2. Effective Nuclear Charge (ZEff)
  3. Screening Effect
  4. Electronic Configuration

Atomic Size: As the size of the atom increases, the magnitude of electron gain enthalpy decreases. This is because, as the size of the atom increases, the force of attraction between the nucleus and the last shell which receives the incoming electron decreases. Therefore, electron gain enthalpy decreases.

Effective Nuclear Charge (ZEff): The greater the effective nuclear charge, the greater the tendency of the atom to attract the incoming electron towards itself. Therefore, electron gain enthalpy increases.

Screening Effect: More the screening effect of the electrons in the completely filled orbitals on the outer electrons, the more will be the electron gain enthalpy. This is because, as the screening effect increases, the force of attraction between the nucleus and the outermost electrons decreases. Therefore. Electron gain enthalpy also decreases.

Electronic Configuration: Elements with half-filled and fully-filled orbitals are more stable. Therefore, adding an electron to a stable electron configuration required more energy. Hence, electron gain enthalpy has a high positive value.

Example:

In the example given above, neon and beryllium have a fully-filled outer orbital and are therefore more stable. Thus, the electron gain enthalpy in the case of neon and beryllium is the maximum. This is followed by nitrogen, which has a half-filled outer electron configuration.

The magnitude of electron gain enthalpies of group 3 elements is greater than the corresponding second period p-block elements.

The order of electron gain enthalpies of group 17 elements follows the order:

In the above order, chlorine has a higher electron gain affinity than fluorine owing to its larger size. Due to the smaller size of the fluorine atom, adding an additional electron will cause interelectronic repulsion and result in instability of the atom. The same is the case with oxygen in group 16.

Chlorine has the highest negative electron gain enthalpy in the periodic table. The electron gain enthalpies of noble gases are positive i.e. it results in an unstable electronic configuration.

Successive Electron Gain Enthalpy: ΔegH for the addition of a second electron to a neutral atom is positive. This is because interelectronic repulsion outweighs nuclear attraction.

Example:

Electron Affinity

Electron affinity is the energy released when an electron is added to the valence shell of an isolated neutral gaseous atom.

Relation Between Electron Affinity and Electron Gain Enthalpy

For comparison purposes in Periodic Trends, we use both the terms interchangeably.

Electronegativity

Electronegativity is the property of an atom in a molecule to attract the shared pair of electrons towards itself. The electronegativity of any atom is not constant, rather it is relative to the element to which it is bonded. Electronegativity helps in predicting the type of bond formed between two atoms. There are various scales of electronegativity. They are as follows:

The Pauling scale is the most commonly used scale of electronegativity.

Factors Affecting Electronegativity

  1. Atomic Size
  2. Effective Nuclear Charge (ZEff)
  3. Magnitude of Positive Charge on the Atom

Atomic Size: As the size of the atom increases, the force of attraction between the nucleus and the outermost electron decreases. Therefore, the electronegativity of the atom decreases.

Effective Nuclear Charge (ZEff): As the effective nuclear charge increases, the force of attraction between the nucleus and the outermost electron increases. Therefore, the electronegativity of the atom increases.

Magnitude of Positive Charge on the Atom: As the magnitude of the positive charge on the atom increases, the force of attraction between the nucleus and the outermost electron increases. Therefore, the electronegativity of the atom increases.

Electronegativity increases as we move from left to right across a period. This is because, on moving from left to right across a period, more electrons are added to the same shell, which increases the force of attraction between the nucleus and the outermost electrons.

Electronegativity decreases as we move down a group. This is because, as we move down a group, subsequent shells are added and therefore, the distance between the nucleus and the outermost electrons increases. Therefore, the force of attraction between the nucleus and the outermost electrons decreases and subsequently electronegativity decreases.

Fluorine has a higher electronegative whereas caesium has the least electronegativity value. Alkali metals have the lowest, and halogens have the highest electronegativity in their respective periods. It also varies among metals and non-metals. Non-metals are more electronegative than metals.

Valency and Valence Electrons

Valency is the number of electrons in the outermost shell of an atom (or) 8- the number of electrons in the outermost shell.

The electrons in the valence shell are called valence electrons. As we move across the period, the number of valence electrons increases as electrons are added to the same shell in a period. Down the group, the number of valence electrons remains the same as the electronic configuration of elements belonging to the same group is the same.

Chemical Reactivity

By seeing how an element interacts with oxygen and halogens, one can clearly understand the chemical reactivity of the element. Oxides are created when elements react with oxygen. According to the periodic table, elements on the far left react with oxygen to generate basic oxides (Na2O), while elements on the far right react with oxygen to form acidic oxides (Cl2O7).

Amphoteric oxides (Al2O3) are formed by the elements in the middle of the periodic table. Amphoteric oxides exhibit both acidic and basic behaviour.

Facts Based on the Period Trends and Periodic Tables

  • Due to their great stability, the majority of noble gases, including helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn), have zero electronegativity. They do not readily lose or absorb electrons because they have a fully-filled valence electronic configuration.
  • The periodic table has 118 elements, of which 90 are found in nature and the remaining 28 are entirely artificial.
  • While oganesson (Og) is the heaviest element in the periodic table, hydrogen (H) is the lightest element (located in the top left corner) (located in the lower right corner)
  • Metals make up over 75% of the elements in the periodic table. There aren't many non-metals, nevertheless.
  • Mercury and bromine are the only two elements that exist as a liquid at room temperature.

Practice Problems

  1. The outermost electronic configuration of the most electronegative element is:
  1. 1s21p5
  2. 1s22p5
  3. 2s22p5
  4. 2s22p4

Answer: C

Solution: Fluorine is the most electronegative element in the periodic table. Its electronic configuration is 2s22p5. By gaining one electron, it can reach the stable electronic configuration of its nearest noble gas i.e. neon (2s22p6).

So, option C is the correct answer.

  1. According to the modern periodic law, the physical and chemical properties of an element are the periodic function of its:
  1. Atomic Volume
  2. Atomic Mass
  3. Atomic Number
  4. Ionic radii

Answer: C

Solution: As per the modern periodic law, the physical and chemical properties of elements are a periodic function of their atomic number.

So, option C is the correct answer.

  1. Why is the ionisation enthalpy of beryllium greater than that of boron?

Solution: The first ionisation energy of beryllium is greater than that of boron. This is because the outermost electron in the case of beryllium is present in the s-orbital which is more penetrated towards the nucleus. In the case of boron, the outermost electron is present in the p-orbital which is less penetrated than the s-orbital. Therefore, removing an electron from a s-orbital required more energy than removing an electron from a p-orbital.

  1. Why does chlorine have a greater electron affinity than fluorine?

Solution: Chlorine has a higher electron gain affinity than fluorine owing to its larger size. Due to the smaller size of the fluorine atom, adding an additional electron will cause interelectronic repulsion and result in instability of the atom. Therefore, the order of electron affinities of group 17 elements follows the order:

Frequently Asked Questions – FAQ

1. Why do atomic radii increase down the group?
Answer:
Down a group, the number of energy levels (n) increases, so there is a greater distance between the nucleus and the outermost orbital. This results in a larger atomic radius.

2. How do periodic trends relate to periodic law?
Answer: 
Periodic trends give similar patterns in the periodic table showing us the various aspects of an element such as electronegativity, atomic radius, or ionising power. The periodic law tells us that when grouped by atomic number, certain properties of elements occur periodically.

3. Why do electronegativity increases across the period and decrease down the group?
Answer: 
The larger the atom, the lesser the electronegativity, since the electrons being farther away from the nucleus, experience a lesser force of attraction. So, the atomic size increases down the group while electronegativity decreases. Similarly, across a period, electronegativity increases from left to right while atomic radius decreases.

4. Why does metallic character increase down the group?
Answer: Metallic nature increases moving down the group because electron shielding causes the atomic radius to increase, thus, the outer electrons ionize more readily than electrons in smaller atoms.

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Related Topics

Mendeleev’s Classification

Ionic Radii

s-block Elements

Electron Affinity and Electron gain enthalpy

Electronegativity

Factors affecting Electron gain Enthalpy

Atomic Radii

Ionization Enthalpy

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