Periodic trends in physical properties
Periodic trends in the modern periodic table are the basics of chemistry. It is of worth importance to know about The variation of different physical and chemical properties in the modern periodic table including groups and periods is known as the periodicity of properties.
Explanation of periodic trends
In the Modern Periodic Table, the elements are arranged in the ascending order of their atomic number. Elements are placed in groups and periods on the basis of similar properties and systematic variation of properties.
We will discuss the following periodic properties in the modern periodic table.
- Atomic radius
- Ionic radius
- Ionization energy or ionization potential
- Electron affinity
- Metallic and nonmetallic character
- Melting and boiling point
- Oxidation State
- Electrical conductivity
1: Atomic radius trends
The average distance between the nucleus of the atom and its outermost shell provided that the atom is assumed to be spherical is called the atomic radius.
Definition of atomic radius
Watch The recommended video to better understand atomic radius trends in periodic table
Explanation of atomic radius
The precise measurement of the size of the atom is difficult because of the following reason.
There are no sharp boundaries for the atomic orbital. The reason is that the probability for finding the electron never becomes zero even if the electron is at a large distance from the nuclear.
The distribution of electrons is affected by the neighboring atoms. For example, if the neighboring atom is electronegative in nature then it will pull the electron more towards itself. Hence, the atomic radius will also be changed.
The techniques have been deployed which can measure the distance between the center of two adjacent atoms of any element. Half of this distance is considered to be the radius of the atom.
The atomic radii are expressed in pm.
1 pm = 10-12 m
1 nm = 103 pm
1 Angstrom = 102 pm
Atomic radius trends in the periodic table
Atomic radii vary systematically in the periodic table.
Atomic radius trends in the period
Atomic radii decrease in a period from left to the right in the periodic table.
Why atomic radius decreases in a period?
Due to increasing nuclear charge, the outermost orbital comes closer to the nucleus and atomic radii are decreased.
- The number of shells remain the same
- The shielding effect remains the same
Atomic radius trends in group
Atomic radii increase in a group from upper to downward direction.
Why atomic radii increases in a group?
Number of shells increases down the group due to which the distance between outermost shell and nucleus is also increased. Therefore radius is increase in a group.
Shielding effect increases down the group.
Atomic radius trends in transition elements
d and f block elements are transition elements. Whenever we go from left to right in transition elements series, then there happened the decrease in atomic radius for the first four members, but after that atomic radius shows systematic behavior.
Ionic radius trends in the periodic table
The distance from the center of charged atom to the periphery of the active sphere is called ionic radius.
Definition of Ionic radius
Watch the video lecture on Trends on ionic radii in periodic table
Explanation of ionic radius
When electrons are removed from the atoms positively charged ions are produced. The radii of positively charged ions are smaller than the neutral atoms.
For example
Atomic radius of Na =157 pm
Ionic radius of cation Na+ = 95 pm
Reason
Due to the Greater attraction of the nucleus charge, the remaining electrons of the ions are drawn closer to the nucleus. The following table shows that how many differences are there between atomic radii and positive ionic radii of the atoms.
Negatively charged ion
When the electron is provided to the atom negatively charged ion is produced, which is bigger in size than the neutral atom. The ionic radii of some negatively charged ions are given as follows.
Ionic radius trends in the periodic table
Ionic radii vary systematically in groups and periods in the periodic table.
ionic radius trends in periods
As we move from left to right in a period, the ionic radii of the isoelectronic positive ions go on decreasing. Anyhow, the values of ionic radii of the negatively charged ions go on increasing from left to right.
ionic radius trends in a group
The value of the positive charge remains the same in a group from top to the bottom. The ionic sizes of the positively charged ion go on increasing from top to bottom.
Similarly, the values of ionic radii of negatively charged ions go on increasing down the group.
Comparative study of ionic radii vs atomic radii
If we plot a graph between atomic numbers on x-axis and atomic and ionic radii in pm on y-axis, then following types of graph are obtained. It is clear that neutral atoms are bigger in sizes than positively charged ions. Negatively charged ions are bigger in sizes than neutral atoms.
Graphical representation of atomic radii
The following graphs can help us to have a clear view of the variation of atomic and ionic radii of I-A and VII-A group elements.
Ionization energy trend in the periodic table
It is the minimum amount of energy which is required to remove the most loosely bounded electron from outermost orbital in its isolated state.
It is also known as ionization potential. It is abbreviated as I.E or I.P.
Explanation of ionization energy
Electrons are bounded to the nucleus in an atom, The electrons in the inner levels ( closer to nucleus) have greater force of attraction than the outermost electrons.
The outermost electrons are valance electrons which require less energy to be removed.
Examples of Ionization energy
First ionization energy
The energy required to remove the first outermost electron is called first ionization energy. This happens in case of group II-A and II-A elements.
The removal of electrons from the outermost orbital of the hydrogen atom, sodium, magnesium, and aluminum along with required energies are as follows.
Second ionization energy
When second electron is removed from the outermost orbital of the unipositive ions, then the higher amount of energy is required. This is called second ionization energy of that element.
For example.
Mg——Mg+ + 1e–
Factors affecting the ionization energy values
The ionization energy of an atom depends upon the following factors.
Atomic size
Greater the atomic size of the atom, smaller the ionization energy value and vice versa.
Nuclear charge
Greater the nuclear charge of the atom greater the ionization energy value
Shielding effect
Shielding effect is decreasing force of attraction between the nucleus and outermost electron due to inner electrons. Greater the shielding effect left the ionization energy of the atom.
Nature of orbital
There are four types of orbital as s, p, d, and f. We know that the S orbital is more penetrating and close to the nucleus of the atom. Therefore the removal of electrons from the S orbital is difficult.
P orbital is in the form of lobes and the electron removal from P orbital is easier. Similarly electron removal from the orbital is easier than p orbital. It is probably the easiest to remove electron from the f orbital.
Variation of ionization energy values in the periodic table
Ionization energy values vary systematically in periods and Groups.
ionization energy trend in the period
As we go from left to right in a period ionization energy value increases. This is due to the sizes and increasing in nuclear charges of atoms. Further the shielding effect also remains the same in a period.
ionization energy trend in a group
Ionization energy value decreases down the group. This is due to the increasing size of atom and increasing shielding effect.
Graphical explanation of ionization energy
Groups
When a graph is plotted between atomic number on x axis and ionization energy values on y axis, then following curve is obtained for alkali metals.
Periods
Rising curves obtained for the first, second, and third periods. Anyhow there are certain abnormalities are which have their own reasons.
Alkali metals like Lithium Sodium and potassium have lowest ionization energy values in respective periods.
Ionization energies of noble gases
Ionization energy values of noble gases have the maximum values in respect to periods because their outermost shells are completely filled.
Electron affinity trend in periodic table
The minimum amount of energy that is released the observed when an electron is added to an isolated neutral gaseous atom in its lowest energy state. It produces an anion is called electron affinity.
Definition of electron affinity
Explanation of electron affinity
Electron affinity is a quantitative measurement of the tendency of an atom of an element to accommodate the electron in its outermost orbital.
If the atom has a tendency to accept the electron then it will release the energy and electron affinity is represented by a negative sign. If the energy required by the electron then the electron affinity is positive.
Units of electron affinity
Electron affinity can be expressed in following units
I) KJ / mole
Ii) Kcal / mol
Second electron affinity
When an electron is added to an anion the process is endothermic due to its repulsion. And energy is supplied to form a dinegative ion. This energy is given positive sign and it is called second electron affinity. The reason is that energy is observed and a positive sign is given to the endothermic process.
Examples of electron affinity
Chlorine accepts the electron leases the 337 KJ / mole of energy. This is the first electron affinity.
Oxygen atom accepts one electron to form O negative anion and related 141 kilo joule per mole of energy. This is first electron affinity.
O-1 can accept one more electron to make over -2. It absorbs 880 KJ / mol of energy. This is called second electron affinity.
Factors affecting the electron affinity
There are three factors which control the electron affinity values.
Atomic sizes
Greater the atomic size of an atom smaller the electron affinity
Nuclear charge
Greater the number of protons in the nucleus greater the electron affinity
Shielding effect
Greater the shielding effect of inner levels lesser electron affinity.
Electron affinity trend in a period
Electron affinity is increased from left to right in a period. This is due to the decreasing atomic sizes and encouraging the Nobel charges. What’s the reason alkali metals have the lowest and halogen have the highest value of electron affinity.
Electron affinity trend in group
Electron affinity decreases from upper to downward direction in a group. This is due to the increasing number of shells increasing shielding effect.
Caution
There is an abnormal behavior in electron affinity values from left to right in group II-A, V-A and VIII-A.
Electron affinities of elements of second period are less than the third period. Actually it should not have been so. The reason is that elements of second period have very small sizes and incoming electron is rather repelled.
Metallic character trend in periodic table
Introduction
The elements of periodic table can be broadly speaking be divided into three categories.
- Metals
- Non-metal
- Metalloid
Metals
Element on the left hand side in the centre and at the bottom of the periodic tables are metals.
The elements like sodium, Potassium are not frequently considered as metals in everyday life but chemically these are true metals.
Metallic character trend in the periodic table
Metallic character increases from top to bottom in a group.
Metallic character decreases from left to right in a period.
All the transition elements are metals and the metallic character increases from left to the right up to the middle of the families in d block elements.
Non-metals
Non metals are present on the top right of the periodic table. Greater the tendency to gain the electrons greater the non metallic character.
Non-metallic character trend in periodic table
Non metallic character decreases down the group
Non metallic character increases from left to right in the period
Element on the top right of the periodic table have the maximum non metallic character. For example F, N, O etc, are highly non metallic element.
Metalloids
The lower element of group III-A, IV-A, V-A have intermediate metallic and nonmetallic character and are known as a metalloid.
Trends of melting and boiling point in periodic table
The melting and boiling point of substances tell us about the strength of forces present in the atoms and molecules. In case of elements the values of melting and boiling points are according to the binding energy is present in them.
Periodic Trends of melting and boiling point
There is an almost systematic variation of melting and boiling points in groups and periods of the periodic table.
Trend of melting and boiling point in periods
As we move from left to right in a period, the melting and boiling point increase up to group IV-A.
The melting and boiling point decrease from group V-A to noble gases.
Reasons
- The elements of group I-A are alkali metal which has only one electron in their outermost shells. They can make only one Bond with other atoms.
- The elements of group II-A have a sufficiently high melting and boiling point than I-A. The reason is that atoms in them provide two binding electrons.
- Elements of group IV-A have the highest melting and boiling points due to the presence of four electrons in the outermost orbital.
- Carbon has the highest melting and boiling point when it is in the form of a diamond.
- In diamond, each carbon is SP3 hybridized and four strong Sigma bonds are produced with other atoms. Diamond exists in the form of a giant molecule.
- The elements of groups V-A, VI-A, and VII-A have low melting and boiling points because these elements exist in the form of diatomic molecules like N2, O2, F2, and Cl2, etc. There are the least intermolecular forces and so their melting and boiling points are very low. So much so that they exist in the form of gas at room temperature.
Graphical explanation melting and boiling point in periods
When we plot a graph between atomic number of the element on x axis and melting point on y axis then we come to know that the elements of group IV-A carbon and silicon show the peaks in the curve.
Variation of melting and boiling point in a group
The melting and boiling points of group I-A, II-A, III-A and IV-A decrease down the group due to lesser binding energies. Atomic size is increased down the group.
The elements of group and zero group elements V-A, VI-A, VII-A show increase of melting and boiling points down the group.
Reason
The binding forces present between large atoms of group I-A, II-A, III-A, and IV-A are less. So their melting points and boiling points for less.
In the case of lower elements of group V-A, VI-A, VII-A, and zero group elements have greater polarizabilities so, their melting and boiling points are higher.
Trend of melting and boiling points in group
The graph has been plotted for melting points of elements of group II-A. It is clear that magnesium has the lowest melting point. There is overall decrease of melting and boiling point except radium.
In case of halogens the melting and boiling point increase from chlorine to Iodine. The upper curve is for the boiling points and the lower curve is for the melting point.
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