Major periodic trends include: electronegativity, ionization energy, electron affinity, atomic radius, melting point, and metallic character. Periodic trends, arising from the arrangement of the periodic table, provide chemists with an invaluable tool to quickly predict an element's properties.

click on the link below to understand the trends in properties of periodic table

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• Trends in Physical Properties

One of the important features of long i‘ettn on periodic table it that several properties of elements vary in regular fashion in going some period or going down the group Such properties on called periodic properties. When the elements are arranged in their increasing order of their atomic numbers, they get arranged in such way that the elements with similar properties reccur at regular intervals, The periodic recurrence of elements having similar properties after regular intervals is called periodicity The distribution of electrons in the valence shell of an atom is most important and it is mainly responsible for physical and chemical properties of the elements For example, all halogens present in group No 1 possess ns2 np5 configuration in their valence shell and show similar properties. In this section the periodic trends in certain physical properties are explained in terms of number of electrons and energy levels

Trends in physical properties

There are numerous physical properties of elements such as boiling point and melting point.Heats of fusion and vapourisation, energy of atomization, etc. which show periodic variations However the periodic trends with respect to atomic and ionic radii, ionization enthalpy, electron gain enthalpy electronegativity, valency and oxidation states discussed here.

• Atomic Radius

One must know that size of atom is small (l 2 x 10*m radius) Size of atom is usually expressed in terms of radius called atomic radius. It is very important property because many physical and chemical properties are related to it.

The atomic radius of a chemical element is a measure of the size of its atoms, usually the mean or typical distance from the center of the nucleus to the boundary of the surrounding cloud of electrons. Since the boundary is not a well-defined physical entity, there are various non-equivalent definitions of atomic radius. Three widely used definitions of atomic radius are: Van der Waals radius, ionic radius, and covalent radius.

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Depending on the definition, the term may apply only to isolated atoms, or also to atoms in condensed matter, covalently bound in molecules, or in ionized and excited states; and its value may be obtained through experimental measurements, or computed from theoretical models. The value of the radius may depend on the atom's state and context.

Electrons do not have definite orbits, or sharply defined ranges. Rather, their positions must be described as probability distributions that taper off gradually as one moves away from the nucleus, without a sharp cutoff. Moreover, in condensed matter and molecules, the electron clouds of the atoms usually overlap to some extent, and some of the electrons may roam over a large region encompassing two or more atoms.

Under most definitions the radii of isolated neutral atoms range between 30 and 300 pm (trillionths of a meter), or between 0.3 and 3 ångströms. Therefore, the radius of an atom is more than 10,000 times the radius of its nucleus (1–10 fm), and less than 1/1000 of the wavelength of visible light (400–700 nm).

The approximate shape of a molecule of ethanol, CH₃CH₂OH. Each atom is modeled by a sphere with the element's Van der Waals radius.

For many purposes, atoms can be modeled as spheres. This is only a crude approximation, but it can provide quantitative explanations and predictions for many phenomena, such as the density of liquids and solids, the diffusion of fluids through molecular sieves, the arrangement of atoms and ions in crystals, and the size and shape of molecules.^{[citation needed]}

Atomic radii vary in a predictable and explicable manner across the periodic table. For instance, the radii generally decrease along each period (row) of the table, from the alkali metals to the noble gases; and increase down each group (column). The radius increases sharply between the noble gas at the end of each period and the alkali metal at the beginning of the next period. These trends of the atomic radii (and of various other chemical and physical properties of the elements) can be explained by the electron shell theory of the atom; they provided important evidence for the development and confirmation of quantum theory. The atomic radii decrease across the Periodic Table because as the atomic number increases, the number of protons increases across the period, but the extra electrons are only added to the same quantum shell. Therefore, the effective nuclear charge towards the outermost electrons increases, drawing the outermost electrons closer. As a result, the electron cloud contracts and the atomic radius decrease.

Factors affecting atomic size

1. Number of shells : The atomic size or radius increases with the increase in the number of electronic shells. Atomic radius is proportional to the number of electronic shells.
2. Nuclear charge : The atomic size or radius decreases with the increase in the nuclear charge. If the nuclear charge is more, then the nucleus attracts the electrons towards it and atomic size decreases. In the other words atomic radius is inversely proportional to the nuclear charge.
3. Screening effect or Shielding effect :We know atomic radius decrease if there is more effective nuclear charge, as electron is found in orbital which surrounds the nucleus so it mean if an inner orbital shields the nucleus, the less force of attraction will be perceived by outer valence electron and this cause atomic radius to increase.

Shielding effect of orbitals decrease like this : S>p>d>f. It mean s orbital have maximum capacity of shielding the nucleus, why?, because s orbital is spherical which means it can act like a screen which covers the nucleus by the angle of 360 degree.

• Electronegativity

Electronegativity can be understood as a chemical property describing an atom's ability to attract and bind with electrons. Because electronegativity is a qualitative property, there is no standardized method for calculating electronegativity. However, the most common scale for quantifying electronegativity is the Pauling scale, named after the chemist Linus Pauling. The numbers assigned by the Pauling scale are dimensionless due to the qualitative nature of electronegativity. Electronegativity values for each element can be found on certain periodic tables. An example is provided below.

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Electronegativity measures an atom's tendency to attract and form bonds with electrons. This property exists due to the electronic configuration of atoms. Most atoms follow the octet rule (having the valence, or outer, shell comprise of 8 electrons). Because elements on the left side of the periodic table have less than a half-full valence shell, the energy required to gain electrons is significantly higher compared with the energy required to lose electrons. As a result, the elements on the left side of the periodic table generally lose electrons when forming bonds. Conversely, elements on the right side of the periodic table are more energy-efficient in gaining electrons to create a complete valence shell of 8 electrons. The nature of electronegativity is effectively described thus: the more inclined an atom is to gain electrons, the more likely that atom will pull electrons toward itself.

• From left to right across a period of elements, electronegativity increases. If the valence shell of an atom is less than half full, it requires less energy to lose an electron than to gain one. Conversely, if the valence shell is more than half full, it is easier to pull an electron into the valence shell than to donate one.
• From top to bottom down a group, electronegativity decreases. This is because atomic number increases down a group, and thus there is an increased distance between the valence electrons and nucleus, or a greater atomic radius.
• Important exceptions of the above rules include the noble gases, lanthanides, and actinides. The noble gases possess a complete valence shell and do not usually attract electrons. The lanthanides and actinides possess more complicated chemistry that does not generally follow any trends. Therefore, noble gases, lanthanides, and actinides do not have electronegativity values.
• As for the transition metals, although they have electronegativity values, there is little variance among them across the period and up and down a group. This is because their metallic properties affect their ability to attract electrons as easily as the other elements.

NOTE

• The ionization energy of the elements within a period generally increases from left to right. This is due to valence shell stability.
• The ionization energy of the elements within a group generally decreases from top to bottom. This is due to electron shielding.
• The noble gases possess very high ionization energies because of their full valence shells as indicated in the graph. Note that helium has the highest ionization energy of all the elements.

Some elements have several ionization energies; these varying energies are referred to as the first ionization energy, the second ionization energy, third ionization energy, etc. The first ionization energy is the energy requiredto remove the outermost, or highest, energy electron, the second ionization energy is the energy required to remove any subsequent high-energy electron from a gaseous cation, etc. Below are the chemical equations describing the first and second ionization energies:

• First Ionization Energy:

X(g)→X+(g)+e−(1.1)(1.1)X(g)→X(g)++e−

Second Ionization Energy:

X+(g)→X2+(g)+e−(1.2)(1.2)X(g)+→X(g)2++e−

Generally, any subsequent ionization energies (2nd, 3rd, etc.) follow the same periodic trend as the first ionization energy.

Electron affinity generally decreases down a group of elements because each atom is larger than the atom above it (this is the atomic radius trend, discussed below). This means that an added electron is further away from the atom's nucleus compared with its position in the smaller atom. With a larger distance between the negatively-charged electron and the positively-charged nucleus, the force of attraction is relatively weaker. Therefore, electron affinity decreases. Moving from left to right across a period, atoms become smaller as the forces of attraction become stronger. This causes the electron to move closer to the nucleus, thus increasing the electron affinity from left to right across a period.

NOTE

• Electron affinity increases from left to right within a period. This is caused by the decrease in atomic radius.
• Electron affinity decreases from top to bottom within a group. This is caused by the increase in atomic radius.