Periodic Trends in Properties of Elements

What are Properties of Elements? Definition and Examples

The periodic table is a tabular display of various chemical elements that are arranged according to their electronic configuration, atomic number, and common chemical properties. There are certain trends that are common throughout all groups and periods.

The rows are termed as periods and the periodic table has 7 periods. The metals are present on the left rows and towards the right-hand side are the nonmetals. Whereas, the columns are termed as groups. In groups, elements have various chemical behaviours. There are 18 groups and the halogens are present under group 17 and noble gases are present under group 18.



Periodic trends

The specific patterns in the properties of chemical elements present in the periodic table are known as periodic trends. The important trends are,
 
  • 1. Ionization energy

  • 2. Metallic character

  • 3. Atomic Radii

  • 4. Electronegativity

  • 5. Ionic radius

  • 6. Electron affinity

  • 7. Chemical reactivity

  • 8. Shielding effect

  • These trends arise due to changes in the structure of atoms of the elements within their groups and periods. A few exceptions exist, for example, the ionization energy of groups 3 and 6.

    Periodic law

    Periodic law forms the basis for periodic trends. According to periodic law, “the chemical elements are listed in an order of increasing atomic number, and main properties thus undergo cyclic changes. Elements having similar chemical properties re-occur in regular intervals”

    This principle was given by Dmitri Mendeleev. He also stated that the periodic table was not just based on the atomic weights, but also based on various physical and chemical properties of elements.
    Later it was also found that the reoccurrence of properties was due to the reoccurrence of similar electronic configurations in the outer shells of atoms.

  • 1. Ionization energy

  • The ionization potential can be defined as,
    “Minimum energy required by an isolated atom to remove one electron in its neutral or gaseous state”
    As one goes across the period, the ionization energy increases. The reason behind is that the nuclear charge across the period increase and thus the electrons are strongly held by the nucleus.
    But as one goes down the group, the ionization energy decreases down the group. The reason behind this is, down the group the valence electrons go farther away from the nucleus, thus the nuclear charge decreases.
    Factors affecting ionization energy

    There are various factors that affect the ionization energy levels 

  • Nuclear charge

  • Lower the nuclear charge lower is the force of attraction between the nucleus and valence electrons, thus low ionization energy. 

  • Shielding effect

  • Shielding effect increases as nuclear charge increases, thus with an increase in shielding effect the ionization energy also increases. 

  • Atomic radius

  • As the atomic radius increases the force of attraction between the nucleus and valence electrons also decreases. Thus, with an increase in atomic radius the ionization decreases.
     
  • Half-filled valence shells

  • Pseudo filled or half-filled valence shells have high ionization energy.

    A simple principle that can be used is that, if the principal quantum number is low, then the ionization number will be high for the electron present in that shell.

    Exceptions

    All the elements in the oxygen and boron family are an exception to the above stated periodic trend. They require a little less energy than the usual trend.

  • 2. Metallic property

  • Metallic property of an element can be defined as their ability to conduct electricity. The metallic properties increase down the group as the nuclear charge decreases down the group. Since the valence electron is loosely bounded by the nuclei, they are able to conduct electricity well.

    But across a period, the metallic character decreases as nuclear charge increases. This causes the force of attraction between the valence electrons and the nuclei increases, thereby inhibiting them from conducting electricity or heat.

  • 3. Atomic Radii

  • The atomic radius is the distance between the atomic nucleus and outermost stable electron orbital of an atom which is at equilibrium. Across a period the atomic radius decreases, as the nuclear charge increases. The reason for the decrease is as nuclear charge increases, the force of attraction between the nucleus and the valence electrons also increases, and the nucleus holds the electron tightly, thereby decreasing the atomic radii.

    In a group, the atomic radius increases down the group. The reason being, new shells are being added and thus the nuclear charge decreases. But the atomic radii also increase diagonally causing some exception.
    Example

    Along the Period – Li> Be > B > C > N > O > F

    Down the Grp – Li < Na < K < Rb < Cs

  • 4. Electronegativity

  • Electronegativity can be defined as the ability of an atom or a molecule to attract a pair of electrons. The bond formed due to this is mainly determined by the difference between the electronegativity of the atoms.
    Across the period, the electronegativity increases as nuclear charge increases. Moving down a group, the electronegativity decreases as nuclear charge decreases. The reason being the distance between the nucleus of the atom and the valence electrons is long and thus the electrons are easily lost.

    Example

    Along the Period- Li < Be < B < C < N < O < F

    Down the Grp - Li > Na > K > Rb > Cs

    Exception

    The group 13 elements are an exception and thus the electronegativity increases from aluminum to thallium. Also, in group 14, the electronegativity of tin is higher than lead.

  • 5. Electron affinity

  • Electron affinity can be defined as the tendency of an atom to accept an electron or an electron pair. This is a characteristic feature of nonmetals as they gain electrons to become anions. Across a period, the electron affinity increase as nuclear charge increases

    Down the group, it decreases, as the nuclear charge decreases. Fluorine has the highest electronegativity and noble gases are not included in this. The reason being they have full valence shell and thus can neither gain nor lose electrons.

  • 6. Shielding effect

  • It can be defined as the repelling of an outer electron by the inner electrons. It can also be used to explain how many nuclei can control the outer electrons. The effective nuclear charge decreases down the group due to increased shielding effect. Across a period, the effective nuclear charge increases as nuclear charge increases.
    To summarize the whole thing, we can make the following conclusions.

    CharacteristicPeriod Group
    Ionization energyIncreases Decreases
    Metallic propertyDecreasesIncreases
    Atomic radiusDecreasesIncreases
    Electron negativityIncreasesDecreases
    Electron affinityIncreasesDecreases
    Shielding effectIncreases Decreases