The ionization energy of an element is the minimum energy required to remove an electron from the valence shell of an isolated gaseous atom to form an ion .
For example: Mg → Mg + e– ΔH = 738kjmol-1
In the gaseous phase, the atoms and ions are isolated and are free from all external influences. Thus, the ionization energy is the qualitative measure of the stability of an atom. The higher ionization energy cause difficult to remove an electron. It is important because it can be used to predict the strength of chemical bonds.
Unit: In chemistry, the unit of ionization energy is expressed as kilojoules per mole (Kj/mol) or kilocalories per mole (kcal/mol). In physics, it is measured in term of electronvolts .
Types of ionization energy
First ionization energy
The first ionization energy is the energy required to remove 1 electron from the valence shell. The first ionization energy of sodium ions is +496 kjmol-1.
Na → Na+ + e– ΔH = +496 kJmol-1.
H → H+ + e– ΔH = -1312.0 kJmol-1.
The first ionization energy of hydrogen is about half in a chemical reaction. When natural gas burnt about 800 kJ of energy is released.
CH4 (g) + 2O2 (g) → CO2 (g) + 2 H2 O (g) ΔH = -802.4 kJmol-1
Second ionization energy
When there is more than one electron in the valence shell, so they can be removed one by one with the help of providing more energy. Like elements of group 2 and 3 they have more than one electron in their outermost shells. That’s why they have more than one ionization energy. Each succeeding energy is larger than the preceding one.
Second ionization energy is higher than the first ionization energy because removing the first electron gives the atom stable electron shell like an alkali metal but removing the second electron involves a new electron shell that is closer and tightly bound to the atomic nucleus .
Mg → Mg+ + e– ΔH = 738 kjmol-1.
Mg → Mg2+ + e– ΔH = 1451 kjmol-1.
To carry out chemical reactions, the role of ionization energy is significant. As it determines the probability of the reactants to form covalent or ionic bonds. For example the ionization energy of sodium is 496 kJ/mol but chlorine first ionization energy is 1251.1 kJ/mol, so due to difference in the energy values the chemically combine and form ionic bonds. Elements that lies very close in the periodic table form covalent bonds due to less difference in ionization energies value. For example carbon and chlorine form CCl4.
Periodic Table Trend
In periodic table ionization energies increases from left to right in a period with the increase in proton number. Each period begins with an element which has one electron in its valence shell and ends with the completion of an electronic shell. Increase in atomic number is associated with increase in nuclear charge which leads to a strong force of attraction.
Elements on the left side of the periodic table have less ionization values as compare to the elements on the right side of the periodic table.
|2nd period elements||3Li||4Be||5B||6C||7N||8O||9F||10Ne|
|Ionization energy (kJ/mol)||520||899||801||1086||1402||1314||1681||2081|
In groups, the ionization energy decreases due to increase in nuclear charge. Due to addition of electronic shells increases the distance between the nucleus and outer electron. That’s why electron are easily removed with less energy. Force of attraction is also decreasing due to increase in shielding effect of intervening electrons.
|1st group elements||Ionization energy (kJ/mol)|
Factors Affecting Ionizing Trend
Ionization potential is an older term used for ionization energy because in history ionization energy was measured through ionizing a sample and accelerating the electron. The electrons are removed through electrostatic potential.
Here are some factors which are affecting ionization energy:
- Nuclear charge
Higher the magnitude of nuclear charge the more closely electrons are held by the nucleus. That’s why more will be the ionization energy.
- Number of electron shells
On the size of atom ionization energy will be measured. If the size of atom is greater less tightly the electrons are held so the ionization energy will be less.
- Effective nuclear charge (Z eff)
Higher the magnitude of electrons and shielding power. The electrons are less tightly present by the nucleus and that’s why there will be less ionization energy.
- Type of orbital ionized
If the atom has stable electronic configuration then it will be very difficult to remove electron from outermost shell which results in higher ionization energy.
- Occupancy of orbital matters
If the orbit is completely filled or half-filled then it is very difficult to remove the electrons.
The ionization energy of (n+1)th is higher than the nth ionization energy because removal of next electron from same shell increases the ionization energy due to increased net charge of the ion. Element also has greater electrostatic force of attraction because of less distance between the nucleus and the electron. These two factors also increase the ionization energy of atom.
However there are certain elements which show irregular trends in period 3. For example ionization energy of Mg and P are higher than those of Al and S. In case of Mg (1s2, 2s2, 2p6, 3s2) it is more difficult to remove electron from completely filled 3s orbital while in case of Al (1s2, 2s2, 2p6, 3s2, 3p1) it is easier to remove the same from partially filled 3p orbital. Since to remove an electron from a 3s-orbital of Mg atom requires more energy than to remove the same from a 3p- orbital of Al atom that is why ionization energy of Mg is higher than that of Al.
The electron affinity of an atom is the energy released when an electron adds to an empty or partially filled orbital of an isolated gaseous atom in its valence energy level to form an anion having a unit negative charge.
Cl (g) + e– → Cl– (g) ΔH=-349kJ/mol
Since energy is released, so electron affinity is given the negative sign. Electron affinity is the measure of the attraction of the nucleus of an atom for the extra electrons.
Factor Influencing Electron Affinity
Like ionization energy, electron affinity is influenced by atomic radius, the nuclear charge and the shielding effects of inner electrons. As the force of attraction between the valence electrons and the nucleus decreases with the increase in the atomic radius, the electron affinities usually decreases.
Variation in the Periodic Table
In a period the atomic radius decreases due to increase in the nuclear charge. Thus the electron affinities of element increases from left to right in the periodic table. That is why the alkali metals have the lowest and the halogens have the highest electron affinities.
In groups atomic radii increase with the increase in the proton number due to successive increase of electronic shells. This also exerts a shielding effect on the force of attraction between nucleus and the valence electron. Thus electron affinities usually decreases from top to bottom.
Prediction of Ionic and covalent bonds
There are many types of bonding and forces are present that bind molecules but two basic types are ionic and covalent bond .
It is defined as complete transfer of valence electron between atoms. In ionic bonding it require one electron donor (metal) and one electron acceptor (non-metal). The metal loses electrons to become a positively charged cation and non-metal accepts that electron to become a negatively charged anion.
For example: Sodium (Na) is donating 1 electron to the chlorine (Cl) atom. The net charge of the resulting compound is zero.
Na + Cl → Na+ Cl–
Overall predicted that ionization energy of metal and electron affinity of non-metal compound is positive. The attraction between these particles release energy to carry out the reaction. Due to opposite charge of both compounds they are mostly dissolve in polar solvents.
It is defined as complete sharing of electrons between atoms. The bonding occurs between atoms having same electronegativity because both have no tendency to donate electrons but they share electrons to complete their outermost shell and become stable.
For example: Carbon share its 4 electrons through single, double and triple bonds.
Electron binding energy
It is also known as ionization potential or energy, it is that amount of energy which is required to free electrons from a molecule or an atom or an ion.
Binding energy of a single proton or neutron in nucleus is million times greater than electron in an atom .
For example: Hydrogen atom has one electron in its outermost shell so it is very harder to remove electron because there is less distance between the nucleus and the electron greater electrostatic force of repulsion is required to remove an electron.
- “Ionization Energy”. ChemWiki. University of California, Davis. 2013-10-02.