Ionization Energy and Electronegativity

Atomic Radius

Below is a chart showing the radius of neutral atoms in picometers (1 pm = 1 x 10-12 m) for the s and p block elements. The situation is a little more complicated for the d and f block elements.

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The volume occupied by an atom mostly depends on the electrons. The latest data on the size of a proton gives it as 0.84 femptometer (1 fm = 1 x 10-15 m). For the hydrogen atom with 1 proton and 1 electron, the radius is of the atom is 37 pm with the nucleus making up only 0.00084 pm of that.The atomic radius increases with each filled shell of electrons. For any column in the periodic table, the size increases down a column. So, for example: He The attraction between the positively charged protons and the negatively charged electrons causes a contraction, or a decrease in size as the number of protons increases. In any row, increasing the number of protons decreases the size of the atom even though the number of protons always equals the number of electrons. So, for example: Na > Mg > Al > Si > P > S > Cl > Ar

Ionization Energy

When atoms are ionized they lose an electron and become positively charged.
Electron configurationIonization ReactionEnergy Required
2s1Li Li+ + e-520 kJ/mol
2s2Be Be+ + e-899 kJ/mol
2s22p1B B+ + e-800 kJ/mol
2s22p2C C+ + e-1090 kJ/mol
2s22p3N N+ + e-1400 kJ/mol
2s22p4O O+ + e-1310 kJ/mol
2s22p5F F+ + e-1680 kJ/mol
2s22p6Ne Ne+ + e-2080 kJ/mol
Ionization always requires energy. The amount of energy required to separate one electron from its atom (first ionization energy) depends on how tightly held the electron is. This depends on the number of protons and on the orbitals that the electron occupies. There is extra stability when a type of orbital is half filled or completely filled. Note that it is easier to remove an electron from the singly occupied 2p orbital of boron than from the filled 2s orbital of beryllium. In nitrogen, the 2p orbitals are half filled (special stability) and so it takes more energy to remove an electron from nitrogen than from oxygen.
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Click here for a chart of ionization energies for elements H-Ar by orbitals.

Electron Affinity

The electron affinity is the heat released (negative energy term) when a free electron combines with a neutral atom to make a negatively charged ion. Below are the electron affinities for some halogen elements.
F + e- F- -328 kJ/mol
Cl + e- Cl- -349 kJ/mol
Br + e- Br- -324 kJ/mol
I + e- I- -295 kJ/mol
There is no particular trend for these values with respect to the number of filled shells or the number of protons. Here is a periodic table that includes the electron affinities. Note that the numbers are in kJ of energy released and should be negative numbers.By convention, energy added to a system has a positive value and energy released from a system has a negative value.
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Electronegativity

Electronegativity is a property of atoms within molecules rather than free atoms. It measures the tendency of that atom to draw bonding electrons towards itself. In HF, the fluorine atom is much more electronegative than the hydrogen atom. The electrons in the H-F bond are not equally distributed between the atoms. The electron density is greater around the fluorine atom. In general, the electronegativity increases from left to right in any row of the periodic table and it increases from bottom to top in any column.

See more: Why Is Baso4 Insoluble In Water, Is Baso4 Insoluble In Water

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Bond Polarity

When the electronegativity difference between atoms is 0.5 or greater, we characterize the bond as polar. The H-F bond is polar but the C-H bond in CH4 is non-polar. We can represent the bond polarity of HF as a vector. The bond is covalent but there is excess electron density around F, giving it a partial negative charge, and a deficiency of electron density around H, giving it a partial positive charge.
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The opposite of electronegative is electropositive. Very electropositive elements, such as Na, typically form salts rather than covalent compounds.BackCompassTablesIndexIntroductionProfessor Patricia Shapley, University of Illinois, 2012