This is a story of two business partners A and B who shared their investments equally among themselves, yet one of them wants all of it . A tries to grab everything he can from the other partner, B. A will be successful in doing so because he is stronger and powerful than B.
This happens even in atoms which share electrons between them. The successful atom which manages to pull electrons towards itself is the atom with high electronegativity and hence more powerful in this case.
But, what is electronegativity? Why do atoms of some elements have high electronegativity while others are less electronegative? We will answer these questions in detail in the following article.
- This article is about electronegativity, which comes under bonding in physical chemistry.
- First, we will define electronegativity and look at the factors affecting it.
- After that, we will look at the electronegativity trends in the periodic table.
- Then, we will look at electronegativity and bonding.
- We will then relate electronegativity and bond polarisation.
- Finally, we will look at the electronegativity formula.
Electronegativity definition
Electronegativity is the ability of an atom to attract the bonding pair of electrons in a covalent bond to itself. This is why its values can be used by chemists in order to predict whether bonds between different types of atoms are polar, non-polar, or ionic. Many factors affect electronegativity within atoms; there are also trends relating the elements in the periodic table to electronegativity.
Electronegativity is the power and ability of an atom to attract and pull a pair of electrons in a covalent bond towards itself.
Which factors affect electronegativity?
In the introduction one of the questions we intended to discuss was- " Why do atoms of some elements have high electronegativity while others are less electronegative? " This question will be answered in the following section where we are going to discuss the factors that affect the electronegativity.
Atomic radius
Atoms do not have a fixed boundary like spheres do, and hence it is difficult to determine and define the radius of an atom. But, if we consider a molecule with a covalent bond between them, half of the distance between the nuclei of the two covalently bonded atoms is considered as the atomic radius of one atom participating in the bond formation. Other types of radii are Vanderwaal's radius, ionic radius and metallic radius.
Not every time the atomic radius is the exact half of the distance between nuclei of the bonded atoms. It depends on the nature of the bond, or to be precise, the nature of the forces between them.
Based on the above explanations,theoretically, we can describe that atomic radius is the distance between the centre of the nucleus and the outermost orbital.
The shorter the distance between the outer electrons and the positive nucleus, the stronger the attraction between them. This means that if the electrons are further away from the nucleus, the attraction will be weaker. Therefore, a decrease in the atomic radius, results in an increase in electronegativity.
As explained above, the covalent radius is half the distance between the nuclei of covalently bonded atoms. Ionic radius is not the exact half, because the cation is smaller than the anion, the size of the cation (ionic radius of the cation) is smaller as compared to that of the anion.
Nuclear charge and Shielding effect
As the name indicates, nuclear charge is the charge of the nucleus felt by the electrons. The nucleus has protons and neutrons,as we already know, with protons carrying positive charge while neutrons are neutral. So, nuclear charge is the pull of the protons felt by the electrons.
The nuclear charge is the attractive force of the nucleus, caused by protons, on the electrons.
As the number of protons increase, the 'pull' felt by the electrons increases. As a result, the electronegativity increases. Hence, in a period from left to right, the increase in electronegativity is attributed to the increase in the nuclear charge.
But, for the outer electrons, to experience this pull, there is a problem called screening effect or shielding effect.
The inner shell electrons repel the outer electrons and will not let the outer electrons experience the love of the nucleus. Thus, as the number of shells increase down the group, the electronegativity decreases due to reduced nuclear charge because of shielding effect.
Beware! Do not confuse nuclear charge with an element or compound having a charge.
Effective Nuclear Charge
Effective nuclear charge, Zeff is the actual pull of the nucleus felt by the outer electrons in the outer shells after cancelling the repulsions experienced by the outer electrons from inner electrons.
This is because the inner electrons shield the nucleus from the outer electrons by repelling them. Hence, the electrons closest to the nucleus experience greater pull while the outer electrons will not due to repulsions from inner electrons.
Fig. 1: Effective nuclear charge and shielding effect | Archana Tadimeti | Vaia Originals
As we move across a period from left to right, the number of inner electrons remain the same, meaning the shielding effect is the same, but the number of valence electrons and the number of protons increase. This will lead to a greater pull of electrons by the nucleus, thus in turn resulting in an increase in the effective nuclear charge. The greater the effective nuclear charge, the greater the attraction of the nucleus towards the valence electrons. Thus, electronegativity also increases across the period from left to right owing to the diminishing shielding effect and increase in Zeff . This is the reason why group 7 elements have high electronegative values and fluorine is the element with the highest electronegativity.
Let us compare the electronegativities of oxygen and nitrogen to understand this concept better.
Nitrogen and oxygen | Inner and outer electrons | Shielding effect | Archana Tadimeti | Vaia
The electronegativity of nitrogen is 3.0, while that of oxygen is 3.5. The increase in electronegativity is due to the increase in Zeff as explained before.
Electronegativity trends in the periodic table
Let's look at some basic trends in electronegativity, which generally hold true in the periodic table.
Electronegativity down a group
Electronegativity decreases going down a group in the periodic table. The nuclear charge increases as protons are added to the nucleus. However, the effect of shielding is also increased as there is an extra filled electron shell in each element going down a group. The atomic radius of the atom increases as you go down the group since you are adding more shells of electrons, which makes the atom larger. This leads to an increase in the distance between the nucleus and the outermost electrons, meaning that there is a weaker force of attraction between them.
Electronegativity across a period
As you go across a period in the periodic table, electronegativity increases. The nuclear charge increases because the number of protons in the nucleus increases. However, shielding remains constant since no new shells are being added to the atoms, and electrons are being added to the same shell each time. As a result of this, the atomic radius decreases because the outermost shell is pulled closer to the nucleus, so the distance between the nucleus and the outermost electrons decreases. This results in a stronger attraction for the bonding pair of electrons.
Fig. 3: The periodic table | Electronegativity trends | Vaia Originals
Electronegativity of elements and bonding
The Pauling scale is a numeric scale of electronegativities that can be used to predict the percentage ionic or covalent character of a chemical bond. The Pauling scale ranges from 0 to 4.
Halogens are the most electronegative elements in the Periodic Table, with fluorine being the most electronegative element of all, with a value of 4.0. The elements that are least electronegative have a value of approximately 0.7; these are caesium and francium.
Single covalent bonds can be formed by the sharing of a pair of electrons between two atoms.
Examples of molecules made up of a single element are diatomic gases, and molecules such as H2, Cl2, and O2. Molecules made up of a single element contain bonds that are purely covalent. In these molecules, the difference in electronegativity is zero since both atoms have the same electronegativity value and, therefore, the sharing of electron density is equal between the two atoms. This means that the attraction towards the bonding pair of electrons is equal, resulting in a non-polar covalent bond.
Fig. 4: Electronegativity- a tug of war between atomic nuclei | Created from image Image by Mohamed Hassan via Pixabay
However, when atoms with different electronegativities form a molecule, the sharing of electron density is not equally distributed between the atoms. This results in the formation of a polar covalent bond. In this case, the more electronegative atom (the atom with the higher value in the Pauling scale) attracts the bonding pair of electrons towards itself. Due to this, partial charges appear on the molecule, since the more electronegative atom gains a partial negative charge, while the less electronegative atom gains a partial positive charge.
An ionic bond is formed when one atom completely transfers its electrons to another atom which gains the electrons. This occurs when there is a large enough difference between electronegativity values of the two atoms in a molecule; the least electronegative atom transfers its electron(s) to the more electronegative atom. The atom which loses its electron(s) becomes a cation which is a positively charged species, whilst the atom which gains the electron(s) becomes an anion, which is a negatively charged species. Compounds such as magnesium oxide (\(MgO\)), sodium chloride( \(NaCl\) ), and calcium fluoride( \(CaF_2\) )are examples of this.
Usually, if the difference in electronegativity exceeds 2.0, the bond is likely to be ionic. If the difference is less than 0.5 then the bond will be a non-polar covalent bond. If there is an electronegativity difference between 0.5 and 1.9, then the bond will be a polar covalent bond.
Difference in Electronegativity | Type of Bond |
\(>2.0\) | ionic |
\(0.5~to~1.9\) | polar covalent |
\(<0.5\) | pure (non-polar) covalent |
It is important to remember that bonding is a spectrum, and some boundaries are not clear-cut. Some sources claim a polar covalent bond to be only until 1.6 in the electronegativity difference. This means that bonding needs to be judged on a case-to-case basis rather than always sticking to the rules above.
Let's have a look at some examples. Take \(LiF\):
The electronegativity difference for this is \(4.0 - 1.0 = 3.0\); therefore this represents an ionic bond.
\(HF\) :
The electronegativity difference for this is \(4.0 - 2.1 = 1.9\); therefore this represents a polar covalent bond.
\(CBr\):
The electronegativity difference for this is \( 2.8 - 2.5 = 0.3\); therefore this represents a non-polar covalent bond.
Note that no bond is 100% ionic. A compound which has more ionic character than covalent is regarded as an ionic bond while the molecule which has more covalent character than ionic is a covalent molecule. For example, \(NaCl\) has 60% ionic character and 40% covalent character. Thus, \(NaCl\) is regarded as an ionic compound. This ionic character arises due to the differences in electronegativity as discussed previously.
Electronegativity formula
As shown above, one can see all the Pauling electronegativity values of the elements from a dedicated Periodic Table. To calculate the bond polarity of a molecule, you have to subtract the smaller electronegativity value from the larger one.
Carbon has an electronegativity value of 2.5, and chlorine has a value of 3.0. So, if we were to find the electronegativity of the \( C-Cl bond\) , we would know the difference between the two.
Therefore, \(3.0 - 2.5 = 0.5\) .
Electronegativity and polarisation
If the two atoms have similar electronegativities, then the electrons sit in the middle of the two nuclei; the bond will be non-polar. For instance, all diatomic gases such as \(H_2\)and \(Cl_2\)have covalent bonds which are non-polar as the electronegativities are equal in the atoms. Therefore, the attraction of electrons to both nuclei is also equal.
If two atoms have different electronegativities, however, the bonding electrons are attracted towards the atom which is more electronegative. Because of the uneven spread of electrons, a partial charge is assigned to each atom as mentioned under the previous heading. As a result, the bond is polar.
A dipole is a difference in charge distribution between two bonded atoms that is caused by a shift in electron density in the bond. The electron density distribution depends on the electronegativity of each atom.
You can read about this in more detail in Polarity.
Fig. 5: Diagram showing the bond dipole. Sahraan Khowaja, Vaia Originals
Thus, a bond is said to be more polar if the difference in electronegativity is larger. Therefore, there is a larger shift in electron density.
Now, you might have grasped the meaning of electronegativity, factors and trends of electronegativity. This topic is a foundation for many aspects of chemistry, particularly organic chemistry. Hence, it is important to get a thorough understanding of the same.
Electronegativity - Key takeaways
- The factors that affect electronegativity are atomic radius, nuclear charge, and shielding.
- The electronegativity decreases as you go down a group in the periodic table and increases as you go across a period.
- The Pauling scale can be used to predict the percentage ionic or covalent character of a chemical bond.
- The more electronegative atom pulls the bonding pair of electrons towards itself.
- A dipole is a difference in charge between two bonded atoms that is caused by a shift in electron density in the bond.
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