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Lesson Video: Electronegativity Chemistry

In this video, we will learn how to explain the chemical property of electronegativity.

17:26

Video Transcript

In this video, we will learn how to explain the chemical property of electronegativity. When atoms of nonmetals react with each other and bond together, covalent bonds are normally formed. These covalent bonds are widely seen in simple molecules and they hold their atoms together.

So what causes the atoms to be attracted to one another? Quite simply, it’s the electrostatic attraction between the shared pair of electrons and the positively charged nucleus of each atom in the covalent bond. The more electrostatic attraction we have, the stronger this covalent bond will be. The electrostatic attraction is largely governed by the size of the positive charges in each nucleus. Since the positive charge in each nucleus is caused by positively charged protons, its size is determined by how many protons there are in each nucleus. It’s not just this simple, though, as the size of each atom is important, too. If the two atoms are large, the distance between the nucleus of each atom is greater. This makes the electrostatic attraction weaker. And we get a weaker covalent bond.

Remember that a shared pair of electrons is just a shared pair of electrons, and the negative charge of these two electrons does not change. If the atoms involved in the covalent bond are identical, the attraction for the shared pair of electrons will also be identical. We could imagine this situation to arise in a molecule of chlorine. The single covalent bond formed in a chlorine molecule contains a shared pair of electrons. One electron comes from each chlorine atom, respectively, shown by a dot and a cross here. So in this chlorine molecule, we can imagine that there is perfect sharing of the pair of electrons in this covalent bond, since the number of protons in each nucleus is identical and the size of each chlorine atom is the same.

This perfect sharing scenario will only arise in covalent bonds formed between two identical atoms in a diatomic molecule. Think of H O N Cl Br I F to remember the most commonly encountered elements that exist as diatomic molecules. We should find perfect sharing of the electron pair in the covalent bonds of hydrogen, oxygen, nitrogen, chlorine, bromine, iodine, and fluorine.

We can imagine the attraction between the shared pair of electrons and the nucleus of each atom in a covalent bond as being similar to a tug of war. In a tug of war, two teams pull on opposite ends of a strong rope with a marker tied in the middle. The team with the biggest pulling power or traction wins by pulling the rope marker over a line on the ground first. Of course, covalent bonds do not involve pieces of rope and teams of people. But imagine that the nucleus of each atom is like the team of people and the rope marker is the shared pair of electrons. It is reasonable to assume that a larger nucleus with more protons will behave like a larger or stronger team of people in a tug of war. A larger nucleus will attract the shared pair of electrons in a covalent bond more strongly than a smaller nucleus with less protons.

When two different atoms are bonded together in a covalent bond, we may see uneven sharing or less-than-perfect sharing of the electron pair in the covalent bond. This could occur if a hydrogen atom is covalently bonded to a chlorine atom. Chlorine has a much larger nucleus than hydrogen. Chlorine has 17 protons or positive charges in its nucleus versus hydrogen’s one. The total amount of positive charge found in a nucleus is often referred to as the nuclear charge. Since the chlorine atom has a bigger nuclear charge than the hydrogen atom, the sharing of the electron pair in the covalent bond is not equal. The electron pair is attracted more strongly to the chlorine atom. The concept of how well electrons are shared in covalent bonds is called electronegativity.

Electronegativity is now defined as the relative attraction that an atom has for the shared pair of electrons in a covalent bond. This concept was developed using a scale by the American scientist Linus Pauling in 1932. Linus Pauling called electronegativity the power of an atom in a molecule to attract electrons to itself. Linus Pauling introduced the scale for electronegativity. He calculated relative values for each element, the calculation involving bond dissociation energies. Bond dissociation energies give an idea of how much energy is needed to break a specific covalent bond. The Pauling electronegativity scale has no units, and the numbers range from 0.79 for cesium and francium to 4.0 for fluorine.

There are definite patterns for the electronegativity values for each element as we move around the periodic table. Electronegativity displays periodic trends. As we move across a period from left to right in the periodic table, electronegativity values generally increase. If we exclude the transition elements where there’s not always a clear trend, electronegativity values increase across each period. Across period two, we see values range from 1.0 for lithium to 4.0 for fluorine, which is in fact the most electronegative element of all. A similar pattern emerges for period three elements sodium to chlorine. Here, the values range from 0.9 for sodium to 3.2 for chlorine.

In any given period, the corresponding halogen in group 17 or seven emerges as the most electronegative element. This is because the number of protons in the nucleus, or nuclear charge, is increasing as we move across each period. There are no new shells being occupied, so no additional screening or shielding of the valence electrons occurs. Screening or shielding is when occupied inner shells reduce the effect of the nuclear charge on the valence shell electrons. A greater nuclear charge attracts bonding pairs of electrons more strongly. The increasing nuclear charge across each period also reduces the size of the atoms, so atomic radii decrease. This also adds to the attraction of valence shell electrons as they’re closer to the nucleus.

All of these factors explain why electronegativity values increase as we move across a period in the periodic table. As we move down each group in the periodic table, electronegativity values decrease. This trend is seen in group one, where values range from 2.2 for hydrogen to 0.7 for francium. This trend is repeated as we move down group 17, where values range from 4.0 for fluorine to 2.2 for astatine. So as we move down each group in the periodic table, electronegativity values decrease. This is because we are adding occupied shells and the atomic radii or the size of the atoms are getting larger.

Although the nuclear charge is increasing as we move down the group, its effect is shielded by the core electrons. These factors combined explain why electronegativity decreases as you move down a group. So on a map of the periodic table, the least electronegative elements are found on the lower-left corner and the most electronegative elements on the upper-right corner. It is worth noting here that since noble gases do not generally form compounds, they are not assigned electronegativity values.

It’s not really too important to learn electronegativity values for each element. We can use them, however, to determine whether specified covalent bonds will exhibit perfect sharing or not. For this purpose, it’s more important to look at differences in electronegativity values. If two elements with similar electronegativity values are covalently bonded together, we can see that there is little to no difference in the electronegativity values. An example of this is a carbon-to-hydrogen covalent bond. Carbon has an electronegativity value of 2.6. Hydrogen has an electronegativity value of 2.2. The difference in electronegativity is 2.6 minus 2.2, which equals 0.4. This difference is considered to be very small and the bond is said to be a nonpolar covalent bond. There is no difference in the sharing of the electron pair between the nucleus of each atom in this covalent bond.

Notice the difference of less than 0.5 is considered to give rise to a nonpolar covalent bond. In contrast, if we examine the hydrogen-to-fluorine covalent bond, we find that there is a difference in electronegativity of 4.0 minus 2.2, which equals 1.8. This difference is considered to be large and will give rise to a highly polar covalent bond. The sharing of the electron pair in this covalent bond is very uneven. Using electronegativity value differences, covalent bonds can be considered to lie on a spectrum from nonpolar, perfect sharing, to highly polar, very uneven sharing and in fact almost ionic.

The differences in electronegativity values are summed up in this table with descriptions for the type of bond expected. We can see that if the difference in electronegativity of the two atoms bonded together lies at the extreme end, for example, rubidium bonded to fluorine, then the bonding is not expected to be covalent at all. In this case, there is electron transfer between rubidium and fluorine, resulting in a fully ionic bond. So it’s more useful to consider differences in electronegativity when determining the nature of the sharing of the electron pair in a covalent bond and, therefore, how polarized that covalent bond would be expected to be. We will now look a question to test your understanding of the concept of electronegativity.

Why is the element potassium less electronegative than the element lithium? (A) Potassium has a greater number of protons and greater nuclear charge to attract the bonding electrons. (B) Lithium has less electrons than potassium. (C) Potassium does not follow the general trend of electronegativity. (D) Bonding electrons are further from the nucleus in atom of potassium causing less attraction. (E) Electronegativity increases as you descend a group in the periodic table.

In this question, we’re being asked to compare the electronegativity of two elements. Remember that electronegativity is the attraction that an atom has for the shared pair of electrons in a covalent bond. To answer this question, we’ll check the validity of each statement against the main factors that influence electronegativity. Electronegativity depends on the number of protons in the nucleus of an atom, sometimes referred to as the nuclear charge. It also depends on the number of shells occupied, which influences the size of an atom and also the amount of shielding or screening taking place in that atom.

As we move across a period in the periodic table, electronegativity values generally increase. This is due to increasing nuclear charge, the same number of shells being occupied, and therefore no additional screening or shielding. As we move down a group, electronegativity values decrease. Although there are more protons in the nucleus of these atoms, the number of shells being occupied is increasing. This increases the amount of shielding or screening. And this factor outweighs the increasing nuclear charge.

We’re comparing potassium, symbol K, with the element lithium, symbol Li. Both metals are located in group one of the periodic table. And since potassium is further down the group than lithium, the element potassium is certainly less electronegative than lithium. Statement (A) suggests that potassium has a greater number of protons and hence nuclear charge than lithium. This part of the statement is true. Lithium has three protons in its nucleus; potassium has 19. This statement could be a good argument to explain why potassium is more electronegative than lithium. This argument would be tempting because electronegativity does depend on the size of the nuclear charge. This statement does not provide any support, though, as to why potassium is less electronegative than lithium. It’s therefore not a correct answer.

Despite the greater nuclear charge in potassium, we find there’s more shielding in this atom, and this outweighs the factor of the increasing nuclear charge as far as electronegativity is concerned. Statement (B) suggests that lithium has less electrons than potassium. On the face of it, this statement is true. Lithium contains three electrons, and potassium contains 19. Simply stating how many electrons each atom has does not attempt to explain how many shells are occupied or the degree of shielding in each atom. There’s also no mention of the nuclear charge in each atom in this statement. And therefore, it doesn’t provide an explanation at all. Statement (B) is not the correct answer.

Statement (C) suggests that potassium doesn’t follow the general trend of electronegativity patterns in the periodic table. In fact, the elements in group one do follow the pattern for electronegativity. As we descend group one, electronegativity values steadily decrease. This pattern is not always observed throughout the transition elements, but that’s not what we’re dealing with here. Statement (C) is therefore not a correct answer.

Statement (D) suggests that the bonding electrons are further from the nucleus in an atom of potassium. This is a true statement so far as a potassium atom having more shells occupied is a larger atom. Since the valence electrons are further from the nucleus, there’ll be less attraction for them too. This will indeed reduce the electronegativity of the atom and it does form part of an explanation. Statement (D) could be the correct answer. Statement (E) offers a full statement. Electronegativity does not increase as you descend a group in the periodic table. Statement (E) does not offer an explanation either. So it’s not the correct answer. Statement (D) is the correct answer.

Now let’s review the key points. Electronegativity is the relative attraction an atom has for the shared pair of electrons in a covalent bond. Electronegativity increases across a period as the nuclear charge increases with no extra shielding. Electronegativity decreases down a group as there are more shells and more shielding. Large differences in electronegativity of covalently bonded atoms leads to highly polar bonds.

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