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If you have ranked bonds by polarity in Class 11 chemistry, you have used Linus Pauling's work. The electronegativity values in your NCERT table, with fluorine at 4.0, are on the scale that bears his name. He also has an unusual record: the Nobel Prize site notes he is the only person to have been awarded two unshared Nobel Prizes (Nobel Prize facts).
What were his two Nobel Prizes for?
Chemistry, 1954: "for his research into the nature of the chemical bond and its application to the elucidation of the structure of complex substances" (Nobel Prize in Chemistry 1954).
Peace, 1962: "for his fight against the nuclear arms race between East and West" (Nobel Peace Prize 1962).
Oregon State University's documentary history of his chemical-bond work says he helped pioneer the field as a postgraduate fellow in 1926-27 and carried it through to his book *The Nature of the Chemical Bond* in 1939 (OSU Special Collections, Pauling and the chemical bond).
How did Pauling put numbers on electronegativity?
The idea that some atoms pull electrons harder than others was old. What Pauling added was a way to put it on a numerical basis. In the spring of 1932, while a visiting professor at MIT, he compared the real strengths of bonds between pairs of elements with what he predicted for purely covalent bonds. The real bonds were always stronger. He took the extra strength as coming from ionic character: the bigger the gap, the more the two atoms differed in their ability to attract electrons (OSU archive, "The Fourth Paper").
By comparing many pairs he built a relative scale. Fluorine sat at the far end and lithium towards the other, so lithium fluoride came out almost 100 percent ionic, while lithium iodide had more covalent character. The archive says he found hydrogen chloride to be both ionic and covalent, in the ratio 20:80. The work appeared as "The Nature of the Chemical Bond. IV. The Energy of Single Bonds and the Relative Electronegativity of Atoms" in the *Journal of the American Chemical Society* in 1932 (Pauling, 1932).
The archive is frank about the scale's weakness: it was "one of his least theoretically well-founded ideas — and became one of his most influential", because chemists could use it to roughly predict bonds without knowing anything about the wave equation (OSU archive, "Electronegativity").
What NCERT says about the scale
The NCERT Class 11 chapter on periodicity calls electronegativity "a qualitative measure of the ability of an atom in a chemical compound to attract shared electrons to itself". It notes that it is "not a measureable quantity", that it varies with the element an atom is bound to, and that of several scales the Pauling scale is the most widely used (NCERT Class 11 Chemistry, Unit 3). Its Table 3.8 gives Pauling values including Li 1.0, C 2.5, N 3.0, O 3.5, F 4.0, Na 0.9, Cl 3.0, Br 2.8 and I 2.5. Values generally rise across a period and fall down a group.
A worked routine: ranking bond polarity
Use the NCERT values to rank these bonds from least to most polar: C–Cl, C–O, Na–Cl, Li–F, Li–I.
Find the difference for each pair. C–Cl: 3.0 − 2.5 = 0.5. C–O: 3.5 − 2.5 = 1.0. Li–I: 2.5 − 1.0 = 1.5. Na–Cl: 3.0 − 0.9 = 2.1. Li–F: 4.0 − 1.0 = 3.0.
Rank by the difference. C–Cl < C–O < Li–I < Na–Cl < Li–F. A bigger gap means more ionic character, as in Pauling's own reasoning.
Mark the dipole. The partial negative charge sits on the more electronegative atom: on Cl in C–Cl, on O in C–O.
Check against a known case. Pauling found LiF almost entirely ionic and LiI more covalent; your ranking should agree.
Remember the limit. These are bond polarities. Whether a whole molecule has a dipole moment also depends on its shape, which you meet in the chemical bonding chapter.
Pauling's scale also led to a bold prediction. The archive records that he expected fluorine to form compounds even with xenon, then thought to be inert; in the summer of 1933 his colleague Don Yost searched for them and found none, but thirty years later another team made xenon compounds (OSU archive, "The Xenon Conundrum").
For more on making periodic trends stick, see learning inorganic chemistry with retrieval, and for seeing bond shapes in 3D, molecule visualisation tools.
If electronegativity, dipoles and bond types feel like separate facts, the Fermi Concept Library gives every concept in chemistry its own lesson and its place in the map, so you can see how they connect.
Try the five-step routine on the next bonding question you meet.
