Periodic Trends: Electronegativity Learning Objectives • Define electronegativity. • Describe how the electronegativity changes within a period. • Describe how the electronegativity changes within a group. • Analyze the importance of electronegativity in determining bond polarity Periodic trends are specific patterns that are present in the periodic table that illustrate different aspects of a certain element, including its size and its electronic properties. Major periodic trends include: electronegativity, ionization energy, electron affinity, atomic radius, melting point, and metallic character. Periodic trends, arising from the arrangement of the periodic table, provide chemists with an invaluable tool to quickly predict an element's properties. These trends exist because of the similar atomic structure of the elements within their respective group families or periods, and because of the periodic nature of the elements. We have already studied metallic character, atomic radius, and ionization energy. We are now going to have a closer look at electronegativity. Electronegativity Trends Electronegativity can be understood as a chemical property describing an atom's ability to attract and bind with electrons. Because electronegativity is a qualitative property, there is no standardized method for calculating electronegativity. However, the most common scale for quantifying electronegativity is the Pauling scale (Table A2), named after the chemist Linus Pauling. The numbers assigned by the Pauling scale are dimensionless due to the qualitative nature of electronegativity. Electronegativity values for each element can be found on certain periodic tables. An example is provided below. Electronegativity measures an atom's tendency to attract and form bonds with electrons. This property exists due to the electronic configuration of atoms. Most atoms follow the octet rule (having the valence, or outer, shell comprise of 8 electrons). Because elements on the left side of the periodic table have less than a half-full valence shell, the energy required to gain electrons is significantly higher compared with the energy required to lose electrons. As a result, the elements on the left side of the periodic table generally lose electrons when forming bonds. Conversely, elements on the right side of the periodic table are more energy-efficient in gaining electrons to create a complete valence shell of 8 electrons. The nature of electronegativity is effectively described thus: the more inclined an atom is to gain electrons, the more likely that atom will pull electrons toward itself. • From left to right across a period of elements, electronegativity increases. If the valence shell of an atom is less than half full, it requires less energy to lose an electron than to gain one. Conversely, if the valence shell is more than half full, it is easier to pull an electron into the valence shell than to donate one. • From top to bottom down a group, electronegativity decreases. This is because atomic number increases down a group, and thus there is an increased distance between the valence electrons and nucleus, or a greater atomic radius. • Important exceptions of the above rules include the noble gases, lanthanides, and actinides. The noble gases possess a complete valence shell and do not usually attract electrons. The lanthanides and actinides possess more complicated chemistry that does not generally follow any trends. Therefore, noble gases, lanthanides, and actinides do not have electronegativity values. • As for the transition metals, although they have electronegativity values, there is little variance among them across the period and up and down a group. This is because their metallic properties affect their ability to attract electrons as easily as the other elements. Patterns of electronegativity in the Periodic Table The distance of the electrons from the nucleus remains relatively constant in a periodic table row, but not in a periodic table column. The force between two charges is given by Coulomb’s law. F=kQ1Q2 r2 In this expression, Q represents a charge, k represents a constant and r is the distance between the charges. When r = 2, then r2= 4. When r = 3, then r2 = 9. When r = 4, then r2 = 16. It is readily seen from these numbers that, as the distance between the charges increases, the force decreases very rapidly. This is called a quadratic change. The result of this change is that electronegativity increases from bottom to top in a column in the periodic table even though there are more protons in the elements at the bottom of the column. Elements at the top of a column have greater electronegativities than elements at the bottom of a given column. The overall trend for electronegativity in the periodic table is diagonal from the lower left corner to the upper right corner. Since the electronegativity of some of the important elements cannot be determined by these trends (they lie in the wrong diagonal), we have to memorize the following order of electronegativity for some of these common elements. F > O > Cl > N > Br > I > S > C > H > metals The most electronegative element is fluorine. If you remember that fact, everything becomes easy, because electronegativity must always increase towards fluorine in the Periodic Table. According to these two general trends, the most electronegative element is fluorine, with 3.98 Pauling units. The least is cesium, at 0.79. Francium is rated as lower, but as it is a radioactive element, its reactivity is not typically a consideration. SUMMARY Trends • Electronegativity refers to the ability of a nucleus to attract electrons or to retain electrons during chemical bonding. • The electronegativity of the elements within a period generally increases from left to right. This is because the nuclear charge is increasing faster than the electron shielding, so the attraction that the atoms have for the valence electrons increases. • The electronegativity of the elements within a group generally decreases from top to bottom. This is because as you go from top to bottom down a group, the atoms of each element have an increasing number of energy levels. The electrons in a bond are thus farther away from the nucleus and are held less tightly. • Atoms with low ionization energies have low electronegativities because their nuclei do not have a strong attraction for electrons. Atoms with high ionization energies have high electronegativities because the nucleus has a strong attraction for electrons. • Although the noble gases possess very high ionization energies, He, Ne and Ar do not have listed electronegativity values, as they do not bond with other elements. Kr and Xe do form QUESTIONS: 1) What is electronegativity? 2) Considering also the periodic trends in atomic radius and ionization energy, explain why fluorine has the highest electronegativity. 3) Why are there no values of EN for He, Ne and Ar? 4) a. What is the trend in EN across a period (row) from left to right? b. What causes this trend? 5) a. What is the trend in EN down a group (column) from top to bottom? b. What causes this trend? 6) In each pair, select the element which has the higher electronegativity: a. N and As b. Mg and Sr c. Na and S d. K and Br 7) Which group would generally have the lowest electronegativity? a. Transition Metals (Groups 3-12) b. Alkali Metals (Group 1) c. Noble Gases (Group 18) d. Alkaline Earth Metals (Group 2) e. Halogens (Group 17) Justify your response. 9) Low electronegativty is considered a property of a. Metals b. Nonmetals Justify your response .