Electronegativity
Electronegativity is a measure of an atom's ability to attract and hold onto electrons in a chemical bond. It is a fundamental concept in understanding the nature of chemical bonds and the reactivity of elements. Electronegativity values range from 0.7 for cesium to 4.0 for fluorine on the Pauling scale.
7 Key excerpts on "Electronegativity"
eBook - ePub- Tony Cox(Author)
- 2004(Publication Date)
- Taylor & Francis(Publisher)
...B1 E LECTRONEGATIVITY AND BOND TYPE Key Notes Definitions Electronegativity is the power of an atom to attract electrons to itself in a chemical bond. Different numerical estimates agree on qualitative trends: electro-negativity increases from left to right along a period, and generally decreases down groups in the periodic table. Elements of low Electronegativity are called electropositive. The bonding triangle Electropositive elements form metallic solids at normal temperatures. Electro-negative elements form molecules or polymeric solids with covalent bonds. Elements of very different Electronegativity combine to form solids that can be described by the ionic model. Bond polarity The polarity of a bond arises from the unequal sharing of electrons between atoms with different electronegativities. There is no sharp dividing line between polar covalent and ionic substances. Related topics Trends in atomic properties (A5) Introduction to solids (D1) Electron pair bonds (C1) Definitions Electronegativity may be defined as the power of an atom to attract electrons to itself in a chemical bond. It is the most important chemical parameter in determining the type of chemical bonds formed between atoms. It is hard to quantify in a satisfactory way, especially as Electronegativity is not strictly a property of atoms on their own, but depends to some extent on their state of chemical combination. Nevertheless several scales have been devised. • Pauling Electronegativity is based on bond energies (see Topic C8), using the empirical observation that bonds between atoms with a large Electronegativity difference tend to be stronger than those where the difference is small...
eBook - ePub- Atef Korchef(Author)
- 2022(Publication Date)
- CRC Press(Publisher)
...Indeed, when an atom has a high ionization energy and a high electron affinity, it strongly retains its electrons, and it can readily attract electrons from another atom. As a result, its Electronegativity is high. For example, metals lose their electrons easily, so they have low electronegativities. However, non-metals tend to attract electrons, so they have high electronegativities. Milliken proposed the following equation to determine the Electronegativity (E N) of an atom: E N = k I + A where k = 0.18 if the ionization energy (I) and the electron affinity (A) of the atom are expressed in eV atom −1. Note that Pauling proposed another equation, based on the bond energy in the diatomic molecule AB, for determining the difference of Electronegativity (ΔE N = E N (A)− E N (B)) between two atoms A and B: Δ E N = k E A − B − (E A − A × E B − B) where E A-B is the bond energy between the atoms A and B, E A-A and E B-B are the bond energies A-A and B-B, respectively and k = 0.0208 if the bond energies are expressed in kcal mol −1. The bond energy is defined as the amount of energy required to break apart a mole of molecules into its component atoms. Note that when the atoms A and B have the same Electronegativity: Δ E N = 0 a n d E A − B = (E A − A × E B − B) This means that, when the atoms A and B have the same Electronegativity, the bond energy E A-B equals the average of the bond energies A-A and B-B. When A and B have different electronegativities, the bond energy E A − B > (E A − A × E B − B). When we move from the left to the right along a period in the periodic table of elements, the Electronegativity increases. Indeed, since there is an increase in the effective nuclear charge, there is a greater attraction of the outer shell electrons to the nucleus. When we move from the top to the bottom along a group in the periodic table of elements, the electrons are further from the nucleus and there is a weaker attraction...
eBook - ePub- Gavin Whittaker, Andy Mount, Matthew Heal(Authors)
- 2000(Publication Date)
- Taylor & Francis(Publisher)
...As more electronic charge accumulates on one atom over the other, this leads to a polar covalent bond. An extreme example of this may be seen as the basis of the ionic bonding in, for example, KCl. Fig. 6 Molecular orbital diagram for a heteronuclear diatomic molecule (CO). The element which most strongly attracts the electrons is referred to as the more electronegative element. Conversely, the element which holds them less strongly is referred to as being more electropositive. Among the main group elements, the more electronegative elements tend to be in the later groups and earlier periods of the periodic table (top right), whilst the electropositive elements tend to be located in the early groups and later periods (bottom left). The strength of the attraction is most commonly measured using the Pauling Electronegativity scale, with values, denoted χ, ranging between 4 for fluorine, the most electronegative element, down to ca. 0.6 for francium, the most electropositive. Dipole moments A polar covalent bond implies that one atom in a molecule will be more positively charged than the other. The positive charge, + q and the negative charge, − q, separated by a distance R, give rise to an electric dipole moment, µ. This is a vector directed from the positive to the negative charge across the molecule, with magnitude qR. This vector is usually represented by an arrow directed from the positive to the negative charge, with a positive sign included to indicate the positive end, thus: In order to generate convenient values, the dipole moment is generally reported in debye, D, where 1 debye is equal to 3.336×10 –30 C m. Water, for example, has a dipole moment of 1.85 D. The size of the dipole moment in debyes between two atoms, A and B may often be estimated from their respective Pauling electronegativities, χ (A) and χ (B): The dipole moments in polyatomic molecules may be calculated by vector addition of the dipole moments for each band (Fig. 7). Fig...
- Kevin Reel, Derrick C. Wood, Scott A. Best(Authors)
- 2014(Publication Date)
- Research & Education Association(Publisher)
...The Electronegativity of an atom exhibits the same trend as electron affinity in the periodic table. You should know the values from Pauling’s scale of Electronegativity for all nonmetallic atoms. This will enable you to determine if the covalent bond is polar or if a molecule is polar. The chart below shows the Electronegativity values for the most common nonmetals. Atom Electronegativity Fluorine 4.0 Oxygen 3.5 Nitrogen 3.0 Chlorine 3.0 Bromine 2.8 Carbon 2.5 Sulfur 2.5 Iodine 2.5 Phosphorus 2.1 Hydrogen 2.1...
eBook - ePub- Subrata Pal(Author)
- 2019(Publication Date)
- Academic Press(Publisher)
...The property was quantified by Linus Pauling based on comparative bond energies of homo- and heteronuclear diatomic molecules. Fluorine (F), the most electronegative element in the periodic table, was assigned an Electronegativity of 4.0. The force of attraction towards one of the two nuclei as experienced by the bonding pair of electrons depends on (a) the number of protons in the nucleus, (b) the distance from the nucleus, and (c) the extent of screening by the inner/core electrons. For example, carbon (electron configuration: 1s 2 2s 2 2p 2) and fluorine (electron configuration: 1s 2 2s 2 2p 5) both have their valence electron in the second shell (n = 2) screened from their respective nuclei by 1s 2 electrons. However, the fluorine nucleus has nine protons compared with six of carbon. Evidently, a bonding pair of electrons will experience a greater force of attraction from the nucleus of fluorine which, therefore, is more electronegative than carbon. Again, considering hydrogen fluoride and hydrogen chloride, one can find that in each case the effective charge pulling the bonding pair of electrons towards the center of fluorine or chorine is + 7. Nevertheless, fluorine has the bonding pair in the second shell (n = 2), while chlorine has it in the third (n = 3). So, fluorine, whose bonding pair is closer to the nucleus experiencing a greater force of attraction, is more electronegative than chlorine. 4.1.4 Polarity of bond—Molecular dipole Let us consider a bond between two atoms, A and B, of equal Electronegativity (e.g., two chlorine atoms) formed by sharing a pair of electrons. The bonding pair of electrons is equally attracted by the two chlorine atoms and, hence, placed exactly halfway between the two atoms. The bond is purely covalent or nonpolar (Fig. 4.1). Fig. 4.1 Pure covalent (nonpolar) bond between two chlorine atoms. In accordance with the octet rule, two atoms can share even more than a single pair of electrons...
eBook - ePubChemistry
Concepts and Problems, A Self-Teaching Guide
- Richard Post, Chad Snyder, Clifford C. Houk(Authors)
- 2020(Publication Date)
- Jossey-Bass(Publisher)
...To be sure, any difference in electronegativities will result in a polar bond. However, it is the degree or extent of the difference that will permit you to tell whether one compound is more or less polar than other compounds, or whether a particular bond is more or less polar than another or more ionic than covalent. The only time we have a 100% covalent bond is when Δen = 0, which usually happens only when identical atoms combine. Using the table shown below and the scheme above, CO would be classified as a polar covalent compound. The Electronegativity value is the number below the symbol in the table, C = 2.5 and O = 3.5 units. The difference (Δen) is 3.5–2.5 = 1.0 unit, therefore CO is polar covalent. Electronegativities of the First 20 Elements How should KF be classified: ionic, polar covalent, or covalent? ____________________________ Answer: Δen = 4.1 (fluorine) − 0.9 (potassium) = 3.2, thus it is an ionic bond Using the table again, how should each of the following be classified? HCl __________ LiCl __________ F 2 ____________ Answer: (a) polar covalent (Δen = 0.8 unit); (b) ionic (Δen = 1.9 units); (c) covalent (Δen = 0.0; they are identical atoms) SHAPES OF MOLECULES The classification scheme discussed in frames 61 and 62 works only with compounds composed of two atoms. If we have compounds of three or more atoms, the shape of the molecule becomes important in determining its polarity. A discussion of the shapes of molecules requires an exhaustive look at atomic and molecular orbitals and their shapes and spatial orientation, along with a discussion of the valence bond electron repulsion theory. Such a presentation is beyond the scope of this book. We will limit our discussion to a few simple molecules that are known to be linear (all the atoms in the compound can be joined by a single straight line through their nuclei) or bent (nuclei cannot be joined by a straight line)...
eBook - ePubFoundations for Teaching Chemistry
Chemical Knowledge for Teaching
- Keith S. Taber(Author)
- 2019(Publication Date)
- Routledge(Publisher)
...A teacher cannot be surprised that students may later talk of atoms ‘stealing’ electrons and the like if a social metaphor is presented as if a satisfactory description. As suggested earlier, it would be better to focus on how the negative electrons can hold the positively charged atomic cores in the structure when there is a balance of attractive and repulsive forces. A student might reasonably ask how the two negative electrons can be considered to behave as a pair if they repel each other – a question that cannot be addressed in any detail without invoking that elephant we may prefer to pretend is not in the room. Covalent bonding is often illustrated with misleading hypothetical schemes for how the molecules came about: thus, methane is often shown as being formed from an isolated carbon atom and four isolated hydrogen atoms – allowing students to acquire the mistaken notion that bonds form in order to allow atoms to follow the octet rule (chemical reaction mechanisms are discussed in Chapter 11). The students may be new enough to the subject not to spot this trick and so to ask where these radical atoms originated: a teacher should know better than to use such deceptive and unscientific devices in explanations. A teacher should only explain a reaction as starting with non-bonded atoms if they can show the students that reaction using reactants in that form – but I doubt any school chemical stores can supply such reagents! The ionic bond is based on the attractions between charges The traditional teaching scheme will often then move on to ionic bonding. This is a more complex situation because it does not involve discrete links between adjacent atomic cores but rather relates to the overall regular structures of myriad cations and anions that allow forces to be balanced by placing positive ions closer to negative ions than other positive ions and vice versa...
