Understanding Electronegativity and Its Impact on Bond Type and Intermolecular Forces | Essays (high school) Chemistry

ELECTRONEGATIVITY

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Whether a bond is nonpolar or polar covalent is determined by a property of the bonding atoms                

called electronegativity. Electronegativity is a measure of the tendency of an atom to attract             

electrons (or electron density) towards itself. It determines how the shared electrons are distributed             

between the two atoms in a bond. The more strongly an atom attracts the electrons in its bonds, the                  

larger its electronegativity



. Electrons in a polar covalent bond are shifted toward the more             

electronegative atom; thus, the more electronegative atom is the one with the partial negative             

charge. The greater the difference in electronegativity, the more polarized the electron distribution            

and the larger the partial charges of the atoms.

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Electronegativity versus Electron Affinity

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We must be careful not to confuse electronegativity and electron affinity. The electron affinity of an              

element is a measurable physical quantity, namely, the energy released or absorbed when an             

isolated gas-phase atom acquires an electron, measured in kJ/mol. Electronegativity

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, on the other            

hand, describes how tightly an atom attracts electrons in a bond. It is a dimensionless quantity that is                 

calculated, not measured. Pauling derived the first electronegativity values by comparing the           

amounts of energy required to break different types of bonds. He chose an arbitrary relative scale               

ranging from 0 to 4.

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Electronegativity and Bond Type

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The absolute value of the difference in electronegativity (ΔEN) of two bonded atoms provides a              

rough measure of the polarity to be expected in the bond and, thus, the bond type. When the                 

difference is very small or zero, the bond is covalent and nonpolar

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. When it is large, the bond is polar                   

covalent or ionic. The absolute values of the electronegativity differences between the atoms in the              

bonds H–H, H–Cl, and Na–Cl are 0 (nonpolar), 0.9 (polar covalent), and 2.1 (ionic), respectively. The               

degree to which electrons are shared between atoms varies from completely equal (pure covalent             

bonding) to not at all (ionic bonding).

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A rough approximation of the electronegativity differences associated with covalent, polar covalent,           

and ionic bonds is shown above . This table is just a general guide, however, with many exceptions.                 

For example, the H and F atoms in HF have an electronegativity difference of 1.9, and the N and H                   

atoms in NH3 a difference of 0.9, yet both of these compounds form bonds that are considered polar                 

covalent. Likewise, the Na and Cl atoms in NaCl have an electronegativity difference of 2.1, and the                

Mn and I atoms in MnI2 have a difference of 1.0, yet both of these substances form ionic                 

compounds.

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Tomboc, MAJS/RN/LPT 2nd Tri/ 2019-2020 [email protected]

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ELECTRONEGATIVITY

Whether a bond is nonpolar or polar covalent is determined by a property of the bonding atoms called electronegativity. Electronegativity is a measure of the tendency of an atom to attract electrons (or electron density) towards itself. It determines how the shared electrons are distributed between the two atoms in a bond. The more strongly an atom attracts the electrons in its bonds, the larger its electronegativity. Electrons in a polar covalent bond are shifted toward the more electronegative atom; thus, the more electronegative atom is the one with the partial negative charge. The greater the difference in electronegativity, the more polarized the electron distribution and the larger the partial charges of the atoms. Electronegativity versus Electron Affinity We must be careful not to confuse electronegativity and electron affinity. The electron affinity of an element is a measurable physical quantity, namely, the energy released or absorbed when an isolated gas-phase atom acquires an electron, measured in kJ/mol. Electronegativity , on the other hand, describes how tightly an atom attracts electrons in a bond. It is a dimensionless quantity that is calculated, not measured. Pauling derived the first electronegativity values by comparing the amounts of energy required to break different types of bonds. He chose an arbitrary relative scale ranging from 0 to 4. Electronegativity and Bond Type The absolute value of the difference in electronegativity (ΔEN) of two bonded atoms provides a rough measure of the polarity to be expected in the bond and, thus, the bond type. When the difference is very small or zero, the bond is covalent and nonpolar. When it is large, the bond is polar covalent or ionic. The absolute values of the electronegativity differences between the atoms in the bonds H–H, H–Cl, and Na–Cl are 0 (nonpolar), 0.9 (polar covalent), and 2.1 (ionic), respectively. The degree to which electrons are shared between atoms varies from completely equal (pure covalent bonding) to not at all (ionic bonding). A rough approximation of the electronegativity differences associated with covalent, polar covalent, and ionic bonds is shown above. This table is just a general guide, however, with many exceptions. For example, the H and F atoms in HF have an electronegativity difference of 1.9, and the N and H atoms in NH3 a difference of 0.9, yet both of these compounds form bonds that are considered polar covalent. Likewise, the Na and Cl atoms in NaCl have an electronegativity difference of 2.1, and the Mn and I atoms in MnI2 have a difference of 1.0, yet both of these substances form ionic compounds. Tomboc, MAJS/RN/LPT 2nd Tri/ 2019-2020 [email protected]

The best guide to the covalent or ionic character of a bond is to consider the types of atoms involved and their relative positions in the periodic table. Bonds between two nonmetals are generally covalent; bonding between a metal and a nonmetal is often ionic. Some compounds contain both covalent and ionic bonds. The atoms in polyatomic ions, such as OH–, NO−3^ , and NH+4^ , are held together by polar covalent bonds. However, these polyatomic ions form ionic compounds by combining with ions of opposite charge. For example, potassium nitrate, KNO3, contains the K+ cation and the polyatomic NO−3 anion. Thus, bonding in potassium nitrate is ionic, resulting from the electrostatic attraction between the ions K+ and NO−3 , as well as covalent between the nitrogen and oxygen atoms in NO−. Intermolecular forces, not to be confused with intramolecular forces, which bind the atoms within a single particle (i.e. covalent bonds, metallic bonds, ionic bonds). Intermolecular forces are the interactions that occur between neighboring particles and have a large effect on a compound’s physical properties such as the melting point, boiling point, viscosity, etc. Ion-Dipole : the interaction between an ion and an oppositely charged dipole. (example: the positive cation of NaCl will be surrounded be water’s oxygens which have a negative dipole). Hydrogen Bonding : requires a hydrogen to be covalently bound to F, O, or N. The large contrast in electronegativities between the hydrogen and these other F, O, N atoms creates large dipoles. Partially-positive hydrogens will interact with partially-negative F,O, or Ns of neighboring molecules. Dipole-Dipole : These can be basically thought of as weaker hydrogen bonds that do not contain F,O, or N as the electronegative atoms. The concept is the same as above except on a smaller scale since the dipoles will not be as large. Just remember, there will be a partially-positive atom that interacts with a partially-negative atom of a neighboring molecule. London Dispersion : All molecules have these. It’s the very brief attraction between neighboring molecules due to the random movement of electrons. At any one snippet of time, the electrons on an atom may be bunched on one side making that side partially-negative while the electron-deficient side is partially-positive. These partially-positive/negative atomic domains interact with the domains of the atoms of neighboring molecules, but on a much smaller scale that the dipole-dipole interactions. However, these forces occur in such large numbers that their summation can’t be ignored. The higher the molecular weight, the stronger the London dispersion forces References: D. Harvey, Modern Analytical Chemistry (McGraw Hill, New York, 2000). D. Harvey, Analytical Chemistry 2.0: An Electronic Textbook for Introductory Courses in Analytical Chemistry (Greencastle, IN, 2009) Tomboc, MAJS/RN/LPT 2nd Tri/ 2019-2020 [email protected]

Understanding Electronegativity and Its Impact on Bond Type and Intermolecular Forces, Essays (high school) of Chemistry

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ELECTRONEGATIVITY