## Electronegativity and Oxidation Number

Electronegativity is the tendency of an atom/molecule to attract electrons; oxidation number is an indicator of its bonding environment.

### Learning Objectives

Apply the rules for assigning oxidation numbers to atoms in compounds

### Key Takeaways

#### Key Points

• An atom ‘s electronegativity is affected by both the element ‘s atomic number and its size.
• The higher its electronegativity, the more an element attracts electrons.
• The atom with higher electronegativity, typically a nonmetallic element, is assigned a negative oxidation number, while metallic elements are typically assigned positive oxidation numbers.

#### Key Terms

• electronegativity: A chemical property that describes the tendency of an atom to attract electrons (or electron density) toward itself.
• oxidation number: The hypothetical charge that an atom in a molecule/compound would have if all bonds were purely ionic. It indicates of the degree of oxidation of an atom in a chemical compound.

### Electronegativity

Electronegativity is a property that describes the tendency of an atom to attract electrons (or electron density) toward itself. An atom’s electronegativity is affected by both its atomic number and the size of the atom. The higher its electronegativity, the more an element attracts electrons. The opposite of electronegativity is electropositivity, which is a measure of an element’s ability to donate electrons.

Electronegativity is not directly measured, but is instead calculated based on experimental measurements of other atomic or molecular properties. Several methods of calculation have been proposed, and although there may be small differences in the numerical values of the calculated electronegativity values, all methods show the same periodic trend among the elements.

Electronegativity, as it is usually calculated, is not strictly a property of an atom, but rather a property of an atom in a molecule. Properties of a free atom include ionization energy and electron affinity. It is expected that the electronegativity of an element will vary with its chemical environment, but it is usually considered to be a transferable property; that is to say, similar values will be valid in a variety of situations.

On the most basic level, electronegativity is determined by factors such as the nuclear charge and the number/location of other electrons present in the atomic shells. The nuclear charge is important because the more protons an atom has, the more “pull” it will have on negative electrons. Where electrons are in space is a contributing factor because the more electrons an atom has, the farther from the nucleus the valence electrons will be, and as a result they will experience less positive charge; this is due to their increased distance from the nucleus, and because the other electrons in the lower-energy core orbitals will act to shield the valence electrons from the positively charged nucleus.

The most commonly used method of calculation for electronegativity was proposed by Linus Pauling. This method yields a dimensionless quantity, commonly referred to as the Pauling scale, with a range from 0.7 to 4. If we look at the periodic table without the inert gases, electronegativity is greatest in the upper right and lowest at the bottom left.

Electronegativity of the elements: Electronegativity is highest at the top right of the table and lowest at the bottom left.

Hence, fluorine (F) is the most electronegative of the elements, while francium (Fr) is the least electronegative.

### Oxidation Numbers

It is common to consider a single value of electronegativity to be valid for most bonding situations a given atom can be in. While this approach has the advantage of simplicity, it is clear that the electronegativity of an element is not an invariable atomic property; rather, it can be thought of as depending on a quantity called ‘the oxidation number’ of the element.

One way to characterize atoms in a molecule and keep track of electrons is by assigning oxidation numbers. The oxidation number is the electric charge an atom would have if the bonding electrons were assigned exclusively to the more electronegative atom, and it can identify which atom is oxidized and which is reduced in a chemical process. Six rules can be used when assigning oxidation numbers:

1. The oxidation number of an element in its natural state (i.e., how it is found in nature) is zero. For example, hydrogen in H2, oxygen in O2, nitrogen in N2, carbon in diamond, etc., have oxidation numbers of zero.
2. In ionic compounds, the ionic charge of an atom is its oxidation number.
3. The sum of the oxidation numbers of all the atoms in an ion or molecule is equal to its net charge.
4. In compounds with nonmetals, the oxidation number of hydrogen is +1. However, when hydrogen is bonded with a metal, its oxidation number reduces to -1 because the metal is a more electropositive, or less electronegative, element.
5. Oxygen is assigned an oxidation number of -2 in most compounds. However, there are certain exceptions. In peroxides (O22-), such as hydrogen peroxide (H2O2), the oxidation number of oxygen is -1. In oxygen difluoride (OF2), the oxidation number of oxygen is +2, while in dioxygen difluoride (O2F2), oxygen is assigned an oxidation number of +1 because fluorine is the more electronegative element in these compounds, so it is assigned an oxidation number of -1.
6. The atom with higher electronegativity, typically a nonmetallic element, is assigned a negative oxidation number, while the other atom, which is often but not necessarily a metallic element, is given a positive oxidation number.

## Bond Polarity

Molecular polarity is dependent on the presence of polar covalent bonds and the molecule’s three-dimensional structure.

### Learning Objectives

Apply knowledge of bond polarity and molecular geometry to identify the dipole moment of molecules

### Key Takeaways

#### Key Points

• When non-identical atoms are covalently bonded, the electron pair will be attracted more strongly to the atom that has the higher electronegativity. This results in a polar covalent bond.
• Polarity refers to a separation of electric charge leading to a molecule or its chemical groups having an electric dipole moment.
• A polar molecule acts as an electric dipole that can interact with electric fields that are created artificially, or that arise from nearby ions or polar molecules.
• The dipole moment $\mu$ that corresponds to an individual bond is given by the product of the quantity of charge, q, and the bond length r: $\mu = qr$.

#### Key Terms

• Bond polarity: A covalent bond is polar if one atom is more electronegative than its bonding partner, resulting in a net dipole moment between the two atoms.
• dipole moment: A measure of the polarity of a covalent bond or of an entire molecule. It is the product of the charge on either pole of the dipole and the distance separating them.
• Molecular polarity: A molecule is polar if it has a net dipole moment, which depends on the existence of polar covalent bonds and the molecule’s three-dimensional structure or geometry.

### Bond vs Molecular Polarity

Polarity refers to the separation of charge that creates permanent positive and negative ‘electric poles.’ This concept can be applied in two contexts:

1. Bond polarity: when atoms from different elements are covalently bonded, the shared pair of electrons will be attracted more strongly to the atom with the higher electronegativity. As a result, the electrons will not be shared equally. Such bonds are said to be ‘polar’ and possess partial ionic character.
2. Molecular polarity: when an entire molecule, which can be made out of several covalent bonds, has a net polarity, with one end having a higher concentration of negative charge and another end having a surplus of positive charge. A polar molecule acts as an electric dipole which can interact with electric fields that are created artificially, or that arise from interactions with nearby ions or other polar molecules.

### Dipole Moment

Dipoles are conventionally represented as arrows pointing in the direction of the negative end. The strength of a dipole’s interaction with an electric field is given by the electric dipole moment of the bond or molecule. The dipole moment is calculated by evaluating the product of the magnitude of separated charge, q, and the bond length, r:

$\mu = q r$

In SI units, q is expressed in coulombs and r in meters, so μ has the dimensions of $C \cdot m$. If two charges of magnitude +1 and -1 are separated by a typical bond length of 100 pm, then:

$\mu = (1.6022 \times 10^{-19} C) \times (10^{-10} m) = 1.6 \times 10^{-29} C \cdot m = 4.8 D$

The Debye unit, D, is commonly used to express dipole moments.

### Determining a Molecule’s Dipole Moment

In molecules containing more than one polar bond, the molecular dipole moment is just the vector addition of the individual bond dipole moments. Being vectors, these can reinforce or cancel each other depending on the geometry of the molecule. Therefore, it is possible for molecules containing polar bonds to be nonpolar overall, as in the example of carbon dioxide.

Molecular dipole moment of carbon dioxide: The linear shape of the CO2 molecule results in the canceling of the dipole moments of the two polar C=O bonds. The net, molecular dipole moment of CO2 is therefore zero, and the molecule is nonpolar.

H2O, by contrast, has a very large molecular dipole moment which results from the two polar H–O bonds forming an angle of 104.5° between them. The water molecule, therefore, is polar.

Dipole moment of a water molecule: Water has a very large dipole moment which results from the two polar H–O bonds oriented at an angle of 104.5° with respect to each other. The bond dipoles add up to create a molecular dipole (indicated by the green arrow).