Electronegativity and bond polarity are foundational concepts in understanding the behaviour of molecules, allowing us to predict how atoms will interact and share electrons. Dive into the concept of electronegativity, the measurement behind it, and the art of deciphering bond polarity.
Electronegativity: A Definition
Electronegativity refers to the ability of an atom in a molecule to attract shared electrons towards itself. This property gives us insight into how strongly an atom desires electrons, which is pivotal in determining how it will interact with other atoms.
- Atom's Electron Desire: Atoms with high electronegativity values are like electron magnets; they have a strong tendency to pull shared electrons close.
- Comparison Tool: By comparing electronegativity values of different atoms, we can predict the nature of their bond: whether it's purely covalent, purely ionic, or somewhere in between.
Image courtesy of Western Oregon University
Measurement of Electronegativity
The electronegativity of an atom is not something we can measure directly. Instead, it's a relative value derived from several observable properties of elements.
- Pauling Scale: Developed by Linus Pauling, this is the most widely used scale for electronegativity. Values on this scale range from around 0.7 (least electronegative, like francium) to 3.98 (most electronegative, like fluorine).
- Other Scales: There are other scales, such as the Mulliken, Allred-Rochow, and Sanderson scales, but the Pauling scale remains the most popular.
- Trends in the Periodic Table: Generally, electronegativity increases as you move from left to right across a period and decreases as you move down a group. Thus, elements like fluorine and oxygen (top right of the table, excluding noble gases) are highly electronegative.
Periodic table of electronegativity using the Pauling scale. It increases as you move from left to right across a period.
Image courtesy of natros
Deduction of Bond Polarity Using Electronegativity
Bond polarity arises when there is an electronegativity difference between the bonded atoms. The greater this difference, the more polar the bond.
- Pure Covalent Bond: When two atoms with the same electronegativity bond, they share electrons equally, resulting in a non-polar bond.
- Polar Covalent Bond: If there's a moderate difference in electronegativity (typically a difference of 0.5 to 1.7 on the Pauling scale), the bond becomes polar covalent. Here, the electrons are not shared equally.
- Ionic Bond: A large electronegativity difference (usually greater than 1.7) leads to an ionic bond. One atom takes electrons from the other.
- A Rule of Thumb: While these values are a good guide, remember that the nature of a bond is a continuum from pure covalent to ionic, and there's often some debate over the exact boundary values.
Representation of Bond Dipoles
A bond dipole moment arises in a polar covalent bond due to the uneven distribution of electrons. This dipole moment can be represented in several ways.
Using Partial Charges
- δ+ and δ-: These symbols represent partial positive and partial negative charges. If an atom in a bond is more electronegative, it will have a partial negative charge (δ-) because it pulls the shared electrons closer. The other atom will have a partial positive charge (δ+).
- Example: In a water molecule, oxygen is more electronegative than hydrogen. Thus, the bond between oxygen and hydrogen will be represented as O-δ H+δ.
Image courtesy of Pradana Aumars
Using Vectors
- Direction of the Vector: A vector arrow can be drawn from the less electronegative atom to the more electronegative atom, indicating the direction of the electron pull. The arrow's head points towards the more electronegative atom.
- Magnitude of the Vector: The length or size of the vector can give an idea about the strength of the bond dipole.
Water molecule showing a vector arrow drawn from the less electronegative atom (H) to the more electronegative atom (O).
Image courtesy of Riccardo Rovinetti
Compound Dipoles
In molecules with multiple bonds, individual bond dipoles can either reinforce each other or cancel out, depending on the molecular geometry. Thus, a molecule can have polar bonds and still be non-polar if the bond dipoles cancel out. This interplay between bond and molecular polarity will be explored in more depth in subsubtopic 2.2.6.
FAQ
Electronegativity is fundamental in predicting how atoms will interact and bond with each other. Atoms with high electronegativities tend to attract electrons, making them likely to form negative ions or participate in polar covalent bonds. Conversely, atoms with low electronegativities are more likely to lose electrons, forming positive ions or sharing electrons almost equally in nonpolar covalent bonds. Understanding electronegativity helps in predicting bond types, bond strengths, and even the reactivity of substances. For instance, elements with greatly differing electronegativities often form ionic compounds, while those with moderate differences usually form polar covalent bonds.
No, two atoms with identical electronegativities will not form a polar bond. When two atoms with the same electronegativity value bond together, they share the bonding electrons equally. Since there's no difference in their electronegativities, there's no region of partial positive or partial negative charge. The bond formed is a pure covalent or nonpolar covalent bond. A classic example of this is the bond between two hydrogen atoms in an H2 molecule.
Generally, electronegativity increases across a period from left to right and decreases down a group. However, there are exceptions. For example, the noble gases, which occupy the far right of the periodic table, have relatively low electronegativities because they already possess a full valence shell and usually don't form compounds. Also, electronegativity doesn't always increase smoothly from one element to the next in a period. Transition metals, for instance, show little variation in electronegativity across each row. It's crucial to use an electronegativity chart or table when predicting properties rather than relying solely on periodic trends.
The concept of electronegativity was introduced by the American chemist Linus Pauling in the 1930s. He developed the Pauling Scale, which is the most widely used scale for measuring electronegativity values. On the Pauling Scale, electronegativity values range from around 0.7 to 3.98, with francium having the lowest and fluorine having the highest electronegativity. The values represent an atom's relative ability to attract and bind with electrons. It's worth noting that there are other scales, such as the Mulliken and Allred-Rochow scales, but the Pauling Scale remains the most popular.
While bond polarity refers to the polarity within a particular bond due to differences in electronegativity, the overall polarity of a molecule depends on both bond polarity and molecular geometry. A molecule can have polar bonds but still be nonpolar if the bond dipoles cancel each other out. For instance, carbon tetrachloride (CCl4) has polar C-Cl bonds, but its tetrahedral symmetry means that the bond dipoles cancel out, making the whole molecule nonpolar. So, both bond dipoles and molecular shape must be considered when determining the overall polarity of a molecule.
Practice Questions
Electronegativity is the ability of an atom in a molecule to attract shared electrons towards itself. It is a measure of how strongly an atom desires electrons. To predict the polarity of a chemical bond, one can compare the electronegativity values of the two bonded atoms. If there is a significant difference in their electronegativity values, the bond is polar. For example, hydrogen has an electronegativity of 2.20 on the Pauling scale, while chlorine has an electronegativity of 3.16. Due to this difference, the bond between hydrogen and chlorine in HCl is polar, with chlorine being slightly negative (δ-) and hydrogen being slightly positive (δ+).
Bond dipoles arise from polar covalent bonds where there's an uneven distribution of electrons due to differences in electronegativity. They can be represented using partial charges, denoted by δ+ for partial positive charge and δ- for partial negative charge. In a water molecule, oxygen is more electronegative than hydrogen. As a result, the shared electrons between oxygen and hydrogen are drawn closer to the oxygen atom, making it slightly negative. On the other hand, the hydrogens become slightly positive. This can be represented as O-δ H+δ for each hydrogen-oxygen bond in the water molecule, indicating the bond's polar nature with uneven electron distribution.
