Chlorine (Cl) is more electronegative than bromine (Br). This difference in chemical behavior is rooted in a fundamental property called electronegativity, which is an atom’s ability to attract electrons when forming a chemical bond. Chlorine exhibits a stronger pull on shared electrons compared to bromine, a distinction that is a direct consequence of their respective positions on the periodic table.
Understanding Electronegativity
Electronegativity is a measure of an atom’s power to attract a shared pair of electrons toward itself within a chemical bond. It is most commonly quantified using the Pauling scale, which assigns a numerical value to each element. A higher number indicates a greater ability to attract electrons. This measurement predicts how electrons will be distributed in a molecule, determining the nature of the chemical bond.
The concept can be visualized as a tug-of-war for electrons between two bonded atoms. If one atom has a significantly higher electronegativity value, it pulls the shared electrons closer to its nucleus. This unequal sharing creates a polar bond, with one end of the molecule becoming slightly negative and the other slightly positive.
Periodic Table Trends and Electron Pull
The arrangement of the periodic table is a powerful predictor of an element’s electronegativity. This property follows two primary trends. First, as you move from left to right across any period, electronegativity increases. This occurs because the increasing number of protons in the nucleus creates a stronger positive charge that pulls electrons inward.
Second, when moving down a group, or column, electronegativity tends to decrease. As you move down, atoms gain additional electron shells, making them progressively larger. The increasing distance between the nucleus and the outermost valence electrons weakens the attractive force on shared electrons.
Why Chlorine Outranks Bromine
Both chlorine and bromine belong to Group 17, the halogens, and both have a strong tendency to gain one electron to achieve a stable configuration. The difference in their electron-attracting power stems from their position: Chlorine is in the third period, while bromine sits directly beneath it in the fourth period.
This vertical separation means chlorine has three occupied electron shells, while bromine has four. The extra electron shell makes the bromine atom physically larger than the chlorine atom. In chlorine, the valence electrons are closer to the positively charged nucleus, allowing them to experience a much stronger effective nuclear charge. This closer proximity results in a more forceful attraction for shared electrons in a bond.
The Pauling values quantify this difference, showing chlorine’s electronegativity at approximately 3.16, while bromine’s is lower at about 2.96. The larger size of the bromine atom means its outermost electrons are shielded by more inner-shell electrons. This shielding effect and greater distance from the nucleus significantly diminish the nucleus’s ability to pull shared electrons toward itself.
Practical Consequences of the Difference
The difference in electronegativity between chlorine and bromine has tangible effects on the chemical compounds they form. This difference directly impacts the polarity of the bonds they create with other elements, such as hydrogen. For instance, in hydrogen chloride (HCl) versus hydrogen bromide (HBr), the chlorine atom pulls the shared electrons more strongly toward itself than the bromine atom does.
This stronger pull in HCl creates a more polar bond, meaning the hydrogen side is more positive and the halogen side is more negative compared to the bond in HBr. The difference in electron-pulling strength also influences the relative chemical reactivity of the two elements. Since chlorine has a stronger desire to gain an electron, it is generally a more reactive element than bromine in specific types of chemical reactions. The greater electronegativity of chlorine means it is a more powerful oxidizing agent.