Question

Draw the Lewis structure of the atom whose electron configuration is \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{6} 4 s^{2} 3 d^{10} 4 p^{5}\). Would you expect this atom to form an ionic, nonpolar covalent, or polar covalent bond with sulfur?

Short answer

The element is bromine (Br), and the Lewis structure consists of the symbol 'Br' with seven dots around it representing its valence electrons. Bromine is expected to form a polar covalent bond with sulfur.

Step-by-step solution

Determine the Element

Based on the given electron configuration, identify that the element is bromine (Br). The electron configuration ends with 4p^5, which is one electron less than the noble gas configuration of krypton (Kr), indicating that it is the element before Kr in the periodic table.

Lewis Structure of Bromine

To draw the Lewis structure, start by writing the symbol 'Br' for bromine. Then distribute seven dots around the symbol representing its valence electrons (since it is in Group 17 and has seven valence electrons).

Determine the Type of Bond with Sulfur

Given that bromine and sulfur (S) are both nonmetals and differ in electronegativity, they are likely to form a polar covalent bond due to the electron sharing and partial charges that will develop.

Key concepts

Electron Configuration

Understanding the electron configuration of an atom provides a fundamental basis for predicting its chemical behavior, including how it will bond with other atoms. Every atom is composed of protons, neutrons, and electrons, with the electrons occupying space around the nucleus in areas called orbitals. The arrangement of electrons within these orbitals is known as the electron configuration, and it follows a specific order defined by quantum mechanics.

The electron configuration is denoted by numbers and letters that indicate the energy level, type of orbital, and the number of electrons in those orbitals. For example, the configuration 1s²2s²2p⁶ implies that there are 2 electrons in the first energy level's s orbital, 2 in the second level's s orbital, and 6 in the second level's p orbital. The electron configuration of an element essentially represents its position in the periodic table and hence the potential to gain, lose, or share electrons during chemical bonding.

Chemical Bonding

Chemical bonding is the force of attraction that holds atoms together in compounds, allowing them to function in a synchronized manner. There are three primary types of chemical bonds: ionic, covalent, and metallic. Ionic bonds are formed between metals and nonmetals when electrons are transferred from one atom to another, resulting in the formation of ions with opposite charges that attract each other.

On the other hand, covalent bonds occur when atoms, typically nonmetals, share one or more pairs of valence electrons, fostering a strong attraction between the atoms while allowing each to attain a more stable electron configuration. If this sharing is equal, the bond is nonpolar covalent; if unequal, it's polar covalent. The distinction lies in the electronegativity difference between the involved atoms, with a polar bond showing a partial charge separation due to uneven sharing. The example of bromine and sulfur points towards a polar covalent bond, which is influenced by their individual electronegativities and their tendency to share electrons unequally.

Valence Electrons

Valence electrons are the electrons in the outermost shell of an atom and they play a crucial role in determining an element's chemical properties, specifically its reactivity and the type of bonds it can form. These electrons can be easily influenced during chemical reactions because they have the most energy compared to electrons in inner shells.

In the case of bromine, with an electron configuration ending in 4p⁵, it has seven valence electrons. These are the electrons used when forming chemical bonds. When examining the reactivity of an element, these valence electrons are what interact with the valence electrons of other atoms to create either ionic or covalent bonds. According to the octet rule, atoms strive to reach the electron configuration of a noble gas, having eight valence electrons. Consequently, bromine will typically either accept an electron to fill its p orbitals and achieve a stable octet or share electrons through covalent bonding, as predicted in its interaction with sulfur.