Question

Write orbital diagrams for the valence electrons and indicate the number of unpaired electrons for each element. (a) Ne (b) \(I\) (c) \(\mathrm{Sr}\) (d) \(\mathrm{Ge}\)

Short answer

Ne has 0 unpaired electrons; I has 1 unpaired electron; Sr has 0 unpaired electrons; Ge has 0 unpaired electrons.

Step-by-step solution

Writing Orbital Diagram for Ne

Neon (Ne) has an atomic number of 10. The electronic configuration is 1s2 2s2 2p6. There are two electrons each in the 1s and 2s orbitals, and six electrons in the 2p orbital, all are paired. The orbital diagram shows no unpaired electrons.

Writing Orbital Diagram for I

Iodine (I) has an atomic number of 53. The valence electrons are in the fifth shell, with the electronic configuration of 5s2 5p5. There are two electrons in the 5s orbital, and five electrons in the 5p orbital. With five electrons in the 5p orbital, there is one unpaired electron in the 5p orbital.

Writing Orbital Diagram for Sr

Strontium (Sr) has an atomic number of 38. The valence electrons are in the fifth shell, with the electronic configuration of 5s2. There are two electrons in the 5s orbital, both are paired. The orbital diagram shows no unpaired electrons.

Writing Orbital Diagram for Ge

Germanium (Ge) has an atomic number of 32. The valence electrons are in the fourth shell, with the electronic configuration of 4s2 3d10 4p2. There are two electrons in the 4s orbital, ten electrons in the 3d orbitals which are all paired, and two electrons in the 4p orbital, both are paired. The orbital diagram shows no unpaired electrons.

Key concepts

Unpaired Electrons

When it comes to understanding chemical reactions and the properties of elements, knowing about unpaired electrons is crucial. Unpaired electrons are the electrons in an atom that are not part of an electron pair. These lone electrons have a strong tendency to pair up with other unpaired electrons during chemical reactions, often forming chemical bonds.

For example, iodine (I) has an electronic configuration of 5s2 5p5, indicating that it has one unpaired electron in the 5p orbital. This single electron plays a significant role in iodine's ability to participate in bonding and its general reactivity. As opposed to iodine, neon (Ne), strontium (Sr), and germanium (Ge) do not exhibit any unpaired electrons in their valence shells, rendering them relatively inert and less likely to form bonds under normal conditions.

Understanding the presence of unpaired electrons can help students predict magnetic properties, reactivity, and the types of bonds an element might form. For example, atoms with unpaired electrons can exhibit paramagnetism, while those with all paired electrons tend to be diamagnetic.

Electronic Configuration

The electronic configuration of an element details how electrons are distributed in an atom's orbitals. It is a foundational concept in chemistry that elucidates the atom's behavior during chemical reactions. The configuration uses the principle of lowest energy fill first, meaning electrons occupy the lowest energy orbitals available.

Let's take neon (Ne) as an example. Its electronic configuration, 1s2 2s2 2p6, reveals how its ten electrons are arranged. This configuration tells us that the 1s orbital is filled first with two electrons, followed by the 2s orbital with another pair, and finally, the 2p orbital with six electrons, which fills it completely. Other elements such as iodine (I), strontium (Sr), and germanium (Ge) have more complicated configurations that denote their position in the periodic table and their chemical properties.

The knowledge of electronic configuration not only helps in understanding the structure of an atom but also guides chemists in predicting the atom’s behavior in terms of reactivity and bonding patterns.

Valence Electrons

Valence electrons are the electrons in the outermost shell of an atom that are involved in forming bonds. These are the electrons that have the most influence on an atom's chemical properties and its reactivity with other atoms to form molecules.

For instance, strontium (Sr) has the electronic configuration 5s2, telling us it has two valence electrons. These two electrons are what Sr will use to form bonds. Similarly, Germanium (Ge) with the electronic configuration 4s2 3d10 4p2 has four valence electrons, which it can share or exchange in chemical reactions. Knowing how many valence electrons are available in an element can aid students in predicting how an element will react and what type of ion it may form, as well as understanding its place in the periodic table.

Atomic Number

The atomic number is a fundamental concept in chemistry that represents the number of protons in an atom's nucleus. It is denoted by the letter Z and also defines the number of electrons surrounding the nucleus of an atom in a neutral state.

Atomic number not only categorizes elements in the periodic table but also is a defining attribute for an element's identity. For example, neon has an atomic number of 10, germanium has 32, strontium comes with 38, and iodine has 53. Each of these atomic numbers tells us exactly what element we're dealing with. Knowing the atomic number allows students to determine the electronic configuration and, subsequently, the chemical behavior of an element in a more straightforward way. It's a simple yet effective tool to start understanding the complexities of elements.

The concept of the atomic number also helps in comprehending isotopes, as isotopes of an element will have the same atomic number but different mass numbers due to variations in the number of neutrons.