Question

How many valence electrons are in each of the following atoms or ions? \(\mathrm{Xe}, \mathrm{Sr}^{2+}, \mathrm{Cl},\) and \(\mathrm{Cl}^{-}.\)

Short answer

Answer: The number of valence electrons for each atom and ion are as follows: Xe has 8 valence electrons, Sr^2+ has 8 valence electrons, Cl has 7 valence electrons, and Cl^- has 8 valence electrons.

Step-by-step solution

Find the electronic configuration of each atom and ion

To find the electronic configuration, we first determine the number of electrons present in each atom according to their atomic numbers, and then arrange the electrons in different orbitals following the aufbau principle. 1. Xe (Atomic number 54): Electronic configuration: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^{10} 5p^6\) 2. Sr (Atomic number 38): Electronic configuration: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2\) 3. Sr^2+ ion: When Sr lost two electrons to form the Sr^2+ ion, its electronic configuration becomes: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6\) 4. Cl (Atomic number 17): Electronic configuration: \(1s^2 2s^2 2p^6 3s^2 3p^5\) 5. Cl^- ion: When Cl gains one electron to form the Cl^- ion, its electronic configuration becomes: \(1s^2 2s^2 2p^6 3s^2 3p^6\)

Identify the valence electrons

After finding the electronic configuration, look for the electrons in the outermost energy level. 1. Xe: The outermost energy level is 5, and there are 2 electrons in the 5s orbital and 6 electrons in the 5p orbital. Total valence electrons = 2 + 6 = 8 2. Sr^2+: The outermost energy level of the Sr^2+ ion is 4, and there are 2 electrons in the 4s orbital and 6 electrons in the 4p orbital. Total valence electrons = 2 + 6 = 8 3. Cl: The outermost energy level is 3, and there are 2 electrons in the 3s orbital and 5 electrons in the 3p orbital. Total valence electrons = 2 + 5 = 7 4. Cl^-: The outermost energy level of the Cl^- ion is 3, and there are 2 electrons in the 3s orbital and 6 electrons in the 3p orbital. Total valence electrons = 2 + 6 = 8 Now we have the number of valence electrons for each of the given atoms and ions: Xe: 8 valence electrons Sr^2+: 8 valence electrons Cl: 7 valence electrons Cl^-: 8 valence electrons

Key concepts

Electronic Configuration

Electronic configuration is a method used to represent the arrangement of electrons within an atom or ion. Electrons occupy specific energy levels, known as shells, and each shell is further divided into subshells, labeled as s, p, d, and f. The number and arrangement of electrons in these subshells determine the chemical properties of an element.
To determine the electronic configuration, we follow the aufbau principle, which states that electrons fill orbitals starting with the lowest energy levels first. The order of filling is further guided by another tool called the "electron configuration chart" or periodic table order. For example:

• Xenon (Xe): With an atomic number of 54, Xe has the electronic configuration: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^{10} 5p^6\).
• Strontium (Sr): At atomic number 38, it’s \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2\). The Sr²⁺ ion loses two electrons: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6\).
• Chlorine (Cl): With atomic number 17, the configuration is \(1s^2 2s^2 2p^6 3s^2 3p^5\). Upon gaining an electron, Cl^- turns into \(1s^2 2s^2 2p^6 3s^2 3p^6\).

Electronic configurations are crucial for understanding how elements will interact with each other during chemical reactions.

Atoms and Ions

Atoms are the basic units of matter, consisting of a nucleus surrounded by electrons. The behavior of atoms is significantly influenced by the arrangement of these electrons, primarily those in the outermost shell, which are called valence electrons.
An ion is an atom that has lost or gained electrons, giving it a net positive or negative charge. When an atom loses electrons, it becomes positively charged and is called a cation. Conversely, when an atom gains electrons, it becomes negatively charged, known as an anion.
Considering examples:

• Sodium ion (Na⁺): Originally an atom with the configuration \(1s^2 2s^2 2p^6 3s^1\), it loses one electron to become stable, leading to a configuration of \(1s^2 2s^2 2p^6\).
• Chloride ion (Cl^-): Starts with \(1s^2 2s^2 2p^6 3s^2 3p^5\) and gains an electron to become \(1s^2 2s^2 2p^6 3s^2 3p^6\).

This gain or loss typically makes ions more stable than their neutral atom forms due to achieving a full outer shell, commonly resembling the electron configuration of noble gases.

Periodic Table

The periodic table is a systematic arrangement of elements according to their atomic number, electron configuration, and recurring chemical properties. It is a powerful tool for predicting the properties of elements and their compounds.
Elements are organized into rows called periods and columns known as groups or families. Each element in a group exhibits similar valence electron configurations, resulting in akin chemical properties.

• Groups: Vertical columns, like Group 1 (alkali metals), where each element has a single electron in its outer shell.
• Periods: Horizontal rows indicate the number of electron shells. Elements across a period gradually fill these shells.

The table also helps identify the metallic character and reactivity of elements. Transition metals, found in the d-block, have varied oxidation states, crucial for forming stable ions and compounds. The periodic table thus serves as a roadmap for understanding elements' behavior in reactions.

Chemical Bonding

Chemical bonding is the force that holds atoms together in compounds. The type and strength of bonds determine the properties of the resulting substances. Understanding the electronic configuration and valence electrons helps in predicting how an atom will bond.
Ionic Bonding: This occurs when atoms transfer electrons between them, resulting in a bond between positive and negative ions. For example, sodium (Na) can donate an electron to chlorine (Cl), forming an ionic compound, NaCl, with a stable electronic structure.
Covalent Bonding: Here, atoms share pairs of electrons. Non-metal elements typically form covalent bonds when they have similar electronegativities. Molecules like \(\text{O}_2\) and \(\text{H}_2\) are common examples of covalent bonding.
A deep understanding of chemical bonding requires examining how atoms approach stability by achieving a complete outer electron shell, often referred to as the "octet rule." In conclusion, bonding is essential in the formation of matter's diverse states and nature.