Question

Use the basic rules for electron configurations to indicate the number of (a) unpaired electrons in an atom of \(\mathrm{P} ;\) (b) \(3 d\) electrons in an atom of \(\mathrm{Br} ;\) (c) \(4 p\) electrons in an atom of \(\mathrm{Ge} ;\) (d) \(6 \mathrm{s}\) electrons in an atom of \(\mathrm{Ba}\) (e) \(4 f\) electrons in an atom of Au.

Short answer

(a) 3 unpaired electrons for P. (b) 0 3d electrons for Br. (c) 2 4p electrons for Ge. (d) 2 6s electrons for Ba. (e) 14 4f electrons for Au.

Step-by-step solution

Unpaired Electrons in Atom of P

Phosphorus (P) has 15 electrons. The electron configuration is \(1s^22s^22p^63s^23p^3\). We can see that there are 3 unpaired electrons in p orbital of the third energy level.

Number of 3d Electrons in Atom of Br

Bromine (Br) is the 35th element with 35 electrons. The electron configuration for bromine ends with \(4p^5\). There are no d electrons in the third energy level because d orbital starts filling from the 4th period.

Number of 4p Electrons in Atom of Ge

Germanium (Ge) is the 32nd element with 32 electrons. The electron configuration ends with \(4p^2\). Therefore, Germanium has 2 electrons in the 4p orbital.

Number of 6s Electrons in Atom of Ba

Barium (Ba) has 56 electrons. The electron configuration ends with \(6s^2\). Hence, the Barium atom has 2 electrons in the 6s orbital.

Number of 4f Electrons in Atom of Au

Gold (Au) has 79 electrons. Its electron configuration ends with \(6s^14f^{14}5d^{10}\). Thus, there are 14 electrons in the 4f orbital.

Key concepts

Unpaired Electrons

When discussing unpaired electrons, we're talking about electrons in an electron orbital that are alone rather than in pairs. The concept of unpaired electrons is crucial in chemistry because these electrons influence an atom's magnetic properties and reactivity. Phosphorus, for example, has an electron configuration of \(1s^22s^22p^63s^23p^3\). Here, the last set of electrons, in the 3p orbitals, includes three unpaired electrons. This is because the 3p subshell can hold up to six electrons, but when filled with only three, each electron occupies a separate orbital. The presence of these unpaired electrons makes phosphorus reactive and influences how it bonds with other elements.

Orbital Filling

Orbital filling is the process of allocating electrons into an atom's orbitals based on increasing energy levels. The specific order is determined by the Aufbau principle, which states that electrons fill the lowest available energy level before moving to higher ones. According to the principle, the filling order follows a specific sequence: 1s, 2s, 2p, 3s, and so on.

• For Bromine, the electron configuration ends in \(4p^5\), indicating that all lower energy levels and the 4p orbital are sequentially filled, leading up to 34 electrons in a specific order before reaching the fifth electron in the 4p orbital.
• In the case of Barium, its configuration ends in \(6s^2\), showing the atoms fill up to the sixth energy level, specifically filling the 6s subshell.

Understanding this filling order helps predict electron configuration and chemical properties of elements.

Energy Levels

Energy levels in an atom describe the distances of electron orbitals from the nucleus. Electrons closer to the nucleus are in lower energy levels and have less energy, while those further away are in higher energy levels with more energy. This concept is significant because it helps explain the stability of electron configurations.In Germanium's configuration, which ends at \(4p^2\), the electrons are at the fourth energy level. Each period on the periodic table represents a different energy level, and elements fill these levels successively as atomic number increases. Higher energy levels have more sublevels (s, p, d, f) and can store more electrons.

Electron Orbitals

Electron orbitals are regions around the nucleus where electrons are most likely to be found. These orbitals come in various shapes, like spherical for s-orbitals and dumbbell-shaped for p-orbitals. The configuration of these orbitals determines many properties of an atom, including its chemical reactivity and location on the periodic table. A practical application is seen in Gold's electron configuration: \(6s^14f^{14}5d^{10}\). Here, the 4f subshell is fully filled with 14 electrons. These filled orbitals contribute to Gold's unique properties as a metal, such as its malleability and conductivity. Understanding the arrangement and filling of these orbitals is key to many chemical phenomena and predictions.