Question

Write electron configurations for each transition metal. (a) \(\mathrm{Zn}\) (b) \(\mathrm{Cu}\) (c) \(\mathrm{Zr}\) (d) \(\mathrm{Fe}\)

Short answer

The electron configurations are: (a) Zn: '1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10}', (b) Cu: '1s^2 2s^2 2p^6 3s^2 3p^6 3d^{10} 4s^1', (c) Zr: '1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^2', (d) Fe: '1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^6'.

Step-by-step solution

Identify the Atomic Number

Look up the atomic number for each transition metal on the periodic table. The atomic number of a chemical element is the number of protons found in the nucleus of an atom of that element.

Understand Electron Configuration Notation

Electron configurations are written using the principal quantum number (n), the sublevel (s, p, d, or f), and the number of electrons in that sublevel, e.g., '1s^2' means two electrons in the 's' sublevel of the first energy level.

Write the Electron Configuration for Zn (Atomic Number 30)

Zinc's electron configuration can be represented by filling up the energy levels and sublevels in order: 1s, 2s, 2p, 3s, 3p, 4s, 3d. Since Zinc has 30 electrons, the configuration is '1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10'. The d sublevel is filled with 10 electrons because it is a transition metal and that is a characteristic property.

Write the Electron Configuration for Cu (Atomic Number 29)

Copper's electron configuration is an exception to the normal order of filling orbitals because of its additional stability from a filled or half-filled d sublevel. Normally, it would be '1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^9', but because a half-filled d10 sublevel is more stable, one electron from the 4s orbital moves to the 3d orbital, making Copper's configuration '1s^2 2s^2 2p^6 3s^2 3p^6 3d^{10} 4s^1'.

Write the Electron Configuration for Zr (Atomic Number 40)

Zirconium's electron configuration can be figured out by filling the orbitals in order. For Zirconium 40 electrons must be placed in the orbitals, thus the configuration is '1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^2'.

Write the Electron Configuration for Fe (Atomic Number 26)

Iron's electrons fill up to the 3d sublevel. The configuration would typically be '1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^6'. Since the d sublevel can hold up to 10 electrons, and there are only 6, there's no need for exceptions in Iron's case.

Key concepts

Atomic Number

The atomic number is a fundamental concept in understanding elements on the periodic table. It is denoted by the symbol 'Z' and represents the number of protons in the nucleus of an atom. For example, iron (Fe) has an atomic number of 26, which means each atom of iron has 26 protons. The atomic number also determines the element's position in the periodic table and hence its chemical behavior. In the context of electron configurations, the atomic number signifies the total number of electrons in a neutral atom, which we need to distribute across the different energy levels and sublevels appropriately.

When writing electron configurations for transition metals, the atomic number allows us to understand how many electrons we’re dealing with. It’s the starting point for arranging electrons in the increasing order of their energy levels and visualizing the electron cloud around the nucleus.

Electron Configuration Notation

Electron configuration notation provides us with a map of where electrons are located around an atom's nucleus. It involves quantum numbers and can be viewed as the 'address' for each electron. The notation includes the principal quantum number (n), which indicates the energy level, followed by a letter (s, p, d, f) that specifies the sublevel, and a superscript number that tells us the number of electrons in that sublevel. For instance, '2p^6' indicates that there are six electrons in the p sublevel of the second energy level.

The notation follows a pattern based on the Aufbau principle, which states that electrons fill orbitals starting with the lowest energy level and move to higher levels progressively. However, transition metals often present complex cases where this straightforward filling order isn't always observed due to additional stability considerations.

d Sublevel Electron Configuration

Transition metals are characterized by their partially filled d sublevels, giving them unique properties and a block of their own on the periodic table. The d sublevel can hold up to 10 electrons. The transition metals fill in the order of '4s' before '3d', a detail that might seem counterintuitive but is predicted by the sequence of energy levels. For example, scandium (Sc) with an atomic number of 21 has the electron configuration '1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^1'.

The arrangement in d orbitals is crucial as it impacts the magnetic and chemical properties of the elements. Furthermore, the d sublevel configuration gives rise to the color and varied oxidation states among transition metals, making them integral in many biological and industrial processes.

Exceptional Electron Configurations

Some transition metals have electron configurations that deviate from the expected patterns due to the relatively higher stability offered by half-filled or completely filled d sublevels. Such exceptions include chromium (Cr) and copper (Cu). Normally, copper would be expected to have the configuration '4s^2 3d^9' but instead follows '3d^{10} 4s^1'. This is because a full d sublevel offers more stability than a partially filled one with just nine electrons.

The anomaly in the electron configurations can often be justified by the specific electron–electron interactions within these sublevels, which lead to a reconfiguration that minimizes the atom's energy. Learning about these exceptions is essential as they offer deep insights into the behavior of elements and influence their placement on the periodic table in the transition metal series.