Question

What is meant by the term ground state? Write the electron configuration of an atom of carbon in the ground state. Write one possible electron configuration of a carbon atom in an excited state (i.e., not in the ground state).

Short answer

The ground state is the lowest energy state of an atom. For carbon, the ground state electron configuration is 1s² 2s² 2p². An example of an excited state configuration for carbon could be 1s² 2s¹ 2p³.

Step-by-step solution

Definition of Ground State

The ground state of an atom refers to the lowest energy state of the atom where all electrons are in the lowest possible energy levels according to the Pauli exclusion principle and Hund’s rule.

Electron Configuration of Carbon in Ground State

Carbon has an atomic number of 6, indicating it has 6 electrons. The electron configuration for carbon in the ground state is written as: 1s² 2s² 2p², where the superscripts represent the number of electrons in each subshell and the numbers indicate the principal energy level.

Electron Configuration of Carbon in Excited State

In an excited state, one or more electrons have absorbed energy and moved to a higher energy level. One possible electron configuration of carbon in an excited state could be 1s² 2s¹ 2p³, where an electron from the 2s orbital has moved to the 2p orbital.

Key concepts

Excited State Electron Configuration

When we talk about an atom's electrons, they can be in different energy states. The ground state is their natural, lowest energy condition. However, when an atom absorbs energy, its electrons can 'jump' to higher levels or orbitals. This state is known as an excited state. Unlike the ground state, which is predictable, the excited state has many variations since electrons can jump to different orbitals based on how much energy they absorb.

For example, consider a carbon atom which usually has the ground state configuration of 1s² 2s² 2p². In an excited state, such as 1s² 2s¹ 2p³, one electron from the 2s subshell has absorbed energy and has moved to a higher energy level, in this case, the 2p subshell. It's important to visualize this process as electrons occupying seats in an auditorium - when one gets excited, it may move to a higher seat (energy level), leaving its lower seat (ground state) behind.

Pauli Exclusion Principle

A key concept when understanding electron configurations is the Pauli Exclusion Principle. This principle, proposed by Wolfgang Pauli in 1925, states that no two electrons in an atom can have the same set of four quantum numbers. This effectively means that each electron must be unique in its characteristics and occupy its own space.

In simpler terms, think of it like this: in any given orbital, only two electrons can exist, and they must have opposite spins. This is like saying in a two-seater car, you can have only two passengers at a time, and they can't be facing the same direction. When writing electron configurations, this is why we have upward and downward arrows (representing electron spins) and why pairs of electrons occupy orbitals individually.

Hund’s Rule

Another fundamental rule for assigning electron configurations is Hund's Rule. This rule states that electrons will fill up orbitals of the same energy (degenerate orbitals) singly first before doubling up. Equally important, the electrons that occupy these orbitals singly will all have the same spin direction initially.

Using the analogy of a bus with empty seats, Hund's Rule would suggest that passengers (electrons) would prefer to sit alone before sharing a seat. And while they sit alone, they all face forward (same spin). Only when there are no empty rows left, they'll start sitting next to each other with one of them facing backward (opposite spin). This is essential to minimize the repulsion between electrons and keep the atom in its most stable arrangement.

Atomic Number

The concept of the atomic number is fundamental in chemistry as it tells us the number of protons in the nucleus of an atom, which in turn reveals the element's identity. For neutral atoms, the atomic number also equals the number of electrons around the nucleus. The atomic number, therefore, determines the chemical properties of the element and its place in the periodic table.

For example, carbon has an atomic number of 6, meaning it has 6 protons and, when neutral, 6 electrons. The atomic number helps us to begin constructing an atom's electron configuration by telling us how many electrons we need to place in orbitals according to the rules we've discussed. Remember, the configuration of electrons based on increasing energy levels (1s, 2s, 2p, etc.) and applying the Pauli Exclusion Principle and Hund's Rule is key to understanding an element's chemical behavior and reactivity.