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Chapter 10: Problem 16

What is the difference between the ground state of an atom and an excited state of an atom?

Short Answer

Expert verified
The ground state is the lowest and most stable energy state of an atom, with electrons in the nearest possible orbitals to the nucleus. The excited state has a higher energy level, with electrons in higher energy orbitals, and is less stable.

Step by step solution

01

Understanding the Ground State

The ground state of an atom is its lowest energy state, where the electrons are in the closest possible orbitals to the nucleus according to the Pauli exclusion principle and Hund's rule. In this state, an atom is at its minimum energy level and is stable.
02

Understanding the Excited State

An excited state of an atom occurs when one or more electrons absorb energy and transition to higher energy orbitals, further away from the nucleus. This state is temporary and the atom is in a higher energy level than its ground state.
03

Comparing Ground and Excited States

The difference between the ground state and an excited state is that in the ground state, the atom has the lowest possible energy, whereas in the excited state, the atom has a higher energy due to the presence of electrons in higher energy orbitals. Excited states are not as stable as the ground state, and the atom will eventually release energy and return to the ground state.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Atomic Energy Levels
The concept of atomic energy levels is fundamental to understanding how electrons exist in atoms. Atoms are structured with a nucleus at the center, surrounded by electrons that occupy specific energy levels, or 'shells.' These shells represent distinct energy levels, and the closest shell to the nucleus is the lowest energy level. As electrons move further from the nucleus, they occupy higher energy levels, which require more energy to maintain.

These energy levels aren't arbitrary, but are quantized, meaning that an electron must gain or lose a certain amount of energy to move between them. This behavior underpins the absorption and emission spectra of elements, with each element having a unique set of energy levels. Whenever an electron jumps from one level to another, it must absorb or release a precise energy amount, which can be observed as specific wavelengths of light.
Electron Orbitals
An electron orbital is a region within an atom where there is a high probability of finding an electron. Each orbital can hold a specific number of electrons, for example, s orbitals can hold 2, p orbitals can hold 6, and so on. These orbitals are not fixed paths like planets around a sun but are more like clouds that represent varying probabilities of where an electron may be at any given time.

Shapes and Sizes

Orbitals have different shapes鈥攕 orbitals are spherical, while p orbitals are dumbbell-shaped. The shape of the orbital affects the electron's probability distribution and energy. As electrons fill these orbitals, they follow a specific order, known as the Aufbau principle, which fills the lower energy orbitals first.
Pauli Exclusion Principle
The Pauli exclusion principle is a quantum mechanical principle that states no two electrons in an atom can have the same set of four quantum numbers. Essentially, this rule means that an orbital can hold a maximum of two electrons, and they must have opposite spins. This rule is a cornerstone of chemistry as it explains the arrangement of electrons in an atom.

Quantum Numbers

Quantum numbers describe the properties of electrons within an atom. There are four types: the principal quantum number (n), which indicates the energy level, the azimuthal quantum number (l), which defines the shape of the orbital, the magnetic quantum number (m), which specifies the orbital's orientation in space, and the spin quantum number (s), which tells the direction of the electron's spin.
Hund's Rule
Hund's rule addresses how electrons are distributed among orbitals of the same energy, known as 'degenerate' orbitals. Hund's rule states that electrons will fill an unoccupied orbital before they pair up in an already occupied one. This is because electrons, being negatively charged, repel each other, and occupying separate orbitals minimizes their repulsion.

This rule helps to determine the electron configuration of atoms in their ground states. Electrons will spread out to occupy as many orbitals as possible before doubling up, which also results in the most stable and lowest energy configuration for an atom, as electron repulsion is minimized.

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Most popular questions from this chapter

Use the electron configuration of oxygen to explain why it tends to form a 2 - ion.

When an electron makes a transition from the \(n=4\) to the \(n=2\) hydrogen atom Bohr orbit, the energy difference between these two orbits \(\left(4.1 \times 10^{-19} \mathrm{~J}\right)\) is emitted as a photon of light. The relationship between the energy of a photon and its wavelength is given by \(E=h c / \lambda\), where \(E\) is the energy of the photon in \(\mathrm{J}, h\) is Planck's constant \(\left(6.626 \times 10^{-34} \mathrm{~J} \cdot \mathrm{s}\right)\), and \(c\) is the speed of light \(\left(3.00 \times 10^{8} \mathrm{~m} / \mathrm{s}\right)\). Find the wavelength of light emitted by hydrogen atoms when an electron makes this transition.

What is the relationship between the wavelength of light and the amount of energy carried by its photons? How are wavelength and frequency of light related?

When atoms lose more than one electron, the ionization energy to remove the second electron is always more than the ionization energy to remove the first. Similarly, the ionization energy to remove the third electron is more than the second and so on. However, the increase in ionization energy upon the removal of subsequent electrons is not necessarily uniform. For example, consider the first three ionization energies of magnesium: \(\begin{array}{ll}\text { First ionization energy } & 738 \mathrm{~kJ} / \mathrm{mol} \\ \text { Second ionization energy } & 1450 \mathrm{~kJ} / \mathrm{mol} \\ \text { Third ionization energy } & 7730 \mathrm{~kJ} / \mathrm{mol}\end{array}\) The second ionization energy is roughly twice the first ionization energy, but then the third ionization energy is over five times the second. Use the electron configuration of magnesium to explain why this is so. Would you expect the same behavior in sodium? Why or why not?

Explain the difference between valence electrons and core electrons.
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