Chapter 8. Electronic Structure
Quantum Numbers for Electrons
Learning Objectives
 Explain what spectra are.
 Learn the quantum numbers that are assigned to electrons.
There are two fundamental ways of generating light: either heat an object up so hot it glows or pass an electrical current through a sample of matter (usually a gas). Incandescent lights and fluorescent lights generate light via these two methods, respectively.
A hot object gives off a continuum of light. We notice this when the visible portion of the electromagnetic spectrum is passed through a prism: the prism separates light into its constituent colours, and all colours are present in a continuous rainbow (part (a) in Figure 8.16 “Prisms and Light”). This image is known as a . However, when electricity is passed through a gas and light is emitted and this light is passed though a prism, we see only certain lines of light in the image (part (b) in Figure 8.16 “Prisms and Light”). This image is called a line spectrum. It turns out that every element has its own unique, characteristic line spectrum.
Why does the light emitted from an electrically excited gas have only certain colours, while light given off by hot objects has a continuous spectrum? For a long time, it was not well explained. Particularly simple was the spectrum of hydrogen gas, which could be described easily by an equation; no other element has a spectrum that is so predictable (Figure 8.17 “Hydrogen Spectrum”). Latenineteenthcentury scientists found that the positions of the lines obeyed a pattern given by the following equation:
Where n = 3, 4, 5, 6,…, but they could not explain why this was so.
In 1913, the Danish scientist Niels Bohr suggested a reason why the hydrogen atom spectrum looked this way. He suggested that the electron in a hydrogen atom could not have any random energy, having only certain fixed values of energy that were indexed by the number n (the same n in the equation above and now called a ). Quantities that have certain specific values are called . Bohr suggested that the energy of the electron in hydrogen was quantized because it was in a specific orbit. Because the energies of the electron can have only certain values, the changes in energies can have only certain values (somewhat similar to a staircase: not only are the stair steps set at specific heights but the height between steps is fixed). Finally, Bohr suggested that the energy of light emitted from electrified hydrogen gas was equal to the energy difference of the electron’s energy states:
This means that only certain frequencies (and thus, certain wavelengths) of light are emitted. Figure 8.18 “Bohr’s Model of the Hydrogen Atom” shows a model of the hydrogen atom based on Bohr’s ideas.
Bohr’s ideas were useful but were applied only to the hydrogen atom. However, later researchers generalized Bohr’s ideas into a new theory called , which explains the behaviour of electrons as if they were acting as a wave, not as particles. Quantum mechanics predicts two major things: quantized energies for electrons of all atoms (not just hydrogen) and an organization of electrons within atoms. Electrons are no longer thought of as being randomly distributed around a nucleus or restricted to certain orbits (in that regard, Bohr was wrong). Instead, electrons are collected into groups and subgroups that explain much about the chemical behaviour of the atom.
In the quantummechanical model of an atom, the state of an electron is described by four quantum numbers, not just the one predicted by Bohr. The first quantum number is called the . Represented by n, the principal quantum number largely determines the energy of an electron. Electrons in the same atom that have the same principal quantum number are said to occupy an of the atom. The principal quantum number can be any nonzero positive integer: 1, 2, 3, 4,….
Within a shell, there may be multiple possible values of the next quantum number, the , represented by ℓ. The ℓ quantum number has a minor effect on the energy of the electron but also affects the spatial distribution of the electron in threedimensional space — that is, the shape of an electron’s distribution in space. The value of the ℓ quantum number can be any integer between 0 and n − 1:
Thus, for a given value of n, there are different possible values of ℓ, as shown in Table 8.2.
If n equals  ℓ can be 

1  0 
2  0 or 1 
3  0, 1, or 2 
4  0, 1, 2, or 3 
and so forth. Electrons within a shell that have the same value of ℓ are said to occupy a in the atom. Commonly, instead of referring to the numerical value of ℓ, a letter represents the value of ℓ (to help distinguish it from the principal quantum number):
If ℓ equals  The subshell is 

0  s 
1  p 
2  d 
3  f 
The next quantum number is called the , represented by mℓ. For any value of ℓ, there are 2ℓ + 1 possible values of m_{ℓ}, ranging from −ℓ to ℓ:
The following explicitly lists the possible values of m_{ℓ} for the possible values of ℓ:
If ℓ equals  The mℓ values can be 

0  0 
1  −1, 0, or 1 
2  −2, −1, 0, 1, or 2 
3  −3, −2, −1, 0, 1, 2, or 3 
The particular value of m_{ℓ} dictates the orientation of an electron’s distribution in space. When ℓ is zero, m_{ℓ} can be only zero, so there is only one possible orientation. When ℓ is 1, there are three possible orientations for an electron’s distribution. When ℓ is 2, there are five possible orientations of electron distribution. This goes on and on for other values of ℓ, but we need not consider any higher values of ℓ here. Each value of m_{ℓ} designates a certain . Thus, there is only one orbital when ℓ is zero, three orbitals when ℓ is 1, five orbitals when ℓ is 2, and so forth. The m_{ℓ} quantum number has no effect on the energy of an electron unless the electrons are subjected to a magnetic field — hence its name.
The ℓ quantum number dictates the general shape of electron distribution in space (Figure 8.19 “Electron Orbitals”). As shown in part (a), any s orbital is spherically symmetric, and there is only one orbital in any s subshell. Part (b) shows that any p orbital has a twolobed, dumbbelllike shape; because there are three of them, we normally represent them as pointing along the x, y, and zaxes of Cartesian space. In part (c), we see that d orbitals are fourlobed rosettes; they are oriented differently in space (the one labelled d_{z}^{2} has two lobes and a torus instead of four lobes, but it is equivalent to the other orbitals). When there is more than one possible value of m_{ℓ}, each orbital is labelled with one of the possible values. It should be noted that the diagrams in Figure 8.19 are estimates of the electron distribution in space, not surfaces electrons are fixed on.
The final quantum number is the , represented by m_{s}. Electrons and other subatomic particles behave as if they are spinning (we cannot tell if they really are, but they behave as if they are). Electrons themselves have two possible spin states, and because of mathematics, they are assigned the quantum numbers +½ and −½. These are the only two possible choices for the spin quantum number of an electron.
Example 8.6
Problems
Of the set of quantum numbers {n, ℓ, m_{ℓ}, m_{s}}, which are possible and which are not allowed?
 {3, 2, 1, +½}
 {2, 2, 0, −½}
 {3, −1, 0, +½}
Solutions
 The principal quantum number n must be an integer, which it is here. The quantum number ℓ must be less than n, which it is. The m_{ℓ} quantum number must be between −ℓ and ℓ, which it is. The spin quantum number is +½, which is allowed. Because this set of quantum numbers follows all restrictions, it is possible.
 The quantum number n is an integer, but the quantum number ℓ must be less than n, which it is not. Thus, this is not an allowed set of quantum numbers.
 The principal quantum number n is an integer, but ℓ is not allowed to be negative. Therefore, this is not an allowed set of quantum numbers.
Test Yourself
Of the set of quantum numbers {n, ℓ, m_{ℓ}, m_{s}}, which are possible and which are not allowed?
 {4, 2, −2, 1}
 {3, 1, 0, −½}
Answers
 Spin must be either +½ or −½, so this set of quantum number is not allowed.
 allowed
Chemistry Is Everywhere: Neon Lights
A neon light is basically an electrified tube with a small amount of gas in it. Electricity excites electrons in the gas atoms, which then give off light as the electrons go back into a lower energy state. However, many socalled “neon” lights don’t contain neon!
Although we know now that a gas discharge gives off only certain colours of light, without a prism or other component to separate the individual light colours, we see a composite of all the colours emitted. It is not unusual for a certain colour to predominate. True neon lights, with neon gas in them, have a reddishorange light due to the large amount of red, orange, and yellowcoloured light emitted. However, if you use krypton instead of neon, you get a whitish light, while using argon yields a bluepurple light. A light filled with nitrogen gas glows purple, as does a helium lamp. Other gases — and mixtures of gases — emit other colours of light. Ironically, despite its importance in the development of modern electronic theory, hydrogen lamps emit little visible light and are rarely used for illumination purposes.
Key Takeaways
 Electrons in atoms have quantized energies.
 The state of electrons in atoms is described by four quantum numbers.
Exercises
Questions
 Differentiate between a continuous spectrum and a line spectrum.
 Under what circumstances is a continuous spectrum formed? Under what circumstances is a line spectrum formed?
 What is the wavelength of light from the hydrogen atom spectrum when n = 3?
 What is the wavelength of light from the hydrogen atom spectrum when n = 5?
 What are the restrictions on the principal quantum number?
 What are the restrictions on the angular momentum quantum number?
 What are the restrictions on the magnetic quantum number?
 What are the restrictions on the spin quantum number?
 What are the possible values for ℓ when n = 5?
 What are the possible values for ℓ when n = 1?
 What are the possible values for m_{ℓ} when ℓ = 3?
 What are the possible values for m_{ℓ} when ℓ = 6?
 Describe the shape of an s orbital.
 Describe the shape of a p orbital.
 Which of these sets of quantum numbers is allowed? If it is not, explain why.
 {4, 1, −2, +½}
 {2, 0, 0, −½}
 Which of these sets of quantum numbers is allowed? If it is not, explain why.
 {5, 2, −1, −½}
 {3, −1, −1, −½}
Answers
 A continuous spectrum is a range of light frequencies or wavelengths; a line spectrum shows only certain frequencies or wavelengths.
 6.56 × 10^{−7} m, or 656 nm
 The principal quantum number is restricted to being a positive whole number.
 The absolute value of m_{ℓ} must be less than or equal to ℓ: m_{ℓ} ≤ ℓ.
 ℓ can be 0, 1, 2, 3, or 4.
 m_{ℓ} can be −3, −2, −1, 0, 1, 2, or 3.
 An s orbital is spherical in shape.

 Because m_{ℓ} must be less than ℓ, this set of quantum numbers is not allowed.
 Allowed.
Media Attributions
 “Prisms and Light” by David W. Ball © CC BYNCSA (AttributionNonCommercialShareAlike)
 “Hydrogen Spectrum” by David W. Ball © CC BYNCSA (AttributionNonCommercialShareAlike)
 “Bohr’s Model of the Hydrogen Atom” by David W. Ball © CC BYNCSA (AttributionNonCommercialShareAlike)
 “Electron Orbitals” by David W. Ball © CC BYNCSA (AttributionNonCommercialShareAlike)
 “Neon Internet Cafe” by Justin Cormack © CC BYSA (AttributionShareAlike)
An image that contains all colours of light.
An index that corresponds to a property of an electron, like its energy.
When a quantity is restricted to having only certain values.
The theory of electrons that treats them as a wave.
The index that largely determines the energy of an electron in an atom.
A term used to describe electrons with the same principal quantum number.
An index that affects the energy and the spatial distribution of an electron in an atom.
A term used to describe electrons in a shell that have the same angular momentum quantum number.
The index that determines the orientation of the electron’s spatial distribution.
The specific set of principal, angular momentum, and magnetic quantum numbers for an electron.
The index that indicates one of two spin states for an electron.