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Introductory Chemistry - 1st Canadian Edition: Quantum Numbers for Electrons

Introductory Chemistry - 1st Canadian Edition
Quantum Numbers for Electrons
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table of contents
  1. Cover
  2. Title Page
  3. Copyright
  4. Table Of Contents
  5. Acknowledgments
  6. Dedication
  7. About BCcampus Open Education
  8. Chapter 1. What is Chemistry
    1. Some Basic Definitions
    2. Chemistry as a Science
  9. Chapter 2. Measurements
    1. Expressing Numbers
    2. Significant Figures
    3. Converting Units
    4. Other Units: Temperature and Density
    5. Expressing Units
    6. End-of-Chapter Material
  10. Chapter 3. Atoms, Molecules, and Ions
    1. Acids
    2. Ions and Ionic Compounds
    3. Masses of Atoms and Molecules
    4. Molecules and Chemical Nomenclature
    5. Atomic Theory
    6. End-of-Chapter Material
  11. Chapter 4. Chemical Reactions and Equations
    1. The Chemical Equation
    2. Types of Chemical Reactions: Single- and Double-Displacement Reactions
    3. Ionic Equations: A Closer Look
    4. Composition, Decomposition, and Combustion Reactions
    5. Oxidation-Reduction Reactions
    6. Neutralization Reactions
    7. End-of-Chapter Material
  12. Chapter 5. Stoichiometry and the Mole
    1. Stoichiometry
    2. The Mole
    3. Mole-Mass and Mass-Mass Calculations
    4. Limiting Reagents
    5. The Mole in Chemical Reactions
    6. Yields
    7. End-of-Chapter Material
  13. Chapter 6. Gases
    1. Pressure
    2. Gas Laws
    3. Other Gas Laws
    4. The Ideal Gas Law and Some Applications
    5. Gas Mixtures
    6. Kinetic Molecular Theory of Gases
    7. Molecular Effusion and Diffusion
    8. Real Gases
    9. End-of-Chapter Material
  14. Chapter 7. Energy and Chemistry
    1. Formation Reactions
    2. Energy
    3. Stoichiometry Calculations Using Enthalpy
    4. Enthalpy and Chemical Reactions
    5. Work and Heat
    6. Hess’s Law
    7. End-of-Chapter Material
  15. Chapter 8. Electronic Structure
    1. Light
    2. Quantum Numbers for Electrons
    3. Organization of Electrons in Atoms
    4. Electronic Structure and the Periodic Table
    5. Periodic Trends
    6. End-of-Chapter Material
  16. Chapter 9. Chemical Bonds
    1. Lewis Electron Dot Diagrams
    2. Electron Transfer: Ionic Bonds
    3. Covalent Bonds
    4. Other Aspects of Covalent Bonds
    5. Violations of the Octet Rule
    6. Molecular Shapes and Polarity
    7. Valence Bond Theory and Hybrid Orbitals
    8. Molecular Orbitals
    9. End-of-Chapter Material
  17. Chapter 10. Solids and Liquids
    1. Properties of Liquids
    2. Solids
    3. Phase Transitions: Melting, Boiling, and Subliming
    4. Intermolecular Forces
    5. End-of-Chapter Material
  18. Chapter 11. Solutions
    1. Colligative Properties of Solutions
    2. Concentrations as Conversion Factors
    3. Quantitative Units of Concentration
    4. Colligative Properties of Ionic Solutes
    5. Some Definitions
    6. Dilutions and Concentrations
    7. End-of-Chapter Material
  19. Chapter 12. Acids and Bases
    1. Acid-Base Titrations
    2. Strong and Weak Acids and Bases and Their Salts
    3. Brønsted-Lowry Acids and Bases
    4. Arrhenius Acids and Bases
    5. Autoionization of Water
    6. Buffers
    7. The pH Scale
    8. End-of-Chapter Material
  20. Chapter 13. Chemical Equilibrium
    1. Chemical Equilibrium
    2. The Equilibrium Constant
    3. Shifting Equilibria: Le Chatelier’s Principle
    4. Calculating Equilibrium Constant Values
    5. Some Special Types of Equilibria
    6. End-of-Chapter Material
  21. Chapter 14. Oxidation and Reduction
    1. Oxidation-Reduction Reactions
    2. Balancing Redox Reactions
    3. Applications of Redox Reactions: Voltaic Cells
    4. Electrolysis
    5. End-of-Chapter Material
  22. Chapter 15. Nuclear Chemistry
    1. Units of Radioactivity
    2. Uses of Radioactive Isotopes
    3. Half-Life
    4. Radioactivity
    5. Nuclear Energy
    6. End-of-Chapter Material
  23. Chapter 16. Organic Chemistry
    1. Hydrocarbons
    2. Branched Hydrocarbons
    3. Alkyl Halides and Alcohols
    4. Other Oxygen-Containing Functional Groups
    5. Other Functional Groups
    6. Polymers
    7. End-of-Chapter Material
  24. Chapter 17. Kinetics
    1. Factors that Affect the Rate of Reactions
    2. Reaction Rates
    3. Rate Laws
    4. Concentration–Time Relationships: Integrated Rate Laws
    5. Activation Energy and the Arrhenius Equation
    6. Reaction Mechanisms
    7. Catalysis
    8. End-of-Chapter Material
  25. Chapter 18. Chemical Thermodynamics
    1. Spontaneous Change
    2. Entropy and the Second Law of Thermodynamics
    3. Measuring Entropy and Entropy Changes
    4. Gibbs Free Energy
    5. Spontaneity: Free Energy and Temperature
    6. Free Energy under Nonstandard Conditions
    7. End-of-Chapter Material
  26. Appendix A: Periodic Table of the Elements
  27. Appendix B: Selected Acid Dissociation Constants at 25°C
  28. Appendix C: Solubility Constants for Compounds at 25°C
  29. Appendix D: Standard Thermodynamic Quantities for Chemical Substances at 25°C
  30. Appendix E: Standard Reduction Potentials by Value
  31. Glossary
  32. About the Authors
  33. Versioning History

Quantum Numbers for Electrons

Learning Objectives

  1. Explain what spectra are.
  2. 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 8.03 “Prisms and Light”). This image is known as a continuous spectrum. 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 8.03 “Prisms and Light”). This image is called a line spectrum. It turns out that every element has its own unique, characteristic line spectrum.

Part (a): A continuous spectrum. Part (b): A line spectrum for mercury.
Figure 8.03 “Prisms and Light.” (a) A glowing object gives off a full rainbow of colours, which are noticed only when light is passed through a prism to make a continuous spectrum. (b) However, when electricity is passed through a gas, only certain colours of light are emitted. Here are the colours of light in the line spectrum of Hg.

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 8.04 “Hydrogen Spectrum”). Late-nineteenth-century scientists found that the positions of the lines obeyed a pattern given by the following equation:

\dfrac{1}{\lambda}=(109,700\text{ cm}^{-1})\left(\dfrac{1}{4}-\dfrac{1}{n^2}\right)

Where n = 3, 4, 5, 6,…, but they could not explain why this was so.

Hydrogen spectrum. Long description needed.
Figure 8.04 “Hydrogen Spectrum.” The spectrum of hydrogen was particularly simple and could be predicted by a simple mathematical expression.

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 quantum number). Quantities that have certain specific values are called quantized. 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:

E_{\text{light}}=h\nu = \Delta E_{\text{electron}}

This means that only certain frequencies (and thus, certain wavelengths) of light are emitted. Figure 8.05 “Bohr’s Model of the Hydrogen Atom” shows a model of the hydrogen atom based on Bohr’s ideas.

Bohr's model of the hydrogen atom. Long description needed.
Figure 8.05 “Bohr’s Model of the Hydrogen Atom.” Bohr’s description of the hydrogen atom had specific orbits for the electron, which had quantized energies.

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 quantum mechanics, 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 quantum-mechanical 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 principal quantum number. 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 electron shell 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 angular momentum quantum number, 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 three-dimensional 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:

\ell=0,1,2,\dots,n-1

Thus, for a given value of n, there are different possible values of ℓ, as shown in Table 8.1.

Table 8.1 Possible Values of ℓ
If n equalsℓ can be
10
20 or 1
30, 1, or 2
40, 1, 2, or 3

and so forth. Electrons within a shell that have the same value of ℓ are said to occupy a subshell 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):

Table 8.2 Atomic Subshells as Defined by the Value of ℓ
If ℓ equalsThe subshell is
0s
1p
2d
3f

The next quantum number is called the magnetic quantum number, represented by mℓ. For any value of ℓ, there are 2ℓ + 1 possible values of mℓ, ranging from −ℓ to ℓ:

\begin{array}{c} -\ell \varleq m_\ell \varleq \ell \\ \\ \text{or} \\ \\ |m\ell | \varleq \ell \end{array}

The following explicitly lists the possible values of mℓ for the possible values of ℓ:

Table 8.3 Possible Values of mℓ
If ℓ equalsThe mℓ values can be
00
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 mℓ dictates the orientation of an electron’s distribution in space. When ℓ is zero, mℓ 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 mℓ designates a certain orbital. 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 mℓ 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 8.06 “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 two-lobed, dumbbell-like shape; because there are three of them, we normally represent them as pointing along the x-, y-, and z-axes of Cartesian space. In part (c), we see that d orbitals are four-lobed rosettes; they are oriented differently in space (the one labelled dz2 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 mℓ, each orbital is labelled with one of the possible values. It should be noted that the diagrams in Figure 8.06 are estimates of the electron distribution in space, not surfaces electrons are fixed on.

Shapes and orientations of electron orbitals.
Figure 8.06 “Electron Orbitals.” (a) The lone s orbital is spherical in distribution. (b) The three p orbitals are shaped like dumbbells, and each one points in a different direction. (c) The five d orbitals are rosette in shape, except for the dz2 orbital, which is a “dumbbell + torus” combination. They are all oriented in different directions.

The final quantum number is the spin quantum number, represented by ms. 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 8.3

Problems

Of the set of quantum numbers {n, ℓ, mℓ, ms}, which are possible and which are not allowed?

  1. {3, 2, 1, +½}
  2. {2, 2, 0, −½}
  3. {3, −1, 0, +½}

Solutions

  1. 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 mℓ 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.
  2. 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.
  3. 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, mℓ, ms}, which are possible and which are not allowed?

  1. {4, 2, −2, 1}
  2. {3, 1, 0, −½}

Answers

  1. Spin must be either +½ or −½, so this set of quantum number is not allowed.
  2. 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 so-called “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 reddish-orange light due to the large amount of red-, orange-, and yellow-coloured light emitted. However, if you use krypton instead of neon, you get a whitish light, while using argon yields a blue-purple 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.

A neon sign that says "Internet Cafe: Open 24 hours."
Figure 8.07 “Neon.” The different colours of these “neon” lights are caused by gases other than neon in the discharge tubes.

Key Takeaways

  • Electrons in atoms have quantized energies.
  • The state of electrons in atoms is described by four quantum numbers.

Exercises

Questions

  1. Differentiate between a continuous spectrum and a line spectrum.
  2. Under what circumstances is a continuous spectrum formed? Under what circumstances is a line spectrum formed?
  3. What is the wavelength of light from the hydrogen atom spectrum when n = 3?
  4. What is the wavelength of light from the hydrogen atom spectrum when n = 5?
  5. What are the restrictions on the principal quantum number?
  6. What are the restrictions on the angular momentum quantum number?
  7. What are the restrictions on the magnetic quantum number?
  8. What are the restrictions on the spin quantum number?
  9. What are the possible values for ℓ when n = 5?
  10. What are the possible values for ℓ when n = 1?
  11. What are the possible values for mℓ when ℓ = 3?
  12. What are the possible values for mℓ when ℓ = 6?
  13. Describe the shape of an s orbital.
  14. Describe the shape of a p orbital.
  15. Which of these sets of quantum numbers is allowed? If it is not, explain why.
    1. {4, 1, −2, +½}
    2. {2, 0, 0, −½}
  16. Which of these sets of quantum numbers is allowed? If it is not, explain why.
    1. {5, 2, −1, −½}
    2. {3, −1, −1, −½}

Answers

  1. A continuous spectrum is a range of light frequencies or wavelengths; a line spectrum shows only certain frequencies or wavelengths.
  1. 6.56 × 10−7 m, or 656 nm
  1. The principal quantum number is restricted to being a positive whole number.
  1. The absolute value of mℓ must be less than or equal to ℓ: |mℓ| ≤ ℓ.
  1. ℓ can be 0, 1, 2, 3, or 4.
  1. mℓ can be −3, −2, −1, 0, 1, 2, or 3.
  1. An s orbital is spherical in shape.
    1. Because |mℓ| must be less than ℓ, this set of quantum numbers is not allowed.
    2. Allowed.

Media Attributions

  • “Prisms and Light” by David W. Ball © CC BY-NC-SA (Attribution-NonCommercial-ShareAlike)
  • “Hydrogen Spectrum” by David W. Ball © CC BY-NC-SA (Attribution-NonCommercial-ShareAlike)
  • “Bohr’s Model of the Hydrogen Atom” by David W. Ball © CC BY-NC-SA (Attribution-NonCommercial-ShareAlike)
  • “Electron Orbitals” by David W. Ball © CC BY-NC-SA (Attribution-NonCommercial-ShareAlike)
  • “Neon Internet Cafe” by Justin Cormack © CC BY-SA (Attribution-ShareAlike)

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Organization of Electrons in Atoms
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Copyright © 2014

                                by Jessie A. Key

            Introductory Chemistry - 1st Canadian Edition by Jessie A. Key is licensed under a Creative Commons Attribution-NonCommercial-ShareAlike 4.0 International License, except where otherwise noted.
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