From Tro's Chemistry: Structure and Properties [1, p. 93]:

2.5 Quantum Mechanics and the Atom

As we have seen, the position and velocity of the electron are complementary properties—if we know one accurately, the other becomes indeterminate. Since velocity is directly related to energy (recall that kinetic energy equals $\frac 1 2 mv^2),$ position and energy are also complementary properties—the more we know about one, the less we know about the other. Many of the properties of an element, however, depend on the energies of its electrons. In the following paragraphs, we describe the probability distribution maps for electron states in which the electron has well-defined energy, but not well-defined position. In other words, for each of these states, we can specify the energy of the electron precisely, but not its location at a given instant. Instead, the electron’s position is described in terms of an orbital, a probability distribution map showing where the electron is likely to be found. Since chemical bonding often involves the sharing of electrons between atoms (see Section 4.2), the spatial distribution of atomic electrons is important to bonding.

The textbook says that electrons can exist in two states simultaneously. As I went further, I saw that an electron could exist as both a particle and a wave. So I thought that this must be the states that I was introduced to earlier, until I was learning about quantum mechanics.

In the text above, the writer was talking about electron states with well defined energy. From what I know the electron state that has well defined energy would be one with well defined velocity. And the velocity of an electron is associated with it's wave nature.

So, aren't we supposed to be dealing with electron as waves only? Or what exactly does the writer mean by the electron states?


  1. Tro, N. J. Chemistry: Structure and Properties, 2nd ed.; Pearson: Hoboken, NJ, 2018. ISBN 978-0-13-429393-6.
  • 1
    $\begingroup$ Keep in mind people are not very careful in using the terminology. You will find this discussion useful chemistry.stackexchange.com/questions/119746/… $\endgroup$
    – AChem
    Aug 22, 2020 at 16:52
  • $\begingroup$ How do you conclude that a defined energy means a defined velocity? That only applies if there is no potential energy term in the Hamiltonian $\endgroup$
    – Andrew
    Aug 22, 2020 at 18:07
  • $\begingroup$ Kinetic energy=1/2mv2 $\endgroup$
    – Taofeek
    Aug 22, 2020 at 18:46
  • 3
    $\begingroup$ Although this question can be easily answered, I'm upvoting because it is a key concept for beginners. $\endgroup$ Aug 22, 2020 at 20:40
  • 1
    $\begingroup$ An electronic state is defined by the electron configuration of the system and by the quantum numbers of each electron contributing to that configuration. Each electronic state corresponds to one of the energy levels of the molecule (source: chemistry libretexts). $\endgroup$ Aug 23, 2020 at 6:24

1 Answer 1


In QM, properties that can be measured (such as energy, spin, position, etc) are each associated with an "operator". Every observation will always return an eigenvalue for the relevant operator. Eigen values are values for $\lambda$ for which the equation $$\hat{O}\Psi = \lambda\Psi$$ is true, where $\hat{O}$ is the operator of interest and $\Psi$ is the wave function of the entity of interest.

For example, let's consider electron "spin", by which we typically mean specifically the component of the spin in the z-direction, given by the operator $\hat{S}_z$. For the electron, this operator has two possible eigenvalues, which we abbreviate as $+1/2$ and $-1/2$. An electron having one of these values of spin is described as being in the eigenstate (or more simply just "state") associated with the given eigenvalue.

So we can say that having $+1/2$ spin and having $-1/2$ spin are two possible states of the electron.

What makes QM different from classical mechanics is that measuring an electron to have $+1/2$ spin does not mean that it was necessarily in the $+1/2$ spin state prior to our taking the measurement. An electron can (and often does) exist in a super-position of states, meaning that it simultaneously has both $+1/2$ and $-1/2$ spin, but our measurement causes it to collapse into one state or the other.

As you described in your question, some pairs of operators have the property that you cannot measure a defined value for both at the same time. Position and momentum are the most famous pair, since they are the example for the Heisenberg uncertainty principle, but many other pairs behave similarly. Energy and position are one such pair. So if we measure energy precisely, we cannot know where the electron with that energy exists. Since energy levels are the basis of the Schrodinger equation and define much of the behavior of atoms, chemists tend to treat electrons as if they are in a state in which the energy has been measured and other properties are unknown.

To address your point of confusion - "wave" and "particle" are descriptions of the behavior of electrons, most useful when understanding the concept of positional uncertainty, since "position" is most associated with particle behavior. When we learn about the Heisenberg uncertainty, we learn about experiments in which electrons behave as waves when their position has not been measured and like particles when their position has been measured, so it's easy to see why you associate wave and particle behavior with velocity and position, respectively, but those are just convenient examples to show the concept of duality. Operators generally should not be thought of as being associated with one description or the other, nor should we think of an electron as switching between particle and wave. It is better just to think of it as having specific position in space or not, and whether or not it has specific position it is neither a wave nor a particle, but rather just a quantum mechanical entity.

The results of any experiments we perform with electrons are predicted very well by the mathematical system that has been developed for quantum mechanics, and some of that math has parallels with classical wave behavior and some with classical particle behavior and some with neither of them. Beyond that, we have very little understanding of what an electron actually is and how its behavior accords with our understanding of the physics of the macroscopic world.

ADDENDUM For completeness, I'll add that in addition to $\hat{S}_z$, two other important operators that measure properties of electrons that can be known at the same time as energy is known are $\hat{L}^2$, the operator for the magnitude of total (quantum) angular momentum and $\hat{L}_z$, the z-component of that angular momentum. Thus, we can theoretically have an electron for which all four properties are fixed at one time, and we define the state described by the four eigenvalues by representing those eigenvalues with the "quantum numbers" (simplified forms of the eigenvalues) with which you are probably familiar:

n is the quantum number which represents the eigenvalues of the energy operator $\hat{H}$ (ie the Hamiltonian, which we associate with a specific shell)

l is the quantum number which represents the eigenvalues of the operator $\hat{L}^2$ (which we equate with s,p, d etc orbitals)

ml is the quantum number which represents the eigenvalues of the operator $\hat{L}_z$ (which we equate with the axial orientation of orbitals)

ms is the quantum number which represents the eigenvalues of the operator $\hat{S}_z$ (which we equate with the above-described spin orientation)


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