Photoelectric Effect — Definition & Overview

Think back for a moment about what you have learned about electron energy levels and how they are associated with light. What do you think will happen if we hit an electron with highly energetic light? Take a moment and think about this on your own.

It turns out that when electrons are hit by highly energetic light they are ejected from their atoms. Let's go into more detail, starting with electron binding energy. Electrons are bound to nuclei through an electrostatic potential. Electrons have some negative potential energy, and in order to escape their atom, they must absorb at least that amount of energy. For example, in its ground state, a hydrogen electron has an energy of -13.6 eV. If the electron absorbs at least 13.6 eV, then the electrostatic potential will not be significant enough to hold on to the electron.

As a side note, you'll notice we use the unit electronvolts ($eV$). An electronvolt is a convenient measure of energies at the atomic scale. It's equivalent to $1.602\times10^{-19}$ joules. From the scientific notation, we can see that it's a very small quantity. So, it's easier to write atomic energies in electronvolts. We'll also encounter this unit when we study nuclear binding energy, mass defects, and radioactivity.

So where does light come into play? Well, light is a type of energy. Specifically, light has energy, $E$, characterized by its frequency, $\nu$, where $E=h\nu$. If we shine light of at least 13.6 eV on hydrogen, we will be able to knock off the electron. This is the principle underlying the photoelectric effect. Below is a diagram depicting this process.

(Video) The Photoelectric Effect

by Ben's Chem Videos

In this quick video, Ben teaches you about the particle behavior of light. He also talks about how Einstein's discovery of the photon changed electromagnetic theory.

Photoelectric Effect: What does it tell us?

It is crucial to note that the intensity of light is not associated with the energy that light carries. And the intensity of light does not vary the speed of emission of electrons.

In the 1900s, many scientists thought that science was reaching its limits. They'd had great success with Newtonian mechanics and classical electromagnetism. But, there were a few remaining experiments they couldn't explain. Most of these experiments were based on the absorption and emission of light. In the previous lesson, we studied Planck's quantum theory. His theory solved the problem of blackbody radiation. The classical theory predicted that a blackbody would give off unlimited ultraviolet radiation. But that disobeys the first law of thermodynamics! Planck realized that at the atomic scale, energy must be quantized. That means it can only have specific values. Ultraviolet light has a lot of energy. So, it's unlikely that a blackbody reaches high intensities of ultraviolet radiation.

In this lesson, we'll look at another seemingly unsolvable experiment. But, this time, Albert Einstein will discover the breakthrough.

Light and Metals

In the 1800s, many scientists observed that shining light on metal produced a current. That means that the light caused electrons to escape from the metal. But there was a catch! Only specific frequencies of light produced emitted electrons. Below that frequency, nothing happened. If we are above that frequency, we generate electrons! As the frequency increased, so did the electron's kinetic energy. Scientists couldn't explain what caused the threshold. Why? The classical theory of electromagnetism correlated intensity and energy. If we shined more light, we should produce electrons regardless of the frequency. Why? In classical electromagnetism, light is a wave. So, the energy should be additive. It's an application of constructive interference. If we used half the frequency for double the time, it should produce the same result. But it didn't!

The Photoelectric Effect

In 1905, Albert Einstein solved the problem. He proposed that the light was not behaving like a wave. Instead, it was acting as a particle. But, the particle's energy still depended on the frequency. We call the particle a photon. To illustrate, let's imagine an atom. The electrons circle the nucleus in their orbits. They are attracted to the nucleus by the protons. How can a photon emit an electron? To discharge an electron, we must overcome the attractive force. But, in this case, light is a particle. We can imagine it like a collision of two billiard balls. It has to collide with an energy greater than or equal to the electrostatic attraction. That explains why intensity is irrelevant. It doesn't matter how many times it gets hit. If the energy is too low, it won't overcome the electrostatic attraction.

How did Einstein come up with this model? He considered Planck's quantization of energy. Emitting an electron demands a minimum amount of energy. Some frequencies of light meet the energy requirement. But some don't, regardless of intensity. So, the energy must be a property of the light itself. That is only true for particles. As we've seen, waves depend on intensity.

Further Problems

Using quantization, scientists solved two of science's most challenging puzzles. But, quantization created new problems. We had a lot of evidence that light acted as a wave. But now, we had evidence that it behaved like a particle. How do we know which is right? And do other things behave as both particles and waves? We'll answer these questions in future explanations on the wave-particle duality and de Broglie's wavelength. These insights would lay the foundation for quantum mechanics and the quantum mechanical model of the atom.