Who Calculated the Charge on an Electron?
Use this calculator to estimate electric charge with the classic oil-drop force-balance formula and compare your result with the accepted elementary charge. The historical answer is Robert A. Millikan, whose oil-drop experiment provided the first precise direct measurement of the electron’s charge.
Electron Charge Calculator
For a suspended oil drop in a uniform electric field, the electric force can balance the droplet’s weight. A simplified estimate uses q = mgd / V, where m is droplet mass, g is gravitational acceleration, d is plate separation, and V is applied voltage.
Results and Chart
Who calculated the charge on an electron?
The scientist most commonly credited with calculating, or more precisely measuring directly, the charge on an electron is Robert Andrews Millikan. His famous oil-drop experiment, performed in the early twentieth century and refined over several years, demonstrated that electric charge occurs in discrete units and allowed him to estimate the size of the smallest unit of charge carried by a single electron. Today that quantity is called the elementary charge, written as e, and its exact SI value is 1.602176634 × 10^-19 coulomb.
That answer, however, becomes richer when you look at the broader history of physics. Millikan did not work in isolation from prior discoveries. Michael Faraday showed through electrolysis that electricity is related to matter in a regular and measurable way. J. J. Thomson then discovered the electron in 1897 and measured its charge-to-mass ratio, proving that cathode rays were made of tiny negatively charged particles. Millikan’s contribution was different: he supplied the direct measurement of the particle’s charge itself. Once physicists had both Thomson’s charge-to-mass ratio and Millikan’s charge, they could determine the electron’s mass as well.
So if you are answering a quiz, article title, or common search question, the short answer is clear: Robert A. Millikan calculated the charge on an electron. If you are answering with full scientific accuracy, the fuller statement is that Millikan directly measured the elementary charge, building on foundational work by Faraday and Thomson.
Why Millikan’s experiment mattered so much
Before Millikan, scientists had strong evidence that matter and electricity were linked, but they did not yet possess a clean, direct number for the charge on one electron. Millikan’s experiment changed that. He used tiny charged oil drops placed between electrically charged plates. By adjusting the voltage across the plates, he created an electric field that could oppose gravity. When the upward electric force matched the downward gravitational force, a droplet could hover almost motionless.
In a simplified form, the balancing idea is straightforward:
Electric force = qE and weight = mg
At balance, qE = mg. Since the electric field E = V/d, this becomes q = mgd / V.
That relation is the basis of the calculator above. The real experiment included corrections for buoyancy, air viscosity, and droplet motion, but the fundamental concept is exactly the same. Millikan repeatedly measured droplets carrying different amounts of charge and found that all of those charges were integer multiples of one smallest value. That smallest repeating unit was the charge of the electron.
- It gave physics a direct numerical value for a fundamental constant.
- It provided persuasive evidence that electric charge is quantized.
- It allowed scientists to combine charge data with Thomson’s charge-to-mass ratio and calculate electron mass.
- It helped establish modern atomic and quantum-era physics on a firmer experimental foundation.
What did Faraday and Thomson contribute?
Michael Faraday’s role
Faraday did not directly measure the charge on a single electron, because the electron had not yet been isolated as a known particle during his earliest work. What he did show was that chemical change in electrolysis occurs in fixed proportional relationships to the quantity of electric charge that passes through a substance. This suggested that electricity is not infinitely divisible in practice but tied to regular atomic-scale processes.
Faraday’s laws of electrolysis introduced a key measurable constant: the Faraday constant, now known as approximately 96485.33212 C/mol. When combined with the modern exact Avogadro constant, the Faraday constant implies the modern elementary charge. Historically, though, that was an indirect route. Faraday set the stage; he did not perform the direct single-electron measurement for which Millikan is known.
J. J. Thomson’s role
Thomson’s 1897 experiments with cathode rays revealed that these rays consisted of negatively charged particles, later called electrons. His major quantitative result was the charge-to-mass ratio of the electron, commonly written as e/m. The modern accepted value is about 1.75882001076 × 10^11 C/kg. Thomson showed that this ratio was much larger than that of known ions, implying that the particles were extraordinarily light.
But Thomson did not yet know either the exact charge or the exact mass separately. Millikan’s later work supplied the missing piece. Once physicists had the value of e, they could rearrange Thomson’s ratio and determine the electron’s mass.
Historical comparison table
| Scientist | Approximate period | Key contribution | Relevant statistic | Why it matters |
|---|---|---|---|---|
| Michael Faraday | 1830s | Established laws of electrolysis | Faraday constant: 96485.33212 C/mol | Showed a regular link between matter and electric charge |
| J. J. Thomson | 1897 | Measured the electron charge-to-mass ratio | Modern accepted e/m: 1.75882001076 × 10^11 C/kg | Proved electrons are tiny charged particles |
| Robert A. Millikan | 1909 to 1913 | Directly measured the elementary charge | Accepted e: 1.602176634 × 10^-19 C | Confirmed charge quantization and enabled electron mass calculation |
| Modern SI definition | 2019 onward | Fixed the elementary charge exactly in the SI | e = 1.602176634 × 10^-19 C exactly | Makes the coulomb and ampere part of a constant-based measurement system |
How the oil-drop experiment worked in practice
Millikan atomized oil into tiny droplets and observed them through a microscope as they moved between metal plates. Some droplets picked up electric charge through friction or ionization. By changing the voltage between the plates, he could vary the electric field and therefore the upward or downward electric force acting on each droplet.
- A tiny oil droplet was introduced into the chamber.
- The droplet’s motion without the field gave information about its size and effective mass.
- An electric field was applied between plates separated by a known distance.
- The voltage was adjusted until the drop nearly hovered.
- At or near equilibrium, the electric force balanced gravity.
- Repeating the experiment across many droplets showed that measured charges came in discrete multiples of one basic amount.
The beauty of the experiment lies not merely in one measurement, but in the pattern across many measurements. If charges had been continuous, the values would have been spread smoothly over all possible numbers. Instead, they clustered around integer multiples of a fundamental unit. That was powerful evidence that charge is quantized.
In classroom presentations, the formula is often written as q = mgd / V. In a more complete laboratory treatment, corrections account for the viscosity of air, buoyant force, and details of the droplet’s terminal velocity. Those refinements are important for precision, yet the core physical insight remains the same: electric force can be tuned until it balances weight, and that balance reveals charge.
Measured value versus modern accepted value
Millikan’s historical values were impressively close to the modern result, especially given the experimental tools of the era. The accepted value today is exact because the SI now defines the elementary charge as a fixed constant. Historically, published early twentieth century values were near 1.59 × 10^-19 C, differing from today’s accepted value by well under 1 percent. That level of agreement was extraordinary for the time and cemented the experiment’s place in physics.
| Quantity | Representative value | Notes |
|---|---|---|
| Millikan-era measured elementary charge | About 1.59 × 10^-19 C | Historical values varied by publication year and corrections |
| Accepted elementary charge today | 1.602176634 × 10^-19 C | Exact SI value |
| Approximate difference | Roughly 0.7% or less | Shows how successful the oil-drop experiment was |
| Faraday constant | 96485.33212 C/mol | Charge per mole of electrons |
| Avogadro constant | 6.02214076 × 10^23 mol^-1 | Exact SI value |
A useful cross-check is this: dividing the Faraday constant by the Avogadro constant gives the elementary charge. That connection links chemistry, electricity, and atomic physics into one elegant quantitative framework.
Why this question still appears in classes and search engines
The question “who calculated the charge on an electron” stays popular because it sits at the intersection of history, physics, and chemistry. Students often learn the milestones separately and then need to sort out who did what. The confusion usually comes from mixing together three different achievements:
- Faraday linked electric charge to chemical change.
- Thomson discovered the electron and measured its charge-to-mass ratio.
- Millikan directly measured the elementary charge of the electron.
Once these roles are separated, the historical picture becomes much clearer. If the task is to name the person associated with the value of the electron’s charge itself, Millikan is the correct answer. If the task is to explain the full scientific path to that answer, then Faraday and Thomson must also be included.
How to interpret the calculator above
The calculator on this page uses the simplified force-balance model. You enter a droplet mass, a plate separation, and an applied voltage. The calculator computes the total charge required to suspend that droplet. If you also enter the number of excess electrons on the droplet, the tool estimates the implied charge per electron and compares it with the accepted value.
What the outputs mean
- Total charge on droplet: the complete charge needed to balance the droplet against gravity.
- Charge per electron: the total charge divided by the number of excess electrons you assume are present.
- Percent error: how far your implied electron charge is from the accepted value.
- Historical interpretation: a reminder that Robert A. Millikan is the scientist credited with this direct measurement.
If your chosen mass, distance, voltage, and electron count combine to produce a result near 1.602176634 × 10^-19 C, you have created a plausible educational example of charge quantization. If the result is far away, the calculator still teaches an important lesson: only certain parameter combinations are physically consistent with a droplet carrying an integer number of electrons.
Common misconceptions
Misconception 1: Millikan discovered the electron
No. J. J. Thomson is credited with discovering the electron. Millikan measured its charge later.
Misconception 2: Faraday measured the electron’s charge directly
No. Faraday’s work was foundational and indirectly related through electrochemistry, but not a direct single-electron measurement.
Misconception 3: The oil-drop experiment only gave one number
Not quite. Its deeper importance was showing that droplet charges came in discrete multiples of a common unit, strongly supporting charge quantization.
Misconception 4: The accepted value has always been experimental
Historically yes, but in the modern SI the elementary charge is fixed exactly by definition, not just measured approximately.
Authoritative sources for further reading
Final answer
If you need the direct answer in one sentence: Robert A. Millikan calculated the charge on an electron through the oil-drop experiment. If you need the expert answer: Faraday laid the groundwork, Thomson discovered the electron and measured its charge-to-mass ratio, and Millikan directly measured the elementary charge itself.