Partial Pressure Calculator

Find one gas’ partial pressure from Dalton, ideal gas or Henry law inputs.

Twelve pressure units, 21 volume units, nine concentration scales and eight tabulated gases.

Updated October 17, 2026

CreatorFrank Zhao

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Introduction

In a mixture, every gas carries its own share of the pressure: that share is the gas's partial pressure. It is the number that decides how much of a gas dissolves in a liquid, how much a sealed vessel holds, and how much of each reactant a gas-phase equilibrium actually sees. This page works out that share from four different starting points, because the data people have available differs from case to case.

Two of the routes start from the mixture itself. Dalton's law takes the total pressure and the gas's mole fraction, and the ideal gas law takes the amount of gas, the temperature and the volume it occupies. The other two start from a liquid: Henry's law links the pressure above a solution to what is dissolved in it, either as a concentration in moles per litre or as a mole fraction. The Henry blocks carry eight tabulated gases — oxygen, hydrogen, carbon dioxide, dinitrogen, helium, neon, argon and carbon monoxide — with a Custom entry for any other gas, solvent or temperature.

Four independent routes

Dalton, ideal gas, and Henry with a concentration or a mole fraction

Twelve pressure units

Pascals through to inches of mercury, on every pressure box

Eight tabulated gases

Henry constants in water at 298 K, plus Custom

The four blocks do not share their numbers

Each block solves its own relation, so typing in one block never moves a box in another. That is deliberate: the four routes describe four different physical situations, and the two Mole fraction boxes are not the same quantity at all. The one in Dalton's law is the composition of a gas mixture; the one in Henry's law is the composition of a liquid solution. Air at sea level has an oxygen mole fraction near 0.21, while the oxygen dissolved in that water sits near $2 \times 10^{-5}$ — same gas, different phase, different fraction. Entering a number in one of those boxes leaves the other alone, and that is what you want.

Quick start

One block answers one question. Pick the block that matches the data you already have, fill in what you know, and read the pressure. Anything the calculator worked out is shown as a computed value; typing over it hands that direction back to you and the calculator solves for a different box instead.

Choose the block

Composition and total pressure → Dalton's law. Amount of gas, temperature and volume → Ideal gas law. Dissolved concentration → Henry's law, method 1. Dissolved mole fraction → Henry's law, method 2. The next section compares them side by side.

Fill in the boxes you already know

Units are yours to pick: 12 pressure units, 21 volume units, three temperature scales and nine concentration scales are available, and changing a menu only rewrites the same number in another unit. If you know the moles of each gas rather than the fraction, the mole fraction calculator turns a parts list into the ratio this box wants.

Pick the gas in a Henry block

The Gas menu loads that gas's tabulated constant at 298 K and holds it fixed. Choose Custom when your gas, solvent or temperature is not in the table, and type the constant yourself.

Read the partial pressure

The result appears in whichever pressure unit that box is set to. If the number is flagged, read the message first — it names the assumption that is being stretched, and the number stays on screen so you can judge it.

A complete run: the oxygen share of air

Take dry air at sea level, treating oxygen as 21 % of the moles, and a total pressure of 101.325 kPa — one standard atmosphere.

p_{\mathrm{O_2}}

=

P_{\mathrm{total}} \times \chi_{\mathrm{O_2}}

=

101.325\ \mathrm{kPa} \times 0.21

=

21.27825\ \mathrm{kPa}

The Partial pressure box prints 21.27825 kPa, and switching that box to atmospheres re-reads the same value as 0.21 atm. Both readings say the same thing: oxygen carries 21 % of the pressure, because a mole fraction transfers straight across into a pressure fraction. Compare it with a measured oxygen partial pressure and the difference is a real difference in composition, not a units problem.

Which relation fits your data

The four relations overlap in what they can produce but not in what they need. Match the left column to your data and the choice is usually immediate.

What you already know	Use this block	What it needs from you
The total pressure of the mixture and the composition of the gas you care about	Dalton's law	Two numbers: the total pressure and one mole fraction between 0 and 1. No temperature or volume is involved.
The amount of the gas, the temperature, and the volume it occupies	Ideal gas law	Moles, temperature and volume of that one gas. The composition of the mixture never enters.
How much of the gas is dissolved, as an amount per litre of solution	Henry's law - method 1	A concentration in mol/L (or a submultiple) and the constant in litre·atm/mol. The Gas menu supplies the constant for eight gases.
How much of the gas is dissolved, as a mole fraction of the solution	Henry's law - method 2	A dissolved mole fraction and the constant expressed as a pressure. The Gas menu supplies that form too.

Two data checks worth making before you type

A mole fraction is only meaningful against a whole. In Dalton's law it is the moles of that gas divided by the total moles of the gas mixture; in Henry's method 2 it is the moles of dissolved gas divided by the total moles of the solution, water included. If your ratio was built on a different basis, convert it before entering it.
A partial pressure belongs to one gas. In the ideal gas law block, enter the moles of the single gas you are asking about. The pressure that comes out is what that amount of gas would exert in that volume; it is not the mixture total unless you entered the mixture total.

What none of the four blocks does

They are equilibrium relations, not reaction models. They will not account for a gas that reacts with the liquid it dissolves in, for a solution that has not finished equilibrating, or for a gas mixture so dense that the molecules attract each other appreciably. The limits section spells out where each one starts to bend.

How each result is calculated

Four relations sit behind the four blocks. Each one is solved for whichever box you leave empty, so every relation runs in both directions.

Dalton's law — a share of the total pressure

p_i = P_{\mathrm{total}} \times \chi_i

The reverse forms are the same statement rearranged, and the block provides both:

P_{\mathrm{total}}

=

\frac{p_i}{\chi_i}

\qquad

\chi_i

=

\frac{p_i}{P_{\mathrm{total}}}

Ideal gas law — a pressure from how much gas is there

p_i = \frac{n_i R T}{V}

Solving for the other three boxes turns the same equation around:

n_i = \frac{p_i V}{R T}

\qquad

T = \frac{p_i V}{n_i R}

\qquad

V = \frac{n_i R T}{p_i}

The temperature box is a real temperature, not a difference, so it converts into kelvin before it reaches the equation: $T_{\mathrm{K}} = T_{\mathrm{{}^\circ C}} + 273.15$ and $T_{\mathrm{K}} = \tfrac{5}{9}\left(T_{\mathrm{{}^\circ F}} - 32\right) + 273.15$ . If you have a mass of gas rather than moles, the grams to moles calculator converts it with the substance's molar mass.

Henry's law, method 1 — from a dissolved concentration

p_i

=

K_{\mathrm{H1}} \times c_i \times 101325\ \mathrm{Pa\,atm^{-1}}

\qquad

c_i

=

\frac{p_i}{K_{\mathrm{H1}} \times 101325}

The constant is tabulated in litre·atm per mole, so $K_{\mathrm{H1}} \times c_i$ is already in atmospheres; the 101,325 is only the step from atmospheres into pascals, the unit the pressure boxes share, so it changes the scale of the answer and nothing else. That is worth keeping in mind, because it means a constant near 769 and a concentration near 1 mmol/L are the ingredients of an answer near one atmosphere — not of something enormous. If your dissolved amount arrives as a mass per volume, the molarity calculator converts it into the mol/L this box expects.

Henry's law, method 2 — from a dissolved mole fraction

p_i = K_{\mathrm{H2}} \times \chi_i

Here the constant itself is a pressure, so no conversion factor appears. Method 1 and method 2 are two tabulations of the same physical tendency, which is why the two constants differ so much in size: dividing one by the other gives roughly the molar concentration of the solvent (about 55 mol/L for water). A mole fraction of a dissolved gas is a small number, so its constant has to be a large one for the product to land at a sensible pressure.

The numbers that are fixed rather than measured

R

8.31446261815324 J mol⁻¹ K⁻¹ — the molar gas constant, exact by definition

N_{\mathrm{A}}

6.02214076 × 10²³ mol⁻¹, one of the two constants R is built from

k_{\mathrm{B}}

1.380649 × 10⁻²³ J K⁻¹, the other one

101325\ \mathrm{Pa}

one standard atmosphere, the factor between atm and the pascal base

K_{\mathrm{H1}}

298 K constants in litre·atm/mol for the eight tabulated gases

K_{\mathrm{H2}}

the same eight gases, expressed as a pressure for a mole fraction

$R$ is not a rounded constant: it is the exact product of the Avogadro constant and the Boltzmann constant, and it terminates at exactly those fifteen significant figures. Both defining constants, and their exact product, are published by NIST as the molar gas constant and the Boltzmann constant. The two Henry constants are a different kind of number: handbook values for water at 298 K, not measurements of your solution, which is exactly what the Custom entry is for.

Where the page pushes back

Two low-pressure limits. A Henry's law result that climbs past about 1000 hPa, and an ideal-gas result past about 100 MPa, are flagged in red. Both are finite answers — they are just answers from outside the range where the relation says anything reliable. The number stays visible, with the reason next to it.
An impossible composition. A zero mole fraction together with a non-zero partial pressure cannot be satisfied by any total pressure, so the total pressure box explains that instead of inventing a number.
A genuinely undefined result. If a division by zero appears — solving for the amount of moles at 0 K, for example — that box goes blank and says which denominator was zero. A blank here means the arithmetic has no answer, not that the answer is zero.

Worked examples

Three short cases, each one a typical reason someone opens this page. The numbers are the calculator's own output for the inputs shown.

A composition check: oxygen in air

A gas analyst has a mixture at 101.325 kPa in which oxygen is 21 % of the moles and wants to know what an oxygen sensor should read. Dalton's law:

p_{\mathrm{O_2}}

=

101.325\ \mathrm{kPa}

\times

0.21

=

21.27825\ \mathrm{kPa}

Reading the same box in atmospheres gives 0.21 atm, and in millimetres of mercury 159.6 mmHg (approximately). If the sensor reads noticeably below that, the mixture is not what the analysis says it is — oxygen was consumed, or air leaked in.

A sealed flask: how much pressure is in there?

Half a mole of a gas sits in a 12 L flask held at 25 °C. The ideal gas law block:

p = \frac{nRT}{V}

=

0.5 \times 8.31446261815324

\times

298.15

\div

0.012

\approx

103.2899\ \mathrm{kPa}

That is 1.019392 atm. Note how the temperature was entered as 25 °C and multiplied as 298.15 K — the scale conversion happens before the relation, so you never have to convert it by hand. Rearranged, the same block answers the practical version of the question: hold the pressure you measured and it returns the moles in the flask, or the volume they must occupy.

The same dissolved gas through both Henry routes

Water at 25 °C in equilibrium with oxygen holds about 1.3 mmol of O₂ per litre, using the tabulated oxygen constant. Method 1 returns:

p_{\mathrm{O_2}}

=

769.23

\times

0.0013

\times

101325

\approx

101.3249\ \mathrm{kPa}

just under one atmosphere, which is the promise of the example: a solution left in contact with roughly one atmosphere of oxygen reaches roughly this concentration. Method 2 should agree, but it needs the dissolved mole fraction rather than the concentration. The matching fraction follows from the two constants:

\chi_{\mathrm{O_2}} = \frac{K_{\mathrm{H1}} \times c}{K_{\mathrm{H2}}}

=

\frac{769.23 \times 0.0013}{42590}

\approx

2.348 \times 10^{-5}

Typing 2.348 × 10⁻⁵ into method 2 returns 101.3263 kPa — the same answer to within 0.002 kPa, from a completely different pair of numbers. That is the useful lesson about the two Henry blocks: the constants look wildly different (769.23 against 42,590) because they pair with quantities that differ by the same factor, the molar concentration of water. Feeding each block the matching input is what makes them agree.

Where this usually leads next

Partial pressures are the raw material of gas-phase equilibrium. If the mixture you are describing is a reacting one, the Kp calculator takes the partial pressures of the species and turns them into the equilibrium constant.

Traps and good habits

Do not expect the two mole fraction boxes to agree

They belong to different phases, so they will normally differ by orders of magnitude. The box in Dalton's law describes the gas mixture; the box in method 2 describes the liquid solution. Nothing is broken when the two numbers look unrelated.

Check whether your concentration is per litre of solution

The concentration box wants an amount per litre of solution — molarity. A molality (per kilogram of solvent) is a different quantity, and the two only coincide for very dilute aqueous solutions near room temperature. For a concentrated solution, convert before typing.

Enter the temperature of the equilibrium

For Henry's law that is the temperature of the liquid and the gas sitting above it. For the ideal gas law it is the temperature of the gas itself. Both blocks accept °C, °F and K, and the Celsius conversion is not an approximation of the kelvin value.

A large Henry constant is not an error

Values in the tens of thousands of atmospheres belong to method 2, where the partner quantity is a dissolved mole fraction typically near 10⁻⁵. Method 1's constants sit in the hundreds because they pair with millimoles per litre. Match the constant to the input you have.

Type the number in the unit you measured it in

Every pressure box can be switched afterwards, and the conversion never changes the underlying value or the result. Working in the measured unit, then switching to the unit your report needs, keeps transcription errors out of the chain.

Custom is the honest option for real samples

The eight tabulated constants describe those gases in water at 298 K. Seawater, a solvent, a cold tank or a different gas all need your own constant — measured, or from the source you trust — entered through Custom. The tabulated row is a starting point, not a substitute.

A warning is information, not a failure

When a result is flagged, the calculator is telling you that the relation has been pushed past the range where it is usually trusted, and it leaves the number there so you can still judge it. Read which boundary was crossed, decide whether your numbers or your expectations are off, and only then act on the result.

Questions

Does typing in one block change the other three?

No. Each block solves its own relation, and the four relations share no boxes. That is intentional: a gas mixture, a sealed volume and a solution are different physical situations, and coupling them would let an unrelated input quietly rewrite a result you were reading. If you want the same gas described from a second starting point, fill the second block in yourself — the two must agree when you use matching values.

Why did a very high ideal-gas pressure get flagged?

Because the ideal gas law treats molecules as points that do not attract one another, and that picture stops holding at high density. Past roughly 100 MPa a real gas occupies a different pressure than the equation predicts, and the gap grows quickly from there. The number is kept on screen because it is what the equation says; the flag is there so you do not read it as a measurement of a real vessel.

What does a zero partial pressure mean, and why is a negative one rejected?

A zero means the gas is absent from the mixture: enter a mole fraction of 0 and Dalton's law returns 0 kPa — a legitimate answer, shown like any other. A negative pressure has no physical meaning, and neither does a mole fraction outside 0 to 1, so those boxes turn red and ask you to check the input rather than carrying a sign error into the rest of the calculation.

How far to trust it

The ideal gas law is a low-density approximation. Its molecules have no volume and no attraction. It is excellent for air and for laboratory gases near room temperature and ordinary pressure, and it drifts badly as pressure and density rise — the page flags results above about 100 MPa for that reason.
Henry's law is a dilute, low-pressure limit. The proportionality between pressure and dissolved amount is a straight line only while the solution is dilute and the pressure is modest — the page warns above about 1000 hPa. The tabulated constants are 298 K values for water; another solvent or temperature needs your own constant.
Gases that react with the solvent break the linear link. Carbon dioxide and ammonia in water are the familiar cases: part of what dissolves is converted into another species, so the total dissolved amount no longer tracks the gas pressure by a single constant.
Every block assumes equilibrium and a consistent basis. Henry's law describes a solution that has finished equilibrating with the gas above it, and none of the blocks checks that your mole fraction, total pressure, volume or concentration describe the same mixture and the same moment.
Do not use it for breathing-gas or clinical decisions. Partial pressures of oxygen and other gases govern oxygen toxicity, narcosis and blood-gas interpretation, and those decisions depend on depth, exposure time, equipment and individual physiology that no single relation captures. This page is an educational and estimation tool; breathing-gas planning and clinical interpretation require qualified supervision and validated instruments.

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