mmHg to Atmospheres

Partial Pressures of the Gases You Breathe

Respiratory physiology is written almost entirely in millimetres of mercury: oxygen at 159 in the air outside, 40 for carbon dioxide leaving the alveoli, 47 for water vapour at body temperature. Physical chemistry, gas laws and most exam questions about mixtures work in atmospheres or fractions instead. Moving between the two is how a partial pressure stops being a number to memorise and starts making sense as a share of the whole gas mixture.

Factor: 1 mmHg = 1/760 atm = 0.0013157895 atm, because one standard atmosphere is defined as exactly 760 mmHg. Oxygen in dry sea-level air exerts about 159.2 mmHg, which converts to 0.2095 atm — precisely its 20.95 % share of the mixture. That identity is Dalton's law in one line.

Reference information, not medical advice. The figures here are textbook teaching values for learning gas physiology. They are not diagnostic criteria, they do not describe any individual, and nothing on this page should guide a clinical decision. Questions about health belong with a qualified professional.

The Ideas Behind the Numbers

Dalton's law

Each gas in a mixture pushes as if it were alone. Multiply its fractional concentration by the total pressure and you have its partial pressure — the quantity that actually drives diffusion.

Humidification comes first

Air is fully saturated by the time it reaches the trachea. Water vapour claims 47 mmHg at 37 °C, leaving only 713 mmHg for everything else to share.

A cascade, not a step

Oxygen falls stagewise — atmosphere, airway, alveolus, arterial blood, tissue. Each drop has a physical cause, and teaching texts call the whole descent the oxygen cascade.

Gradients, not absolutes

Gas moves because of a difference in partial pressure, not because of concentration. That is why a value in atm and the same value in mmHg carry identical meaning.

Converting a Partial Pressure for a Physiology Problem

Most study questions give you mmHg and then ask for something in fractions, atmospheres or SI — or the other way round when the source is a chemistry text.

Put the mmHg figure in the left field

Enter the partial pressure you are working with — 100, 47, 159.2 — and the atmosphere value builds as you type. Decimal commas are understood, so notes typed on a European keyboard paste in unchanged.

Read the result as a fraction of the mixture

At sea level the atm value is the gas fraction directly: 0.2095 atm of oxygen is 20.95 % of the air. That trick only works when total pressure is one atmosphere, which is exactly why altitude problems are harder.

Carry the value into your working

The copy control on each field gives the number with no unit attached, ready for a calculation sheet or a set of revision notes. Ctrl + C from inside a field does the same job.

Reverse it when the source is in atmospheres

Press the swap control (↔) to run atm → mmHg, which multiplies by 760. A chamber described as 0.5 atm therefore holds a total pressure of 380 mmHg, and every partial pressure inside it halves too.

Water vapour does not scale with altitude. Because it depends on body temperature and not on ambient pressure, the 47 mmHg term stays put while everything else shrinks — so it takes a bigger and bigger slice of the mixture as you climb.

Gas Partial Pressures from Room Air to Blood

Standard teaching values at sea level, with a total pressure of 760 mmHg. Real figures vary between texts and between individuals; these are the reference numbers used to illustrate the cascade.

Location	Gas	Partial pressure (mmHg)	Partial pressure (atm)
Dry atmospheric air	Nitrogen	593.4 mmHg	0.7808 atm
Dry atmospheric air	Oxygen	159.2 mmHg	0.2095 atm
Humidified tracheal air	Water vapour at 37 °C	47 mmHg	0.0618 atm
Humidified tracheal air	Oxygen	149.4 mmHg	0.1966 atm
Alveolar gas	Oxygen	100 mmHg	0.1316 atm
Alveolar gas	Carbon dioxide	40 mmHg	0.0526 atm
Arterial blood	Oxygen	95 mmHg	0.1250 atm
Mixed venous blood	Oxygen	40 mmHg	0.0526 atm
Mixed venous blood	Carbon dioxide	46 mmHg	0.0605 atm

Follow the oxygen column downwards and the story is complete: 0.2095 atm outside, 0.1966 after humidification, 0.1316 in the alveolus once carbon dioxide has taken its place, and 0.0526 by the time blood returns from the tissues. Two thirds of the starting partial pressure is spent before the gas ever reaches a cell.

What Makes This Workable for Study

Step through a cascade in one sitting

Both boxes accept input, so you can enter each stage in turn and watch the atmosphere column fall without ever resetting the page.

Enough decimals for small terms

Results run to eight decimal places, which keeps the trace-gas end of the mixture — carbon dioxide in room air at roughly 0.3 mmHg — from vanishing into zero.

Switch to kPa when the syllabus does

Both unit lists carry all 26 pressure units in 8 groups, so a course that teaches partial pressures in kilopascals is one search away on the same screen.

Reverse for chemistry-style problems

The swap control turns the page into atm → mmHg, the direction you need when a gas-law question hands you a fraction of an atmosphere.

Respiratory Gas Questions

Why is alveolar oxygen near 100 mmHg when the air outside carries 159?

Two dilutions happen in sequence. Humidification adds 47 mmHg of water vapour, cutting the usable pressure from 760 to 713 and dropping inspired oxygen to about 149 mmHg. Then carbon dioxide arriving from the blood occupies roughly 40 mmHg of the alveolar mixture, and oxygen is continuously being absorbed. What is left is the familiar figure near 100 mmHg, or 0.1316 atm.

Why is water-vapour pressure always given as 47 mmHg?

Saturated vapour pressure depends on temperature alone, and core temperature is held near 37 °C, so the value is effectively a constant in these calculations — 47 mmHg, about 0.0618 atm. It is not affected by how high you are or how fast you breathe. Cool the gas and the figure falls, which is why a sample analysed at room temperature is corrected back to body conditions.

Do the alveolar gas equation's terms have to be in the same unit?

Yes — every pressure term must share one unit, and the respiratory exchange ratio has none of its own. Written in mmHg the classic worked example reads 0.21 × (760 − 47) − 40 ÷ 0.8, giving about 100 mmHg. The identical calculation in atmospheres uses 1 and 0.0618 in place of 760 and 47 and produces 0.1316 atm. Mixing a 760 with a 0.21 fraction and a kPa carbon dioxide term is the usual source of a wrong answer.

What happens to inspired oxygen partial pressure at 3 000 m?

The fraction of oxygen in the air does not change — it is still 20.95 % — but total pressure does. At 3 000 m the standard atmosphere gives about 526 mmHg, that is 0.69 atm. Subtract the unchanged 47 mmHg of water vapour and multiply by 0.2095 and inspired oxygen lands near 100 mmHg, roughly where the alveolar value sits at sea level. Textbooks use exactly this arithmetic to explain why altitude matters even though the air composition does not change.

Should the partial pressures of a mixture add up to the total?

They should, and it is a good way to check your working. Alveolar gas at sea level is roughly 100 mmHg oxygen, 40 carbon dioxide, 47 water vapour and about 573 nitrogen — 760 mmHg in total, or 1.000 atm when each term is converted. If your four numbers do not reach the total, one of them was taken from a dry-air table and never corrected for humidification.