Skip to main content

Partial Pressure Calculator

Ready to calculate
4 Standard Methods.
8 Gases Preset.
7 Pressure Units.
100% Free.
No Data Stored.

How it Works

01Pick the Method

Dalton (gas-mix), Ideal gas (n,T,V), Henry method 1 (K in L·atm/mol), or method 2 (K in atm).

02Dalton or Ideal Gas

P_i = x_i · P_total (Dalton) or P = nRT/V (ideal gas) — for known mixtures or pure conditions.

03Henry's Law (Dissolved Gases)

P = K_H · C (concentration form) or P = K_H · x (mole-fraction form) for gas-liquid equilibria.

04Get Partial Pressure

Output in your chosen unit (kPa / atm / bar / mmHg / psi / torr / Pa) with conversion to all 7.

What is a Partial Pressure Calculator?

Partial pressure is the pressure that a single component of a gas mixture would exert if it alone occupied the entire container at the same temperature. It is the central concept in Dalton's law (1801), the ideal gas law applied to mixtures, and Henry's law for gas-liquid equilibria. Our Partial Pressure Calculator implements four standard methods — each appropriate for a different physical situation. Dalton's law (P_i = x_i · P_total) computes partial pressure of a component in a gas mixture from the mole fraction. Ideal gas law (P = nRT/V) computes partial pressure of a single gas from moles, temperature, and volume. Henry's law method 1 (P = K_H · C) computes the equilibrium gas-phase pressure above a solution of known dissolved-gas concentration, with K_H in L·atm/mol. Henry's law method 2 (P = K_H · x) is the same equilibrium expressed with K_H in atm and the dissolved-gas mole fraction.

The calculator includes built-in Henry's law constants for 8 common gases at 25 °C: Oxygen, Nitrogen, Carbon dioxide, Hydrogen, Helium, Methane, Argon, and Carbon monoxide — sourced from the CRC Handbook of Chemistry and Physics. Output gives the partial pressure in 7 different pressure units simultaneously (Pa, kPa, atm, bar, mmHg, psi, torr) so you can copy whichever unit your downstream calculation, instrument, or report requires.

Designed for chemistry students learning gas laws, atmospheric scientists modelling gas exchange (CO₂ flux at the ocean surface, water-vapour saturation), respiratory physiologists analysing alveolar O₂ / CO₂, divers calculating partial pressures at depth (oxygen toxicity at > 1.4 atm; nitrogen narcosis depth limits), HVAC engineers spec'ing humidity control, and any researcher needing to convert between concentration and equilibrium gas pressure, the tool runs entirely in your browser — no account, no data stored.

Pro Tip: Pair this with our Molarity Calculator for stock preparation, our Serial Dilution Calculator for concentration series, or our Rate Constant Calculator for kinetics analysis.

How to Use the Partial Pressure Calculator?

Pick the Method: 4 collapsible sections — Dalton, Ideal gas, Henry method 1, Henry method 2. Click a section header to expand its inputs and select that method.
Dalton's Law (Gas Mixtures): Use when you know total pressure of a gas mixture and the mole fraction of the component. Inputs: P_total in any unit, mole fraction (0-1). Output: P_i = x_i · P_total.
Ideal Gas Law (P = nRT/V): Use when you know amount (moles), temperature, and container volume of a single gas. Inputs: n in mol, T in °C / K / °F, V in m³ / L / mL / cm³. R = 8.314 J/(mol·K) used internally.
Henry's Law Method 1 (P = K_H · C): Use for the equilibrium pressure above a solution of dissolved gas. Pick gas (auto-fills K_H in L·atm/mol), enter concentration in M / mM / µM. Output: gas-phase partial pressure at equilibrium.
Henry's Law Method 2 (P = K_H · x): Same physical situation as method 1, but using the mole-fraction form of Henry's constant (K_H in atm). Pick gas (auto-fills K_H), enter mole fraction of dissolved gas (typically very small for sparingly soluble gases like O₂ ~ 4×10⁻⁶ at saturation in water).
Read Result in 7 Pressure Units: Hero card shows partial pressure in the cleanest unit (Pa for tiny values, kPa for moderate, atm for typical lab conditions, bar for elevated pressure). All-unit grid shows simultaneous conversion to Pa, kPa, atm, bar, mmHg, psi, torr.

How is partial pressure calculated?

Four equations for partial pressure, each from a different chapter of physical chemistry. All assume ideal gas behaviour; Henry's law adds the dilute-solution assumption.

References: Atkins' Physical Chemistry, IUPAC Quantities Units and Symbols, CRC Handbook of Chemistry and Physics for Henry's law constants.

1. Dalton's Law of Partial Pressures (1801)

P_i = x_i · P_total

where P_i is the partial pressure of component i, x_i is its mole fraction (n_i / n_total), and P_total is the total pressure of the mixture. Equivalently P_total = Σ P_i — the total pressure equals the sum of partial pressures.

Standard example: dry atmospheric air has x_N₂ = 0.781, x_O₂ = 0.209, x_Ar = 0.0093, x_CO₂ ≈ 4.2×10⁻⁴. At sea level (P_total = 101.325 kPa): P_O₂ = 0.209 × 101.325 = 21.18 kPa.

2. Ideal Gas Law for a Single Component

P = n · R · T / V

where n is moles, R = 8.31446 J/(mol·K) is the universal gas constant, T is absolute temperature in K, V is volume in m³. Result: P in Pa.

Useful when you know amount of substance and container conditions — common in vacuum-line gas handling, cylinder pressure calculations, balloon experiments. For a gas mixture, apply this to each component individually with its own n_i to get P_i.

3. Henry's Law — Method 1 (Concentration Form)

P_gas = K_H · C_dissolved

where C is the concentration of dissolved gas in mol/L (= M), and K_H has units of L·atm/mol. P_gas comes out in atm. The Henry's law constant K_H is gas-specific and temperature-dependent.

For O₂ in water at 25 °C: K_H = 769.2 L·atm/mol. So 1 mM dissolved O₂ corresponds to gas-phase pressure 769.2 × 0.001 = 0.7692 atm = 78 kPa above the solution at equilibrium.

4. Henry's Law — Method 2 (Mole-Fraction Form)

P_gas = K_H · x_dissolved

where x is the mole fraction of dissolved gas, and K_H has units of atm (the mole-fraction form of Henry's constant). Same physical situation as method 1, but uses dimensionless mole fraction instead of molar concentration.

For O₂ in water at 25 °C: K_H = 42,590 atm. So if x_O₂ in water = 4×10⁻⁶ (typical saturation at 1 atm O₂), P_O₂ = 42,590 × 4×10⁻⁶ = 0.17 atm. Method 2 is preferred in chemical engineering and physical chemistry texts; method 1 is preferred in environmental and biological chemistry.

Henry's Law Constants at 25 °C (CRC Handbook):

  • Oxygen (O₂): K_H1 = 769.2 L·atm/mol; K_H2 = 42,590 atm.
  • Nitrogen (N₂): K_H1 = 1639 L·atm/mol; K_H2 = 86,920 atm. (Less soluble than O₂.)
  • Carbon dioxide (CO₂): K_H1 = 29.4 L·atm/mol; K_H2 = 1605 atm. (Far more soluble than O₂ — the chemistry of carbonated drinks.)
  • Hydrogen (H₂): K_H1 = 1282 L·atm/mol; K_H2 = 71,420 atm.
  • Helium (He): K_H1 = 2700 L·atm/mol; K_H2 = 144,300 atm. (Lowest solubility of common gases.)
  • Methane (CH₄): K_H1 = 714 L·atm/mol; K_H2 = 41,470 atm.
  • Argon (Ar): K_H1 = 715 L·atm/mol; K_H2 = 39,930 atm.
  • Carbon monoxide (CO): K_H1 = 1052 L·atm/mol; K_H2 = 56,440 atm.

Pressure Unit Conversions

  • 1 atm = 101,325 Pa = 101.325 kPa = 1.01325 bar = 760 mmHg = 760 torr = 14.696 psi.
  • 1 bar = 100,000 Pa = 100 kPa = 0.987 atm = 750.06 mmHg = 14.504 psi.
  • 1 mmHg = 1 torr = 133.322 Pa.
  • 1 psi = 6,894.76 Pa.

SI standard pressure unit is the pascal (Pa), but practical chemistry uses atm or kPa most commonly. Vacuum work uses torr (= mmHg) at typical 10⁻³ to 10⁻⁹ ranges. Engineering and meteorology often use bar or kPa. Tire pressure and US plumbing use psi.

Real-World Example

Partial Pressure – Worked Examples

Example 1 — Atmospheric Oxygen at Sea Level (Dalton). Total pressure 101.325 kPa, x_O₂ = 0.209.
  • P_O₂ = 0.209 × 101.325 = 21.18 kPa = 0.209 atm = 159 mmHg.
  • Equivalent to 21.18 kPa O₂ partial pressure entering alveoli — the basis of pulmonary gas-exchange physiology.
  • At altitude (lower P_total), P_O₂ scales proportionally — at 5,000 m where P_total ≈ 54 kPa, P_O₂ drops to ~11 kPa, half of sea-level value (the basis of altitude sickness).

Example 2 — Helium Tank (Ideal Gas). 0.5 mol He in a 10 L tank at 25 °C.

  • T = 25 + 273.15 = 298.15 K. V = 10 L = 0.01 m³.
  • P = nRT/V = 0.5 × 8.314 × 298.15 / 0.01 = 123,924 Pa = 124 kPa = 1.22 atm.
  • Compare to ideal-gas-law verification: at 1 atm, 25 °C, 0.5 mol fills 22.4 × 0.5 × (298/273) ≈ 12.2 L. Our 10 L gives slightly higher pressure (1.22 atm), consistent.

Example 3 — Dissolved O₂ in Water (Henry Method 1). Lake water analysis: dissolved O₂ measured at 0.27 mM (typical for cold well-aerated water).

  • K_H1 (O₂ in water at 25 °C) = 769.2 L·atm/mol.
  • P_O₂ = 769.2 × 0.27×10⁻³ = 0.208 atm = 21.0 kPa.
  • Matches the atmospheric O₂ partial pressure (Example 1) — the lake is at equilibrium with the atmosphere, as expected for an open well-aerated water body.
  • At the water surface, dissolved O₂ saturation drops with temperature (warmer water holds less O₂) — this is the basis of summer fish kills in eutrophic lakes.

Example 4 — Carbonated Beverage (Henry Method 1). Coca-Cola has dissolved CO₂ at ~0.1 M (sweet-spot for pleasant carbonation).

  • K_H1 (CO₂ in water at 25 °C) = 29.4 L·atm/mol (CO₂ is far more soluble than O₂).
  • P_CO₂ above the liquid at equilibrium = 29.4 × 0.1 = 2.94 atm = 297 kPa.
  • This is the CO₂ pressure inside the sealed bottle. Open it: P_CO₂ above the liquid drops to atmospheric ~4×10⁻⁴ atm; the dissolved CO₂ supersaturated by ~7,000-fold rapidly bubbles out — the famous fizzing.

Example 5 — Same Equilibrium via Method 2 (Henry Mole-Fraction). O₂ saturation in water at 1 atm O₂.

  • Inverse calculation: at equilibrium with P_O₂ = 1 atm, what's x_O₂ in water?
  • x_O₂ = P / K_H2 = 1 / 42,590 = 2.35 × 10⁻⁵ mole fraction (extremely small — O₂ is sparingly soluble).
  • In molar terms (using density of water 55.5 mol/L): C_O₂ = 2.35×10⁻⁵ × 55.5 = 1.30 × 10⁻³ M = 1.30 mM. Cross-check method 1: P = 769.2 × 1.30×10⁻³ = 1.00 atm ✓.
  • The two Henry methods give identical results when the conversion factor (water density 55.5 mol/L for dilute aqueous solutions) is properly applied.

Who Should Use the Partial Pressure Calculator?

1
Chemistry Students: Learning Dalton, ideal gas law, and Henry's law side-by-side; converting between the four standard partial-pressure formulations.
2
Atmospheric Scientists: Modelling P_O₂ / P_CO₂ at altitude, ocean-surface CO₂ flux, water-vapour saturation, gas exchange between atmosphere and biosphere.
3
Respiratory Physiologists: Alveolar P_O₂ / P_CO₂ calculations at sea level (~100 / 40 mmHg) vs altitude; pulmonary gas-exchange efficiency assessment.
4
Diving Medicine / Hyperbaric Engineers: Oxygen toxicity calculations (P_O₂ > 1.4 atm causes CNS toxicity); nitrogen narcosis depth limits; mixed-gas (heliox, trimix) safe-depth charts.
5
Environmental Engineers: Dissolved-O₂ saturation in lakes / rivers; CO₂ uptake from atmosphere; aeration system design for wastewater treatment.
6
Beverage Industry: Carbonation level calculations (P_CO₂ vs dissolved CO₂); bottle-pressure safety margins; nitrogen-flushed packaging headspace.
7
Brewers and Cellar Operators: Beer carbonation, sparkling-wine bottle pressures, kegging system pressure calculations.

Technical Reference

Dalton's Law — Historical and Practical. John Dalton published the law of partial pressures in 1801 in his "New System of Chemical Philosophy". The law states that gases in a mixture behave independently — each component contributes a partial pressure equal to what it would exert alone in the container. Dalton's law is exact for ideal gases and very accurate (better than 1%) for real gases at low to moderate pressures. It breaks down at high pressures where intermolecular forces become significant — for high-pressure gas mixtures, use the Lewis-Randall fugacity rule or full equation-of-state mixing rules.

Ideal Gas Law and Real-Gas Corrections. The ideal gas law PV = nRT (and its rearrangement P = nRT/V for partial pressure) assumes: (1) gas molecules occupy negligible volume; (2) intermolecular forces are negligible; (3) collisions are perfectly elastic. These approximations are excellent at low pressures (P ≤ 10 atm), high temperatures (T ≥ 100 K), and for small / non-polar gases (He, Ne, Ar, H₂, N₂, O₂, CH₄). They break down for: high-pressure compressed gases, gases near critical points (CO₂ at > 73 bar, 304 K), polar molecules (NH₃, H₂O, HCl), and at low temperatures near liquefaction. Real-gas corrections via:

  • Van der Waals (1873): (P + a·n²/V²)(V − n·b) = nRT, where a corrects for intermolecular attraction and b for molecular volume. Gives qualitatively correct PVT behaviour.
  • Redlich-Kwong (1949): empirical improvement over Van der Waals; better accuracy for many gases.
  • Peng-Robinson (1976): the modern industrial standard for petroleum / natural-gas / refinery process design; handles both gases and liquids.
  • Compressibility factor Z = PV/(nRT): the empirical correction. For ideal gas Z = 1; real gases at high P typically have Z > 1 (repulsion-dominated); at moderate P near critical, Z < 1 (attraction-dominated).

Henry's Law — Origin and Conventions. William Henry published his law in 1803: at constant temperature, the amount of a gas dissolved in a liquid is proportional to the partial pressure of the gas above the liquid. The proportionality constant K_H comes in MULTIPLE conventions that are widely confused in literature:

  • K_H,cp = C / P (concentration / pressure): solubility form; units mol/(L·atm) or M/atm. Larger K_H,cp = MORE soluble gas. Used in environmental chemistry.
  • K_H,pc = P / C = 1 / K_H,cp: volatility form; units L·atm/mol. Larger K_H,pc = LESS soluble gas. This is the calculator's "method 1" form.
  • K_H,px = P / x: mole-fraction form; units atm. Larger K_H,px = LESS soluble gas. This is the calculator's "method 2" form.
  • K_H,xp = x / P = 1 / K_H,px: reciprocal of the above; units atm⁻¹. Common in engineering.

Critical caveat: when reading published Henry's law constants, ALWAYS check which convention is used. The same gas (e.g. O₂ in water) has K_H values of 769.2 (method 1) or 42,590 (method 2) — same physics, different units. The CRC Handbook explicitly tabulates which form it uses; many papers leave it ambiguous.

Temperature Dependence of K_H — van't Hoff Equation. Henry's law constants are strongly temperature-dependent — gas solubility DECREASES with increasing temperature for most gases (an unusual entropy-driven effect). The van't Hoff form: d(ln K_H) / d(1/T) = −ΔH_sol / R, where ΔH_sol is the enthalpy of solution (typically negative — gas dissolution is exothermic). For O₂ in water: ΔH_sol ≈ −12.5 kJ/mol; warming from 25 °C to 35 °C decreases O₂ solubility by ~15%. This is why warmer water holds less dissolved O₂ — a critical issue in thermal pollution and summer fish kills.

Reference: O₂ Saturation in Water vs Temperature (at 1 atm air = P_O₂ = 0.209 atm above the surface):

  • 0 °C: 14.6 mg/L = 0.456 mM dissolved O₂.
  • 10 °C: 11.3 mg/L = 0.353 mM.
  • 20 °C: 9.1 mg/L = 0.284 mM.
  • 25 °C: 8.3 mg/L = 0.260 mM.
  • 30 °C: 7.6 mg/L = 0.236 mM.
  • 40 °C: 6.4 mg/L = 0.200 mM.

Aquatic-life thresholds: salmonids need ≥ 6 mg/L O₂; most warm-water fish need ≥ 4 mg/L; below 2 mg/L hypoxic / "dead zone" conditions develop.

Atmospheric Composition Reference (Dry Air, Sea Level):

  • N₂: 78.084% (mole fraction 0.78084) → P_N₂ = 79.1 kPa.
  • O₂: 20.946% (0.20946) → P_O₂ = 21.2 kPa.
  • Ar: 0.934% (9.34×10⁻³) → P_Ar = 0.946 kPa.
  • CO₂: 0.042% (~420 ppm in 2024, rising ~2.5 ppm/yr) → P_CO₂ = 0.043 kPa.
  • Trace: Ne, He, CH₄, Kr, H₂, N₂O, Xe — all sub-ppm.
  • Total: 101.325 kPa = 1 atm = 760 mmHg.

Diving / Hyperbaric Partial Pressure Limits:

  • P_O₂ ≥ 1.4 atm: CNS oxygen toxicity risk (the "Paul Bert effect") — convulsions, blackouts. Recreational diving max 1.4 atm; technical max 1.6 atm during decompression.
  • P_O₂ ≤ 0.16 atm: hypoxic (insufficient O₂); equivalent to altitude > 5,500 m on land. Loss of consciousness within minutes.
  • P_N₂ ≥ ~3 atm: nitrogen narcosis ("the rapture of the deep") — judgment impairment, similar to alcohol intoxication. Equivalent to depth ~30 m breathing air.
  • P_He has no narcosis effect at recreational pressures — basis of helium-oxygen (heliox) and helium-nitrogen-oxygen (trimix) breathing gases for technical / commercial diving.
  • P_CO₂ ≥ 0.05 atm: hypercapnia symptoms (headache, breathlessness); equivalent to ~5% CO₂ at 1 atm — the OSHA short-term exposure limit.

Respiratory Physiology Reference:

  • Alveolar P_O₂ at sea level: ~100 mmHg (lower than inspired 159 mmHg due to humidification, dilution by alveolar CO₂).
  • Alveolar P_CO₂ at sea level: ~40 mmHg (set by ventilation rate; the body's primary acid-base regulator).
  • Arterial P_O₂ (PaO₂): normal 80-100 mmHg; mild hypoxemia < 80; moderate < 60; severe < 40 (medical emergency).
  • Arterial P_CO₂ (PaCO₂): normal 35-45 mmHg; hypercapnia > 45 (respiratory acidosis); hypocapnia < 35 (respiratory alkalosis).
  • Mixed venous P_O₂: ~40 mmHg after tissues extract O₂.

When Henry's Law Breaks Down. Henry's law is the limiting law for very dilute solutions. It breaks down when:

  • Gas concentration is high (typically > 1 M for sparingly soluble gases, > 0.1 M for highly soluble gases) — non-ideal interactions between dissolved gas molecules.
  • Solvent is non-aqueous with strongly different polarity than water — the gas-solvent interactions can be specific (e.g. O₂ in fluorocarbons is anomalously soluble; NH₃ in water has acid-base chemistry that breaks the simple Henry framework).
  • Chemical reaction occurs: CO₂ in water partially converts to carbonic acid, bicarbonate, carbonate; the "effective" solubility differs from physical Henry's law solubility.
  • Near phase transitions: close to liquefaction or solidification of the dissolved gas, simple proportionality between concentration and pressure breaks down.

For these cases, use empirical solubility tables (NIST, IUPAC Solubility Data Series, CRC Handbook), activity-coefficient corrections, or full thermodynamic modelling (NRTL, UNIQUAC, equation-of-state methods like SRK, PR for non-aqueous systems).

Key Takeaways

Partial pressure can be computed by four standard methods, each appropriate for different physical situations. Dalton's law (1801): P_i = x_i · P_total — for known gas mixtures with mole fraction. Ideal gas law: P = n·R·T/V — for a single component with known moles, temperature, volume. Henry's law method 1: P = K_H · C (K in L·atm/mol) — gas-liquid equilibrium with concentration form. Henry's law method 2: P = K_H · x (K in atm) — same equilibrium with mole-fraction form. Pressure unit conversions: 1 atm = 101,325 Pa = 101.325 kPa = 1.01325 bar = 760 mmHg = 760 torr = 14.696 psi. Henry's law constants at 25 °C from CRC: O₂ 769.2 L·atm/mol, N₂ 1639, CO₂ 29.4 (most soluble of common gases — basis of carbonation chemistry), H₂ 1282, He 2700 (least soluble), CH₄ 714, Ar 715, CO 1052. All methods assume ideal gas behaviour; valid for low-pressure (≤ 10 atm), high-temperature (≥ 100 K) conditions for small/non-polar gases. Henry's law adds the dilute-solution assumption — valid below ~1 M dissolved-gas concentration. For non-ideal conditions use Van der Waals or Peng-Robinson EOS; for high gas-loading use activity coefficients or empirical solubility data.

Frequently Asked Questions

What is the Partial Pressure Calculator?
It implements four standard methods for computing partial pressure of a gas: (1) Dalton's law (P_i = x_i · P_total) for known gas mixtures; (2) Ideal gas law (P = nRT/V) for a single component with known moles, temperature, volume; (3) Henry's law method 1 (P = K_H · C, K in L·atm/mol) for gas-liquid equilibrium with concentration; (4) Henry's law method 2 (P = K_H · x, K in atm) for the same equilibrium with mole-fraction form. Built-in Henry's law constants for 8 common gases (O₂, N₂, CO₂, H₂, He, CH₄, Ar, CO) at 25 °C; output in 7 pressure units simultaneously (Pa, kPa, atm, bar, mmHg, psi, torr).

Pro Tip: Pair this with our Molarity Calculator for stock preparation.

What's Dalton's law of partial pressures?
Dalton's law (1801): each gas in a mixture exerts a pressure as if it alone occupied the container — the total pressure is the sum of partial pressures. P_i = x_i · P_total, where x_i is the mole fraction of component i. Equivalently P_total = Σ P_i. Example: dry air has x_N₂ = 0.781, x_O₂ = 0.209. At sea level (P_total = 101.325 kPa): P_N₂ = 79.1 kPa, P_O₂ = 21.2 kPa, summing to total. Exact for ideal gases; very accurate (≤ 1% error) for real gases at low to moderate pressures.
How does the ideal gas law give partial pressure?
Apply P = nRT/V to each component individually using its own moles n_i. Each component fills the entire volume at the system temperature; its partial pressure is P_i = n_i · R · T / V. For a gas mixture in a container: P_total = (n_total · R · T) / V = Σ (n_i · R · T) / V = Σ P_i — the same conclusion as Dalton's law from a different starting point. Useful when you've prepared a gas mixture by manometric methods (adding known amounts of each gas to a vacuum line) rather than by mixing known mole fractions.
What's Henry's law and why are there two methods?
Henry's law (1803): at constant temperature, the equilibrium pressure of a gas above a solution is proportional to the concentration of dissolved gas in the solution. The proportionality constant K_H is gas-specific and temperature-dependent. The TWO methods reflect TWO different conventions for K_H: Method 1 uses K_H in L·atm/mol with P = K_H · C (concentration in mol/L); Method 2 uses K_H in atm with P = K_H · x (mole fraction). They describe the same physical equilibrium and give the same answer when properly converted (using the molarity-to-mole-fraction relation through solvent density). Method 1 is more common in environmental and biological chemistry; method 2 in physical chemistry and chemical engineering. Always verify which form a published K_H value uses — confusion between methods is a classic source of order-of-magnitude errors.
What's the partial pressure of O₂ in air?
At sea level (P_total = 101.325 kPa = 1 atm), with x_O₂ = 0.20946 (dry-air composition): P_O₂ = 0.20946 × 101.325 = 21.22 kPa = 0.209 atm = 159.2 mmHg. This is the inspired O₂ partial pressure entering the lungs. After humidification by airway mucosa (P_H₂O = 47 mmHg at body temp) and dilution by alveolar CO₂ (P_CO₂ = 40 mmHg), alveolar P_O₂ drops to ~100 mmHg — the equilibrium pressure that drives O₂ diffusion into pulmonary capillary blood. At altitude, P_O₂ scales proportionally with P_total (~10.6 kPa at 5,500 m, half of sea level — the basis of altitude sickness onset).
How does temperature affect Henry's law constants?
Strongly! Gas solubility in water DECREASES as temperature increases for most gases (an entropy-driven effect — heating raises the entropic cost of confining gas molecules in liquid). The van't Hoff form: d(ln K_H) / d(1/T) = −ΔH_sol / R, where ΔH_sol is the enthalpy of solution (typically −10 to −20 kJ/mol). For O₂ in water (ΔH_sol ≈ −12.5 kJ/mol): warming 25 °C → 35 °C decreases solubility ~15%. O₂ saturation in water at 1 atm air vs T: 0 °C 14.6 mg/L; 20 °C 9.1; 25 °C 8.3; 30 °C 7.6; 40 °C 6.4 mg/L. The calculator's built-in K_H values are at 25 °C; for other temperatures use the van't Hoff correction or look up tabulated K_H(T) data.
Why is CO₂ so much more soluble than O₂?
K_H for CO₂ in water is ~26× lower than for O₂ (29.4 vs 769.2 L·atm/mol — lower K_H_method1 means MORE soluble). Three reasons: (1) CO₂ is a polar molecule with a quadrupole moment; it interacts more favorably with water dipoles than non-polar O₂; (2) CO₂ undergoes acid-base chemistry in water — partially converting to carbonic acid (H₂CO₃), bicarbonate (HCO₃⁻), and carbonate (CO₃²⁻); the chemical sink amplifies apparent solubility; (3) O₂ is hydrophobic — its non-polar nature makes water cluster around it (hydrophobic hydration), entropically unfavorable. This solubility difference is the foundation of carbonated beverages: a sealed Coke at 2.94 atm CO₂ holds ~0.1 M dissolved CO₂; an O₂ bottle at the same pressure would hold only ~0.004 M dissolved O₂.
What's the partial pressure of CO₂ in a Coke bottle?
Typical carbonation level: ~0.1 M dissolved CO₂ (pleasant carbonation level for cola). Henry's law method 1 with K_H1 = 29.4 L·atm/mol: P_CO₂ = 29.4 × 0.1 = 2.94 atm = 297 kPa = 43 psi. This is the CO₂ pressure inside the sealed bottle at equilibrium with the dissolved CO₂. When opened, P_CO₂ above the liquid drops to atmospheric ~0.0004 atm — a 7,000-fold drop — and the dissolved CO₂ becomes massively supersaturated, driving rapid bubble formation (the famous fizzing). The bottle's plastic + cap is designed for ~5 atm sustained pressure; champagne bottles routinely hold ~6 atm and use a wire cage on the cork to prevent unwanted ejection.
What partial pressure is dangerous to breathe?
Oxygen: P_O₂ ≥ 1.4 atm sustained causes CNS oxygen toxicity (convulsions, blackouts). Recreational diving max 1.4 atm; technical max 1.6 atm during decompression. P_O₂ ≤ 0.16 atm causes hypoxia — equivalent to altitude > 5,500 m. Nitrogen: P_N₂ ≥ 3 atm causes nitrogen narcosis (judgment impairment like alcohol intoxication) — equivalent to ~30 m depth on air. Carbon dioxide: P_CO₂ ≥ 0.05 atm (5% CO₂) causes hypercapnia symptoms; OSHA short-term exposure limit. Helium: no narcosis at any recreational depth — basis of heliox / trimix diving. Carbon monoxide: even tiny P_CO ≥ 25 ppm causes headache; ≥ 100 ppm dangerous; ≥ 800 ppm rapidly fatal due to hemoglobin binding (240× tighter than O₂).
What pressure unit should I use?
Pa (pascal): SI standard; most-used in physics, scientific papers. kPa: the practical SI unit for atmospheric and biological pressures. atm: historically standard in chemistry; 1 atm = 101.325 kPa. bar: close to 1 atm (1 bar = 100 kPa); used in engineering and meteorology. mmHg = torr: standard in clinical medicine (blood pressure!), vacuum technology; 760 mmHg = 1 atm. psi: US engineering, tire pressure (1 atm = 14.7 psi). The calculator outputs all 7 simultaneously so you don't have to pick — copy whichever your downstream calculation needs.
When does the ideal gas approximation break down?
Ideal gas law assumes molecules occupy negligible volume and have no intermolecular forces. Fails when: (1) Pressure is high (typically > 10 atm) — molecules are forced close enough that volume and forces matter. (2) Temperature is low (typically < 100 K for most gases, < critical T) — kinetic energy no longer dwarfs intermolecular attraction. (3) Gas is polar (NH₃, H₂O vapour, HCl) — strong dipole-dipole interactions even at moderate conditions. (4) Near phase transitions — within ~10% of critical pressure / temperature, real-gas effects dominate. (5) High molecular weight gases (alkanes > C₄, halogenated compounds) — larger molecular volume becomes significant. For these cases use Van der Waals (qualitative), Peng-Robinson (industrial standard for petroleum / natural gas), or compressibility-factor charts. For typical chemistry conditions (1 atm, room temperature, common small gases), ideal gas error is < 0.5%.

Author Spotlight

The ToolsACE Team - ToolsACE.io Team

The ToolsACE Team

Our ToolsACE physical-chemistry team built this calculator to handle the four standard methods for computing the partial pressure of a gas — each appropriate for different physical situations. <strong>Dalton's law</strong> (P_i = x_i · P_total) is the workhorse for known gas mixtures: enter total pressure and mole fraction. <strong>Ideal gas law</strong> (P = nRT/V) computes partial pressure of a single component from its moles, temperature, and volume — useful when you've prepared a gas mixture by manometric methods. <strong>Henry's law method 1</strong> (P = K_H · C) computes the equilibrium gas-phase pressure above a solution of known dissolved-gas concentration, using K_H in L·atm/mol units. <strong>Henry's law method 2</strong> (P = K_H · x) uses K_H in atm units with the dissolved-gas mole fraction. Both Henry methods compute the same physical quantity from the same gas-liquid equilibrium but use different conventions for the Henry's law constant. The calculator includes built-in K_H values for 8 common gases (O₂, N₂, CO₂, H₂, He, CH₄, Ar, CO) at 25 °C in both unit conventions, sourced from the CRC Handbook.

Dalton's Law of Partial PressuresIUPAC Ideal Gas Constant RCRC Handbook of Chemistry and Physics

Disclaimer

All four methods assume ideal gas behaviour. Approximation valid for low pressures (≤ 10 atm), high temperatures (≥ 100 K), and small / non-polar gases. Breaks down at high pressures, low temperatures, polar molecules, near phase transitions. Henry's law adds the dilute-solution assumption (typically valid below ~1 M). Henry's law constants are temperature-dependent; built-in values are at 25 °C. References: CRC Handbook of Chemistry and Physics, IUPAC Quantities Units and Symbols, Atkins' Physical Chemistry. For non-ideal conditions use Van der Waals or Peng-Robinson EOS; for high-loading solutions use activity-coefficient corrections.