Molar Mass of Carbon Dioxide (CO₂)
Learn how chemists calculate the molar mass of Carbon Dioxide (CO₂), with a clear formula breakdown, worked steps, and study notes · IUPAC name: Carbon dioxide.
Quick answer
The molar mass of Carbon Dioxide (CO₂) is
44.009g/mol
One mole of Carbon Dioxide therefore has a mass of 44.009 grams—the value you use for stoichiometry and laboratory preparation.
Reviewed for educational accuracy · Accuracy policy
- CAS Registry Number
- 124-38-9
- PubChem CID
- 280
- SMILES
- O=C=O
Step-by-step calculation
Let's find the molar mass of Carbon Dioxide (CO₂) together—step by step, as if you are seeing the formula for the first time.
Step 1 — Look at the chemical formula
The formula is CO₂. Each letter stands for an element. The little number after a letter (the subscript) tells you how many atoms of that element are in one molecule or formula unit.
- 1 Carbon atom (C)
- 2 Oxygen atoms (O)
Step 2 — Look up each atomic mass
Atomic mass comes from the periodic table. It is the average mass of one mole of atoms of that element, in grams per mole (g/mol). Think of it as the "price tag" for one mole of that element.
- Carbon (C) = 12.011 g/mol
- Oxygen (O) = 15.999 g/mol
Step 3 — Multiply atoms × atomic mass
Why multiply? If one oxygen atom "costs" about 16 g/mol, then two oxygen atoms cost twice as much. Each element's contribution is: number of atoms × atomic mass.
- 1 × 12.011 = 12.011 g/mol (Carbon)
- 2 × 15.999 = 31.998 g/mol (Oxygen)
Step 4 — Add the contributions
Why add? The molar mass of the whole compound is simply the total mass of every atom in the formula. Add each element's contribution:
12.011 + 31.998 = 44.009 g/mol
Step 5 — Final answer
Molar mass of Carbon Dioxide = 44.009 g/mol
That means one mole of Carbon Dioxide (CO₂) has a mass of about 44.01 grams.
Quick summary
Read the formula → count atoms → look up atomic masses → multiply → add → report g/mol. For CO₂, the total is 44.009 g/mol.
Common beginner mistakes
- Writing CO² instead of CO₂ (subscript vs superscript confusion).
- Assuming CO₂ supports combustion because it contains oxygen.
- Forgetting that CO₂ is heavier than air when discussing gas collection methods (downward displacement).
Memory trick
Link 44 g/mol to the ideal gas law: 44 g at STP ≈ 22.4 L.
Mini practice
Without looking above, list the atoms in CO₂ and write one multiplication line for the heaviest element. Then check your work against Step 3.
Real-world example
If a recipe asks for 0.100 mol of Carbon Dioxide, mass needed = 0.100 × 44.009 = 4.401 g. That is how chemists turn a mole amount into a weighable sample.
Atomic contribution table
Each row shows how much mass one element contributes to the total for CO₂.
| Element | Atoms | Atomic mass | Contribution | Mass % |
|---|---|---|---|---|
| C | 1 | 12.011 | 12.011 g/mol | 27.3% |
| O | 2 | 15.999 | 31.998 g/mol | 72.7% |
| Total molar mass | 44.009 g/mol | 100% | ||
Mass contribution chart
Linear molecule · 180°
Download study sheets
Save a printable summary, revision sheet, practice worksheet, or laboratory reference for Carbon Dioxide (CO₂).
Practice this calculation
Without looking above, write the atom count for CO₂, then compute the molar mass. Check your answer against 44.009 g/mol.
Next challenge: how many grams are in 0.250 mol of Carbon Dioxide? Multiply 0.250 × 44.009 to get 11.002 g.
Physical and chemical properties
Physical properties
| Appearance | Colorless gas; solid (dry ice) is white opaque material |
| Color | Colorless (gas and liquid) |
| Odor | Odorless at normal concentrations (high concentrations smell sharp/acidic due to carbonic acid in mucous membranes) |
| State (STP) | Gas |
| Density | 1.977 g/L at 0 °C, 1 atm (gas); 1.101 g/mL (liquid at −37 °C); 1.562 g/mL (solid dry ice) |
| Melting point | −56.6 °C (triple point); solid sublimes at −78.5 °C at 1 atm |
| Boiling point | Sublimes at −78.5 °C at 1 atm (no liquid phase at 1 atm below triple point pressure) |
| Solubility | 0.145 g/100 mL water at 25 °C (decreases with temperature, unlike most gases — relevant to ocean CO₂ uptake) |
| Crystal structure | Cubic (solid CO₂, Pa-3) |
Chemical properties
| Classification | Acidic oxide (non-metal oxide) |
| Family | Group 14 dioxide / acidic oxide |
| Polarity | Non-polar (dipoles cancel) |
| Geometry | Linear |
| Bond angle | 180° |
| Oxidation states | C: +4, O: −2 |
Applications
Industrial uses
- Supercritical CO₂ extraction of caffeine from coffee and hops from beer
- Enhanced oil recovery by injecting CO₂ into depleted wells
- Welding shield gas (CO₂/argon mixtures for MIG welding)
- Refrigerant R-744 in commercial heat pumps and transcritical cycles
Laboratory uses
- Inert atmosphere for air-sensitive syntheses
- Preparation of metal carbonates by bubbling through hydroxide solutions
- Calibration gas for infrared gas analyzers
Primary anthropogenic greenhouse gas; monitored via Mauna Loa Keeling Curve; target of carbon capture and storage (CCS) technologies.
Essential substrate for C3 and C4 photosynthesis; regulated in blood as part of the bicarbonate buffer system (Henderson–Hasselbalch); elevated levels trigger the medullary respiratory center to increase breathing rate.
Preparation and production
Laboratory CO₂ is generated by reacting calcium carbonate with dilute hydrochloric acid: CaCO₃ + 2 HCl → CaCl₂ + H₂O + CO₂↑. It is also produced by fermenting sugar with yeast or by subliming dry ice.
Industrial CO₂ is captured from ammonia synthesis, hydrogen production, natural gas processing, and fermentation. It is purified, compressed, and liquefied or solidified for commercial distribution.
Important reactions of Carbon Dioxide
CO₂(g) + H₂O(l) ⇌ H₂CO₃(aq)
- Reaction type
- Hydration
- Conditions
- Aqueous solution, ambient conditions
- Explanation
- CO₂ hydrates reversibly to form carbonic acid, the basis of the bicarbonate buffer in natural waters and blood.
- Products
- Carbonic acid (equilibrium mixture with HCO₃⁻ and CO₃²⁻)
- Why it matters
- Ocean acidification studies, blood gas analysis, carbonated beverages
Related ideas: Weak acids · Buffer systems · Henry's law
CO₂(g) + 2 NaOH(aq) → Na₂CO₃(aq) + H₂O(l)
- Reaction type
- Acid–base (with base)
- Conditions
- Excess NaOH, aqueous
- Explanation
- CO₂ reacts with strong base to form carbonate; with limited base, bicarbonate forms first.
- Products
- Sodium carbonate and water
- Why it matters
- CO₂ scrubbing, air purification in submarines and spacecraft
Related ideas: Acidic oxides · Gas absorption · Stoichiometry
6 CO₂ + 6 H₂O → C₆H₁₂O₆ + 6 O₂
- Reaction type
- Photosynthesis (light-driven)
- Conditions
- Chlorophyll, sunlight, chloroplasts
- Explanation
- Plants convert CO₂ and water into glucose, storing solar energy in chemical bonds and releasing O₂.
- Products
- Glucose and oxygen
- Why it matters
- Agriculture, carbon sequestration, global carbon cycle
Related ideas: Redox · Bioenergetics · Carbon fixation
CaCO₃(s) + 2 HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g)
- Reaction type
- Acid–carbonate (reverse direction for CO₂ production)
- Conditions
- Room temperature, dilute acid
- Explanation
- Carbonates and acids produce CO₂; this is the standard lab preparation and the test for carbonate ions.
- Products
- Calcium chloride, water, and carbon dioxide gas
- Why it matters
- Laboratory CO₂ generation, carbonate identification
Related ideas: Gas evolution · Limiting reagent · Qualitative analysis
CO₂(g) + RNH₂(aq) ⇌ RNHCOO⁻(aq) + RNH₃⁺(aq)
- Reaction type
- Reversible amine capture (carbon capture chemistry)
- Conditions
- Aqueous amine solution (e.g., monoethanolamine), moderate temperature; reversed by heating
- Explanation
- CO₂ reversibly reacts with amine solvents to form a carbamate salt at low temperature, then releases concentrated CO₂ gas again on heating, the core chemistry of industrial post-combustion carbon capture.
- Products
- Ammonium carbamate salt (captured form); releases concentrated CO₂ on regeneration
- Why it matters
- Power plant flue-gas carbon capture, direct air capture technology
Related ideas: Carbon capture and storage · Reversible reactions · Industrial gas separation
History and discovery
Joseph Black identified "fixed air" (CO₂) in the 1750s by showing that it was distinct from atmospheric air and could extinguish flames. Lavoisier later recognized it as an oxide of carbon. The role of CO₂ in the greenhouse effect was proposed by Fourier (1824), experimentally demonstrated by Tyndall (1859), and linked to fossil fuel emissions by Arrhenius (1896).
Joseph Black, 1754 — discovered carbon dioxide as a gas released from heating limestone and fermenting solutions.
Interesting facts
- Solid CO₂ is called dry ice because it sublimes without leaving liquid at 1 atm.
- CO₂ makes soda fizzy; approximately 3–4 volumes of CO₂ dissolve under pressure in soft drinks.
- Venus's atmosphere is ~96.5% CO₂, creating a runaway greenhouse effect with surface temperatures near 465 °C.
- The critical point of CO₂ is 31.1 °C and 73.8 bar, enabling supercritical fluid applications.
Comparison with similar compounds
CO (carbon monoxide, 28.01 g/mol) is toxic and binds hemoglobin; CO₂ (44.01 g/mol) is non-toxic at low levels but acts as an asphyxiant and is a greenhouse gas. The two carbon oxides differ fundamentally in bonding (CO has a triple bond, CO₂ has two double bonds) and in the oxidation state of carbon (+2 in CO versus +4 in CO₂), reflecting CO's origin in incomplete combustion versus CO₂'s origin in complete combustion.
Storage, handling, and safety
Store compressed gas cylinders upright, secured, away from heat. Solid CO₂ (dry ice) requires vented insulated containers; never seal in airtight vessels as sublimation pressure can rupture containers.
Ensure adequate ventilation. CO₂ is an asphyxiant at high concentrations (>5% can cause dizziness; >10% can be fatal). Use CO₂ monitors in basements, breweries, and dry-ice storage areas. Handle dry ice with insulated gloves to prevent frostbite.
Low toxicity at ambient atmospheric levels but dangerous as an asphyxiant in enclosed spaces; dry ice causes cold burns.
- Asphyxiation in poorly ventilated areas
- Frostbite from contact with dry ice or liquid CO₂
- Rupture hazard if solid CO₂ is stored in sealed containers
Classification: GHS: Compressed gas; may displace oxygen
Exam notes and student tips
Exam notes
- Molar mass CO₂ = 12.01 + 2(16.00) = 44.01 g/mol.
- 1 mole of CO₂ at STP occupies 22.4 L and contains 1 mole of C and 2 moles of O atoms.
- In photosynthesis: 6 CO₂ + 6 H₂O → C₆H₁₂O₆ + 6 O₂ (memorize for biology crossover questions).
- Linear geometry with 180° bond angle → non-polar molecule.
Student tips
- Link 44 g/mol to the ideal gas law: 44 g at STP ≈ 22.4 L.
- Use the lime water test (CO₂ + Ca(OH)₂ → CaCO₃ white precipitate) as a identification reaction.
- Draw the two double bonds and note why dipoles cancel for polarity questions.
Common mistakes
- Writing CO² instead of CO₂ (subscript vs superscript confusion).
- Assuming CO₂ supports combustion because it contains oxygen.
- Forgetting that CO₂ is heavier than air when discussing gas collection methods (downward displacement).
Misconceptions
- CO₂ is not poisonous like CO; it harms by oxygen displacement, not by binding hemoglobin.
- Plants do not " breathe out" only oxygen; respiration releases CO₂ continuously.
- Carbon dioxide is not the same as carbon monoxide (CO, molar mass 28.01 g/mol).
Practice questions
1. Calculate the molar mass of CO₂.
Show answer
12.01 + 2(16.00) = 44.01 g/mol
2. What mass of CO₂ contains the same number of oxygen atoms as 36 g of water?
Hint: Compare moles of O atoms, not molecules.
Show answer
36 g H₂O = 2 mol H₂O = 4 mol O atoms. CO₂ has 2 O per molecule, so need 2 mol CO₂ = 2 × 44.01 = 88.02 g
3. Why is CO₂ collected by downward displacement of air?
Show answer
CO₂ (M = 44 g/mol) is denser than air (~29 g/mol average), so it sinks and displaces air upward.
4. How many liters does 88 g of CO₂ occupy at STP?
Show answer
88 g ÷ 44.01 g/mol = 2.0 mol; 2.0 × 22.4 L/mol = 44.8 L
Frequently asked questions about Carbon Dioxide
44.01 g/mol, from 12.01 (C) + 2 × 16.00 (O).
Chemistry of Carbon Dioxide
The sections above give the number you need for calculations. Here we look more closely at how Carbon Dioxide (CO₂) behaves chemically—so the molar mass connects to real reactions, properties, and laboratory practice.
Carbon dioxide (CO₂) is a linear triatomic molecule with molar mass 44.01 g/mol, formed from one carbon atom (12.01 g/mol) and two oxygen atoms. At standard temperature and pressure it is an invisible, odorless gas denser than air (ρ ≈ 1.98 g/L vs. air ~1.29 g/L), which is why it accumulates in low-lying areas and can displace breathable air in confined spaces.
CO₂ plays a central role in Earth's carbon cycle. Plants fix atmospheric CO₂ into carbohydrates via photosynthesis, while respiration, combustion, and volcanic outgassing return it to the atmosphere. Since the Industrial Revolution, fossil fuel burning has raised atmospheric CO₂ from roughly 280 ppm to over 420 ppm, driving radiative forcing through the greenhouse effect: CO₂ absorbs infrared radiation emitted by Earth's surface, trapping thermal energy in the troposphere.
Because CO₂ is both a natural biogeochemical participant and the dominant anthropogenic greenhouse gas, it occupies a unique dual role in chemistry education: it is simultaneously the harmless-seeming product of respiration and combustion that students first encounter in balanced equations, and the central molecule of one of the most consequential scientific and policy debates of the modern era. Understanding its simple linear structure and acidic-oxide chemistry is inseparable, in a full chemistry education, from understanding its outsized role in climate science, carbon capture engineering, and the global carbon budget.
The formula CO₂ shows a 1:2 carbon-to-oxygen ratio. Carbon forms two double bonds to oxygen (O=C=O), leaving no lone pairs on carbon and satisfying the octet rule for all atoms. The empirical formula equals the molecular formula because the molecule is already the simplest whole-number ratio.
CO₂ is a non-polar molecule despite polar C=O bonds because the dipoles cancel in the linear geometry. It is an acidic oxide: dissolving in water forms weak carbonic acid (H₂CO₃). CO₂ is not combustible and does not support combustion; instead, it extinguishes fires by displacing oxygen. At −78.5 °C and 1 atm it sublimes directly from solid (dry ice) to gas.
The greenhouse effect: how CO2 traps heat
CO₂ absorbs infrared radiation because its bending and stretching vibrational modes are excited by photons in the thermal infrared range emitted by Earth's warm surface; symmetric linear CO₂ has no permanent dipole, but its asymmetric stretch and bend vibrations create transient dipole changes that let it interact with infrared light despite being a nonpolar molecule overall. This absorbed energy is partly re-radiated back toward the surface rather than escaping directly to space, which is the mechanistic core of the greenhouse effect. Since pre-industrial times, atmospheric CO₂ has risen from about 280 ppm to over 420 ppm, and this single number is one of the most closely tracked metrics in climate science, measured continuously since 1958 at Mauna Loa Observatory (the Keeling Curve).
Photosynthesis: CO2 as the carbon source for all plant biomass
Virtually all carbon in plant biomass, and by extension in the food chains and fossil fuels derived from ancient plants, originates from atmospheric CO₂ fixed during photosynthesis: 6 CO₂ + 6 H₂O → C₆H₁₂O₆ + 6 O₂, powered by light energy captured in chlorophyll. This reaction is the biological reverse of respiration and combustion, closing a cycle in which carbon moves between the atmosphere and living matter on timescales ranging from days (leaf-level gas exchange) to millions of years (fossil fuel formation and eventual combustion).
Dry ice: sublimation chemistry without a liquid phase
Solid CO₂ (dry ice) does not melt into a liquid at ordinary atmospheric pressure; instead it sublimes directly from solid to gas at −78.5 °C because the triple point of CO₂ (−56.6 °C at 5.11 atm) lies above 1 atm, meaning liquid CO₂ can only exist at pressures higher than normal atmospheric pressure. This unusual phase behavior makes dry ice a uniquely 'dry' refrigerant that leaves no liquid residue as it warms, valuable for shipping temperature-sensitive goods, creating fog effects in theatrical productions, and even blasting/cleaning applications that exploit the rapid volume expansion during sublimation.
The global carbon cycle: reservoirs and residence times
Carbon cycles between the atmosphere, oceans, terrestrial biosphere, and geological reservoirs (carbonate rocks, fossil fuels) on vastly different timescales — atmospheric CO₂ exchanges with surface ocean water and plant biomass over years to decades, but the deep ocean and sedimentary rock reservoirs turn over on timescales of centuries to millions of years. Human fossil fuel combustion releases carbon that has been sequestered in geological reservoirs for tens to hundreds of millions of years, injecting it into the fast atmospheric-biological cycle far more rapidly than natural processes can re-sequester it, which is the fundamental reason atmospheric CO₂ concentration is currently rising rather than remaining in steady-state balance.
Carbon capture and storage (CCS) technologies
Carbon capture and storage technologies aim to intercept CO₂ from power plant flue gas or directly from ambient air (direct air capture) and sequester it underground in geological formations or convert it into stable mineral carbonates, preventing its release to the atmosphere. Common capture chemistries include amine scrubbing (where CO₂ reversibly binds to amine solutions and is later released by heating for concentrated storage) and emerging solid sorbent and mineralization approaches; despite significant engineering progress, CCS remains energy-intensive and costly to deploy at the gigaton scale needed to meaningfully offset global emissions.
Recalculate any formula with the molar mass calculator, compare atoms on the periodic table, or browse more compounds in the gas library.
References and further reading
- IPCC AR6: Radiative forcing of CO₂ and climate projections
- NIST Chemistry WebBook: Thermophysical data for CO₂
- PubChem CID 280: Compound identifiers and properties