Molar Mass of Carbon Monoxide (CO)
Learn how chemists calculate the molar mass of Carbon Monoxide (CO), with a clear formula breakdown, worked steps, and study notes · IUPAC name: Carbon monoxide.
Quick answer
The molar mass of Carbon Monoxide (CO) is
28.010g/mol
One mole of Carbon Monoxide therefore has a mass of 28.010 grams—the value you use for stoichiometry and laboratory preparation.
Reviewed for educational accuracy · Accuracy policy
- CAS Registry Number
- 630-08-0
- PubChem CID
- 281
- SMILES
- [C-]#[O+]
Step-by-step calculation
Let's find the molar mass of Carbon Monoxide (CO) 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 CO. 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)
- 1 Oxygen atom (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)
- 1 × 15.999 = 15.999 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 + 15.999 = 28.010 g/mol
Step 5 — Final answer
Molar mass of Carbon Monoxide = 28.010 g/mol
That means one mole of Carbon Monoxide (CO) has a mass of about 28.01 grams.
Quick summary
Read the formula → count atoms → look up atomic masses → multiply → add → report g/mol. For CO, the total is 28.010 g/mol.
Common beginner mistakes
- Confusing carbon monoxide (CO, 28.01 g/mol, toxic) with carbon dioxide (CO₂, 44.01 g/mol, non-toxic at ambient levels) — different formula, bonding, and hazard profile.
- Assuming CO is detectable by smell — it is completely odorless, unlike gases such as ammonia or hydrogen sulfide.
- Writing the Lewis structure of CO with formal charges reversed from what molecular orbital theory predicts; the correct structure places a negative formal charge on carbon.
Memory trick
Remember CO's molar mass (28.01) is very close to N₂ (28.02) — a useful memory link since both are strong triple-bonded diatomics of similar mass.
Mini practice
Without looking above, list the atoms in CO 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 Monoxide, mass needed = 0.100 × 28.010 = 2.801 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 CO.
| Element | Atoms | Atomic mass | Contribution | Mass % |
|---|---|---|---|---|
| C | 1 | 12.011 | 12.011 g/mol | 42.9% |
| O | 1 | 15.999 | 15.999 g/mol | 57.1% |
| Total molar mass | 28.010 g/mol | 100% | ||
Mass contribution chart
Count every atom in this formula, multiply by atomic mass, then add. That total is the molar mass used in lab weighing.
Download study sheets
Save a printable summary, revision sheet, practice worksheet, or laboratory reference for Carbon Monoxide (CO).
Practice this calculation
Without looking above, write the atom count for CO, then compute the molar mass. Check your answer against 28.010 g/mol.
Next challenge: how many grams are in 0.250 mol of Carbon Monoxide? Multiply 0.250 × 28.010 to get 7.002 g.
Physical and chemical properties
Physical properties
| Appearance | Colorless gas |
| Color | Colorless |
| Odor | Odorless |
| State (STP) | Gas |
| Density | 1.145 g/L at 25 °C (slightly less dense than air) |
| Melting point | −205.0 °C |
| Boiling point | −191.5 °C |
| Solubility | 0.0026 g/100 mL water at 20 °C (poorly soluble) |
| Crystal structure | Not applicable at standard conditions (gas) |
Chemical properties
| Classification | Inorganic oxide (non-metal oxide) / reducing gas |
| Family | Carbon oxides (lower oxidation state oxide of carbon) |
| Polarity | Weakly polar (dipole moment 0.11 D) |
| Geometry | Linear (diatomic) |
| Bond angle | 180° (diatomic, trivially linear) |
| Oxidation states | C: +2, O: −2 |
Applications
Industrial uses
- Reducing agent in blast furnace iron and steel production
- Key component of syngas for Fischer–Tropsch synthesis of liquid fuels
- Feedstock for methanol and acetic acid manufacture (carbonylation processes)
- Metal carbonyl formation for purification (Mond process for nickel refining)
Laboratory uses
- Ligand in organometallic and coordination chemistry research
- Reducing atmosphere for certain high-temperature metallurgical experiments
- Model system for studying molecular orbital theory and bond order
A regulated air pollutant from vehicle exhaust and combustion sources; though not a major direct greenhouse gas, atmospheric CO indirectly affects climate by consuming hydroxyl radicals that would otherwise break down methane.
Endogenously produced in trace amounts during heme breakdown and studied as a signaling molecule in vascular biology, but at higher concentrations acts as a potent respiratory toxin by displacing oxygen from hemoglobin.
Preparation and production
Laboratory CO can be generated by dehydrating formic acid with concentrated sulfuric acid: HCOOH → CO + H₂O (H₂SO₄ catalyst, heat). Industrially, it is produced via the water-gas reaction over hot coke, partial combustion of carbon-based fuels, or steam reforming of natural gas as part of syngas production.
Global CO production is largely captured in situ within syngas and blast-furnace gas streams rather than isolated as a standalone commodity, though purified CO is sold for specialty carbonylation chemistry and metal refining applications.
Important reactions of Carbon Monoxide
2 CO(g) + O₂(g) → 2 CO₂(g)
- Reaction type
- Combustion (oxidation)
- Conditions
- Ignition source, presence of oxygen
- Explanation
- Carbon monoxide burns with a characteristic blue flame, releasing substantial heat as it is fully oxidized to carbon dioxide.
- Products
- Carbon dioxide
- Why it matters
- Complete combustion in engines and furnaces, catalytic converter chemistry
Related ideas: Combustion · Oxidation states · Thermochemistry
Fe₂O₃(s) + 3 CO(g) → 2 Fe(l) + 3 CO₂(g)
- Reaction type
- Redox (reduction of metal oxide)
- Conditions
- Blast furnace, ~1,200–2,000 °C
- Explanation
- Carbon monoxide reduces iron(III) oxide to molten iron metal while itself being oxidized to CO₂, the central reaction of primary steel production.
- Products
- Molten iron and carbon dioxide gas
- Why it matters
- Blast furnace ironmaking, extractive metallurgy
Related ideas: Reducing agents · Extractive metallurgy · Redox
CO(g) + 2 H₂(g) → CH₃OH(g)
- Reaction type
- Catalytic synthesis (carbonylation/hydrogenation)
- Conditions
- Cu/ZnO/Al₂O₃ catalyst, ~250 °C, 50–100 atm
- Explanation
- Syngas (CO and H₂) is catalytically combined to form methanol, a major industrial route to this key feedstock chemical.
- Products
- Methanol
- Why it matters
- Industrial methanol production, syngas chemistry
Related ideas: Catalysis · Syngas · Industrial synthesis
HCOOH(l) → CO(g) + H₂O(l)
- Reaction type
- Dehydration (acid-catalyzed decomposition)
- Conditions
- Concentrated H₂SO₄, heat
- Explanation
- Formic acid is dehydrated to produce carbon monoxide gas, a common laboratory method for generating small quantities of CO.
- Products
- Carbon monoxide and water
- Why it matters
- Laboratory CO generation for controlled experiments
Related ideas: Dehydration reactions · Gas generation · Organic decomposition
History and discovery
Carbon monoxide's toxic properties were documented as early as antiquity through observations of charcoal-burning fumes causing death, but its chemical identity as a distinct gas was established in the late 18th century. French chemist Claude Louis Berthollet showed it was composed of carbon and oxygen in the 1780s, and its formula and bonding were clarified through 19th- and 20th-century advances in molecular orbital theory that explained its unusual electronic structure.
Recognized as a distinct toxic gas in the late 18th century; its molecular composition was established by Berthollet and later chemists studying combustion gases.
Interesting facts
- CO binds hemoglobin roughly 200–250 times more strongly than oxygen, which is why even small amounts can be dangerous.
- The triple bond in CO (bond order 3) is one of the strongest chemical bonds known between two atoms, even stronger than the N≡N bond in nitrogen gas.
- The body produces small amounts of CO naturally during heme catabolism, and it functions as a minor signaling molecule in blood vessel regulation.
- Home carbon monoxide detectors became mandatory in many jurisdictions only in recent decades, despite CO poisoning being recognized as a hazard for over a century.
Comparison with similar compounds
CO (28.01 g/mol, triple-bonded, acutely toxic via hemoglobin binding) contrasts sharply with CO₂ (44.01 g/mol, double-bonded, non-toxic at ambient levels but a major greenhouse gas) — both carbon oxides but with almost opposite hazard profiles.
Storage, handling, and safety
Store compressed CO gas cylinders in well-ventilated areas equipped with CO detectors, secured upright and away from heat sources. Because CO is both toxic and flammable, storage areas require both gas-leak monitoring and fire-safety precautions.
Handle only in well-ventilated areas or fume hoods with continuous CO monitoring. Never operate combustion engines, generators, or unvented fuel-burning appliances in enclosed spaces. Install and regularly test carbon monoxide detectors in any residence with gas appliances, fireplaces, or attached garages.
Highly toxic by inhalation due to strong, competitive binding to hemoglobin; colorless and odorless, giving no warning of dangerous exposure; also flammable and forms explosive mixtures with air.
- Acute poisoning from hemoglobin binding, causing hypoxia even at low blood oxygen displacement
- No sensory warning (colorless, odorless) prior to dangerous exposure
- Flammable gas; can form explosive mixtures with air (12.5–74% by volume)
- Chronic low-level exposure causing headache, fatigue, and cognitive impairment
Classification: GHS: Flam. Gas 1, Acute Tox. 3 (inhalation), Repr. 1A
Exam notes and student tips
Exam notes
- Molar mass CO = 12.01 + 16.00 = 28.01 g/mol.
- CO has a triple bond and bond order 3 — contrast with CO₂'s two double bonds and O₂'s double bond when discussing bond strength trends.
- Complete combustion of hydrocarbons produces CO₂; incomplete combustion (limited O₂) produces CO — a key exam distinction in combustion stoichiometry.
- CO is a reducing agent in blast furnace reactions: Fe₂O₃ + 3 CO → 2 Fe + 3 CO₂ — memorize for extraction of metals topics.
Student tips
- Remember CO's molar mass (28.01) is very close to N₂ (28.02) — a useful memory link since both are strong triple-bonded diatomics of similar mass.
- Use the blast furnace equation as your go-to example when asked for a reaction where CO acts as a reducing agent.
- Link CO's toxicity mechanism (hemoglobin binding) to real-world safety advice about detectors and ventilation for a memorable exam answer.
Common mistakes
- Confusing carbon monoxide (CO, 28.01 g/mol, toxic) with carbon dioxide (CO₂, 44.01 g/mol, non-toxic at ambient levels) — different formula, bonding, and hazard profile.
- Assuming CO is detectable by smell — it is completely odorless, unlike gases such as ammonia or hydrogen sulfide.
- Writing the Lewis structure of CO with formal charges reversed from what molecular orbital theory predicts; the correct structure places a negative formal charge on carbon.
Misconceptions
- CO poisoning is not the same mechanism as CO₂ asphyxiation — CO actively displaces oxygen from hemoglobin rather than simply diluting available oxygen in the air.
- CO is not a significant direct contributor to the greenhouse effect at typical concentrations, unlike CO₂; its main environmental role is as an air pollutant and indirect atmospheric chemistry participant.
- A working smoke detector does not detect carbon monoxide — dedicated CO detectors use different sensing technology and must be installed separately.
Practice questions
1. Calculate the molar mass of carbon monoxide, CO.
Show answer
12.01 + 16.00 = 28.01 g/mol
2. How many grams of CO are produced from the incomplete combustion of 24 g of carbon with limited oxygen (2 C + O₂ → 2 CO)?
Show answer
24 g C ÷ 12.01 g/mol = 2.0 mol C → 2.0 mol CO produced = 2.0 × 28.01 = 56.0 g
3. What bond order does carbon monoxide have, and how does it compare to N₂?
Show answer
CO has a bond order of 3 (triple bond), comparable to and even slightly stronger than the triple bond in N₂, explaining CO's high bond dissociation energy.
4. Why is a carbon monoxide detector necessary even if a smoke detector is present?
Show answer
CO is colorless and odorless and does not produce smoke particles, so smoke detectors cannot sense it; a dedicated CO detector uses a different sensing mechanism.
Frequently asked questions about Carbon Monoxide
28.01 g/mol, from 12.01 (C) + 16.00 (O).
Chemistry of Carbon Monoxide
The sections above give the number you need for calculations. Here we look more closely at how Carbon Monoxide (CO) behaves chemically—so the molar mass connects to real reactions, properties, and laboratory practice.
Carbon monoxide (CO) is a diatomic molecule with a molar mass of 28.01 g/mol, formed from one carbon and one oxygen atom joined by a strong triple bond. Despite its simple two-atom structure, CO carries a formal negative charge on carbon and positive charge on oxygen in its best Lewis structure — an unusual polarity reversal explained by molecular orbital theory, where the highest-occupied orbital is concentrated on carbon rather than the more electronegative oxygen. This electronic arrangement is directly responsible for carbon monoxide's most notorious property: an extraordinarily strong affinity for binding transition-metal centers, including the iron atom at the heart of hemoglobin.
CO forms whenever carbon-containing fuel burns without enough oxygen for complete combustion. Complete combustion of a hydrocarbon yields CO₂ and H₂O, but oxygen-starved (incomplete) combustion — in poorly ventilated furnaces, idling car engines, charcoal grills used indoors, or faulty gas heaters — instead produces the more dangerous CO. Because carbon monoxide is colorless, odorless, and non-irritating, it gives victims no sensory warning before dangerous concentrations accumulate, making it responsible for the majority of accidental poisoning deaths from gas inhalation worldwide and earning its reputation as "the silent killer."
Despite its toxicity, CO plays an essential and controlled role in industrial chemistry. It is the key building block of syngas (synthesis gas, a CO/H₂ mixture) used in Fischer–Tropsch synthesis to make liquid hydrocarbons and in the industrial production of methanol. In blast furnaces, CO generated from coke serves as the reducing agent that strips oxygen from iron ore, making it indispensable to global steel production despite the careful containment its toxicity demands.
The formula CO represents a 1:1 ratio of carbon to oxygen joined by a formal triple bond (one sigma and two pi bonds), giving a bond order of 3 — among the strongest bonds known for a diatomic molecule (bond dissociation energy ≈ 1072 kJ/mol, even greater than N₂). The molecule carries a small dipole moment (0.11 D) with an unusual polarity: carbon bears a slight negative charge despite oxygen's higher electronegativity, a consequence of the molecular orbital structure rather than simple electronegativity-based reasoning.
Carbon monoxide is a reducing agent, readily donating its lone pair on carbon to metal centers (forming metal carbonyl complexes) and reducing metal oxides to free metals at high temperature. It burns in air to form CO₂ (2 CO + O₂ → 2 CO₂), releasing significant heat, and this combustibility is itself a secondary hazard in enclosed spaces where CO accumulates. CO is also a ligand in organometallic chemistry, forming stable complexes such as Ni(CO)₄ and Fe(CO)₅ through back-donation from metal d-orbitals into CO's empty pi-antibonding orbitals.
Toxicity mechanism: why CO is so much more dangerous than CO2
Carbon monoxide poisons the body by binding to the iron atom in hemoglobin roughly 200–250 times more tightly than oxygen does, forming carboxyhemoglobin (COHb) that cannot carry oxygen and also shifts the oxygen-hemoglobin dissociation curve, impairing oxygen release to tissues even from unaffected hemoglobin. This is fundamentally different from CO₂, which is not directly toxic at moderate levels and instead causes harm mainly through simple oxygen displacement (asphyxiation) at very high concentrations. Because CO binds so tightly and reversibly, treatment for poisoning involves administering high-flow or hyperbaric oxygen to competitively displace CO from hemoglobin.
Hemoglobin binding and the carboxyhemoglobin curve
The affinity constant of hemoglobin for CO relative to O₂ is described by the Haldane relationship, showing why even low ambient CO concentrations (as little as 0.1% by volume) can produce dangerous carboxyhemoglobin saturation over time. Chronic low-level exposure — from a poorly vented furnace or an idling car in an attached garage — can raise COHb to symptomatic levels (headache, confusion) well before life-threatening acute poisoning, which is why CO detectors are calibrated to alarm at low sustained concentrations rather than only at acutely lethal levels.
Industrial synthesis: water-gas shift, syngas, and blast furnaces
Carbon monoxide is produced industrially through the water–gas reaction (C + H₂O → CO + H₂, at high temperature over coke) and the partial oxidation of hydrocarbons or coal to make syngas, a CO/H₂ mixture central to Fischer–Tropsch synthesis of liquid fuels and to methanol manufacture (CO + 2 H₂ → CH₃OH). In blast furnaces, CO generated in situ from coke and air reduces iron ore (Fe₂O₃ + 3 CO → 2 Fe + 3 CO₂), making CO chemistry fundamental to virtually all primary steel production.
CO vs. CO2: contrasting the two carbon oxides
CO and CO₂ are both products of carbon combustion but differ dramatically in bonding, toxicity, and environmental role. CO has a triple bond, a small dipole, and forms from incomplete combustion; CO₂ has two double bonds, is nonpolar due to its linear symmetric geometry, and forms from complete combustion. CO is acutely toxic through hemoglobin binding but is not a significant greenhouse gas at typical atmospheric concentrations, whereas CO₂ is essentially non-toxic at ambient levels but is the dominant anthropogenic driver of the greenhouse effect. In the atmosphere, CO is eventually oxidized to CO₂ by hydroxyl radicals, linking the two compounds in atmospheric chemistry.
Sources, detectors, and everyday exposure risks
Common sources of hazardous CO exposure include idling vehicles in enclosed garages, malfunctioning or unvented gas water heaters and furnaces, charcoal or propane grills used indoors, blocked chimneys, and tobacco smoke. Because CO is undetectable by human senses, electrochemical or biomimetic CO detectors are essential safety devices in homes, functioning by sensing CO's reducing chemistry (electrochemical sensors) or its effect on a hemoglobin-like pigment (biomimetic sensors) rather than relying on smell or visible smoke.
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
- NIST Chemistry WebBook: Thermophysical and spectroscopic data for carbon monoxide
- PubChem CID 281: Structural and toxicological identifiers
- CDC/NIOSH: Carbon monoxide poisoning health and safety guidance