Molar Mass of Sulfuric Acid (H₂SO₄)
Learn how chemists calculate the molar mass of Sulfuric Acid (H₂SO₄), with a clear formula breakdown, worked steps, and study notes · IUPAC name: Sulfuric acid.
Quick answer
The molar mass of Sulfuric Acid (H₂SO₄) is
98.072g/mol
One mole of Sulfuric Acid therefore has a mass of 98.072 grams—the value you use for stoichiometry and laboratory preparation.
Reviewed for educational accuracy · Accuracy policy
- CAS Registry Number
- 7664-93-9
- PubChem CID
- 1118
- SMILES
- OS(=O)(=O)O
Step-by-step calculation
Let's find the molar mass of Sulfuric Acid (H₂SO₄) 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 H₂SO₄. 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.
- 2 Hydrogen atoms (H)
- 1 Sulfur atom (S)
- 4 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.
- Hydrogen (H) = 1.008 g/mol
- Sulfur (S) = 32.060 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.
- 2 × 1.008 = 2.016 g/mol (Hydrogen)
- 1 × 32.060 = 32.060 g/mol (Sulfur)
- 4 × 15.999 = 63.996 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:
2.016 + 32.060 + 63.996 = 98.072 g/mol
Step 5 — Final answer
Molar mass of Sulfuric Acid = 98.072 g/mol
That means one mole of Sulfuric Acid (H₂SO₄) has a mass of about 98.07 grams.
Quick summary
Read the formula → count atoms → look up atomic masses → multiply → add → report g/mol. For H₂SO₄, the total is 98.072 g/mol.
Common beginner mistakes
- Treating H₂SO₄ as monoprotic in stoichiometry (needs 2 equivalents of base for complete neutralization).
- Adding water to concentrated acid instead of acid to water.
- Forgetting that concentrated H₂SO₄ has a high boiling point (~337 °C) due to hydrogen bonding.
Memory trick
Memorize 98 g/mol — close to 100 — for quick mental stoichiometry.
Mini practice
Without looking above, list the atoms in H₂SO₄ 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 Sulfuric Acid, mass needed = 0.100 × 98.072 = 9.807 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 H₂SO₄.
| Element | Atoms | Atomic mass | Contribution | Mass % |
|---|---|---|---|---|
| H | 2 | 1.008 | 2.016 g/mol | 2.1% |
| S | 1 | 32.060 | 32.060 g/mol | 32.7% |
| O | 4 | 15.999 | 63.996 g/mol | 65.3% |
| Total molar mass | 98.072 g/mol | 100% | ||
Mass contribution chart
In water, many acids transfer H⁺ (a proton) to form hydronium. Molar mass still comes from the undissociated formula used to prepare the solution.
Download study sheets
Save a printable summary, revision sheet, practice worksheet, or laboratory reference for Sulfuric Acid (H₂SO₄).
Practice this calculation
Without looking above, write the atom count for H₂SO₄, then compute the molar mass. Check your answer against 98.072 g/mol.
Next challenge: how many grams are in 0.250 mol of Sulfuric Acid? Multiply 0.250 × 98.072 to get 24.518 g.
Physical and chemical properties
Physical properties
| Appearance | Colorless to slightly yellow oily liquid (concentrated); hygroscopic |
| Color | Colorless (pure) |
| Odor | Odorless (pure); pungent when heated or contaminated |
| State (STP) | Liquid (high boiling point prevents gaseous state at STP under normal conditions) |
| Density | 1.84 g/cm³ (98% solution, 25 °C) |
| Melting point | 10.4 °C (98% acid) |
| Boiling point | 337 °C (98% acid, decomposes at higher temperatures) |
| Solubility | Miscible with water in all proportions (highly exothermic) |
| Crystal structure | Forms hydrates; pure H₂SO₄ has hydrogen-bonded network in liquid state |
Chemical properties
| Classification | Mineral acid / strong diprotic acid |
| Family | Oxoacids of sulfur (highest oxidation state) |
| Acidity | Strong acid (pKa₁ ≈ −3, pKa₂ = 1.99) |
| Polarity | Highly polar |
| Geometry | Tetrahedral about sulfur |
| Bond angle | ~109.5° (approximate, tetrahedral S center) |
| Oxidation states | S: +6, O: −2, H: +1 |
Applications
Industrial uses
- Manufacture of phosphate fertilizers (e.g., converting rock phosphate to superphosphate)
- Petroleum refining alkylation (e.g., isooctane production)
- Lead-acid car batteries (dilute ~30–50% H₂SO₄ as electrolyte)
- Metal pickling and steel cleaning
- Rayon and synthetic fiber production
Laboratory uses
- Strong acid for titrations and pH adjustment
- Dehydrating agent in organic preparations (with caution)
- Electrolyte in electrolysis demonstrations
Acid rain contains dilute sulfuric acid from atmospheric oxidation of SO₂; scrubbers in power plants neutralize flue gas SO₂ before emission.
Not biological in beneficial sense; corrosive to tissue. Sulfate ion (SO₄²⁻) from dissociation is a normal metabolic anion in the body.
Preparation and production
The contact process dominates industrial production: burn sulfur to SO₂, catalytically oxidize to SO₃ (V₂O₅ catalyst, 400–450 °C), then absorb SO₃ into 98% H₂SO₄ to form oleum (H₂S₂O₇), which is diluted to desired concentration. Laboratory small amounts: react Na₂SO₄ with concentrated HCl (not preferred) or purchase reagent grade.
Global production via contact process from elemental sulfur, pyrite (FeS₂), or recovered SO₂ from smelting. Spent acid is often regenerated by decomposition and re-concentration.
Important reactions of Sulfuric Acid
H₂SO₄(aq) + 2 NaOH(aq) → Na₂SO₄(aq) + 2 H₂O(l)
- Reaction type
- Acid–base neutralization
- Conditions
- Aqueous, stoichiometric base
- Explanation
- Both protons of sulfuric acid are neutralized by strong base, forming sulfate salt and water.
- Products
- Sodium sulfate and water
- Why it matters
- Acid–base titrations, wastewater neutralization
Related ideas: Diprotic acids · Equivalence point · Stoichiometry
2 SO₂(g) + O₂(g) ⇌ 2 SO₃(g)
- Reaction type
- Catalytic oxidation (contact process step)
- Conditions
- V₂O₅ catalyst, 450 °C, 1–2 atm
- Explanation
- Sulfur dioxide is oxidized to sulfur trioxide, the key intermediate in industrial H₂SO₄ manufacture.
- Products
- Sulfur trioxide
- Why it matters
- Industrial sulfuric acid production
Related ideas: Le Chatelier's principle · Heterogeneous catalysis · Industrial chemistry
Cu(s) + 2 H₂SO₄(conc) → CuSO₄(aq) + SO₂(g) + 2 H₂O(l)
- Reaction type
- Redox (hot concentrated acid)
- Conditions
- Hot concentrated H₂SO₄
- Explanation
- Copper does not release H₂ with cold dilute acid but is oxidized by hot concentrated acid to Cu²⁺ with SO₂ as reduced sulfur product.
- Products
- Copper(II) sulfate, sulfur dioxide, water
- Why it matters
- Demonstrating oxidizing power of concentrated H₂SO₄
Related ideas: Redox · Activity series exceptions · Oxidation state changes
C₁₂H₂₂O₁₁(s) → 12 C(s) + 11 H₂O(l) [catalyzed by H₂SO₄]
- Reaction type
- Dehydration
- Conditions
- Concentrated H₂SO₄, room temperature
- Explanation
- Sucrose is dehydrated to elemental carbon and water; the reaction is dramatic and exothermic, producing a rising carbon column.
- Products
- Carbon and water
- Why it matters
- Classic demonstration of dehydrating power
Related ideas: Dehydration · Concentration effects · Organic chemistry demonstrations
History and discovery
Alchemists knew "oil of vitriol" from roasting green vitriol (FeSO₄·7H₂O). Joshua Ward produced it commercially in England (~1740). The contact process replaced the lead chamber process in the 20th century, improving purity and enabling higher concentrations for industrial demand.
Known since antiquity in impure form; pure composition established by Lavoisier's systematic chemistry and later confirmed by molecular structure studies showing tetrahedral sulfate.
Interesting facts
- More sulfuric acid is produced worldwide than any other chemical.
- A car battery contains roughly 4–5 kg of sulfuric acid electrolyte.
- Concentrated H₂SO₄ can dissolve a paper towel by dehydrating cellulose to carbon.
- The molar mass ~98 g/mol makes 1 L of concentrated acid (~18 mol/L) contain about 18 moles.
Comparison with similar compounds
H₂SO₄ (98.08 g/mol) is diprotic and a dehydrating agent; HCl (36.46 g/mol) is monoprotic and lacks significant dehydrating or oxidizing character when dilute. Unlike SO₂ (64.07 g/mol, S in the +4 state), sulfuric acid holds sulfur in its fully oxidized +6 state, making H₂SO₄ the thermodynamic 'end point' of atmospheric and industrial sulfur oxidation chemistry.
Storage, handling, and safety
Store in corrosion-resistant containers (glass, polyethylene, or passivated steel for concentrated acid). Keep tightly closed, dry, and separated from bases, organic materials, and metals. Secondary containment is mandatory for bulk storage.
Wear acid-resistant gloves, face shield, and apron. Work in fume hood for concentrated acid. Always add acid to water slowly with stirring. Never add water to concentrated H₂SO₄. Have spill neutralization materials (sodium bicarbonate, vermiculite) available.
Highly corrosive; causes severe skin burns, eye damage, and respiratory irritation. Exothermic dilution poses splashing hazard.
- Corrosive to skin, eyes, and respiratory tract
- Violent reaction with water if water is added to concentrated acid
- Dehydrating and charring of organic matter
- Reacts with metals to produce flammable hydrogen gas
Classification: GHS: Skin Corr. 1A, Eye Dam. 1
Exam notes and student tips
Exam notes
- Molar mass H₂SO₄ = 2(1.008) + 32.07 + 4(16.00) = 98.08 g/mol.
- Diprotic: 1 mol H₂SO₄ can provide up to 2 mol H⁺ if fully neutralized (e.g., with 2 mol NaOH).
- S oxidation state in H₂SO₄ is +6; in SO₂ it is +4; track changes in redox problems.
- For titration: H₂SO₄ + 2 NaOH → Na₂SO₄ + 2 H₂O (2:1 base-to-acid mole ratio).
Student tips
- Memorize 98 g/mol — close to 100 — for quick mental stoichiometry.
- Link to contact process equations for industrial chemistry questions.
- Use oxidation state +6 for S in sulfate to balance redox half-reactions.
Common mistakes
- Treating H₂SO₄ as monoprotic in stoichiometry (needs 2 equivalents of base for complete neutralization).
- Adding water to concentrated acid instead of acid to water.
- Forgetting that concentrated H₂SO₄ has a high boiling point (~337 °C) due to hydrogen bonding.
Misconceptions
- Dilute H₂SO₄ does not dehydrate sugar the way concentrated acid does — concentration matters enormously.
- Sulfuric acid is not the same as sulfurous acid (H₂SO₃, sulfur in +4 state).
- Not all metals react with H₂SO₄ the same way; nitric acid reactions differ due to oxidizing ability.
Practice questions
1. Calculate the molar mass of H₂SO₄.
Show answer
2(1.008) + 32.07 + 4(16.00) = 98.08 g/mol
2. How many moles of NaOH are needed to neutralize 49 g of H₂SO₄ completely?
Hint: H₂SO₄ is diprotic — 2 mol NaOH per mol acid.
Show answer
49 g ÷ 98.08 g/mol = 0.50 mol H₂SO₄; needs 2 × 0.50 = 1.0 mol NaOH
3. What is the oxidation state of sulfur in H₂SO₄?
Show answer
+6 (balance: 2(+1) + x + 4(−2) = 0 → x = +6)
4. Why must acid be added to water, not water to acid?
Show answer
The hydration of H₂SO₄ releases large amounts of heat; adding water to acid can cause localized boiling and violent spattering of corrosive acid.
Frequently asked questions about Sulfuric Acid
98.08 g/mol.
Chemistry of Sulfuric Acid
The sections above give the number you need for calculations. Here we look more closely at how Sulfuric Acid (H₂SO₄) behaves chemically—so the molar mass connects to real reactions, properties, and laboratory practice.
Sulfuric acid (H₂SO₄) is the world's most produced industrial chemical, with annual output exceeding 230 million tonnes. Its molar mass of 98.08 g/mol comes from two hydrogen atoms, one sulfur (32.07 g/mol), and four oxygen atoms. A viscous, colorless, odorless liquid when pure, concentrated H₂SO₄ is 98% by mass and has a density of about 1.84 g/cm³ — nearly twice that of water.
H₂SO₄ is a diprotic strong acid: the first proton dissociates completely in dilute aqueous solution (Ka₁ ≈ 10³, effectively infinite for most calculations), while the second is weak (Ka₂ = 1.2 × 10⁻², corresponding to HSO₄⁻ ⇌ H⁺ + SO₄²⁻). Concentrated sulfuric acid is a powerful dehydrating agent, removing water from organic compounds as H and O atoms, leaving carbonaceous char — a property exploited in the contact process and feared in laboratory accidents. Its affinity for water is so extreme that dilution must always add acid to water, never the reverse, to avoid violent spattering from the heat of hydration.
H₂SO₄ contains sulfur in the +6 oxidation state, bonded to four oxygen atoms (two as S=O double bonds and two as S–O–H groups in the tetrahedral sulfate geometry). The formula indicates two ionizable protons, though only the first fully dissociates in typical strong-acid contexts.
Sulfuric acid is a strong Brønsted acid, oxidizing agent when hot and concentrated, and dehydrating agent. It protonates alcohols to alkyl sulfates, nitrates aromatics indirectly by generating HNO₃ in mixed acid, and converts metals below hydrogen in the activity series (except Cu, Hg, Ag, Pt, Au cold) to sulfates with H₂ evolution. Hot concentrated H₂SO₄ oxidizes copper to CuSO₄ with SO₂ evolution.
The contact process: how the world's top industrial chemical is made
The contact process converts elemental sulfur (or sulfide ore) into sulfuric acid through three carefully controlled stages: combustion of sulfur to SO₂, catalytic oxidation of SO₂ to SO₃ over a vanadium pentoxide catalyst at 400–450 °C (chosen as a compromise between reaction rate and the exothermic equilibrium's temperature sensitivity), and absorption of SO₃ into concentrated sulfuric acid to form oleum, which is then diluted with water to the desired final concentration. This carefully staged process — rather than direct SO₃ absorption into water, which would produce a dangerous, uncontrollable acid mist — exemplifies how industrial chemistry manages both thermodynamics and practical engineering safety simultaneously.
Battery acid: sulfuric acid in the lead-acid cell
In a lead-acid car battery, dilute sulfuric acid (roughly 30–50% by mass) serves as both electrolyte and active reactant: during discharge, lead reacts with sulfate ions and the positive plate's lead dioxide is reduced, both processes consuming sulfuric acid and producing lead sulfate and water, which is why a battery's electrolyte density (and thus its measured specific gravity) drops as it discharges. Recharging reverses this reaction, regenerating sulfuric acid and metallic lead — making the state of charge of a lead-acid battery directly and quantitatively linked to the concentration of sulfuric acid in its electrolyte.
Dehydration power: charring sugar and organic matter
Concentrated sulfuric acid is one of the most powerful chemical dehydrating agents known, capable of removing water directly from organic molecules rather than simply absorbing ambient moisture. Its affinity for water is so strong that it can strip hydrogen and oxygen atoms out of carbohydrates like sucrose in the exact 2:1 ratio needed to form water molecules, leaving behind a rising black column of nearly pure carbon — a visually dramatic classroom demonstration that also illustrates why concentrated H₂SO₄ causes severe chemical burns on skin, since it dehydrates living tissue by the same mechanism.
Corrosion and materials chemistry
Sulfuric acid's corrosivity toward metals, concrete, and biological tissue arises from a combination of its strong Brønsted acidity (attacking basic or reactive materials directly) and, when hot and concentrated, its oxidizing power (attacking even relatively unreactive metals like copper). This dual corrosive mechanism is why materials selection for sulfuric acid storage and transport — often specialized alloys, glass-lined steel, or certain plastics — must account for both acid concentration and temperature, since dilute cold H₂SO₄ and hot concentrated H₂SO₄ present quite different corrosion challenges to the same piece of equipment.
Industrial production scale and economic importance
Sulfuric acid production, at over 230 million tonnes annually, exceeds that of nearly every other manufactured chemical, and its output is often used by economists as a rough proxy indicator of a country's industrial and manufacturing activity, since it feeds into fertilizer production, metal processing, petroleum refining, and countless other downstream industries. This makes sulfuric acid a foundational 'workhorse' chemical whose production statistics are tracked globally as an economic and industrial health indicator in their own right.
Recalculate any formula with the molar mass calculator, compare atoms on the periodic table, or browse more compounds in the acid library.
References and further reading
- USGS Mineral Commodity Summaries: Global sulfuric acid production statistics
- NIST Chemistry WebBook: Thermodynamic and physical properties
- PubChem CID 1118: Structural and safety data