Molar Mass of Potassium Chloride (KCl)
Learn how chemists calculate the molar mass of Potassium Chloride (KCl), with a clear formula breakdown, worked steps, and study notes · IUPAC name: Potassium chloride.
Quick answer
The molar mass of Potassium Chloride (KCl) is
74.548g/mol
One mole of Potassium Chloride therefore has a mass of 74.548 grams—the value you use for stoichiometry and laboratory preparation.
Reviewed for educational accuracy · Accuracy policy
- CAS Registry Number
- 7447-40-7
- PubChem CID
- 4873
- SMILES
- [K+].[Cl-]
Step-by-step calculation
Let's find the molar mass of Potassium Chloride (KCl) 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 KCl. 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 Potassium atom (K)
- 1 Chlorine atom (Cl)
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.
- Potassium (K) = 39.098 g/mol
- Chlorine (Cl) = 35.450 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 × 39.098 = 39.098 g/mol (Potassium)
- 1 × 35.450 = 35.450 g/mol (Chlorine)
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:
39.098 + 35.450 = 74.548 g/mol
Step 5 — Final answer
Molar mass of Potassium Chloride = 74.548 g/mol
That means one mole of Potassium Chloride (KCl) has a mass of about 74.55 grams.
Quick summary
Read the formula → count atoms → look up atomic masses → multiply → add → report g/mol. For KCl, the total is 74.548 g/mol.
Common beginner mistakes
- Assuming KCl tastes identical to NaCl — it has a distinct bitter, metallic aftertaste that limits its use as a full 1:1 salt substitute.
- Confusing potassium chloride (KCl) with potassium chlorate (KClO₃) — chlorate is an oxidizer used in different applications and has very different hazards.
- Underestimating the cardiac danger of concentrated or rapid IV potassium administration — this is a well-known clinical error to avoid, not just a laboratory chemistry detail.
Memory trick
Draw a direct comparison table between KCl and NaCl (molar mass, ions released, flame color, biological role) to master alkali halide comparison questions.
Mini practice
Without looking above, list the atoms in KCl 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 Potassium Chloride, mass needed = 0.100 × 74.548 = 7.455 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 KCl.
| Element | Atoms | Atomic mass | Contribution | Mass % |
|---|---|---|---|---|
| K | 1 | 39.098 | 39.098 g/mol | 52.4% |
| Cl | 1 | 35.450 | 35.450 g/mol | 47.6% |
| Total molar mass | 74.548 g/mol | 100% | ||
Mass contribution chart
Teal and blue circles alternate like positive and negative ions in a crystal lattice. Formula mass is the mass of one formula unit, not a single molecule.
Download study sheets
Save a printable summary, revision sheet, practice worksheet, or laboratory reference for Potassium Chloride (KCl).
Practice this calculation
Without looking above, write the atom count for KCl, then compute the molar mass. Check your answer against 74.548 g/mol.
Next challenge: how many grams are in 0.250 mol of Potassium Chloride? Multiply 0.250 × 74.548 to get 18.637 g.
Physical and chemical properties
Physical properties
| Appearance | White or colorless crystalline solid |
| Color | White (pure); may show pink/red tint from mineral impurities in natural sylvite |
| Odor | Odorless |
| State (STP) | Solid |
| Density | 1.98 g/cm³ (25 °C) |
| Melting point | 770 °C |
| Boiling point | 1,420 °C |
| Solubility | 342 g/L water at 20 °C (increases notably with temperature) |
| Crystal structure | Face-centered cubic (rock salt structure, same as NaCl) |
Chemical properties
| Classification | Ionic salt / alkali metal halide |
| Family | Group 1 halide (alkali metal chloride) |
| Polarity | Ionic (no molecular dipole in the crystal lattice) |
| Oxidation states | K: +1, Cl: −1 |
Applications
Industrial uses
- Muriate of potash (MOP) fertilizer, the dominant global source of agricultural potassium
- Feedstock for producing potassium hydroxide and other potassium chemicals
- Component of drilling fluids in the oil and gas industry
- De-icing agent (less common than NaCl or CaCl2 due to higher cost)
Laboratory uses
- Standard reference electrolyte for reference electrodes (e.g., saturated KCl salt bridge)
- Source of K⁺ ion in qualitative analysis and flame test demonstrations
- Isotonic and physiological solution preparation in biochemistry
Potash mining generates saline tailings that require careful management to avoid contaminating nearby soil and waterways with excess chloride and salinity.
Essential electrolyte for nerve conduction and cardiac muscle function; used clinically to treat hypokalemia and, in tightly controlled dosing, certain cardiac conditions; critical macronutrient for plant growth as fertilizer potassium.
Preparation and production
Industrially, KCl is separated from mined sylvinite ore (a mixture of KCl and NaCl) by exploiting the differing solubility–temperature behavior of the two salts, using selective crystallization, flotation, or electrostatic separation. It can also be recovered from potassium-rich brines by solar evaporation.
Global potash production is dominated by Canada, Russia, and Belarus, which together supply the majority of the world's potassium fertilizer; total production exceeds 40 million tonnes of KCl-equivalent annually.
Important reactions of Potassium Chloride
KCl(aq) + AgNO₃(aq) → AgCl(s) + KNO₃(aq)
- Reaction type
- Precipitation (double displacement)
- Conditions
- Aqueous, room temperature
- Explanation
- Silver ion combines with chloride to form insoluble silver chloride, the same qualitative chloride test used for any soluble chloride salt including KCl.
- Products
- Silver chloride precipitate and potassium nitrate solution
- Why it matters
- Qualitative analysis for chloride ion, gravimetric chloride determination
Related ideas: Solubility rules · Precipitation reactions · Qualitative analysis
2 KCl(l) → 2 K(l) + Cl₂(g)
- Reaction type
- Electrolysis (molten salt)
- Conditions
- Molten KCl, direct current, inert electrodes
- Explanation
- Electrolysis of molten potassium chloride yields potassium metal at the cathode and chlorine gas at the anode, analogous to the Downs process for sodium.
- Products
- Potassium metal and chlorine gas
- Why it matters
- Alternative potassium metal production route, electrolysis teaching example
Related ideas: Electrolysis · Redox · Ionic compounds
KCl(aq) + NaOH(aq) → no reaction
- Reaction type
- Non-reaction (spectator ion demonstration)
- Conditions
- Aqueous, room temperature
- Explanation
- Because both possible product combinations (KOH and NaCl) are soluble strong electrolytes, no precipitate or gas forms — a useful negative example when teaching solubility rules and net ionic equations.
- Products
- No net reaction; all ions remain in solution as spectators
- Why it matters
- Teaching solubility rules, contrasting reactive vs. non-reactive ion combinations
Related ideas: Spectator ions · Solubility rules · Net ionic equations
KOH(aq) + HCl(aq) → KCl(aq) + H₂O(l)
- Reaction type
- Acid–base neutralization
- Conditions
- Aqueous, room temperature
- Explanation
- Strong base and strong acid neutralize completely to form the neutral salt potassium chloride and water, a textbook example of a strong acid–strong base titration.
- Products
- Potassium chloride and water
- Why it matters
- Acid–base titration standardization, teaching neutral salt formation
Related ideas: Neutralization · Strong acid–strong base titration · Salt hydrolysis (absence of)
History and discovery
Potassium salts have been used since antiquity in the form of potash (crude potassium carbonate leached from wood ashes), the source of the element's name. Pure potassium chloride was characterized chemically in the 19th century as mineralogists identified sylvite deposits and industrial chemists began separating it from associated sodium chloride. Its importance grew enormously in the 20th century with the rise of intensive agriculture and demand for concentrated potassium fertilizers.
Potash-derived potassium compounds were used since antiquity; pure KCl was characterized as 19th-century mineralogy and industrial separation techniques matured alongside the discovery and mining of sylvite deposits.
Interesting facts
- KCl and NaCl look and taste almost identical as crystals but differ enough in flavor perception that most people notice KCl's bitter aftertaste at high concentration.
- Potassium's flame test color (lilac) is much fainter than sodium's intense orange, requiring a blue glass filter to observe clearly when sodium is also present.
- Canada's Saskatchewan potash deposits are among the largest in the world, supplying a major share of global fertilizer potassium.
- KCl solutions are used in reference electrodes (e.g., saturated calomel and silver/silver chloride electrodes) because K⁺ and Cl⁻ diffuse through salt bridges at closely matched rates, minimizing junction potential errors.
Comparison with similar compounds
KCl (74.55 g/mol) and NaCl (58.44 g/mol) share the same rock-salt crystal structure and 1:1 ionic ratio, but KCl's larger, bitterer-tasting K⁺ ion makes it a physiologically distinct nutrient used for potassium fertilization and supplementation rather than as a direct 1:1 seasoning replacement.
Storage, handling, and safety
Store in dry, sealed containers, as KCl is mildly hygroscopic and can cake if exposed to humid air. Keep fertilizer-grade and pharmaceutical-grade material separated to avoid contamination given their very different purity requirements.
Low acute hazard for handling the solid; avoid dust inhalation and eye contact with concentrated solutions. Medical-grade KCl for injection must always be diluted and administered under strict clinical protocols due to cardiac risk from concentrated or rapid dosing.
Low hazard as a solid salt for industrial and agricultural handling, but concentrated or rapidly administered medical KCl solutions pose serious, potentially fatal cardiac risk if mismanaged.
- Eye and mild skin irritation from concentrated solutions or dust
- Serious cardiac arrhythmia risk from rapid or concentrated intravenous administration (medical context)
- Ingestion of large doses can cause gastrointestinal upset and, in extreme cases, hyperkalemia
Classification: Not classified as hazardous under GHS for the pure solid (agricultural/industrial grade)
Exam notes and student tips
Exam notes
- Molar mass KCl = 39.10 + 35.45 = 74.55 g/mol.
- KCl is a strong electrolyte: 1 mol dissolved → 1 mol K⁺ + 1 mol Cl⁻ (van't Hoff factor i ≈ 2, same pattern as NaCl).
- Flame test: potassium gives a lilac/violet flame, distinct from sodium's orange-yellow — useful for qualitative cation identification.
- KCl does not hydrolyze in water (salt of strong acid HCl and strong base KOH), so its aqueous solution is neutral (pH ≈ 7).
Student tips
- Draw a direct comparison table between KCl and NaCl (molar mass, ions released, flame color, biological role) to master alkali halide comparison questions.
- Remember 'K stays inside, Na stays outside' as a mnemonic for potassium's intracellular vs. sodium's extracellular physiological dominance.
- Link the lilac flame test result to the general pattern that heavier alkali metals show progressively different emission colors down the group.
Common mistakes
- Assuming KCl tastes identical to NaCl — it has a distinct bitter, metallic aftertaste that limits its use as a full 1:1 salt substitute.
- Confusing potassium chloride (KCl) with potassium chlorate (KClO₃) — chlorate is an oxidizer used in different applications and has very different hazards.
- Underestimating the cardiac danger of concentrated or rapid IV potassium administration — this is a well-known clinical error to avoid, not just a laboratory chemistry detail.
Misconceptions
- Potassium chloride is not more dangerous than sodium chloride in normal dietary or agricultural use — its serious risks arise specifically from concentrated medical/IV administration, not typical salt-substitute consumption.
- KCl and table salt are not simply interchangeable 1:1 in cooking due to taste differences, even though both are simple alkali halides.
- Muriate of potash (KCl fertilizer) is not the same as 'potash' in the broadest historical sense, which originally referred to potassium carbonate leached from wood ash.
Practice questions
1. Calculate the molar mass of potassium chloride, KCl.
Show answer
39.10 + 35.45 = 74.55 g/mol
2. How many grams of KCl are needed to prepare 250 mL of a 0.100 M solution?
Show answer
0.100 mol/L × 0.250 L = 0.0250 mol; 0.0250 × 74.55 = 1.86 g
3. What color flame does a potassium salt like KCl produce, and why might it be hard to observe if sodium is also present?
Show answer
A pale lilac flame; sodium's much more intense orange-yellow emission can mask the fainter potassium color unless viewed through a cobalt blue filter.
4. Why must concentrated KCl never be injected rapidly as an undiluted IV push?
Show answer
A sudden spike in extracellular K⁺ disrupts the resting membrane potential of cardiac cells, halting normal electrical conduction and potentially stopping the heart.
Frequently asked questions about Potassium Chloride
74.55 g/mol, calculated as 39.10 (K) + 35.45 (Cl).
Chemistry of Potassium Chloride
The sections above give the number you need for calculations. Here we look more closely at how Potassium Chloride (KCl) behaves chemically—so the molar mass connects to real reactions, properties, and laboratory practice.
Potassium chloride (KCl) is an ionic salt with a molar mass of 74.55 g/mol, composed of K⁺ (39.10 g/mol) and Cl⁻ (35.45 g/mol) in a 1:1 ratio. It crystallizes, like sodium chloride, in a face-centered cubic rock-salt lattice, and the two compounds share many physical similarities — both are white, water-soluble, cubic crystalline solids — yet they diverge sharply in biological role and taste. While Na⁺ is the dominant cation of extracellular fluid, K⁺ is the dominant intracellular cation, and this fundamental physiological difference underlies both KCl's importance as a fertilizer nutrient and its serious risks when used medically or nutritionally without care.
Potassium is one of the three primary macronutrients (alongside nitrogen and phosphorus) required for plant growth, and KCl — known in the fertilizer industry as muriate of potash (MOP) — is the most widely used potassium fertilizer worldwide, supplying the majority of potassium applied to global cropland. It is mined from ancient evaporite deposits, most notably as the mineral sylvite, often found interlayered with sodium chloride (halite) in massive potash deposits formed from evaporated inland seas.
In human physiology, potassium ion is essential for nerve impulse transmission and cardiac muscle function, maintained within a narrow serum range (roughly 3.5–5.0 mmol/L). This narrow safe range makes KCl a double-edged compound: as an oral or IV supplement it corrects potassium deficiency (hypokalemia) and treats certain cardiac arrhythmias, but rapid intravenous administration or gross overdose disrupts cardiac electrical activity and can be fatal — a property that has made concentrated KCl infusion a component of lethal injection protocols, underscoring how the same ion essential to a heartbeat can also stop one if mismanaged.
The formula KCl reflects a simple 1:1 ionic ratio between the alkali metal cation K⁺ and the halide anion Cl⁻, directly parallel to NaCl. As with all alkali halides, there is no discrete covalent molecule in the solid state — the formula unit describes the smallest whole-number ratio of ions in an extended ionic lattice held together by electrostatic (Coulombic) attraction.
KCl is chemically similar to NaCl: it is a strong electrolyte that dissociates completely in water, does not hydrolyze appreciably (a salt of a strong acid and strong base), and participates in standard precipitation and metathesis reactions typical of alkali halides. Molten KCl can be electrolyzed to produce potassium metal and chlorine gas, analogous to the Downs process for sodium, though potassium's higher reactivity and lower melting point make this route less commercially favored than other potassium extraction methods.
KCl as fertilizer: muriate of potash and global agriculture
Potassium chloride, marketed in agriculture as muriate of potash (MOP), supplies the majority of the world's fertilizer potassium, a nutrient essential for plant water regulation, enzyme activation, and disease resistance. Unlike nitrogen, which plants incorporate into organic molecules, potassium remains as the free K⁺ ion within plant tissue, functioning primarily in osmotic and electrochemical roles. Global potash mining and trade — concentrated in Canada, Russia, and Belarus — makes KCl one of the most strategically important bulk commodities in world agriculture.
KCl vs. NaCl: the low-sodium salt substitute
Because K⁺ and Na⁺ are both monovalent alkali cations of similar ionic size, KCl can substitute for NaCl in many chemical contexts, including as a seasoning for people needing to reduce dietary sodium (marketed as 'lite salt,' often as a 50:50 KCl/NaCl blend). However, KCl has a distinctly more bitter, metallic aftertaste than NaCl due to differences in how the two ions interact with taste receptors, which limits the concentration at which it can fully replace table salt without altering flavor.
Medical potassium supplementation and cardiac risk
KCl is a critical medical electrolyte replacement for hypokalemia (low blood potassium), which can arise from diuretic use, vomiting, diarrhea, or poor dietary intake, and hypokalemia itself can cause dangerous cardiac arrhythmias. Yet potassium's narrow therapeutic window means intravenous KCl must always be diluted and infused slowly under monitoring — undiluted or rapid IV push of concentrated KCl disrupts the resting membrane potential of cardiac cells, halting the heart's electrical conduction. This dual life-saving/lethal nature makes KCl a frequently cited case study in pharmacology and toxicology courses on the importance of concentration and administration rate.
Flame test signature: the lilac potassium color
When a potassium salt like KCl is introduced into a flame, excited potassium electrons emit a characteristic pale lilac (light violet) color as they fall back to lower energy levels, a signature distinct from sodium's intense orange-yellow flame. Because sodium contamination is common and its flame color is so much more intense, chemists often view potassium flame tests through a cobalt blue glass filter to block interfering sodium emission and reveal the fainter lilac potassium signal underneath.
Sylvite, sylvinite, and evaporite potash deposits
KCl occurs naturally as the mineral sylvite, frequently intermixed with halite (NaCl) in a rock called sylvinite, both formed by evaporation of ancient inland seas and salt lakes over geological time. Massive potash deposits — such as those in Saskatchewan, Canada, one of the world's largest — are mined and processed to separate KCl from co-occurring NaCl using differences in solubility and crystallization behavior at different temperatures, a large-scale industrial application of fractional crystallization principles.
Recalculate any formula with the molar mass calculator, compare atoms on the periodic table, or browse more compounds in the salt library.
References and further reading
- USGS Mineral Commodity Summaries: Global potash mining and production statistics
- NIST Chemistry WebBook: Thermophysical properties of potassium chloride
- PubChem CID 4873: Structural and safety identifiers