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Freezing Point Depression Calculator

Ready to calculate
ΔT_f = K_f · m · i.
10 Solvent Presets.
°C · °F · K Units.
100% Free.
No Data Stored.

How it Works

01Enter Molality

Moles of solute per kg of solvent — the colligative concentration unit (does NOT depend on temperature)

02Pick a Solvent Preset

Auto-fills K_f and pure-solvent T_f for water, benzene, ethanol, chloroform, and 6 other common solvents

03Set Van't Hoff Factor

i = 1 for non-electrolytes; i = 2 for NaCl; i = 3 for CaCl₂ (number of dissolved ions per formula unit)

04ΔT_f = K_f · m · i

Get the depression magnitude and the new solution freezing point in your chosen unit (°C / °F / K)

What is a Freezing Point Depression Calculator?

Freezing point depression is one of the four classical colligative properties — properties of solutions that depend only on the number of dissolved particles, not their chemical identity. The relation ΔTf = Kf × m × i was published by François-Marie Raoult in 1882 and remains the textbook tool for cryoscopic molar-mass determination, antifreeze formulation, ice-cream chemistry, road de-icing science, and biological cryoprotection. Our Freezing Point Depression Calculator implements this equation with 10 pre-loaded laboratory solvents (water, benzene, ethanol, chloroform, diethyl ether, nitrobenzene, acetic acid, camphor, naphthalene, cyclohexane), full van't Hoff factor support for non-electrolytes and ionic compounds, and three temperature systems (°C, °F, K) with proper delta-vs-absolute conversion handling.

Just enter the molality (moles solute per kg solvent), select a solvent (or enter custom values), and the calculator instantly returns the depression magnitude ΔTf and the new solution freezing point Tf(solution) = Tf(pure) − ΔTf. Our cryoscopic-constant database uses standard CRC Handbook values: water Kf = 1.86 K·kg/mol, benzene 5.12, ethanol 1.99, chloroform 4.68, camphor 39.7 (the highest among common solvents — the historical reason camphor was used by Beckmann for cryoscopic molar-mass determination). The van't Hoff factor i lets you correctly model dissolved ionic compounds: i ≈ 1 for non-electrolytes (sugar, urea, ethylene glycol), i ≈ 2 for 1:1 salts (NaCl, KBr), i ≈ 3 for 2:1 or 1:2 salts (CaCl₂, K₂SO₄, MgCl₂).

Designed for general chemistry students learning colligative properties, biochemistry students working with cryoprotection of cells (DMSO, glycerol), automotive engineers formulating coolants, food scientists controlling ice crystal formation in frozen desserts, environmental chemists studying brine de-icing of roads, and pharmaceutical scientists using cryoscopy to verify drug molecular weights, the tool runs entirely in your browser — no data is stored or transmitted.

Pro Tip: Pair this with our Molar Mass Calculator to convert grams to moles for the molality calculation, or our Molarity Calculator if you need to convert between molarity and molality for dilute aqueous solutions.

How to Use the Freezing Point Depression Calculator?

Enter Molality (m): Moles of solute per kilogram of solvent — the colligative concentration unit. Unlike molarity, molality is temperature-independent (mass doesn't change with T). Example: 1.0 mol of glucose dissolved in 1.0 kg of water = 1.0 mol/kg.
Pick a Solvent Preset: Choose from water (Kf = 1.86), benzene (5.12), ethanol (1.99), chloroform (4.68), diethyl ether (1.79), nitrobenzene (7.00), acetic acid (3.90), camphor (39.7), naphthalene (6.94), or cyclohexane (20.0). The cryoscopic constant Kf and pure-solvent freezing point Tf auto-fill. Pick "Custom" to override either value.
Verify or Override Kf: The cryoscopic constant has units of temperature × kg / mol. Choose its temperature unit (°C, °F, or K) — magnitudes per kelvin equal magnitudes per °C, but °F values are 9/5 larger. Editing Kf automatically switches the solvent to "Custom".
Verify or Override Tf (pure): Pure-solvent freezing point — water is 0 °C, benzene 5.5 °C, ethanol −114.6 °C, etc. Choose any of three temperature units; the calculator converts internally to kelvin.
Set Van't Hoff Factor (i): Defaults to 1. Use 1 for non-electrolytes (sugar, urea, glycerol, ethylene glycol). Use 2 for 1:1 strong electrolytes (NaCl → Na⁺ + Cl⁻). Use 3 for CaCl₂, K₂SO₄, MgBr₂. Real values are slightly less than ideal due to ion pairing.
Press Calculate: Get ΔTf in your chosen output unit, the solution freezing point Tf(solution) = Tf(pure) − ΔTf, a 5-band magnitude classification (negligible → extreme), and a full calculation breakdown.

How is freezing point depression calculated?

Freezing point depression is one of four colligative properties (the others: boiling point elevation, vapor pressure lowering, osmotic pressure) that all share the same fundamental physics — solute particles disrupt the ideal solid-liquid equilibrium of a pure solvent. Here's the complete derivation:

In 1882 François-Marie Raoult discovered that a wide variety of solutes lowered the freezing point of water by an amount proportional to molality, with the proportionality constant being a property of the solvent alone. This result earned him the Davy Medal of the Royal Society and laid the foundation of cryoscopy.

The Formula

For dilute solutions:

ΔTf = Kf × m × i

where ΔTf is the magnitude of freezing-point depression (always positive — depression means the new Tf is LOWER), Kf is the cryoscopic constant (also called molal freezing point depression constant) of the solvent, m is molality of the solute (mol/kg), and i is the van't Hoff factor (effective number of dissolved particles per formula unit).

The Cryoscopic Constant Kf

Kf is a thermodynamic property of the SOLVENT (not the solute) given by:

Kf = R · Tf² · M / (1000 · ΔHfus)

where R is the gas constant, Tf is the pure-solvent freezing point in K, M is solvent molar mass in g/mol, and ΔHfus is solvent enthalpy of fusion in J/mol. For water: Tf = 273.15 K, M = 18.015, ΔHfus = 6010 J/mol → Kf = 8.314 × 273.15² × 18.015 / (1000 × 6010) = 1.86 K·kg/mol ✓.

The Van't Hoff Factor i

For non-electrolytes that dissolve as intact molecules: i = 1. For strong electrolytes that fully dissociate:

  • NaCl → Na⁺ + Cl⁻: i = 2 (ideal); ~1.87 measured at 0.1 M (some ion pairing)
  • CaCl₂ → Ca²⁺ + 2 Cl⁻: i = 3 ideal; ~2.7 measured at 0.1 M
  • K₂SO₄ → 2 K⁺ + SO₄²⁻: i = 3 ideal; ~2.3 measured
  • Sugar, urea, ethylene glycol: i = 1 (non-electrolytes)
  • Acetic acid (weak): i ≈ 1.01-1.05 (only ~1% dissociation in water)

Why Solute Lowers the Freezing Point

At equilibrium between solid and liquid solvent, the chemical potentials must match: μ(solid) = μ(liquid in solution). Adding solute lowers μ(liquid) (it dilutes the solvent) but doesn't change μ(solid) (the solid is pure crystalline solvent — solute is excluded). To restore equilibrium, the system must lower temperature, which raises μ(liquid) more than μ(solid) (because liquid has more entropy). The new equilibrium temperature is below Tf(pure) by ΔTf.

Solution Freezing Point

Tf(solution) = Tf(pure) − ΔTf

Cryoscopic Molar Mass Determination

Rearranging to solve for solute molar mass M:

M = (1000 · Kf · wsolute) / (ΔTf · wsolvent · i)

where wsolute is grams of solute and wsolvent is grams of solvent. This was THE classical method for determining unknown molar masses before mass spectrometry — Ernst Beckmann's apparatus (1888) using camphor as the solvent (high Kf = 39.7) was the gold standard for organic-compound analysis through ~1960.

Limits of the Equation

  • Dilute regime only. Strictly valid for m → 0; works well for m < 0.5 mol/kg in most solvents.
  • Solute insolubility in solid solvent. Assumes solute does not enter the crystal lattice (true for most cases — pure ice forms even from salty seawater).
  • Ideal-solution behavior. No specific solute-solvent interactions, no ion pairing, no association.
  • Constant pressure (1 atm). The formula doesn't account for pressure effects on Tf.
Real-World Example

Freezing Point Depression Calculator – Worked Examples

Consider the classic textbook problem: add 100 g of ethylene glycol (HOCH₂CH₂OH, M = 62.07 g/mol) to 500 g of water. Find the new freezing point. Inputs: m = (100/62.07)/(0.500) = 3.222 mol/kg, Kf(water) = 1.86, Tf(pure) = 0 °C, i = 1 (ethylene glycol is non-electrolyte).
  • Step 1 — Compute molality: moles ethylene glycol = 100 / 62.07 = 1.611 mol. Molality = 1.611 / 0.500 = 3.222 mol/kg.
  • Step 2 — Apply the formula: ΔTf = Kf × m × i = 1.86 × 3.222 × 1 = 5.99 K = 5.99 °C.
  • Step 3 — New freezing point: Tf(solution) = 0 − 5.99 = −5.99 °C = 21.2 °F.
  • Comparison: Real automotive coolant is 50:50 ethylene glycol:water by volume — about 8 mol/kg — giving ~36 °C depression and freezing around −36 °C, which is why your car radiator stays liquid in winter. Pure water would freeze at 0 °C and burst the engine.

Now consider road de-icing with NaCl: 100 g of salt in 500 g of water. M(NaCl) = 58.44 g/mol. Inputs: m = (100/58.44)/0.500 = 3.422 mol/kg, Kf = 1.86, i = 2 (NaCl → Na⁺ + Cl⁻).

  • ΔTf = 1.86 × 3.422 × 2 = 12.73 K.
  • Tf(solution) = 0 − 12.73 = −12.73 °C.
  • Note that using twice as much salt by mass as ethylene glycol gives only twice the depression — but salt is much cheaper. That's why road salt (NaCl, MgCl₂, CaCl₂) is the universal de-icer on highways.
  • CaCl₂ for the same mass: 100 g / 110.98 = 0.901 mol → m = 1.802; with i = 3, ΔTf = 1.86 × 1.802 × 3 = 10.06 K. CaCl₂ is preferred for very cold conditions because it's still effective at lower T (its eutectic is −51 °C vs −21 °C for NaCl).

Cryoscopic molar mass example: dissolve 1.50 g of an unknown organic compound in 25 g of camphor (Kf = 39.7, Tf = 178.4 °C). Measured Tf(solution) = 174.0 °C. Find M.

  • ΔTf = 178.4 − 174.0 = 4.4 °C.
  • m = ΔTf / Kf = 4.4 / 39.7 = 0.1108 mol/kg.
  • moles solute = m × kg solvent = 0.1108 × 0.025 = 0.00277 mol.
  • M = 1.50 g / 0.00277 mol = 541 g/mol. Likely a small natural product or polypeptide. This is exactly how Beckmann's apparatus was used pre-MS to characterize organic compounds.

Who Should Use the Freezing Point Depression Calculator?

1
General Chemistry Students: Solve textbook colligative-property problems on freezing-point depression — the canonical first introduction to chemical thermodynamics.
2
Automotive Engineers: Formulate coolants — calculate exact ethylene glycol:water ratios for desired freezing protection (typical: 50:50 → −36 °C; 70:30 → −51 °C).
3
Highway / Airport De-Icing: Estimate NaCl, CaCl₂, MgCl₂ application rates needed to keep pavement liquid at given low temperatures; comparing salt cost vs effectiveness.
4
Food Scientists: Control ice crystal formation in ice cream, sorbet, and gelato (sugar concentration determines softness — higher m = more depression = softer texture at freezer T).
5
Cryobiologists: Design cryoprotection protocols for cells (DMSO, glycerol, ethylene glycol) — the right freezing-point depression prevents intracellular ice crystallization that would rupture membranes.
6
Environmental Chemists: Predict freezing behavior of natural brines (sea ice forms at −1.9 °C because salinity ~35 ppt = ~0.6 mol/kg salts), study Antarctic dry valleys with hyper-saline lakes that don't freeze.

Technical Reference

Raoult's Original Work: F.-M. Raoult, "Loi de congélation des solutions aqueuses" (Law of freezing of aqueous solutions), Comptes Rendus Acad. Sci. Paris 94, 1517 (1882). Raoult systematically measured ΔTf for hundreds of organic and inorganic solutes in water and discovered the now-famous proportionality. He won the Davy Medal in 1892 and the Faraday Medal in 1899 for this and related work on vapor pressure (Raoult's law).

Beckmann's Apparatus (1888). Ernst Beckmann developed the high-precision differential thermometer (Beckmann thermometer, ±0.001 K) and the cryoscopic apparatus for molar-mass determination. Camphor (Kf = 39.7) was preferred because its huge Kf means even very small molality gives easily measurable ΔTf — perfect for high-molar-mass natural products. The Beckmann method dominated organic-compound molecular weight determination from 1888 until reliable mass spectrometry emerged in the 1950s-60s.

Cryoscopic Constants Kf for Common Solvents (CRC Handbook, K·kg/mol):

  • Water: 1.86 (Tf = 0.00 °C)
  • Benzene: 5.12 (Tf = 5.50 °C)
  • Ethanol: 1.99 (Tf = −114.6 °C)
  • Chloroform: 4.68 (Tf = −63.5 °C)
  • Diethyl ether: 1.79 (Tf = −116.3 °C)
  • Nitrobenzene: 7.00 (Tf = 5.7 °C)
  • Acetic acid: 3.90 (Tf = 16.6 °C)
  • Cyclohexane: 20.0 (Tf = 6.5 °C)
  • Naphthalene: 6.94 (Tf = 80.2 °C)
  • Camphor: 39.7 (Tf = 178.4 °C — highest among common solvents)

Connection to Other Colligative Properties. Freezing point depression is mathematically parallel to boiling point elevation (ΔTb = Kb·m·i) and osmotic pressure (Π = MRTi). All three quantify the same underlying effect: solute lowers the chemical potential of the solvent. Vapor pressure lowering (Raoult's law: ΔP = P°·xsolute·i) is the underlying cause. For water at 1 m solute (i = 1): ΔTf ≈ 1.86 K, ΔTb ≈ 0.51 K, Π ≈ 24 atm — the osmotic pressure is enormous compared to the temperature shifts, which is why osmometry is the most sensitive cryometric technique.

Real-World Applications:

  • Automotive coolant (50/50 ethylene glycol/water): ΔTf ≈ 36 °C → freezes at −36 °C; protects engine to about −20 °C in routine use.
  • Road salt (NaCl): Effective to about −10 °C; below that, CaCl₂ or MgCl₂ are used (eutectic of CaCl₂ + water = −51 °C).
  • Sea ice: Seawater (~3.5% salt by mass) freezes at −1.9 °C; sea ice itself is much fresher than seawater because salt is excluded from the ice lattice.
  • Ice cream: Sugar (sucrose ~30% by mass) and milk solids depress freezing to ~−2.5 °C, keeping the dessert soft and scoopable in a typical home freezer (−18 °C).
  • Cryopreservation: 10% DMSO in cell culture medium → ΔTf ≈ 2.4 K, which slows ice nucleation enough to prevent intracellular crystallization during freezing.

Sources of Error. (1) Ion pairing — real i is less than ideal i for concentrated electrolytes (NaCl 0.1 M: i_observed ≈ 1.87 vs ideal 2.00). (2) Activity coefficients — molality should strictly be replaced with molal activity for non-ideal solutions. (3) Solvent purity — trace impurities can shift Tf(pure) by 0.1-1 °C. (4) Solute insolubility in solid solvent — most solutes don't enter the crystal lattice (good), but if they do (solid solution), the formula breaks down. (5) Volatile solute — the formula assumes the solute stays dissolved; volatile solutes that partially evaporate will show smaller ΔTf.

Key Takeaways

Freezing point depression is the colligative property captured by the simple equation ΔTf = Kf × m × i. It depends only on the NUMBER of dissolved particles, not their identity. Three things to remember: Kf is a property of the SOLVENT (water = 1.86, benzene = 5.12, camphor = 39.7); m is molality (mol solute per kg solvent, NOT volume-based); i is the van't Hoff factor (1 for non-electrolytes, 2 for NaCl, 3 for CaCl₂). The new freezing point is Tf(pure) − ΔTf. Use the ToolsACE Freezing Point Depression Calculator with 10 pre-loaded solvents, full van't Hoff handling, and three temperature units (°C, °F, K) to solve any cryoscopy problem instantly. Bookmark it for chemistry homework, antifreeze formulation, road-salt calculations, food and cryobiology applications, and any time you need to know how much a solute lowers the freezing point.

Frequently Asked Questions

What is the Freezing Point Depression Calculator?
It computes the freezing-point depression ΔTf = Kf × m × i — a colligative property that depends only on the number of solute particles, not their chemical identity. Inputs: molality m (mol/kg), solvent (10 presets — water, benzene, ethanol, chloroform, ether, nitrobenzene, acetic acid, camphor, naphthalene, cyclohexane), cryoscopic constant Kf (auto-fills from solvent), pure-solvent freezing point Tf (auto-fills), and van't Hoff factor i (1 for non-electrolytes, 2 for NaCl, 3 for CaCl₂).

Output: ΔTf in your chosen unit (°C, °F, or K), the new solution freezing point Tf(solution) = Tf(pure) − ΔTf, a 5-band magnitude classification (negligible → extreme), full step-by-step calculation breakdown, and a comparison table of cryoscopic constants for common solvents. Designed for general chemistry students, automotive engineers, food scientists, cryobiologists, and pharmaceutical scientists. Runs entirely in your browser — no data stored.

Pro Tip: Use our Molar Mass Calculator to convert grams to moles for the molality calculation.

What's the formula for freezing point depression?
ΔTf = Kf × m × i, where ΔTf is the magnitude of depression (always positive — depression means LOWER freezing point), Kf is the cryoscopic constant (a property of the SOLVENT, units K·kg/mol or °C·kg/mol — same magnitude), m is molality (mol solute per kg solvent), and i is the van't Hoff factor (number of dissolved particles per formula unit). The new solution freezing point is Tf(solution) = Tf(pure) − ΔTf.
What is the cryoscopic constant K_f and where do values come from?
Kf is a thermodynamic property of the solvent given by Kf = R·Tf²·M / (1000·ΔHfus), where R is the gas constant, Tf is the pure solvent's freezing point in K, M is solvent molar mass (g/mol), and ΔHfus is enthalpy of fusion (J/mol). For water this gives 1.86 K·kg/mol. The biggest Kf values come from solvents with high Tf, low ΔHfus, and high M — camphor (Kf = 39.7) is the standard example, which is why Beckmann's pre-MS molar-mass determinations used camphor.
What's the van't Hoff factor and why does it matter?
The van't Hoff factor i is the effective number of dissolved particles produced when one formula unit of solute dissolves. For non-electrolytes (sugar, urea, ethylene glycol): i = 1. For 1:1 strong electrolytes (NaCl → Na⁺ + Cl⁻): i = 2. For salts that produce 3 ions (CaCl₂ → Ca²⁺ + 2Cl⁻; K₂SO₄ → 2K⁺ + SO₄²⁻): i = 3. Real measured i is slightly less than ideal because of ion pairing in solution — for 0.1 M NaCl, observed i ≈ 1.87 (not 2.00); for 0.1 M CaCl₂, i ≈ 2.7 (not 3.0).
Why molality and not molarity?
Because molality (mol/kg solvent) is temperature-independent. Mass doesn't change with T, but volume does — so molarity (mol/L solution) drifts as the system cools toward freezing. The freezing-point depression formula needs a measure of solute concentration that's stable across the temperature change of interest. For dilute aqueous solutions at room temperature molarity ≈ molality (within ~5%), but for non-aqueous solvents or concentrated solutions the difference matters significantly.
Can I use this calculator for ice melting or de-icing?
Absolutely — that's one of the most common applications. To find how much salt to add to lower the freezing point of an ice patch by, say, 5 °C: ΔTf = 5 K = 1.86 × m × 2 (for NaCl, i = 2). Solve for m = 1.34 mol/kg = 1.34 × 58.44 / 1000 = 0.078 g salt per g of water = 78 g per kg. To lower further, use CaCl₂ (i = 3): m = 5/(1.86 × 3) = 0.90 mol/kg → only 53 g per kg. CaCl₂ is more effective per mole because it produces 3 particles instead of 2.
How is this used to determine an unknown molar mass?
The cryoscopic method (Beckmann's apparatus, 1888): dissolve a known mass of unknown solute in a known mass of solvent with high Kf (camphor is best). Measure ΔTf precisely with a Beckmann thermometer (±0.001 K). Calculate molality: m = ΔTf / (Kf · i). Then moles solute = m × kg solvent, and molar mass = grams solute / moles solute. Pre-mass-spectrometry, this was THE standard method for molecular weight determination — Hermann Staudinger used it to confirm that polymers were really long-chain molecules, work that earned him the 1953 Nobel Prize.
Why does adding solute lower the freezing point at all?
Thermodynamically: at the freezing point, chemical potential of pure solid equals that of liquid. Adding solute lowers liquid chemical potential (it dilutes the solvent) but doesn't change solid potential (solute is excluded from the crystal lattice). To restore equilibrium, T must decrease, which raises liquid potential more than solid (because liquid has higher entropy). The new equilibrium temperature is below Tf(pure). Intuitively: solute particles get in the way of solvent molecules trying to organize into a crystal lattice, requiring lower temperature to overcome that disorder.
Does the type of solute matter?
Surprisingly, no — that's the meaning of "colligative". Equal moles of any solute produce the same ΔTf as long as the van't Hoff factor is the same. 1 mol/kg of glucose gives the same depression as 1 mol/kg of urea or 1 mol/kg of ethylene glycol — all i = 1, so ΔTf = 1.86 K in water. NaCl at 0.5 mol/kg gives the same depression as glucose at 1.0 mol/kg, because i = 2 doubles the particle count. The solute identity matters only through (a) molar mass (to convert grams to moles) and (b) van't Hoff factor (to count dissolved particles).
What are the limits of this formula?
Strictly valid only in the dilute, ideal regime — typically m < 0.5 mol/kg for accurate predictions (errors < 5%). For concentrated solutions, real ΔTf deviates because: (1) activity coefficients depart from 1 (need molal activity, not molality); (2) ion pairing lowers effective i for strong electrolytes; (3) solvent-solute interactions (H-bonding, hydration) become non-ideal. For practical use up to m ~ 5 mol/kg the formula gives values within ~30% — fine for engineering estimates (antifreeze, road salt) but not for analytical-grade work. Use Debye-Hückel theory or experimental tabulations for high precision in concentrated solutions.
Can the calculator handle weak electrolytes like acetic acid?
Yes — set i to the effective dissociation factor. Acetic acid in dilute aqueous solution (~0.1 M) is only ~1% dissociated, so i ≈ 1.01 (essentially 1). Ammonia in water has Kb ~ 1.8 × 10⁻⁵, also only slightly dissociated → i ≈ 1.02. For weak electrolytes you need to compute the actual dissociation fraction α from K_a or K_b at the relevant concentration: i = 1 + α(ν − 1), where ν is the ideal-dissociation factor. For most weak acids/bases at typical lab concentrations, treating them as i = 1 introduces < 5% error.

Author Spotlight

The ToolsACE Team - ToolsACE.io Team

The ToolsACE Team

Our chemistry tools team implements the colligative freezing-point-depression equation ΔT_f = K_f · m · i — the textbook relation François-Marie Raoult published in 1882 that founded the entire field of cryoscopy. The calculator handles 10 common laboratory solvents (water, benzene, ethanol, diethyl ether, chloroform, nitrobenzene, acetic acid, camphor, naphthalene, cyclohexane) with their CRC-Handbook cryoscopic constants and pure-solvent freezing points pre-loaded, plus a Custom mode for any other solvent. The van't Hoff factor i is exposed so you can correctly handle non-electrolytes (i = 1: sugar, urea, ethylene glycol), 1:1 strong electrolytes (i = 2: NaCl, KBr), and higher-charge salts (i = 3: CaCl₂, K₂SO₄). All three temperature systems are supported: Celsius, Fahrenheit, and Kelvin, with separate unit selection for the cryoscopic constant K_f and the absolute freezing point T_f. Output includes ΔT_f, the new solution freezing point T_f(solution) = T_f(pure) − ΔT_f, a 5-band classification of depression magnitude, and a complete step-by-step calculation breakdown.

Colligative PropertiesSolution ThermodynamicsSoftware Engineering Team

Disclaimer

The equation ΔT_f = K_f · m · i strictly applies only to dilute, ideal solutions. For concentrated solutions, real depression can deviate 10-30% (need Debye-Hückel corrections). Real van't Hoff factors for strong electrolytes are slightly less than ideal due to ion pairing (NaCl 0.1 M: i_observed ≈ 1.87 vs 2.00). Cryoscopic constants are CRC Handbook reference values; real solvents may show small variation due to purity. The formula assumes solute is excluded from the solid solvent lattice and solvent doesn't participate in chemical reactions with solute.