Solute Calculators

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A solute is the substance that is dissolved in a solvent to form a solution. In aqueous solutions (the most common in biology and chemistry), water is the solvent and the dissolved substance (salt, sugar, gas, or other compound) is the solute. The concentration of a solute in solution can be expressed in many ways: molarity (mol/L), molality (mol/kg solvent), mass percent, parts per million (ppm), or normality. Solute concentration determines osmotic pressure, colligative properties (boiling point elevation, freezing point depression), reaction rates, and biological function.

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Concentration Units for Solutes

  • Molarity (M = mol/L): moles of solute per liter of solution; most common in chemistry and biology; changes with temperature
  • Molality (m = mol/kg solvent): moles of solute per kilogram of solvent; temperature-independent; used for colligative properties
  • Mass percent (w/v or w/w): grams of solute per 100 mL or 100 g of solution; common in biochemistry (e.g., 1% agarose)
  • Parts per million (ppm): mg solute per kg (or L) of solution = μg/mL for dilute aqueous solutions; used for trace concentrations in environmental science
  • Normality (N): equivalents of solute per liter; used for acid-base and redox reactions

Saturation and Solubility

Saturated solution: maximum concentration of solute dissolved at a given temperature. Solubility (S): amount of solute that dissolves in a fixed amount of solvent at equilibrium. Factors affecting solubility: temperature (most solids: solubility ↑ with T; gases: solubility ↓ with T); polarity ('like dissolves like' — polar solutes dissolve in polar solvents); pH (weak acids and bases more soluble at appropriate pH); pressure (for gases: Henry's Law — S ∝ partial pressure).

Colligative Properties

Properties dependent on solute concentration (not identity): boiling point elevation (ΔTb = i × Kb × m); freezing point depression (ΔTf = i × Kf × m); osmotic pressure (π = i × M × R × T); vapor pressure lowering. i = van't Hoff factor (= 1 for non-electrolytes; > 1 for electrolytes).

Glossary

Solute
The substance dissolved in a solvent to form a solution; usually present in smaller amount; concentration expressed as molarity (M), molality (m), mass percent, or ppm.
Molarity (M)
Moles of solute per liter of solution (mol/L); most common concentration unit in chemistry; temperature-dependent; used for stoichiometry and solution preparation.
Solubility
The maximum amount of solute that dissolves in a fixed solvent amount at a given temperature and pressure; increases with T for most solids; decreases with T for gases; affected by polarity and pH.

Frequently Asked Questions

In a solution: Solvent: the component present in the larger amount (the 'dissolving medium'); usually liquid; water is the most common biological solvent. Solute: the substance dissolved in the solvent; usually present in smaller amount; can be solid, liquid, or gas. Solution: a homogeneous mixture of solute and solvent at the molecular level. Examples: saline solution (salt + water): NaCl (solute) + H₂O (solvent). Carbonated water: CO₂ gas (solute) + H₂O (solvent). Ethanol in water: when ethanol content < 50%, water is the solvent; when > 50%, ethanol is the solvent. Blood: plasma (mostly water + dissolved proteins, ions, glucose) is the solution; blood cells are suspended (not truly dissolved).

Molarity (M = mol/L): most common; 1 M = 1 mol per liter of solution. Example: dissolve 58.44 g NaCl (MW=58.44) in water to make 1 liter → 1.00 M NaCl. Molality (m = mol/kg solvent): 1 m = 1 mol per 1 kg of solvent (not solution). Used for colligative properties because it's temperature-independent. Mass/volume percent (w/v, %): g solute per 100 mL solution × 100. Example: 1% agarose = 1 g agarose in 100 mL buffer. Parts per million (ppm): mg/kg or mg/L (for dilute aqueous solutions, ppm ≈ μg/mL ≈ mg/L). Used for trace pollutants: drinking water nitrate limit = 10 ppm = 10 mg/L.

Temperature: Most solid solutes: solubility increases with temperature (dissolving is endothermic — adding heat favors dissolution). Example: KNO₃ solubility: 13 g/100 mL at 0°C → 247 g/100 mL at 100°C. Gas solutes: solubility decreases with temperature — warm soda goes flat; warm water has less dissolved oxygen. Polarity ('like dissolves like'): polar and ionic solutes dissolve in polar solvents (water). Nonpolar solutes dissolve in nonpolar solvents (hexane, chloroform). pH: weak acids more soluble in basic solution (deprotonated form is more soluble); weak bases more soluble in acidic solution. Pressure (gases only): Henry's Law: C = k × P. Higher partial pressure of gas → higher dissolved concentration. Important for blood CO₂/O₂ transport and carbonation.

Osmotic pressure (π) is the pressure required to prevent osmosis (water movement from lower to higher solute concentration across a semipermeable membrane). π = i × M × R × T. i = van't Hoff factor (number of particles per formula unit: NaCl: i = 2; glucose: i = 1). M = molarity; R = 0.08206 L·atm/mol/K; T = temperature (K). Example: physiological saline 0.9% NaCl ≈ 0.154 M NaCl × i = 2 → 0.308 osmol/L: π = 2 × 0.154 × 0.08206 × 310 K = 7.84 atm ≈ 5,950 mmHg. This is isotonic with blood plasma (~290 mOsm/L). Hypertonic solutions (more solute than plasma) → draw water out of cells → shrinkage. Hypotonic solutions → water enters cells → swelling → lysis. Osmolarity = total solute concentration in mOsm/L.