What is molarity and how is it calculated?

Molarity (M) is the number of moles of solute dissolved per litre of solution: M = n(solute) / V(solution in litres). For example, dissolving 58.44 g of NaCl (molar mass = 58.44 g/mol → 1 mol) in enough water to make 1.00 L of solution gives a 1.00 M NaCl solution. Molarity is temperature-dependent because liquid volume changes with temperature. It is the most common concentration unit in analytical and general chemistry.

What is molality and how does it differ from molarity?

Molality (m) is the number of moles of solute per kilogram of solvent: m = n(solute) / m(solvent in kg). Unlike molarity, molality uses mass of solvent (not volume of solution) and is therefore temperature-independent and pressure-independent. This makes molality the preferred concentration unit for calculations involving colligative properties (boiling point elevation, freezing point depression, osmotic pressure) that depend on the number of solute particles, not on volume.

What is percent by volume (v/v)?

Percent by volume (volume/volume, v/v%) is the volume of solute divided by the total volume of solution, multiplied by 100: % volume = (V_solute / V_solution) × 100. It is most commonly used for liquid-in-liquid solutions where volumes are most convenient to measure. Note: volumes of mixed liquids are not always additive (they can expand or contract slightly on mixing), so the total volume of solution should be measured after mixing, not calculated by addition. Example: beer at 5% v/v contains 5 mL of ethanol per 100 mL of beer.

What does ppm mean in chemistry?

Parts per million (ppm) means one part of solute per million parts of solution by mass: ppm = (m_solute / m_solution) × 10⁶. For dilute aqueous solutions (density ≈ 1 kg/L), 1 ppm ≈ 1 mg/L ≈ 1 mg/kg. PPM is used to express trace-level concentrations in environmental monitoring (pollutants in water or air), food safety (additives, contaminants), and materials science (impurities in metals and semiconductors). Example: the World Health Organisation guideline for arsenic in drinking water is 10 ppb (10 μg/L).

What is the molar mass and why is it needed?

Molar mass (M_w) is the mass of one mole of a substance, in grams per mole (g/mol). It equals the sum of atomic masses of all atoms in the formula unit. Examples: NaCl = 23.0 + 35.5 = 58.5 g/mol; H₂SO₄ = 2(1.0) + 32.1 + 4(16.0) = 98.1 g/mol; glucose C₆H₁₂O₆ = 6(12.0) + 12(1.0) + 6(16.0) = 180.0 g/mol. Molar mass is needed to convert between mass of solute (in grams) and moles of solute for molarity and molality calculations.

What are colligative properties?

Colligative properties are physical properties of a solution that depend on the number of dissolved solute particles (moles), not on the identity of the solute. The four main colligative properties are: (1) Vapour pressure lowering: ΔP = X_solute × P°_solvent (Raoult's Law); (2) Boiling point elevation: ΔTb = Kb × m, where Kb is the ebullioscopic constant (~0.512 °C·kg/mol for water); (3) Freezing point depression: ΔTf = Kf × m, where Kf is the cryoscopic constant (~1.853 °C·kg/mol for water); (4) Osmotic pressure: π = MRT, where R = 8.314 J/(mol·K). Because all these formulas use molality m or molarity M, accurately calculating solution concentration is the essential first step.

What is the difference between solute, solvent, and solution?

A solution is a homogeneous mixture of two or more substances. The solvent is the substance present in the larger amount (or the substance that does the dissolving) — water is the universal solvent. The solute is the substance dissolved in the solvent — it is usually present in smaller amount. Example: in saline (salt water), water is the solvent and NaCl is the solute; together they form the solution. The total mass of the solution = mass of solute + mass of solvent.

Why is molality preferred over molarity for colligative property calculations?

Molality is preferred because it is temperature-independent. Since molality is defined in terms of mass of solvent (not volume of solution), it does not change when temperature changes — unlike molarity, which changes as solution volume expands or contracts with temperature. When measuring boiling point elevation or freezing point depression, the experiment involves temperature change by definition — so using a concentration unit that is stable across temperatures (molality) gives more reliable results than using one that changes (molarity).

Solution Concentration Calculator — Molarity, Molality, % Mass, % Volume, Mole Fraction, PPM & PPB

Chemistry Molarity Molality Concentration PPM/PPB Mole Fraction

Whether you are preparing a buffer solution in a chemistry lab, calculating drug concentrations in a pharmacy, measuring pollutant levels in a water sample, or studying for a university exam, understanding and calculating solution concentration is one of the most essential skills in all of the chemical sciences. Yet with seven different concentration units in common use — each suited to a different context — the terminology can be confusing.

The HeLovesMath Solution Concentration Calculator handles all seven units in one tool: Molarity (M), Molality (m), Percent by Mass (% w/w), Percent by Volume (% v/v), Mole Fraction (X), Parts Per Million (ppm), and Parts Per Billion (ppb). Solve for any variable — concentration, moles, mass, or volume — with built-in unit conversion and a full step-by-step solution. This guide explains every concept and formula from first principles, with all mathematical expressions properly rendered.

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Solutions, Solutes & Solvents — The Building Blocks

A solution is a homogeneous mixture of two or more chemically distinct substances that are uniformly distributed at the molecular or ionic level. The two main components are:

Solvent — the component present in the largest amount, or the one into which others dissolve. Water (H₂O) is called the "universal solvent" because it dissolves more substances than any other liquid.
Solute — the dissolved substance, present in smaller amount. Solutes can be solid (e.g., table salt NaCl), liquid (e.g., ethanol), or gas (e.g., CO₂ in carbonated water).

✦ Fundamental Mass Relationship

\[m_{\text{solution}} = m_{\text{solute}} + m_{\text{solvent}}\]

This basic mass balance underlies every concentration calculation. Note that volumes are NOT always additive (mixing 500 mL ethanol + 500 mL water gives ~953 mL, not 1000 mL, due to intermolecular interactions), but masses always are.

💧 Aqueous Solutions

Solutions in which water is the solvent. Most biochemistry, medicine, and environmental chemistry involves aqueous solutions. Water's high polarity makes it excellent at dissolving ionic compounds (salts) and polar molecules (sugars, alcohols).

🧫 Non-Aqueous Solutions

Solutions using organic solvents: ethanol, acetone, hexane, dichloromethane. Used in organic synthesis, extraction, and industrial processes. "Like dissolves like" — polar solvents dissolve polar solutes; non-polar solvents dissolve non-polar solutes.

🔬 Electrolytic Solutions

Solutions in which the solute dissociates into ions (electrolytes). Strong electrolytes (NaCl, HCl, NaOH) fully dissociate; weak electrolytes (CH₃COOH, NH₃) partially dissociate. The degree of dissociation affects colligative properties and conductivity.

Molarity — The Universal Lab Concentration Unit

Molarity (symbol M, unit mol/L or M) is defined as the number of moles of solute dissolved per litre of total solution. It is by far the most widely used concentration unit in laboratory chemistry, biochemistry, and medicine because solution volumes are easy to measure accurately with volumetric glassware (flasks, pipettes, burettes).

✦ Molarity

\[M = \frac{n_{\text{solute}}}{V_{\text{solution}}} \qquad \text{(mol/L)}\]

M = molarity (mol/L, also written molar or simply M) | n_solute = moles of solute (mol) = mass (g) / molar mass (g/mol) | V_solution = volume of the total solution (L), not just the solvent

✦ Rearrangements of the Molarity Formula

\[n_{\text{solute}} = M \times V \qquad V = \frac{n_{\text{solute}}}{M} \qquad m_{\text{solute}} = M \times V \times M_w\]

M_w = molar mass of solute (g/mol) | These rearrangements allow you to find moles, volume, or mass given the other two quantities.

To prepare a 1.00 M NaCl solution (molar mass = 58.44 g/mol) in 500 mL: calculate the required mass → m = M × V × M_w = 1.00 × 0.500 × 58.44 = 29.22 g. Dissolve 29.22 g NaCl in about 400 mL of water in a 500 mL volumetric flask, then add water to the 500 mL mark and mix.

⚠️ Critical Note: Always dissolve the solute in a smaller volume of solvent first, then dilute to the final volume mark in a volumetric flask. Never fill to the mark with solvent and then add solute — the volume change on dissolving would give the wrong concentration.

Molality — Temperature-Independent Concentration

Molality (symbol m, unit mol/kg) is the number of moles of solute per kilogram of solvent. Unlike molarity (which uses volume of solution and changes with temperature), molality uses mass of solvent — making it temperature-independent and pressure-independent.

✦ Molality

\[m = \frac{n_{\text{solute}}}{m_{\text{solvent (kg)}}}\]

m = molality (mol/kg, also written molal) | n_solute = moles of solute | m_solvent = mass of the solvent in kilograms (NOT the total solution mass) | Note: solute mass is NOT included — only the mass of the dissolving medium (solvent).

Molality is used whenever temperature changes are involved (boiling point elevation, freezing point depression) because it does not depend on volume. The molality of a 1 M aqueous NaCl solution can be calculated if we know the solution density ρ:

✦ Converting Molarity M to Molality m

\[m = \frac{M}{\rho - \dfrac{M \cdot M_w}{1000}}\]

ρ = density of solution in g/mL | M = molarity in mol/L | M_w = molar mass in g/mol | The denominator gives the mass of solvent (in kg) per litre of solution.

Percent by Mass (% w/w) — Simple, Scale-Independent

✦ Percent by Mass (w/w)

\[\%_{\text{mass}} = \frac{m_{\text{solute}}}{m_{\text{solution}}} \times 100\%\]

m_solute = mass of solute | m_solution = total mass of solution = m_solute + m_solvent | Both masses must be in the same unit (g, kg, mg — they cancel). Result is a dimensionless percentage.

Percent by mass is temperature-independent (masses do not change with temperature), making it ideal for industrial specifications, food labelling (e.g., 10% w/w acetic acid in vinegar), pharmaceutical manufacturing, and reagent bottle labels (concentrated acids are often labelled by % mass and density).

✦ Converting % Mass to Molarity

\[M = \frac{10 \times \rho \times \%_{\text{mass}}}{M_w}\]

ρ = solution density in g/mL | %_mass = percentage (e.g., 36 for 36%) | M_w = molar mass in g/mol | The factor of 10 converts mL to L and % to fraction simultaneously.

Percent by Volume (% v/v) — Liquid Mixtures

✦ Percent by Volume (v/v)

\[\%_{\text{vol}} = \frac{V_{\text{solute}}}{V_{\text{solution}}} \times 100\%\]

V_solute = volume of solute | V_solution = total volume of the final solution (measured after mixing, not the sum of component volumes) | Both volumes in the same unit (mL, L). Result is a dimensionless percentage.

Percent by volume is most naturally used for liquid-in-liquid solutions. Alcoholic beverages use % v/v: wine at 12% v/v contains 12 mL of pure ethanol per 100 mL of wine. Hand sanitiser at 70% v/v ethanol contains 70 mL ethanol per 100 mL product.

📐 Non-Additivity of Volumes: When two miscible liquids mix, the total volume of the mixture is generally slightly different from the sum of the individual volumes. For example, 500 mL of ethanol + 500 mL of water ≠ 1000 mL (it gives about 953 mL due to hydrogen bonding). This is why % v/v should always be calculated from the measured volume of the final solution, not from adding solute and solvent volumes.

Mole Fraction — Dimensionless Thermodynamic Concentration

✦ Mole Fraction

\[X_{\text{solute}} = \frac{n_{\text{solute}}}{n_{\text{solute}} + n_{\text{solvent}}} = \frac{n_{\text{solute}}}{n_{\text{total}}}\] \[X_{\text{solvent}} = 1 - X_{\text{solute}}\]

X = mole fraction (dimensionless, 0 ≤ X ≤ 1) | n_solute, n_solvent = moles of solute and solvent respectively | All mole fractions in a solution always sum to 1: X_solute + X_solvent = 1 (for a two-component system).

Mole fraction is fundamental in thermodynamics. Raoult's Law for ideal solutions states that the vapour pressure of a component equals the mole fraction times the pure vapour pressure:

✦ Raoult's Law (Vapour Pressure Lowering)

\[P_{\text{solvent}} = X_{\text{solvent}} \times P^{\circ}_{\text{solvent}}\] \[\Delta P = P^{\circ}_{\text{solvent}} - P_{\text{solvent}} = X_{\text{solute}} \times P^{\circ}_{\text{solvent}}\]

P°_solvent = vapour pressure of pure solvent | P_solvent = vapour pressure of solvent in solution | ΔP = vapour pressure lowering (always positive for a non-volatile solute)

Mole fraction is also used in Dalton's Law of Partial Pressures for gas mixtures: the partial pressure of each gas in a mixture equals its mole fraction times the total pressure.

Parts Per Million (ppm) & Parts Per Billion (ppb) — Trace Concentrations

When solute concentrations are extremely low — trace pollutants in water, heavy metals in soil, impurities in semiconductor materials — the familiar units like molarity become inconveniently small (e.g., 1 × 10⁻⁶ mol/L). Parts per million (ppm) and parts per billion (ppb) provide more intuitive expressions of these tiny concentrations.

✦ Parts Per Million

\[\text{ppm} = \frac{m_{\text{solute}}}{m_{\text{solution}}} \times 10^6 = \frac{\text{mg solute}}{\text{kg solution}}\]

For dilute aqueous solutions (density ≈ 1 g/mL ≈ 1 kg/L): 1 ppm ≈ 1 mg/L ≈ 1 mg/kg ≈ 1 μg/mL. 1 ppm = 0.0001% = 10,000 parts per billion.

✦ Parts Per Billion

\[\text{ppb} = \frac{m_{\text{solute}}}{m_{\text{solution}}} \times 10^9 = \frac{\mu\text{g solute}}{\text{kg solution}}\]

For dilute aqueous solutions: 1 ppb ≈ 1 μg/L ≈ 1 μg/kg. | 1 ppb = 0.001 ppm. | Regulatory limits: WHO arsenic in water: 10 ppb; EPA lead in drinking water action level: 15 ppb.

💡 Intuitive Scale: 1 ppm is 1 second in 11.6 days. 1 ppb is 1 second in 31.7 years. 1 ppt (part per trillion, 10⁻¹²) is 1 second in 31,700 years. Environmental chemistry routinely measures chemicals at ppb or even ppt levels.

Dilution Formula — C₁V₁ = C₂V₂

Dilution is the process of reducing concentration by adding more solvent. Since moles of solute are conserved during dilution:

✦ Dilution Formula

\[C_1 V_1 = C_2 V_2\]

C₁ = initial (stock) concentration | V₁ = initial (stock) volume | C₂ = final (diluted) concentration | V₂ = final (diluted) volume | C must be in the same units for both sides. Works for molarity, % mass, ppm, and any other concentration unit.

This formula is critical in laboratory work. To prepare 100 mL of 0.1 M HCl from concentrated HCl (12 M stock): V₁ = C₂V₂/C₁ = (0.1 mol/L × 0.100 L) / 12 mol/L = 8.33 × 10⁻⁴ L = 0.833 mL of stock acid, diluted to 100 mL.

⚠️ Safety Rule — "Add Acid to Water": When diluting concentrated acids (HCl, H₂SO₄, HNO₃), ALWAYS add the acid slowly to water, never add water to the acid. The intense exothermic heat of dilution can cause water to boil and spray corrosive acid if water is added to concentrated acid.

Colligative Properties — Why Concentration Matters

Colligative properties are physical properties of solutions that depend on the number of dissolved solute particles, not on the chemical identity of the solute. All four major colligative properties use concentration (molality m or molarity M) in their formulas.

✦ Boiling Point Elevation

\[\Delta T_b = K_b \times m \times i\]

Kb = ebullioscopic constant (water: 0.512 °C·kg/mol) | m = molality of solution | i = van't Hoff factor (number of particles per formula unit: 1 for glucose, 2 for NaCl → Na⁺ + Cl⁻, 3 for CaCl₂ → Ca²⁺ + 2Cl⁻)

✦ Freezing Point Depression

\[\Delta T_f = K_f \times m \times i\]

Kf = cryoscopic constant (water: 1.853 °C·kg/mol; benzene: 5.12 °C·kg/mol) | This is why salt lowers the freezing point of water (roads in winter) and antifreeze lowers the freezing point of engine coolant.

✦ Osmotic Pressure

\[\pi = iMRT\]

π = osmotic pressure (Pa or atm) | M = molarity of solution | R = 8.314 J/(mol·K) | T = temperature in Kelvin | Osmotic pressure drives water across cell membranes and is the basis for reverse osmosis water purification.

Comparison of All Seven Concentration Units

Unit	Symbol	Formula	Temperature-Independent?	Best Used For
Molarity	M	n_solute / V_solution (L)	No (volume changes)	Lab work, titrimetry, reactions
Molality	m	n_solute / m_solvent (kg)	Yes	Colligative properties, thermodynamics
% Mass (w/w)	%w/w	(m_sol/m_soln) × 100	Yes	Industrial, food, pharmacy specs
% Volume (v/v)	%v/v	(V_sol/V_soln) × 100	No (volume changes)	Liquid-liquid solutions, beverages
Mole Fraction	X	n_i / n_total	Yes	Raoult's Law, gas mixtures, thermodynamics
PPM	ppm	(m_sol/m_soln) × 10⁶	Yes	Environmental, trace analysis, regulations
PPB	ppb	(m_sol/m_soln) × 10⁹	Yes	Ultra-trace analysis, contamination levels

Worked Examples — Step by Step

Example 1 — Molarity: Preparing 250 mL of 0.500 M NaOH

Molar mass of NaOH: Na(23.0) + O(16.0) + H(1.0) = 40.0 g/mol.

Moles required: n = M × V = 0.500 mol/L × 0.250 L = 0.125 mol NaOH.

Mass required: m = n × M_w = 0.125 mol × 40.0 g/mol = 5.00 g NaOH.

Weigh 5.00 g NaOH, dissolve in ~150 mL water in a 250 mL volumetric flask, cool to room temperature, then add water to the 250 mL graduation mark.

✅ 0.500 M NaOH: weigh 5.00 g and dilute to 250.0 mL. Enter into the calculator: Type = Molarity, Solve for = Mass of Solute, M = 0.500, V = 250 mL, Molar Mass = 40.0 g/mol.

Example 2 — Molality: Antifreeze Solution for Freezing Point Depression

We want to freeze-protect to −10 °C. ΔTf = Kf × m → 10 = 1.853 × m → m = 10/1.853 = 5.40 mol/kg (for non-electrolyte like ethylene glycol, i=1).

Target: 5.40 molal ethylene glycol (M_w = 62.07 g/mol) in 1.00 kg of water.

Mass of ethylene glycol = n × M_w = 5.40 mol × 62.07 g/mol = 335 g.

Mix 335 g of ethylene glycol with 1,000 g (1.000 kg) of water: this gives a 5.40 molal solution, protecting to −10 °C.

✅ 335 g of ethylene glycol per kg of water gives −10 °C freeze protection.

Example 3 — % Mass to Molarity Conversion: Concentrated HCl

Bottle label: 36.5% w/w HCl, density ρ = 1.18 g/mL. Molar mass of HCl = 36.46 g/mol.

Use formula: \(M = \frac{10 \times \rho \times \%_{\text{mass}}}{M_w} = \frac{10 \times 1.18 \times 36.5}{36.46}\)

M = (10 × 1.18 × 36.5) / 36.46 = 430.7 / 36.46 = 11.81 mol/L ≈ 11.8 M.

✅ Concentrated HCl (36.5% w/w, ρ=1.18 g/mL) ≈ 11.8 M. Now use C₁V₁ = C₂V₂ to prepare laboratory dilutions.

Example 4 — PPM: Lead Contamination in Drinking Water

A 500 mL water sample is found to contain 7.5 μg of lead (Pb). Calculate the lead concentration in ppb.

For dilute aqueous solution, density ≈ 1 kg/L, so 500 mL ≈ 0.500 kg of solution.

\(\text{ppb} = \frac{m_{\text{solute (μg)}}}{m_{\text{solution (kg)}}} = \frac{7.5\ \mu\text{g}}{0.500\ \text{kg}} = 15\ \text{ppb}\).

The US EPA action level for lead in drinking water is 15 ppb. This sample is exactly at the action level — trigger corrective action.

✅ Lead concentration = 15 ppb. At or above the EPA action level of 15 ppb.

Frequently Asked Questions

What is solution concentration?＋

Solution concentration is a measure of the amount of solute dissolved in a given amount of solution or solvent. It can be expressed in many units: molarity (mol/L) for lab work, molality (mol/kg) for temperature-sensitive calculations, percent mass or volume for industrial use, mole fraction for thermodynamics, and parts per million/billion (ppm/ppb) for trace-level environmental and food safety applications.

What is molarity and how do I calculate it?＋

Molarity M = n_solute / V_solution (in litres). Steps: (1) convert mass of solute to moles: n = mass(g) / molar mass(g/mol); (2) convert volume to litres if needed; (3) divide moles by litres. Example: 58.44 g NaCl (1.000 mol) in 2.000 L of solution → M = 1.000/2.000 = 0.5000 M. Molarity is temperature-dependent because solution volume changes with temperature.

What is the difference between molarity and molality?＋

Molarity (M) = moles of solute per litre of solution — uses solution volume, temperature-dependent. Molality (m) = moles of solute per kilogram of solvent — uses mass of solvent only, temperature-independent. For dilute aqueous solutions near room temperature, M ≈ m numerically. For concentrated solutions or non-aqueous solvents, they differ significantly. Use molality for colligative property calculations (boiling/freezing point changes) because those involve temperature changes.

How do I convert between concentration units?＋

Most interconversions require two pieces of information beyond what one unit provides: (1) molar mass M_w to convert between mass and moles; (2) solution density ρ to convert between volume of solution and mass of solution. Key formulas: Molarity ↔ % mass: M = (10ρ × %mass)/M_w. % mass ↔ ppm: 1% = 10,000 ppm. Molarity ↔ molality: m = 1000M/(1000ρ − M·M_w). Mole fraction ↔ molality: m = 1000·X_sol/(M_w_solvent·X_solv).

What is percent by mass (w/w)?＋

% mass = (mass of solute / mass of solution) × 100. It is dimensionless and temperature-independent. The mass of solution = mass of solute + mass of solvent. Example: 25 g NaCl dissolved in 75 g water gives 100 g solution → 25% w/w NaCl. Used in food labelling, pharmaceutical batch records, industrial reagent specifications, and concentrated acid/base labels.

What is mole fraction and when is it used?＋

Mole fraction X_i = n_i / n_total. It is dimensionless and always between 0 and 1. All mole fractions in a solution sum to 1. Applications: Raoult's Law (vapour pressure: P_i = X_i × P°_i), Dalton's Law for gas mixtures (partial pressure: P_i = X_i × P_total), thermodynamic activity, phase diagrams, and distillation calculations. Mole fraction is preferred in theoretical and thermodynamic treatments because it is symmetric — it treats all components equally.

What does ppm mean in water quality?＋

In water quality, 1 ppm = 1 mg of solute per litre of water (because water density ≈ 1 g/mL ≈ 1 kg/L). Regulatory limits in drinking water: WHO arsenic limit = 10 ppb; EPA lead action level = 15 ppb; EPA nitrate limit = 10 ppm. For very clean or ultra-pure water in semiconductor manufacturing, impurity specs are often in ppt (parts per trillion = μg per 1000 L).

What is the dilution formula and how do I use it?＋

C₁V₁ = C₂V₂ — moles of solute are conserved during dilution. C₁ and V₁ are the stock (concentrated) values; C₂ and V₂ are the target (diluted) values. Example: prepare 500 mL of 0.200 M HCl from 12.0 M stock: V₁ = C₂V₂/C₁ = (0.200 × 0.500)/12.0 = 8.33 × 10⁻³ L = 8.33 mL. Measure 8.33 mL of 12.0 M HCl and dilute carefully to 500 mL total volume.

What is the van't Hoff factor i?＋

The van't Hoff factor i accounts for the number of particles produced when a solute dissolves. For non-electrolytes (glucose, sucrose): i = 1. For strong electrolytes: NaCl → Na⁺ + Cl⁻, i = 2; CaCl₂ → Ca²⁺ + 2Cl⁻, i = 3; AlCl₃ → Al³⁺ + 3Cl⁻, i = 4. For weak electrolytes, i is between 1 and the maximum theoretical value depending on degree of ionisation. Colligative property formulas (ΔTb, ΔTf, π) multiply by i because colligative effects depend on the total number of particles.

Why does adding salt lower the freezing point of water?＋

Adding a solute (like NaCl) increases the number of particles in solution, which disrupts the regular crystal lattice of ice trying to form. The solution must be cooled to a lower temperature (below 0 °C) before the water molecules slow down enough to form an ordered ice lattice despite the interfering solute particles. Formula: ΔTf = Kf × m × i. A 1 molal NaCl solution (i=2) lowers the freezing point by 1.853 × 1 × 2 = 3.706 °C, giving a freezing point of −3.71 °C.

What is osmotic pressure and how is it related to concentration?＋

Osmotic pressure π = iMRT is the pressure required to prevent osmosis (flow of solvent through a semi-permeable membrane from low to high concentration). At 25 °C (298 K), a 1 M glucose solution has π = 1 × 1 × 0.08206 × 298 = 24.4 atm. Osmotic pressure is why cells shrivel in hypertonic salt solutions (water flows out) and swell in hypotonic pure water (water flows in). It is also the principle behind reverse osmosis water purification (applying pressure greater than π to force water through a membrane against the concentration gradient).

What is the molar mass and how do I find it?＋

Molar mass M_w (g/mol) = sum of atomic masses of all atoms in the chemical formula (from the periodic table). Examples: NaCl = 23.0 + 35.5 = 58.5 g/mol; H₂SO₄ = 2(1.008) + 32.065 + 4(15.999) = 98.08 g/mol; Glucose C₆H₁₂O₆ = 6(12.011) + 12(1.008) + 6(15.999) = 180.16 g/mol; NaHCO₃ = 22.990 + 1.008 + 12.011 + 3(15.999) = 84.007 g/mol. Molar mass is needed to convert between mass (grams) and moles: n = m/M_w.