Name: electrolyte ions and the van't hoff factor
Uploaded: 2020-09-24T15:06:50.000Z
Duration: 3 min 8 s
Channel: JoVE

K12 Standards

Tennessee Academic Standards

High School

Chemistry

Matter and Its Interactions

CHEM1.PS1.15

Electrolytes: van't Hoff Factor

Colligative properties of a solution depend on the actual number of dissolved solute particles. 
For a non-electrolyte, each solute molecule that dissolves yields one dissolved solute molecule. 
However, ionic electrolytes, such as sodium chloride, dissociate into ions when dissolved, while ammonia gas, a covalent electrolyte, reacts with water to release ammonium and hydroxide ions in solution. 
So for electrolytes, each solute molecule that dissolves yields more than one dissolved solute particle. 
Thus, a 1 M solution of a non-electrolyte such as dextrose, will have 1 mole of dextrose molecules in a 1-liter solution, whereas a 1 M solution of an electrolyte such as potassium chloride will dissociate into nearly 1 mole of potassium ions and 1 mole of chloride ions-- a total of 2 moles of ions in a 1-liter solution.
With double the number of solute particles, the osmotic pressure of a 1 M potassium chloride solution will be twice that of a 1 M dextrose solution.
The ratio between the moles of particles that a dissolving solute forms in solution and the moles of solute added to make a solution, is called the van&#8217;t Hoff factor represented by i.
It is calculated by dividing the measured

electrolytes: van't hoff factor

Colligative Properties of ElectrolytesThe colligative properties of a solution depend only on the number, not on the identity, of solute species dissolved. The concentration terms in the equations for various colligative properties (freezing point depression, boiling point elevation, osmotic pressure) pertain to all solute species present in the solution. Nonelectrolytes dissolve physically without dissociation or any other accompanying process. Each molecule that dissolves yields one dissolved solute molecule. The dissolution of an electrolyte, however, is not this simple, as illustrated by the two common examples below:<img alt="Kirchhoff's circuit laws diagram; &#931;I=0, &#931;V=0; electrical circuit analysis; educational schematic." src="/files/ftp_upload/11371/11371_Equation_1.png" style="text-align: center; height: 90px;" data-altupdatedat="1761830439000" data-alt="Eq1" />Considering the first of these examples, and assuming complete dissociation, a 1.0 m aqueous solution of NaCl contains 2.0 moles of ions (1.0 mol Na+ and 1.0 mol Cl&#8722;) per each kilogram of water, and its freezing point depression is expected to be<img alt="static equilibrium; &#931;Fx=0; block diagram; demonstrates mechanical balance on an inclined plane" src="/files/ftp_upload/11371/11371_Equation_2.png" style="opacity: 0.9; text-align: center; height: 90px;" data-altupdatedat="1761830439000" data-alt="Eq2" />However, when this solution is actually prepared and its freezing point depression is measured, a value of 3.4 &#176;C is obtained. Similar discrepancies are observed for other ionic compounds, and the differences between the measured and expected colligative property values typically become more significant as solute concentrations increase. These observations suggest that the ions of sodium chloride (and other strong electrolytes) are not completely dissociated in solution.To account for this and avoid the errors accompanying the assumption of total dissociation, an experimentally measured parameter named in honor of Nobel Prize-winning German chemist Jacobus Henricus van&#8217;t Hoff is used. The van&#8217;t Hoff factor (i) is defined as the ratio of solute particles in solution to the number of formula units dissolved:<img alt="Synthetic division process, polynomial division formula. Diagram illustrates step-by-step method." src="/files/ftp_upload/11371/11371_Equation_3.png" style="text-align: center; height: 90px;" data-altupdatedat="1761830439000" data-alt="Eq3" />In 1923, the chemists Peter Debye and Erich H&#252;ckel proposed a theory to explain the apparent incomplete ionization of strong electrolytes. They suggested that although interionic attraction in an aqueous solution is very greatly reduced by solvation of the ions and the insulating action of the polar solvent, it is not completely nullified. The residual attractions prevent the ions from behaving as totally independent particles. In some cases, a positive and negative ion may actually touch, giving a solvated unit called an ion pair. Thus, the activity&#8212;or the effective concentration&#8212;of any particular kind of ion is less than that indicated by the actual concentration. Ions become more and more widely separated as the solution becomes more dilute and the residual interionic attractions become less and less. Thus, in extremely dilute solutions, the effective concentrations of the ions (their activities) are essentially equal to the actual concentrations. For 0.05 m solutions, the value of i for NaCl is 1.9, as opposed to an ideal value of 2.This text is adapted from <a href="https://openstax.org/books/chemistry-2e/pages/11-4-colligative-properties?query=vapor%20pressure&#38;target={">Openstax, Chemistry 2e, Section 11.4: Colligative Properties.</a>

Colligative Properties of Electrolytes
The colligative properties of a solution depend only on the number, not on the identity, of solute species ...

Education

JoVE Core

Solutions and Colloids

KEY TERMS AND CONCEPTS

electrolyte ions and the van't hoff factor

Electrolyte ions affect colligative properties because these properties depend on the number of dissolved particles, not their identity. The concentration terms used for freezing point depression, boiling point elevation, and osmotic pressure include every solute species present in the solution. Nonelectrolytes dissolve without dissociation, so each molecule that enters solution counts as one particle.

Electrolytes behave differently because they can form ions in water. If sodium chloride were fully dissociated, a 1.0 m aqueous NaCl solution would contain 2.0 moles of ions per kilogram of water, with 1.0 mol of Na+ and 1.0 mol of Cl−. That complete dissociation would predict a certain freezing point depression, but the measured value for this solution is 3.4 °C.

Similar differences show up with other ionic compounds, and the gap between expected and measured colligative properties often grows at higher concentrations. These results show that strong electrolytes do not always act as if they are completely broken into independent ions in solution. To correct for this, chemists use the van't Hoff factor, written as i.

The van't Hoff factor is the ratio of solute particles in solution to the number of formula units dissolved. It gives an experimental way to describe how many particles a dissolved electrolyte produces. In 1923, Peter Debye and Erich Hückel предложed a theory to explain why strong electrolytes seem to ionize incompletely.

Their idea was that interionic attraction is greatly reduced in water because the ions are solvated and the polar solvent acts as an insulator, but the attraction is not removed completely. The remaining attraction keeps ions from behaving as fully independent particles. In some cases, a positive ion and a negative ion can touch and form a solvated ion pair.

This means the activity, or effective concentration, of an ion is lower than its actual concentration. As a solution becomes more dilute, the ions are farther apart and these residual attractions get weaker. In very dilute solutions, the effective concentration is nearly the same as the actual concentration, and for 0.05 m NaCl the value of i is 1.9 instead of the ideal value of 2.

Electrolyte Ions and the van't Hoff Factor

electrolyte-ions-and-the-van-t-hoff-factor

solution formation

Why Some Substances Dissolve Better

Some substances dissolve easily in one solvent but not in another. A quick way to predict solubility is the idea that "like dissolves like." Polar substances, such as salt and sugar, dissolve well in polar solvents like water. Non-polar substances are more soluble in non-polar solvents such as carbon tetrachloride. This selective solubility depends on intermolecular forces, which are the attractions within and between molecules. Stronger forces between solute and solvent molecules lead to...

intermolecular forces in solutions

How Solutions Form Without Energy Change

A solution can form as a spontaneous process. That means it happens under the right conditions without energy from an outside source. When the attraction between solute and solvent particles is about the same as the attraction in the separate substances, the solution forms with no energy change. A solution like this is called an ideal solution. Mixtures of ideal gases are one example, and helium and argon behave close to that model. These gases do not have significant intermolecular...

enthalpy of solution

Dissolution and Heat Changes

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aqueous solutions and heats of hydration

How Water Dissolves Ionic Compounds

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solution equilibrium and saturation

Saturation Limits in Solutions

Saturation limits in solutions describe how much solute can dissolve in a solvent under given conditions. As sugar is added to water and stirred, it dissolves at first. More sugar can be added until the solution reaches its natural limit. That limit depends mainly on the relative strengths of three attractions: solute-solute, solute-solvent, and solvent-solvent forces. When no more sugar dissolves, even after long stirring, undissolved sugar stays behind. The concentration of sugar in the...

physical properties affecting solubility

Gas Solubility in Liquids

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expressing solution concentration

Measuring Solution Strength

Solution strength can be described with terms like dilute and concentrated, and it can also be measured with several scientific units. A solute is the part of a solution present at a lower amount than the solvent. Different units are useful in different situations, especially in chemistry. Molarity, written as M, is a common unit for concentration. It tells how many moles of solute are in one liter of solution. Because solution volume changes with temperature, molarity can change too when the...

vapor pressure lowering

Raoult's Law and Lower Vapor Pressure

Raoult’s law explains why a solution with a nonvolatile solute has a lower vapor pressure than the pure liquid. Vapor pressure is the pressure from a liquid’s gas phase when vaporization and condensation happen at equal rates. When a nonvolatile substance dissolves in a volatile liquid, the liquid’s equilibrium vapor pressure drops. One reason for this drop is at the surface of the solution. Solvent molecules must be at the surface to vaporize, and added solute molecules reduce the surface...

ideal solutions

Raoult's Law in Real Solutions

Raoult's law describes how vapor pressure changes in a solution. For a solvent, the partial vapor pressure equals the vapor pressure of the pure solvent multiplied by its mole fraction in the solution. This relationship works best for ideal solutions. An ideal solution is one in which solvent-solute attractions are just as strong as solvent-solvent and solute-solute attractions. In that case, both components need about the same amount of energy to escape into the vapor phase as they do in...

freezing point depression and boiling point elevation

freezing-point-depression-and-boiling-point-elevation

Colligative Effects in Solutions

Colligative effects in solutions include boiling point elevation and freezing point depression. These properties depend on how many solute particles are present, not on the solute’s chemical identity. Because of that, a solution can behave differently from the pure solvent even when only a nonvolatile solute is added. Boiling point elevation happens because the vapor pressure of a solution is lower than the vapor pressure of the pure solvent. The boiling point is the temperature at which vapor...

osmosis and osmotic pressure of solutions

osmosis-and-osmotic-pressure-of-solutions

Semipermeable Membranes and Reverse Osmosis

Semipermeable membranes let some molecules pass while blocking others. They are found in biological cell membranes and in dialysis tubing used to remove metabolic wastes from blood. These examples show how selective permeation depends on factors such as size, shape, polarity, and charge. Osmosis is the net movement of solvent through a semipermeable membrane. In a U-shaped setup, pure solvent and a solution are separated by a membrane that allows only solvent molecules to pass. Solvent...

Electrolyte ions affect colligative properties because these properties depend on the number of dissolved particles, not their identity. The concentration terms used for freezing point depression, boiling point elevation, and osmotic pressure include every solute species present in the solution. Nonelectrolytes dissolve without dissociation, so each molecule that enters solution counts as one particle. Electrolytes behave differently because they can form ions in water. If sodium chloride were...

colloids

Colloids and Everyday Mixtures

Colloids are mixtures with particles between the size of a solution and a suspension. They stay dispersed instead of settling out, even though their particles are larger than most simple molecules. This makes colloids different from cloudy suspensions like mud and water, flour and water, or tempera paint. It also makes them different from solutions, where dissolved molecules or ions are invisible and do not settle. Colloidal systems can be made in two main ways. In dispersion methods, larger...

key terms and concepts

Learn how electrolyte ions change colligative properties, why NaCl does not act as a fully ideal electrolyte, and how the van't Hoff factor helps measure particle count.

Electrolyte Ions and the van't Hoff Factor | JoVE K12

chemistry

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Virginia  Standards of Learning

HIgh School

CH.5

CH.5.c

CH.5.e

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Formula Mass and Mole Concepts of Compounds

Formula Mass of Covalent Compounds

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CHEM1.PS1.1

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Analyze solutions to identify solutes and solvents, quantitatively analyze concentrations (molarity, percent composition, and ppm), and perform separation methods such as evaporation, distillation, and/or chromatography and show conceptual understanding of distillation. Construct argument to justify the use of certain separation methods under different conditions. EXPLANATION: In seventh grade, students developed an understanding that chemical and physical properties of a sample of matter could be used to identify a substance. Building on this understanding, students now explore the use of these physical and chemical properties to separate mixtures. Analysis of the substances is performed to aid in further discussion and facilitate further quantitative analysis.

CHEM1.PS1.7

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Lewis Acids and Bases

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Acid/Base Strengths and Dissociation Constants

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Identify acids and bases as a special class of compounds with a specific set of properties. EXPLANATION: The concept of pH is first introduced in seventh grade as a chemical property of matter. Students should be introduced to multiple explanations of acid and base behavior to permit classification of common substances (e.g. baking soda, ammonia, carbon dioxide) as acids or bases. To further develop discussions of solubility, qualitative discussions of the differences between strong and weak acids are appropriate. (Calculations of pH or ion concentrations are beyond the scope of this standard. Discussions of pOH are beyond the scope of this standard. Conjugate pairs are beyond the scope of this standard.)

CHEM1.PS1.8

Electrolyte Ions and the van't Hoff Factor

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