Name: colligative effects in solutions
Uploaded: 2020-09-24T15:06:50.000Z
Duration: 3 min 12 s
Channel: JoVE

K12 Standards

Tennessee Academic Standards

High School

Chemistry

Matter and Its Interactions

CHEM1.PS1.15

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

Freezing Point Depression and Boiling Point Elevation

The temperature at which the vapor pressure of a liquid equals the atmospheric pressure is known as its boiling point. 
Since adding a non-volatile solute lowers the vapor pressure of a solvent, a solution requires a higher temperature to increase its vapor pressure to a point that equals the atmospheric pressure. Thus, the boiling point of a solution is greater than that of a pure solvent.
These changes in vaporization can be examined over a range of temperatures and pressure using a phase diagram.
A solution has a lower vapor pressure than the pure solvent at all temperatures. So, the vaporization curve of the solution would lie below that of the solvent. 
At 1 atm, the curve corresponds to a temperature higher than the boiling point of the pure solvent. 
The increase in the boiling point of the solution compared to that of the pure solvent is known as boiling point elevation.
The boiling point of a solution is a colligative property. The temperature increase, or &#916;Tb, is directly proportional to the concentration of solute and can be calculated by multiplying the molality of the solute and the molal boiling point elevation constant.
The boiling point elevation constant has the units

freezing point depression and boiling point elevation

Boiling Point Elevation
The boiling point of a liquid is the temperature at which its vapor pressure is equal to ambient atmospheric pressure. Since the vapor pressure of a solution is lowered due to the presence of nonvolatile solutes, it stands to reason that the solution&#8217;s boiling point will subsequently be increased. Vapor pressure increases with temperature, and so a solution will require a higher temperature than will pure solvent to achieve any given vapor pressure, including one equivalent to that of the surrounding atmosphere. The increase in boiling point observed when a non-volatile solute is dissolved in a solvent, &#916;Tb, is called boiling point elevation and is directly proportional to the molal concentration of solute species:
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where Kb is the boiling point elevation constant, or the ebullioscopic constant and m is the molal concentration (molality) of all solute species. Boiling point elevation constants are characteristic properties that depend on the identity of the solvent.
Freezing Point Depression
Solutions freeze at lower temperatures than pure liquids. This phenomenon is exploited in &#8220;de-icing&#8221; schemes that use salt, calcium chloride, or urea to melt ice on roads and sidewalks, and in the use of ethylene glycol as an &#8220;antifreeze&#8221; in automobile radiators. Seawater freezes at a lower temperature than freshwater, and so the Arctic and Antarctic oceans remain unfrozen even at temperatures below 0 &#176;C (as do the body fluids of fish and other cold-blooded sea animals that live in these oceans).
The decrease in freezing point of a dilute solution compared to that of the pure solvent, &#916;Tf, is called the freezing point depression and is directly proportional to the molal concentration of the solute
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where m is the molal concentration of the solute and Kf is called the freezing point depression constant (or cryoscopic constant). Just as for boiling point elevation constants, these are characteristic properties whose values depend on the chemical identity of the solvent.
Determination of Molar Masses
Osmotic pressure and changes in freezing point, boiling point, and vapor pressure are directly proportional to the number of solute species present in a given amount of solution. Consequently, measuring one of these properties for a solution prepared using a known mass of solute permits determination of the solute&#8217;s molar mass.
For example, a solution of 4.00 g of a nonelectrolyte dissolved in 55.0 g of benzene is found to freeze at 2.32 &#176;C. Assuming ideal solution behavior, what is the molar mass of this compound?
To solve this problem, first, the change in freezing point from the observed freezing point and the freezing point of pure benzene is calculated:
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Then, the molal concentration is determined from Kf, the freezing point depression constant for benzene, and &#916;Tf:
<img alt="static equilibrium &#931;Fx=0 &#931;Fy=0 diagram physics force system analysis" src="/files/ftp_upload/11369/11369_Equation_4.png" style="text-align: center; height: 115px;" data-altupdatedat="1754904545000" data-alt="Chemical structure diagram of fexofenadine, an antihistamine; showing carbon, hydrogen, oxygen, and nitrogen atoms connected by bonds." />
Next, the number of moles of the compound in the solution is found from the molal concentration and the mass of solvent that was used to make the solution.
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And, finally, the molar mass from the mass of the solute and the number of moles in that mass is determined.
<img alt="Hydrogen atom wavefunction graph; &#936;(r,&#952;,&#966;) spherical harmonics with quantum numbers n, l, m." src="/files/ftp_upload/11369/11369_Equation_6.png" style="text-align: center; height: 70px;" data-altupdatedat="1754904545000" data-alt="Chemical equilibrium derivation with equations; thermodynamics diagram; static equilibrium principle." />
This text is adapted from <a href="https://openstax.org/books/chemistry-2e/pages/11-4-colligative-properties">Openstax, Chemistry 2e, Section 11.4: Colligative Properties.</a>

Boiling Point Elevation
The boiling point of a liquid is the temperature at which its vapor pressure is equal to ambient atmospheric pressure. Since the ...

Education

JoVE Core

Solutions and Colloids

KEY TERMS AND CONCEPTS

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 pressure matches ambient atmospheric pressure. Since a solution needs a higher temperature to reach that pressure, its boiling point rises. The increase in boiling point, ΔTb, is directly proportional to the molal concentration of solute species. The constant Kb is the boiling point elevation constant, also called the ebullioscopic constant, and it depends on the solvent.

Freezing point depression is the lowering of a liquid’s freezing temperature compared with the pure solvent. Solutions freeze at lower temperatures than pure liquids. This effect is used in de-icing with salt, calcium chloride, or urea on roads and sidewalks. It also explains why ethylene glycol is used as antifreeze in car radiators. Seawater also freezes below 0 °C, which helps keep the Arctic and Antarctic oceans unfrozen, and it supports the body fluids of fish and other cold-blooded sea animals living there. The decrease in freezing point, ΔTf, is directly proportional to molal concentration. The freezing point depression constant, Kf, or cryoscopic constant, also depends on the solvent.

These colligative properties can be used to find molar mass. Osmotic pressure and changes in freezing point, boiling point, and vapor pressure all depend on the number of solute species in a given amount of solution. That means measuring one of these properties for a solution made with a known mass of solute can reveal the solute’s molar mass. In the benzene example, 4.00 g of a nonelectrolyte dissolves in 55.0 g of benzene, and the solution freezes at 2.32 °C. The calculation begins by finding ΔTf from the freezing point of pure benzene and the observed freezing point. Next, molality is found using Kf for benzene. Then the moles of solute are calculated from the solvent mass, and the solute’s molar mass is found from its mass and moles.

Colligative Effects in Solutions

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

Dissolution and heat changes help explain why some solutions form on their own. Two factors can favor the spontaneous formation of a solution. One is a decrease in internal energy, which happens in an exothermic change. The other is increased dispersal of matter, which means the solute spreads evenly through the solvent. When a substance dissolves, the internal energy often changes as heat is absorbed or released. This does not happen in every case, but it is common. A solution forms when...

aqueous solutions and heats of hydration

How Water Dissolves Ionic Compounds

Water dissolves ionic compounds because its polar molecules attract ions. This attraction between an ion and a molecule with a dipole is called an ion-dipole attraction. It helps explain why many ionic solids can separate in water and form aqueous solutions. When an ionic compound dissolves, its ions separate and spread out evenly through the solution. Water molecules surround the ions and solvate them, which lowers the strong electrostatic forces between the ions. This separation is a...

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

Gas solubility in liquids depends on the attractive forces between the gas and the liquid. When a gas dissolves, there is no strong solute-solute attraction to break apart first. Gas particles are already far apart, so their internal interactions are negligible. That makes solute-solvent interactions the main energetic factor that controls how much gas can dissolve. Examples show how these interactions matter. Oxygen is about three times more soluble in water than helium because oxygen has...

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...

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...

electrolytes: van't hoff factor

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...

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 boiling point elevation and freezing point depression change solution behavior, with real examples and a method for finding molar mass.

Colligative Effects in Solutions | JoVE K12

chemistry

JoVE Core Chemistry brings general chemistry concepts to life using concise and easy-to-understand animated video lessons. Plus, in the scientist-in-action videos, we demonstrate classical and original research experiments performed in today's laboratories worldwide.

Virginia  Standards of Learning

HIgh School

CH.5

CH.5.b

CH.5.e

Georgia Standards of Excellence

9-12

SC6.

Boiling Point and Freezing Point Changes

formula-mass-and-mole-concepts-of-compounds

Formula Mass and Mole Concepts of Compounds

Formula Mass of Covalent Compounds

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Determining the Mass Percent Composition in an Aqueous Solution

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Stoichiometry, Product Yield, and Limiting Reactants

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...

Understand and be prepared to use values specific to chemical processes: the mole, molar mass, molarity, and percent composition. EXPLANATION: During the early and mid-1800s, a tremendous amount of early research poured into building an understanding of the way that atoms combine. This work required analysis of the percent composition of different substances. Joseph Proust's work with tin oxides led him to publish his law of definite proportions noting that a given pure substance will have a constant ratio for the combination of elements making up the substance. Building on Proust's work, John Dalton published his law of multiple proportions and atomic theory noting that multiple patterns exist for the way that atoms combine, yet the combining atoms are invariable. Meanwhile Amedeo Avogadro hypothesizes that given equal pressures and temperatures, two equal volumes of gas will contain equal numbers of atoms. Acknowledging that the exact number of atoms was unknown in these gaseous volumes, the Latin word "mole" was used. "Mole" translates to "lump" recognizing the ambiguous quantity of gas. A number of attempts have been undertaken historically to quantify the mole. Robert Millikan's oil drop experiment revealed the change of a single electron, which could then be divided into Faraday's Constant for the charge of a mole of electrons. Dividing the charge of a mole of electrons by the charge of a single electron yields the number of electrons in a mole, $6.022\times10^{\wedge}23$. Students should be familiar with how to translate between the given quantities. Student's first use atomism/Dalton's atomic theory in third grade noting that matter is made up of particles too small to be seen. This the first time where students will undertake the process of counting these invisible particles. Electrolysis of water provides a tangible experience to demonstrate the constituent parts of water, though the definite proportions of water may not be easily seen due to the solubility of the gasses produced. The concept of relative mass can be explored by comparing small groups of like items. For example, relative masses of beans can be compared with pinto, white, and black beans. If bags are created with equal numbers of like beans, students can develop a concept of why the smallest bean will have a relative mass of 1, analogous to hydrogen.

CHEM1.PS1.1

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The Atomic Theory of Matter

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Demonstrate that atoms, and therefore mass, are conserved during a chemical reaction by balancing chemical reactions. EXPLANATION: Standard 7.PS1.4 explores the idea that chemical reactions are rearrangements of atoms, and students undertake the process of balancing chemical reactions with that standard in seventh grade. The introduction of molar masses in high school makes an explicit connection between the number of atoms, as previously investigated, and the actual masses of the atoms, permitting the collection of more concrete evidence for the conservation of mass in chemical reactions.

CHEM1.PS1.2

Reaction Yield

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Chemical Equilibria: Systematic Approach to Equilibrium Calculations

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Perform stoichiometric calculations involving the following relationships: mole- mole; mass-mass; mole-mass; mole- particle; and mass- particle. Show a qualitative understanding of the phenomenon of percent yield, limiting, and excess reagents in a chemical reaction through pictorial and conceptual examples. (states of matter liquid and solid; excluding volume of gasses). EXPLANATION: This standard builds on the idea that balancing chemical reactions provides evidence for conservation of mass and that the behavior of atoms follows predictable patterns. Students now have the opportunity to utilize this understanding as they perform and evaluate chemical reactions. (The concepts of limiting reagent and percent yield are presented; however, students are not expected to determine these values algebraically. Use of molar volume in stoichiometric processes is beyond the scope of this standard.)

CHEM1.PS1.3

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Oxidation&ndash;Reduction Reactions

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types-of-chemical-reactions-exchange-and-reversible

Types of Chemical Reactions: Exchange and Reversible

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Use the reactants in a chemical reaction to predict the products and identify reaction classes (synthesis, decomposition, combustion, single replacement, double replacement). EXPLANATION: Chemical reactions are rearrangements of atoms that follow predicable patterns. There are patterns both at the macroscopic level in the behavior of some of the reaction classes, as well as patterns in the rearrangements of the atoms underlying the reaction. Students should be able to predict the products of the reactions, which also requires the ability to recognize the general patterns for each type of reaction.

CHEM1.PS1.4

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Kinetic Molecular Theory and Gas Laws Explain Properties of Gas Molecules

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Basic Postulates of Kinetic Molecular Theory: Particle Size, Energy, and Collision

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Ideal Gas Law

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Gas Laws: Boyle's, Gay-Lussac, Charles', Avogadro's, and Ideal Gas Law

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Kinetic Molecular Theory: Molecular Velocities, Temperature, and Kinetic Energy

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Applications of the Ideal Gas Law: Molar Mass, Density, and Volume

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Conduct investigations to explore and characterize the behavior of gasses (pressure, volume, and temperature), develop models to represent this behavior, and construct arguments to explain this behavior. Evaluate the relationship (qualitative and quantitatively) at STP between pressure and volume (Boyle's Law), temperature and volume (Charles' Law), temperature and pressure (Gay-Lussac Law) and moles and volume (Avogadro's Law), and evaluate and explain these relationships with respect to kinetic molecular theory. Be able to understand, establish, and predict the relationships between volume, temperature, and pressure using combined gas law both qualitatively and quantitatively. EXPLANATION: When exploring the behavior of gasses, it is important to consider experimental design. Experiments used to show the relationships between these sets of variables should include one independent and one dependent variable. Other variables should be held constant. Pressure should serve as the dependent variable because it cannot be manipulated directly. Avogadro's Law can be used to relate the amount of gas in two different containers in lieu of knowing actual number of gas particles in an individual container. Individual demonstrations can be performed to explore each of the different gas laws.

CHEM1.PS1.5

Ideal Gas Equation

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Use the ideal gas law, PV = nRT, to algebraically evaluate the relationship among the number of moles, volume, pressure, and temperature for ideal gases. EXPLANATION: If students have collected data during investigations for CHEM1.PS1.5, these data can then be used to linearize. With pressure assigned as the dependent variable and plotted against (moles) (temperature)/volume, the slope of the graph will be the gas constant, yielding the ideal gas law. After developing the mathematical model, this model can be used to make predictions for other systems.

CHEM1.PS1.6

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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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CHEM1.PS1.8

Colligative Effects in Solutions

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