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Concepts of Solutions for JEE Main Chemistry
In Chemistry, when you have a mix of two or more things where the particles are tinier than 1nm and are evenly spread out, it's called a solution. Usually, we talk about this in relation to liquids, but it can also be for gases or solids. For instance, the air we breathe is a solution because it mainly has oxygen and a bit of nitrogen and other tiny bits. Another example is alloys, where different metals are mixed together in a solid form.
Even though a solution is a blend of different things, it looks like one uniform substance. In simpler terms, everything in a solution is spread out evenly. Think of it like when you're sipping a soft drink – the taste stays the same from the first sip to the last, showing that the stuff in the drink is evenly mixed, or homogeneous.
In this article, we will learn what are solutions, the types of solutions, properties of solutions and so on. So, read on and explore all about solutions in Chemistry.
JEE Main Chemistry Chapters 2024
What is Solution?
In chemistry, a solution is a homogeneous mixture of two or more substances. The components of a solution are distributed uniformly throughout the mixture, and they cannot be separated by physical means. The substance that is present in the largest amount is called the solvent, while the other substances are called solutes.
Characteristics of Solutions
A solution is a mixture of two or more different components. Out of these, one substance is always a component in which the other component dissolves. That is, one substance is a solvent and others are called solutes which dissolve in the solvent. So, the solvent is the substance which dissolves other substances in it, such as water. Water is also called the universal solvent because it dissolves most of the other particles in nature. The quantity of solvent is larger than the quantity of solute.
Solute, on the other hand, is the substance or component which is usually present in lower quantities than the solvent and it gets dissolved in the solvent.
Examples Of Solutions
Given below are some of the solution examples:
1.Brass is an alloy of copper and zinc.
2. The air around us is an example of a gaseous form of solution.
3. Sugar syrup is a solution of sugar in water. Water is the solvent and sugar are the solutes in this case.
4. Coffee or tea is also an example of a solution.
5. Carbonated drinks are solutions of water as a solvent and carbon dioxide and other ingredients as solutes.
6. Tincture of iodine has alcohol as the solvent and iodine as the solute.
Solution Examples in Your Everyday Life
Solutions are ubiquitous in our everyday lives. Here are a few examples of solutions:
Air: Air is a solution of gases, primarily nitrogen and oxygen.
Water: Water is a solution of various substances, including salts and minerals.
Blood: Blood is a solution of cells, proteins, and other substances.
Gasoline: Gasoline is a solution of hydrocarbons.
Sugar Water: Sugar water is a solution of sugar and water.
Types Of Solutions
Liquid solutions with water as the solvent are the most often found solutions around us. But gaseous and solid solutions are also quite abundant in nature. A solute or a solvent may be in any state of matter- gases, liquids, or solids. So, depending upon the physical state, solutions can be classified into various types.
Expressing Concentration of Solutions and Vapor Pressure
In the world of chemistry, expressing the concentration of a solution is crucial for understanding its properties and behaviors. There are several methods to quantify this concentration, each offering unique insights. Additionally, the vapor pressure of solutions is a vital parameter, and Raoult's Law provides insights into its behavior.
Methods for Expressing Concentration:
Molality (m): Molality is defined as the number of moles of solute per kilogram of solvent. It is an excellent choice for expressing concentration when dealing with changes in temperature since it is independent of temperature.
Molality = Moles of Solute / Mass of Solvent in Kg
Molarity (M): Molarity represents the number of moles of solute per liter of solution. Molarity is a convenient way to prepare solutions with a specific volume in mind, but it can change with temperature due to volume changes.
Molarity = Moles of Solute / Volume of Solution in Liter
To learn more about Molality and Molarity you can check Vedantu’s page on Molarity and Molality - Important Concepts for JEE.
Mole Fraction (X): Mole fraction is the ratio of the number of moles of solute to the total number of moles in the solution. It is a dimensionless quantity and provides information about the solute's contribution to the overall properties of the solution.
Mole Fraction of a Component = Number of Moles of the Component / Total Number of Moles of all Components of the Solution
Percentage (by Mass and by Volume): Percentage by mass is the mass of the solute divided by the total mass of the solution, multiplied by 100. Percentage by volume is the volume of the solute divided by the total volume of the solution, multiplied by 100.
Mass by Volume Percentage = Mass of Solute Dissolved in 100mL of the Solution.
Vapor Pressure of Solutions and Raoult's Law:
It was given by French Chemist Francois Marte Raoult in 1886. Raoult’s law states that for a solution of volatile liquids, the partial vapor pressure of each component of the solution is directly proportional to its mole fraction present in the solution.
The vapor pressure of a solution is influenced by the presence of solute particles. Raoult's Law describes this relationship for ideal solutions. It states that the vapor pressure of a component in a solution is directly proportional to its mole fraction.
If we take a binary solution of two volatile liquids and those two components are denoted by components A and B. Then, for component A –
PA ∝ xA
pA = pA° xA
Where PA = partial vapor pressure of the component of A, xA = mole fraction of component A, pA° = vapor pressure of the pure component A at the same temperature.
For component B –
pB = pB° xB
Where PB = partial vapor pressure of the component of B, xB = mole fraction of component B, PB° = vapor pressure of the pure component B at the same temperature.
Total Pressure
According to Dalton’s law of partial pressures, the total pressure (ptotal) over the solution phase in the container will be the sum of the partial pressures of the components of the solution. So, it can be written as –
Ptotal = pA + pB
On substituting the values of pA and pB –
Ptotal = pA° xA + pB° xB
Ptotal = (1 - xB) pA° + pB° xB
Ptotal = pA° + (pB° - pA°)xB
Following Conclusions can be Drawn from the Above Equation
Total vapor pressure over the solution can be related to the mole fraction of any one component of the solution.
Total vapor pressure over the solution varies linearly with the mole fraction of component B.
Depending on the vapor pressures of the pure components A and B, total vapor pressure over the solution decreases or increases with the increase of the mole fraction of component A.
Raoult’s Law can Also be Stated as Follows
The partial pressure of each volatile component (or gas) in the solution is directly proportional to its mole fraction.
Based on Raoult’s law, the liquid-liquid solution can be classified into the following two types –
Ideal solutions
Non – ideal solutions
If you want to know more about the importance, deviations, and limitations of Rasult’s Law, check out Vedantu’s page on Rault’s law.
Ideal and Non-Ideal Solutions:
Ideal Solutions: In ideal solutions, the interactions between solute-solute, solvent-solvent, and solute-solvent molecules are similar. The vapor pressure follows Raoult's Law precisely.
Non-Ideal Solutions: In non-ideal solutions, the interactions between molecules differ from those in ideal solutions. This leads to deviations from Raoult's Law. Non-ideal behavior can manifest as positive or negative deviations, indicating stronger or weaker interactions, respectively, in the mixture.
Do you want to know more about how Ideal and Non-ideal solutions related to Rault’s law? Check out Vedantu’s page on Ideal and Non-Ideal Solutions Raoult's Law For JEE.
Vapor Pressure-Composition Plots:
Vapor pressure-composition plots illustrate how the vapor pressure of a solution changes with varying solute concentrations. In ideal solutions, the plot is a straight line, while in non-ideal solutions, it deviates from linearity.
Colligative Properties of Dilute Solutions:
Colligative properties are properties of solutions that depend on the number of solute particles rather than their nature. These include:
Relative Lowering of Vapor Pressure: The presence of solute particles lowers the vapor pressure of the solvent. This is used in techniques like osmometry to determine molecular masses.
Depression of Freezing Point: Adding a solute depresses the freezing point of a solvent. This is utilized in antifreeze solutions for vehicles.
Elevation of Boiling Point: The boiling point of a solvent is raised when a solute is added. This is evident in the preparation of concentrated sulfuric acid.
Osmotic Pressure: Osmotic pressure is the pressure required to prevent the flow of solvent into a more concentrated solution through a semipermeable membrane. It's a key parameter in osmosis and has applications in cell biology and chemical separation processes.
Relative Lowering of Vapor Pressure
Raoult established that lowering of vapor pressure does not depend on the identity of solute particles; instead , it only depends on the concentration of solute particles. We can write –
P1 = x1 p1°
The Vapor pressure of the pure solvent will be more than that of the solvent. So, we can write change in vapor pressure as follows –
∆p1 = p1° - p1
= p1° - p1° x1
= p1° (1 - x1)
As we know x2 = 1 - x1 so, we can write –
∆p1 = x2 p1°
If a solution contains many non – volatile solutes, then the lowering of the vapor pressure depends on the sum of mole fraction of different solutes. So, the above equation can be written as –
∆p1/p1° = (p1°- p1)/p1° = x2 ---------(1)
In the above equation, the left-hand side equation is called relative lowering of vapor pressure which is equal to the mole fraction of the solute.
As x2 = n2\(n1+n2)so, the equation (1) can be written as follows –
(p1°- p1)/p1° = n2\(n1+n2)
n1 = number of moles of solvent in the solution , n2 = number of moles of solute in the solution
For a Highly Diluted Solution
For a highly diluted solution, n1 > > n2, so n2 can be neglected as it’s a very small value. Thus, we can write –
(p1°- p1)/p1° = n2/n1
As we know number of moles = mass/molar mass so, we can write –
(p1°- p1)/p1° = (w2/M2)/(w1/M1)
(p1°- p1)/p1° = (w2/M2) x (M1/w1)
Where w1 and M1 are the mass and molar mass of solute while w2 and M2 are the mass and molar mass of solute.
Elevation of Boiling Point
Elevation of boiling point also depends only on the number of solute particles instead of the nature of the solute particles.
If the boiling point of the pure solvent is Tb° and the boiling point of solution is Tb. Elevation of boiling point will be –
∆Tb = Tb - Tb°
According to the results of the experiments, for dilute solutions, the elevation of the boiling point is directly proportional to the molal concentration of the solute in a solution. Thus, we can write –
∆Tb ∝ m
On removing the proportionality –
∆Tb = Kbxm
Where, m = molality, Kb = Boiling point elevation constant or molal elevation constant or Ebullioscopic constant. Its unit is K kg/ mol.
If w2 and M2 are the mass and molar mass of solute which are dissolved in the w1 gram of solvent.
m = (w2/M2)/(w1/1000) = (1000 x w2)/(M2 x w1)
∆Tb = (Kb x 1000 x w2)/(M2 x w1)
M2 = (Kb x 1000 x w2)/(∆Tbx w1)
Depression of Freezing Point
The solution shows depression of freezing point compared to the pure solvent.
Freezing Point
The freezing point can be defined as the temperature at which the vapor pressure of the substance in its liquid phase is equal to its vapor pressure in the solid phase.
When we add some non – volatile solids to the solvent, its vapor pressure decreases (Raoult’s law). Due to decrease in vapor pressure, it becomes equal to solid at lower temperature. Therefore, the freezing point of the solvent decreases.
Depression in Freezing Point
∆Tf= Tf* - Tf
Where Tf* = freezing point of pure solvent, Tf = freezing point of the solvent when a non - volatile solute is dissolved in it.
Depression of freezing point for dilute solution is directly proportional to molality of the solution. It can be expressed as –
∆Tf ∝ m
∆Tf = Kf m
Where, ∆Tf = depression of freezing point, m = molality, Kf = Freezing point depression constant or molal depression constant or cryoscopic constant.
If w2 and M2 are the mass and molar mass of solute which are dissolved in the w1 gram of solvent.
m = (w2/M2)/(w1/1000) = (1000 x w2)/(M2 x w1)
∆Tf = (Kf x 1000 x w2)/(M2 x w1)
M2 = (Kf x 1000 x w2)/(∆Tfx w1)
Osmosis and Osmotic Pressure
The process of movement of the solvent across a semipermeable membrane towards a higher concentration of solute is called osmosis. Osmotic pressure is the minimum pressure required or applied to a solution to halt the flow of its pure solvent across a semipermeable membrane. Osmotic pressure is also a colligative property. It depends on the concentration of the solute in the solution. It is expressed as follows –
π = iCRT
Where, π = osmotic pressure
i = van’t Hoff factor
C = molar concentration of the solute in the solution
R = universal gas constant
T = temperature
When the pressure is applied more than that of osmotic pressure, then pure solvent starts flowing out of the solution through the semipermeable membrane. This phenomenon is called reverse osmosis.
Experimental values of molar masses sometimes differ from the theoretical values of molecular masses (calculated from the colligative properties of solutions). These values are known as abnormal molar masses.
Determination of Molecular Mass using Colligative Properties:
Colligative properties, like relative lowering of vapor pressure and osmotic pressure, can be used to determine the molecular mass of a solute. By measuring the colligative property and applying mathematical relationships, the molecular mass can be calculated.
Abnormal Value of Molar Mass, Van’t Hoff Factor, and Its Significance:
The van't Hoff factor (i) represents the number of particles into which a solute dissociates in a solution. Sometimes, the van't Hoff factor is higher or lower than expected, indicating the presence of associated or aggregated particles. Understanding these deviations is essential for accurately predicting a solution's behavior and properties. It is particularly important in areas such as pharmaceuticals and industrial chemistry. Check Vedantu’s page to know more about Van't Hoff Factor Equation and Abnormal Molar Mass.
JEE Main Chemistry Solutions Study Materials
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JEE Main Chemistry Study and Practice Materials
Explore an array of resources in the JEE Main Chemistry Study and Practice Materials section. Our practice materials offer a wide variety of questions, comprehensive solutions, and a realistic test experience to elevate your preparation for the JEE Main exam. These tools are indispensable for self-assessment, boosting confidence, and refining problem-solving abilities, guaranteeing your readiness for the test. Explore the links below to enrich your Chemistry preparation.
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Conclusion
The "Solutions" chapter plays a crucial role in the JEE curriculum, offering essential insights into how solutions behave. It covers various concentration methods, vapor pressure, colligative properties, and Raoult's Law, forming the foundation of physical chemistry. Mastering this chapter is vital for JEE aspirants, enhancing problem-solving and analytical skills. This knowledge extends beyond exam success, proving invaluable for academic and professional pursuits. A solid understanding of this chapter prepares students to navigate competitive exams effectively, making it an indispensable part of JEE preparation. Clearly, the "Solutions" chapter serves as a fundamental stepping stone to success.
Solutions Chapter - Chemistry JEE Main
ShareFAQs on Solutions Chapter - Chemistry JEE Main
1. For JEE Main preparation, what is the key difference between Molarity (M) and Molality (m), and why is one preferred over the other in certain calculations?
The key difference lies in their definitions and temperature dependence. Molarity (M) is the number of moles of solute per litre of solution, whereas Molality (m) is the number of moles of solute per kilogram of solvent. Molality is preferred for studies involving temperature changes (like elevation in boiling point or depression in freezing point) because the mass of the solvent does not change with temperature, unlike the volume of the solution, which can expand or contract.
2. How does Raoult's Law define the vapour pressure of a solution containing a non-volatile solute?
For a solution with a non-volatile solute, Raoult's Law states that the partial vapour pressure of the solvent above the solution (P₁) is directly proportional to its mole fraction (x₁) in the solution. The presence of the non-volatile solute lowers the vapour pressure of the solvent. The mathematical expression is P₁ = P₁°x₁, where P₁° is the vapour pressure of the pure solvent.
3. What are the two essential conditions for a solution to be considered an 'ideal solution' as per Raoult's Law?
A solution is considered ideal if it obeys Raoult's Law over the entire range of concentration. The two primary thermodynamic conditions for an ideal solution are:
- The enthalpy of mixing (ΔHmix) is zero. This means no heat is evolved or absorbed when the components are mixed.
- The volume of mixing (ΔVmix) is zero. This means the total volume of the solution is exactly equal to the sum of the volumes of the individual components.
4. Why do real solutions exhibit positive or negative deviations from Raoult's Law?
Real solutions deviate from Raoult's Law because the intermolecular forces of attraction between the components are different from those in the pure components.
- Positive Deviation: Occurs when the new solute-solvent (A-B) interactions are weaker than the original solute-solute (A-A) and solvent-solvent (B-B) interactions. This makes it easier for molecules to escape into the vapour phase, resulting in a higher vapour pressure than predicted.
- Negative Deviation: Occurs when the new solute-solvent (A-B) interactions are stronger than the original interactions. This holds the molecules more tightly in the solution, making it harder for them to vaporise and resulting in a lower vapour pressure than predicted.
5. What is the significance of the van't Hoff factor (i) in the context of colligative properties for JEE Main?
The van't Hoff factor (i) is crucial as it corrects the colligative property formulas for solutes that behave abnormally in solution. It is the ratio of the observed colligative property to the calculated colligative property.
- For dissociation (e.g., NaCl → Na⁺ + Cl⁻), i > 1, as the number of particles in the solution increases.
- For association (e.g., ethanoic acid dimerising in benzene), i < 1, as the number of effective particles decreases.
- For non-electrolytes (e.g., glucose, urea), i = 1.
6. How do the colligative properties 'Elevation in Boiling Point' and 'Depression in Freezing Point' relate to the concentration of the solute?
Both properties are directly proportional to the molal concentration (m) of the solute in a dilute solution.
- Elevation in Boiling Point (ΔTb): The increase in boiling point is given by ΔTb = i * Kb * m, where Kb is the Ebullioscopic Constant.
- Depression in Freezing Point (ΔTf): The decrease in freezing point is given by ΔTf = i * Kf * m, where Kf is the Cryoscopic Constant.
7. What is an azeotropic mixture, and why can't it be separated by fractional distillation?
An azeotropic mixture, or azeotrope, is a binary mixture of liquids that has a constant boiling point and whose vapour has the same composition as the liquid. Because the liquid and vapour phases have identical compositions, the mixture boils at a single temperature without any change in composition. Therefore, its components cannot be separated by fractional distillation, a technique that relies on differences in boiling points and compositions.
8. How is osmotic pressure conceptually different from other colligative properties like vapour pressure lowering?
While all colligative properties depend on the number of solute particles, osmotic pressure (π) is distinct in its mechanism and measurement. It is not a property of the solution itself (like boiling point) but is an equilibrium pressure that must be applied to a solution to prevent the inward flow of pure solvent across a semipermeable membrane. Its magnitude is generally much larger than other colligative properties for the same concentration, making it a more sensitive method for determining the molar masses of macromolecules like proteins and polymers.
9. What is the principle of reverse osmosis, and how does it relate to osmotic pressure?
Reverse Osmosis (RO) is a process where the natural flow of a solvent is reversed. If an external pressure greater than the solution's osmotic pressure (π) is applied to the solution side of a semipermeable membrane, the solvent molecules are forced to move from the region of higher solute concentration (the solution) to the region of lower solute concentration (the pure solvent). This principle is widely used in water purification and desalination.
10. How can colligative properties be used to determine if an unknown solute undergoes association or dissociation in a solvent?
By comparing the experimentally determined molar mass (Mobserved) with the theoretically calculated normal molar mass (Mnormal), one can determine the solute's behaviour. First, calculate the van't Hoff factor, i = Mnormal / Mobserved.
- If i > 1, it implies Mobserved < Mnormal, indicating that the solute has dissociated into a greater number of particles.
- If i < 1, it implies Mobserved > Mnormal, indicating that the solute has associated to form larger particles.
