Understanding Solutions: Types, Properties, and Laws

• Solutions are homogeneous mixtures of two or more components and their utility depends on their composition.
• Body processes occur in liquid solutions.
• The properties of brass, which contains copper and zinc, differ from German silver, which contains copper, zinc, and nickel, and bronze, which contains copper and tin.
• Fluoride ions at 1 ppm prevent tooth decay; 1.5 ppm causes mottling; higher concentrations are poisonous.
• Injections are dissolved in water with specific salt concentrations to match blood plasma.

Objectives of Studying Solutions

• Describe different solutions' formation.
• Express solution concentration using different units.
• Understanding Henry's and Raoult's Laws.
• Differentiate ideal from non-ideal solutions.
• Explain real solution deviations from Raoult's Law.
• Relate colligative properties to molar masses.
• Explain solutes' abnormal colligative properties.

Types of Solutions

• Solutions are homogeneous mixtures with uniform composition and properties.
• The solvent, present in the largest quantity, determines the solution's physical state.
• Solutes are components other than the solvent.
• In this unit, only binary solutions consisting of two components are mostly considered.

Types of Solutions Overview

• Gaseous solutions include gas in gas (oxygen and nitrogen mixture)
• Liquid in gas (chloroform in nitrogen)
• Solid in gas (camphor in nitrogen).
• Liquid solutions include gas in liquid (oxygen in water)
• Liquid in liquid (ethanol in water)
• Solid in liquid (glucose in water).
• Solid solutions include gas in solid (hydrogen in palladium)
• Liquid in solid (mercury in sodium amalgam)
• Solid in solid (copper in gold).

Expressing Concentration of Solutions

• Solution composition can be described qualitatively (dilute or concentrated) or quantitatively.
• Quantitative descriptions reduce confusion.

Mass Percentage (w/w)

• Mass percentage is the mass of a component in the solution divided by the total mass of the solution, multiplied by 100.
• A 10% glucose solution has 10 g of glucose in 100 g of solution.
• This percentage is commonly used in industrial applications, ex: bleaching solutions containing 3.62% sodium hypochlorite.

Volume Percentage (V/V)

• Volume percentage is the volume of the component divided by the total solution volume, multiplied by 100.
• A 10% ethanol solution contains 10 mL of ethanol in 100 mL of solution.
• Commonly used for solutions containing liquids.
• For example, a 35% ethylene glycol solution is used as antifreeze.

Mass by Volume Percentage (w/V)

• Defined as the mass of solute in 100 mL of solution.
• Commonly used in medicine and pharmacy.

Parts Per Million (ppm)

• Used for solutes in trace amounts.
• Calculated as the number of parts of the component divided by the total parts, multiplied by 10^6.
• Can be expressed as mass to mass, volume to volume, or mass to volume.
• Seawater contains about 5.8 ppm of dissolved oxygen and pollutant concentrations in water or atmosphere.

Mole Fraction

• Mole fraction (x) with a subscript denotes the component and is defined as moles of the component divided by total moles in the solution.
• In a binary mixture of A and B, the mole fraction of A is nA / (nA + nB).
• The sum of all mole fractions in a solution is unity: x1 + x2 + ... + xi = 1.
• Useful in relating vapor pressure with solution concentration and in describing gas mixtures.

Molarity (M)

• Molarity is the number of solute moles per liter of solution.
• A 0.25 M NaOH contains 0.25 mol NaOH per liter.

Molality (m)

• Molality is the solute moles per kilogram of solvent.
• A 1.00 m solution of KCl contains 1 mol of KCl per kg of water.
• Each concentration expression has merits and demerits.

Temperature Dependence of Concentration

• Mass %, ppm, mole fraction, and molality are temperature-independent.
• Molarity varies with temperature due to volume changes.

Solubility

• Solubility refers to the maximum solute amount dissolved in a solvent at a specific temperature.
• Depends to the solute, solvent, temperature, and pressure.

Solubility of a Solid in a Liquid

• Not every solid dissolves in a liquid with sodium chloride and sugar readily dissolving in water, unlike naphthalene and anthracene.
• Polar solutes dissolve in polar solvents; nonpolar solutes dissolve in nonpolar solvents, which known as "like dissolves like."
• Dissolution is when a solid solute is added to a solvent and its concentration increases.
• Crystallization is when solute particles collide with solid solutes, separating them.
• Dynamic equilibrium is reached when dissolution and crystallization rates are same.

Saturated and Unsaturated Solutions

• Saturated solutions cannot dissolve any more solute at a given temperature and pressure.
• Unsaturated solutions can dissolve more solute at the same temperature.
• A solution in dynamic equilibrium with undissolved solute is saturated and contains the maximum solute amount.

Temperature and Pressure Effect on Solid Solubility

• Temperature significantly impacts solubility.
• Pressure has negligible effect on solid and liquid solubility due to their incompressibility.
• Le Chatelier's Principle says that endothermic dissolution (Δsol H > 0) increases solubility with temperature, and exothermic dissolution (Δsol H < 0) decreases solubility.

Solubility of a Gas in a Liquid

• Gases dissolve in water to varying degrees, with oxygen sustaining aquatic life and hydrogen chloride being highly soluble.
• Solubility increases with pressure and decreases with temperature.

Henry's Law

• State that at a constant temperature, the solubility of a gas in a liquid is directly proportional to the gas's partial pressure above the liquid.
• The partial pressure of the gas in vapor phase (p) is proportional to the mole fraction of the gas (x) in the solution: p = KHx.
• KH is Henry's law constant which varies among gases and with temperature.
• Higher KH indicates lower solubility.

Applications of Henry's Law

• Used to increase CO2 solubility in soft drinks by sealing bottles under high pressure.
• Explains why scuba divers cope with high dissolved gas concentrations underwater.
• Rapid pressure decreases as divers ascend, producing nitrogen bubbles in the blood in a condition known as "bends" which can be avoided by using helium-diluted air.
• Explains why people at high altitudes have low blood oxygen, leading to anoxia.

Temperature Effect on Gas Solubility

• Gas solubility in liquids decreases with temperature.
• Because dissolution can be considered a type of condensation
• It is exothermic and follows Le Chatelier's Principle.

Vapor Pressure of Liquid Solutions

• Liquid solutions form when the solvent is a liquid and include solutions of gases, liquids, or solids in a liquid.
• Solutions contain one or more volatile components and the solvent is generally volatile and the solute may or may not be.
• The focus is mainly on binary liquid solutions with (i) liquids in liquids and (ii) solids in liquids.

Vapor Pressure of Liquid-Liquid Solutions

• Defined as binary solution is considered with two volatile liquids, denoted as 1 and 2.
• Both components evaporate in a closed vessel, establishing equilibrium between the vapor and liquid phases.
• The mole fractions are represented by x1 and x2 and partial vapor pressures are p1 and p2 and total vapor pressure is ptotal.

Raoult's Law

• Raoult's Law states that the partial vapor pressure of each solution component is directly proportional to its mole fraction in the solution.
• Thus, for component 1: p1 α x1 and p1 = p1°x1, where p1° is the vapor pressure of pure component 1.
• The vapor pressure of the pure component 2 is expressed by p2 = p2°x2.

Dalton's Law of Partial Pressures

• Dalton's Law of Partial Pressures states that the total pressure over the solution (Ptotal) is the sum of the partial pressures of the components: Ptotal = p1 + p2.
• Leading to Ptotal = x1p1° + x2p2° = (1 – x2)p1° + x2p2° = p1° + (p2° – p1°)x2

Conclusions from Raoult's and Dalton's Laws

• Total vapor pressure can be related to the mole fraction of one component.
• Total vapor pressure varies linearly with the component 2 mole fraction.
• Increase/decrease depends on pure components' vapor pressures and the increase of mole fraction of component 1.
• Composition of vapor phase is defined by partial pressures of components.
• The component mole fractions are are y1 and y2 respectively.

Ideal vs. Non-Ideal Solutions

• Ideal solutions are solutions that obey Raoult's law over the entire concentration range.
• They have two other properties: AmixH = 0 which means no heat is absorbed or evolved and AmixV = 0 meaning the solution volume equals the sum of the components' volumes.
• These solutions are explained by considering the types of intermolecular interactions.

Interactions in Ideal Solutions

• In pure components, intermolecular interactions are A-A and B-B.
• In binary solutions, A-B interactions are also present.
• The solutions behave ideally when the forces between A-A and B-B are nearly equal to A-B.
• n-hexane and n-heptane, bromoethane and chloroethane, and benzene and toluene fall into this category.

Non-Ideal Solutions

• Fail to obey Raoult's law over the entire range of concentration.
• Positive deviation happens if solution exhibits a higher vapor pressure.
• Negative deviation happens if the vapor pressure is lower.

Causes of Deviations in Non-Ideal Solutions

• Positive deviation is caused by A-B interactions weaker than those between A-A or B-B which means the molecules escape easier in the pure state than in solutions.
• In ethanol-acetone mixtures, adding acetone breaks some hydrogen bonds and causes the positive deviation.
• Negative deviations occur with A-A and B-B attractions weaker than those of A-B, leading to decreased vapor pressures.
• Chloroform and acetone mixtures form hydrogen bonds, resulting in negative deviation.

Explore the composition and utility of solutions, crucial in various applications like body processes and material science. Understand solution formation, concentration units, and the distinctions between ideal and non-ideal solutions. Delve into Henry's and Raoult's Laws, colligative properties, and real solution deviations.