Types of Solutions

• A solution happens when a homogenous mixture includes two or more than two components.
• Properties, including utility, depend on composition.

Types of Solutions

• Solutions are homogeneous mixtures.
• Homogenous mixtures are uniform in composition and properties.
• The component in the largest quantity is the solvent; it determines the solution's physical state.
• Solutes are other components in the solution.
• Only binary solutions, i.e., consisting of two components, are considered.
• Components can be solids, liquids, or gases.
• Gas in gas solution example: Mixture of oxygen and nitrogen gases.
• Liquid in gas solution example: Chloroform mixed with nitrogen gas.
• Solid in gas solution example: Camphor in nitrogen gas.
• Gas in liquid solution example: Oxygen dissolved in water.
• Liquid in liquid solution example: Ethanol dissolved in water.
• Solid in liquid solution example: Glucose dissolved in water.
• Gas in solid solution example: Solution of hydrogen in palladium.
• Liquid in solid solution example: Amalgam of mercury with sodium.
• Solid in solid solution example: Copper dissolved in gold.

Expressing Solution Concentration

• Concentration can be expressed qualitatively or quantitatively.
• "Dilute" describes a relatively small quantity of solute.
• "Concentrated" describes a relatively large quantity of solute.
• Quantitative descriptions are more precise.

Mass Percentage (w/w)

• Mass percentage is the mass of a component in a solution divided by the total mass of the solution, multiplied by 100.
• 10% glucose in water by mass contains 10 g of glucose in 90 g of water.
• Mass percentage is commonly used in industrial chemical applications.
• Commercial bleaching solutions contain 3.62 mass percentage of sodium hypochlorite in water.

Volume Percentage (V/V)

• Volume percentage is the volume of the component divided by the total volume of the solution, multiplied by 100.
• 10% ethanol solution in water means that 10 mL of ethanol is dissolved in water for a total volume of 100 mL.
• Solutions containing liquids are commonly expressed in this unit.
• 35% (v/v) ethylene glycol solution is used as an antifreeze in cars, lowering the freezing point of water.

Mass by Volume Percentage (w/V)

• Mass by volume percentage is the mass of solute dissolved in 100 mL of solution, commonly used in medicine and pharmacy.

Parts Per Million (ppm)

• Parts per million expresses concentration when a solute is present in trace quantities.
• ppm is defined as the number of parts of the component divided by the total number of parts of all components in the solution, multiplied by 10^6.
• Concentration in parts per million can be expressed as mass to mass, volume to volume, and mass to volume.
• A liter of seawater weighs 1030 g and contains about 6 × 10^-3 g of dissolved oxygen (O2) or 5.8 ppm of sea water.
• Pollutant concentration in water/atmosphere is often expressed in µg mL⁻¹ or ppm.

Mole Fraction

• Mole fraction (x) is the number of moles of the component divided by the total number of moles of all components.
• Formula for mole fraction (x) in a binary mixture.
• The sum of all the mole fractions in a given solution is unity.
• Mole fraction is useful in relating physical properties like vapor pressure to solution concentration and in describing gas mixture calculations.

Molarity

• Molarity (M) is the number of moles of solute dissolved in one liter (or one cubic decimeter) of solution.
• A 0.25 mol L⁻¹ (or 0.25 M) solution of NaOH means that 0.25 mol of NaOH has been dissolved in one liter.

Molality

• Molality (m) is the number of moles of solute per kilogram (kg) of the solvent.
• A 1.00 mol kg¯¹ (or 1.00 m) solution of KCl means that 1 mol of KCl is dissolved in 1 kg of water.
• Mass %, ppm, mole fraction, and molality are independent of temperature.
• Molarity is a function of temperature; volume depends on temperature while mass does not.

Solubility

• Solubility measures the maximum amount of substance that can dissolve in a specified amount of solvent at a specific temperature.
• Solubility depends on the natures of solute and solvent, temperature, and pressure.

Solubility of a Solid in a Liquid

• Not every solid dissolves in a given liquid.
• Polar solutes dissolve in polar solvents, and non-polar solutes dissolve in non-polar solvents.
• "Like dissolves like" says a solute dissolves in a solvent if the intermolecular interactions are similar.
• Dissolution is when a solid solute is added to a solvent, some solute dissolves, and its concentration increases.
• Crystallization is when some solute particles in solution collide with solid solute particles and separate.
• A stage is reached when dissolution and crystallization occur at the same rate, establishing dynamic equilibrium and a saturated solution.
• Unsaturated solution vs saturated solution.
• Solubility of one substance into another depends on the natures of the substances plus temperature and pressure.

Effect of Temperature

• Temperature changes significantly affect solid solubility in a liquid.
• Approaching a nearly saturated solution, the dynamic equilibrium follows Le Chatelier's Principle.
• Endothermic dissolution (∆sol H > 0) increases solubility with rising temperature.
• Exothermic dissolution (∆sol H < 0) decreases solubility with rising temperature.

Effect of Pressure

• Pressure has no significant effect on the solubility of solids in liquids.
• Solids and liquids are highly incompressible and practically unaffected by pressure changes.

Solubility of a Gas in a Liquid

• Many gases dissolve in water, but to varying extents.
• Solubility of gases in liquids is greatly affected by pressure and temperature.
• Solubility of gases increases with increasing pressure.

Henry's Law

• Quantitative relationship between pressure and gas solubility in a solvent.
• At constant temperature, gas solubility in a liquid is directly proportional to the partial pressure of the gas above the liquid or solution.
• The mole fraction of gas in the solution is proportional to the partial pressure of the gas over the solution.
• The partial pressure of the gas in the vapor phase (p) is proportional to the mole fraction of the gas (x) in the solution, expressed as p = KHx.
• KH is the Henry's law constant.
• Different gases have different KH values at the same temperature, indicating that KH is a function of the gas's nature.
• Higher KH value at a given pressure means lower gas solubility in that liquid.
• KH values for N₂ and O₂ increase with increasing temperature, which means that gas solubility decreases.
• Aquatic species are more comfortable in colder waters due to increased oxygen.

Applications of Henry's Law

• To increase CO2 solubility in soft drinks and soda water, the bottle is sealed under high pressure.
• Scuba divers cope with dissolved gases at high pressures underwater.
• Increased pressure raises atmospheric gas solubility in blood as divers breathe compressed air underwater.
• When divers ascend, decreasing pressure releases dissolved gases and leads to nitrogen bubbles in the blood. The blocks capillaries and creates a medical condition known as bends, which are painful and dangerous to life.
• To avoid bends and toxic effects of high nitrogen concentrations, scuba divers use tanks filled with air diluted with helium.
• High altitudes have lower partial oxygen pressures than at ground level, leading to lower oxygen concentrations in blood/tissues.
• Low blood oxygen causes climbers to become weak and unable to think clearly: a condition known as anoxia.

Effect of Temperature

• With rising temperature, the solubility of gases in liquids decreases, similar to condensation.
• Because dissolution involves dynamic equilibrium, it follows Le Chatelier's Principle.
• As dissolution of gases tends to be an exothermic process, solubility decreases as temperature increases.

Vapor Pressure of Liquid Solutions

Vapor Pressure of Liquid-Liquid Solutions

• Liquid solutions are formed when the solvent is a liquid.
• The solute can be a gas, a liquid, or a solid.
• Consider a binary solution of two volatile liquids (components 1 and 2) in a closed vessel, where both components evaporate and eventually reach equilibrium between the vapor and liquid phases.
• The total vapor pressure at this stage is ptotal, and the partial vapor pressures are p1 and p2.
• These partial pressures relate to x1 and x2 (mole fractions of the two components).

Raoult's Law

• For a solution of volatile liquids, the partial vapor pressure of each component in the solution is directly proportional to its mole fraction in the solution.
• The mathematical expressions of this law are outlined.
• Dalton's law of partial pressures states that the total pressure over the solution phase equals to the sum of the partial pressures of the components of the solution.
• Conclusions can be drawn from the equations regarding total vapor pressure.

Graphical Representation

• Plots and conclusions about their meaning, outlined.

Types of Solutions

• Liquid-liquid solutions are classified into ideal and non-ideal solutions on the basis of Raoult's law.

Ideal Solutions

• Solutions obeying Raoult's law over the entire concentration range.
• Ideal solutions have zero enthalpy and volume of mixing.
• Intermolecular attractive forces between the components are nearly equal.
• A perfectly ideal solution is rare, bur some solutions are nearly ideal in behavior.

Non-Ideal Solutions

• A solution that does not obey Raoult's law.
• The vapor pressure of such a solution is either higher (positive deviation) or lower (negative deviation) than that predicted by Raoult's law.
• The cause of these result from the nature of interactions at the molecular level.
• If A-B interactions are weaker than A-A or B-B, molecules of A (or B) will find it easier to escape than in pure state. This will increase the vapor pressure and result in positive deviation.

Azeotropes

• Some liquids, when mixing, form azeotropes, or binary mixtures with the same liquid/vapor phase composition that boil at a constant temperature.
• Fractional distillation cannot separate the components in these cases.
• Minimum and maximum boiling azeotropes vs Raoult's law solutions.

Colligative Properties and Determination of Molar Masses

• Decreased vapor pressure of solutions has many connected properties
• They are: relative lowering of vapor pressure of the solvent, depression of freezing point of the solvent, elevation of boiling point of the solvent, and osmotic pressure of the solution.
• They depend on number of solute particles relative to the total number of particles present in the solution.

Relative Lowering of Vapor Pressure

• Vapor pressure of a solvent in solution is less than that of the pure solvent, depending on the concentration of solute particles and is independent of their identity.
• Raoult established equation of the above.
• The reduction in the vapor pressure of solvent is outlined with formulas
• The left-hand side expression is called relative lowering of vapor pressure and is equal to the solute's mole fraction.
• This allows the molar mass of a solute to be calculated.

Elevation of Boiling Point

• The vapor pressure of a liquid increases with increasing temperature.
• It boils at the temperature at which its vapor pressure equals atmospheric pressure.
• Adding a non-volatile solute decreases the vapor pressure of the solvent.
• The elevation of boiling point depends on the number of solute molecules.
• Let To be the boiling point of pure solvent and To be the boiling point of solution.
• The difference between them known as elevation of boiling point.
• For dilute solutions the elevation of boiling point is directly proportional to the molal concentration of the solute in a solution.
• This has been formalized in formals listing constants
• Thus, molar mass of the solute can be determined

Depression of Freezing Point

• Lowering of vapor pressure causes a lowering of the freezing point compared to the pure solvent.
• At freezing point solid phase is in dynamic equilibrium with the liquid phase.
• Vapor pressure will reach will now drop equal at lower temperatures, thus freezing point decreases
• Solvent freezing point (T^o_f) and T_f are referenced with formulas and explanations as before

Osmosis and Osmotic Pressure

• Various natural phenomena depend on solute-solvent membranes: semipermeable membranes (SPM).
• Solvent molecules flow through membrane from pure solvent to solution: Osmosis.
• This has been formalized in formulas.
• Excess pressure on the solution side to achieve equilbirum is called osmotic pressure
• The solute only effects the number of molecules not identity
• This method is widely utilized determine molar masses of proteins, polymers and other
• Measure using osmotic pressure methods: better temperature measurement
• Two solutions exerting the same osmotic pressure = isotonic solutions.

Reverse Osmosis and Water Purification

• If applied to solution the osmosis direction reverses.
• This called, reverse osmosis, applied via device.
• Quite simple in theory, application is harder, especially for saltwater, but membrane are getting better.

Abnormal Molar Masses

• We can determine compound ion makeup by measuring this. This called degree of dissociation when value is measured.
• Sometimes the reverse occurs in benzene (low dielectric), in this case dimerzation has occurred and also can be quantified
• Molar masses can be experimentally different than the real because of the above.
• This also done utilizing the van't Hoff factor which is calculated for any reaction.