PHT 231 LECTURE TU6-Ideal and real solutions.pptx

Real and ideal solutions

1. PHYSICAL CHEMISTRY I TU6: IDEAL AND REAL SOLUTIONS, SOLUTIONS OF GASES IN LIQUIDS, COLLIGATIVE PROPERTIES, PH: determination of PH buffers, theory of buffers.

2. GENERAL INTRODUCTION Solute -Is a substance that can be dissolved in a solvent to produce a solution. Solvent -It is the part that the solute is dissolved in. Solution -Is a homogeneous mixture of one or more solutes dissolved in a solvent.

3. Solutions - Can be divided into three types: gaseous, liquid and solid solution 1. Gaseous solution -solutions in which solvent is present in gaseous state. This can be divided into three types based on the phases of solute and solvent . a. Gas-gas solution-Solutions in which solute and solvent are both gases. For example-solution (mixture) of nitrogen and oxygen. b. Liquid-gas solution-Solutions in which solute is in liquid state and solvent is in the gaseous state. For example, solution of chloroform in nitrogen gas. c. Solid-gas solution-Solutions in which solute is in its solid state and solvent is in gaseous state. For example, solution of camphor in nitrogen gas.

4. 2. Liquid solution Solutions in which solvent is present in liquid state. a. Gas-liquid solution -Solutions in which solute is in in gaseous state and solvent in liquid state. For example-solution (mixture) of oxygen in water. b. Liquid-liquid solution-Solutions in which solute and solvent are both in the liquid state. For example, solution of ethanoic acid and water (vinegar). c. Solid-liquid solution-Solutions in which solute is in its solid state and solvent is in liquid state. For example, solution of glucose in water.

5. 3. SOLID SOLUTION Solutions in which solvent is present in solid state. a. Gas-solid solution-Solutions in which solvent is in solid state and solute in gaseous state. For example-solution (mixture) of hydrogen and palladum. b. Liquid-solid solution-Solutions in which solvent is in liquid state and solute is in the liquid state. For example, solution of amalgum of mercury with sodium.. c. Solid-solid solution-Solutions in which solute and solvent are both in the solid state. For example, solution of gold and copper.

6. RAOULT’S Law In 1986, a French chemist (Francois Marte Raoult) Proposed a relationship between partial pressure and mole fraction of volatile liquids. According to the law, ‘the mole fraction of the solute component is directly proportional to its partial pressure’. (PV=nRT) On the basis of this law, liquid solutions can be of two types. 1. Ideal solutions 2. Non-ideal solutions

7. Ideal solutions Solution which obey Raoult’s law under all standard temperature and concentration. No heat is evolved or absorbed during the mixing process. (i.e. It satisfies the ) The solution of two components A and B in which the A---B interactions are of same magnitude as A---A and B---B interaction. Only solutions with low concentration of solute behave ideally. Example: benzene+toluene, chlorobenzene+bromobenzene, n-hexane+n-heptane When the intermolecular forces of attraction between A—A, B—B and A—B are nearly equal.  

8. Non-ideal solutions The solution which do not follow Raoult’s law. Hmixing It is the solution in which solute and solvent molecules interact with one another with a different force than the forces of interaction between the molecules of the pure compounds. Ex. Sulphuric acid (solute) and water (solvent) the amount of heat is evolved is large and thus change in volume is also seen. Non-ideal solutions are of two types: non ideal positive and negative deviations  

9. Deviations Positive deviations In mixtures showing a positive deviation from Raoult’s law, the vapour pressure of the mixture is always higher than you would expect from an ideal mixture. This means that molecules are breaking away more easily than they do in the pure liquids. Negative deviations Here the vapour pressure of the mixture is less than would be expected by the Raoult’s law. This means that molecules break away from the mixture less easily than they do from the pure liquids.

10. SOLUBILITY OF GASES IN LIQUIDS Solubility of gases in liquids is the concentration of dissolved gas in the liquid when it is in equilibrium with the pure gas above the solution. Examples include: effervescent preparations containing dissolved CO2, ammonia and hydrochloride gas. NH3 ( aq )+H2O (l) NH4+( aq ) + OH-( aq ) Aerosol products containing nitrogen or carbon dioxide as propellant The gas solubility is greatly affected by temperature and pressure as well as the nature of the solute and solvent.

11. Factors that influence solubility generally include Temperature Nature of solvent Pressure pH Particle size Crystal structure Molecular structure Solute-solvent interactions Melting and boiling points Addition of substituent Solubilizing agents The two major factors that influence solubility of gases in liquid are: Pressure and temperature

12. 1. Effect of Pressure The solubility of a gas in a liquid is directly proportional to the partial pressure of the gas present above the surface of the liquid or solution. Henry’s law- the partial pressure of a gas above a solution is proportional to the mole fraction of the gas in the solution. P= K H x, where P= partial pressure of the gas X=mole fraction of the gas in solution K H =Henry’s law constant Effect of pressure-increase in pressure increase in solubility

13. 2. Effect of temperature Temperature is a measure of the average kinetic energy. Effect of temperature-increase in temperaturedecrease in solubility Why? As temperature increases, kinetic energy increases. This greater kinetic energy results in greater molecular motion of the gas particles. As a result, the gas particles dissolved in the liquid are more likely to escape to the gas phase and the existing gas particles are less likely to be dissolved.

14. Physical properties of substances Classified as additive, constitutive and colligative A. Additive properties- depend on the total contribution of the atoms in the molecule or on the sum of the properties of the constituents in a solution. - An example of additive property of a compound is the molecular weight, i.e. the sum of the masses of the constituent atoms. - The masses of the components of a solution are also additive. The total mass of the solution being the sum of the masses of the individual components.

15. B. Constitutive properties Depend on the arrangement and to a lesser extent on the number and kind of atoms within a molecule. Many physical properties may be partly additive and partly constitutive. The refraction of light, electric properties, surface and interfacial characteristics, and the solubility of drugs are at least in part constitutive and in part additive properties.

16. C. Colligative properties: depend ONLY on the number of dissolved particles (molecules or ions, small or large) in solution and not on their identity. The properties include: ( i ) Osmotic pressure (ii) Vapour pressure lowering (iii) Boiling point elevation (IV) Freezing point depression

17. (I) OSMOTIC PRESSURE Diffusion in liquids-Substances tend to move or diffuse from regions of higher concentration to region of lower concentration so the differences in concentration disappear. Osmosis is the passage of the solvent molecules from pure solvent into a solution through a semipermeable membrane. Solvent molecules move through a semipermeable membrane (SPM) from a dilute solution to a concentrated solution. The flow of solvent towards the solution side across the SPM can be stopped by applying extra pressure on the solution.

18. Osmotic pressure Is the external pressure that must be applied to the solution in order to prevent it being diluted by the entry of solvent via osmosis. Osmotic pressure depends on the concentration of the solution . Depends on the number of solute molecules and NOT their identity. The osmotic pressure is proportional to the reduction in vapour pressure brought about by the concentration of solute present.

19. Van’t Hoff’s equation:   is the osmotic pressure in atm , V is the volume of the solution in liters , n is the number of moles pf solute, R is the gas constant, equal to 0.082 liter atm /mole degree, T is the absolute temperature (Kelvin) Example: 1 g of sucrose, molecular weight=342 g/mol, is dissolved in 100 mL of solution at 25 ⁰C. What is the osmotic pressure of the solution? Solution: Step 1: calculate the moles of sucrose Moles of sucrose= = 0.0029 mol Step 2: Using Van’t Hoff’s equation, = 0.71 atm  

20. Osmotic pressure of non-electrolytes (not ionized) Osmotic pressure . This means twice concentration, twice osmotic pressure number of molecules . This means osmotic pressure of two solutions having the same molal concentration are identical. Example: Explain why solution contains 34.2 g sucrose (mol wt 342) in 1000 g water has the same osmotic pressure as dextrose solution (mol wt 180) contains 18 g/1000 g water? Solution: No of moles of sucrose= wt /mol wt (=34.2/342=0.1 molal) - No of moles of dextrose =18/180=0.1 molal. Therefore, the two solutions are iso-osmotic .  

21. Important terms to note Isotonic solutions-Solutions that have the same osmotic pressure at a given temperature. Hypotonic solutions-Solutions that have a lower osmotic pressure than other solutions. Hypertonic solutions- Solutions that have higher osmotic pressure than other solutions.

22. (II) Vapour pressure lowering When a non-volatile solute is dissolved in solvent, the vapour pressure of the solvent is lowered . Solvent molecules on the surface which can escape into vapour is replaced by solute molecules and have little (if any) vapour pressure. The lowering of vapour pressure is  

23. (III) BOILING POINT (BP) ELEVATION The boiling point of a liquid is defined as the temperature at which the vapour pressure of that liquid equals the atmospheric pressure (760 mmHg). For a solution, the vapour pressure of the solvent is lower at any given temperature. Therefore, a higher temperature is required to boil the solution than the pure solvent. Elevation of the BP- The bp of a solution of a non-volatile solute is higher than that of a pure solvent. This is due to the fact that the solute lowers the vapor pressure of the solvent. The more of the solute that is dissolved, the greater is the effect.

24. Addition of solute will decrease the vapour pressure and so will increase the boiling point. The elevation of the boiling point is given as m-molality(molality is the moles of solute per kilograms of solvent) Kb : boiling point elevation constant that depends on the particular solvent being used. ( Kb water=0.51 ⁰C/m)  

25. (IV) FREEZING POINT DEPRESSION Freezing (melting) point is the temperature at which solid and liquid are in equilibrium under 1 atm. Addition of solute will decrease the vapour pressure and so will decrease the freezing point. For a liquid to freeze, It MUST achieve a very ordered state that results in the formation of a crystal. Depression of the freezing point If a solute is dissolved in the liquid at the triple point, the escaping tendency or vapour pressure of the liquid solvent is lowered below that of the pure solid solvent. The temperature must drop to reestablish equilibrium between the liquid and the solid. Due to this, the freezing point of a solution is always lower than that of the pure solvent. The more concentrated the solution, the farther apart are the solvent and the greater is the freezing point depression.  

26. As per the law, the freezing point for a given dilute solution stays directly proportional to molality of the solute, same as the boiling elevation point change in temperature  

27. Example: What is the freezing point of a solution containing 3.42 g of sucrose and 500 g of water? (The molecular weight of sucrose is 342. In this relatively dilute solution, Kf is approximately equal to -1.86 ⁰C/m) Solution: = =- 0.037 ⁰C Therefore the freezing point of the aqueous solution is -0.037 ⁰C  

28. Buffer solutions A buffer solution is a solution that can resist changes in pH only slightly when small amounts of strong acid or strong bases are added. A buffer contains significant concentrations of both -weak acid and its conjugate base Weak base and its conjugate acid Examples of buffers: Sodium acetate buffer-this buffer relies on the dissociation reaction of acetic acid. CH3COOH  ch3coo- + H+

29. Buffer capacity and range - There is a limit to the ability of a buffer solution to neutralize added acid or base. The buffer capacity is reached before either buffer component is consumed. As a rule, a buffer is most effective if the concentrations of the buffer acid and its conjugate base are equal or nearly so

30. PH: determination of PH buffers, theory of buffers. The pH of a buffer is determined by two factors: The equilibrium constant Ka of the weak acid Different weak acids have different equilibrium constants (Ka). Ka tells us what proportion of HA will be dissociated into H+ and A – in solution. The more H+ ions that are created, the more acidic and lower the pH of the resulting solution. 2. The ratio of weak base [A-] and weak acid [HA] in solution. The ratio of weak base [A-] to weak acid [HA] in a buffer also affects the pH. -If a buffer has more base than acid, more OH- ions are likely to be present and the pH will rise. -If a buffer has more acid than base, more H+ ions are present and the pH will fall. -When the concentrations of A- and HA are equal, the concentration of H+ is equal to Ka, (or equivalently pH=pKa )

31. Handerson-Hasselbalch equation In predicting or determining the pH of a buffer solution, the HH equation is used given below pH=pKa+log( ) Since pKa is a property of the weak acid used in selecting our buffer, we can control the pH by manipulating the proportion of weak base (A-) and weak acid (HA) in solution. When remains close to 1, the pH remains close to pKa. As the goes from 1/10 to 10, the pH changes from pKa+1 to pKa-1