1. CHEMICAL EQUILIBRIUM 2

2. Ionic Equilibrium

3. Acid & Base Ionization For weak acids like acetic acid there will be an equilibrium according to its ionization in water: The equilibrium constant known as acid constant is given by:

4. From this equilibrium constant the Henderson- Hasselbalch equation is derived: which is applied in buffer / pH calculations. In the same way the base equilibrium constant is: The form of Henderson-Hasselbalch equation will then be:

5. Solubility Product The solubility of a salt like AgCl will be having the solubility product The solubility product can be written as:

6. Stability or Formation Constant When complexes are formed their formation constant indicates the stability of the complex. The K f or K stab may be written

7. Distribution Coefficient In the fields of organic and medicinal chemistry, a partition (P) or distribution coefficient (D) is the ratio of concentration of a compound in the two phases of a mixture of two immiscible solvents at equilibrium. Hence these coefficients are a measure of differential solubility of the compound between these two solvents.

8. Normally one of the solvents chosen is water while the second is hydrophobic such as octanol. Hence both the partition and distribution coefficient are measures of how hydrophilic ("water loving") or hydrophobic ("water fearing") a chemical substance is. Partition coefficients are useful for example in estimating distribution of drugs within the body.waterhydrophobicoctanolhydrophilichydrophobicdistribution

9. Hydrophobic drugs with high partition coefficients are preferentially distributed to hydrophobic compartments such as lipid bilayers of cells while hydrophilic drugs (low partition coefficients) preferentially are found in hydrophilic compartments such as blood serumlipid bilayersblood serum

10. The distribution coefficient can be written as: Hence the value indicates the solubility is s specific phase.

11. Factors Affecting the Equilibrium Le Châtelier’s Principle “If you disturb an equilibrium, it will shift to undo the disturbance”.

12. Change of concentration In the equation: 2NO(g) + O 2 (g) 2NO 2 (g) If you add more NO(g) the equilibrium shifts to the right producing more NO 2 (g) If you add more O 2 (g) the equilibrium shifts to the right producing more NO 2 (g) If you add more NO 2 (g) the equilibrium shifts to the left producing more NO(g) and O 2 (g)

13. Consider the Haber process If H 2 is added while the system is at equilibrium, the system must respond to counteract the added H 2 (by Le Châtelier). That is, the system must consume the H 2 and produce products until a new equilibrium is established. Therefore, [H 2 ] and [N 2 ] will decrease and [NH 3 ] increases

14. Change in Reactant or Product Concentrations

15. Common ion effect The common ion effect is the shift in equilibrium caused by the addition of a compound having an ion in common with the dissolved substance. The presence of a common ion suppresses the ionization of a weak acid or a weak base. Consider mixture of CH3COONa (strong electrolyte) and CH3COOH (weak acid). –CH 3 COONa (s) ↔ Na + (aq) + CH3COO - (aq) –CH 3 COOH (aq) ↔ H + (aq) + CH3COO - (aq) CH3COO - common ion

16. What is the pH of a solution containing 0.30 M HCOOH and 0.52 M HCOOK? –HCOOH (aq) ↔H+ (aq) + HCOO- (aq) Initial (M) 0.30 0.00 0.52 Change (M) - x +x+x Equilibrium (M) 0.30-x x 0.52 + x 0.30 – x ≈ 0.30 0.52 + x ≈ 0.52

17. Pressure/volume change In the equation 2SO 2 (g) + O 2 (g) 2SO 3 (g), an increase in pressure will cause the reaction to shift in the direction that reduces pressure, that is the side with the fewer number of gas molecules. Therefore an increase in pressure will cause a shift to the right, producing more product. (A decrease in volume is one way of increasing pressure.)

18. Consider the production of ammonia As the pressure increases, the amount of ammonia present at equilibrium increases.

19. Temperature Change The temperature equilibrium constant is governed by van’t Hoff equation: Where K is the equilibrium constant, T is temperature and ∆H is the enthalpy.

20. Effect of Temperature Changes The equilibrium constant is temperature dependent. For an endothermic reaction,  H > 0 and heat can be considered as a reactant. For an exothermic reaction,  H < 0 and heat can be considered as a product. Adding heat (i.e. heating the vessel) favors away from the increase: –if  H > 0, adding heat favors the forward reaction, –if  H < 0, adding heat favors the reverse reaction.

21. Effect of Temperature Changes Removing heat (i.e. cooling the vessel), favors towards the decrease: –if  H > 0, cooling favors the reverse reaction, –if  H < 0, cooling favors the forward reaction.

22. Using a catalyst A catalyst increases the speed in which a reaction takes place, however it never has any effect on the equilibrium.