and base. This compound partially dissociates in water. Thus, the symbol () is used to show that the reaction is reversible. When a Bronsted –Lowry acid is. ions in aqueous solutions which is called as ionic equilibrium. EQUILIBRIUM. IN. PHYSICAL. PROCESSES. The characteristics of system at equilibrium. THE KEY. Fundamentals of Acids, Bases & Ionic Equilibrium. Acids & Bases. When dissolved in water, acids release H+ ions, base release OH– ions. Arrhenius.
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Ionic Equilibria Notes - Download as PDF File .pdf), Text File .txt) or read online. HCI Lecture notes. Ionic Equilibria. Modern Theories of Acids, Bases, and Salts. Species Concentration as a Function of pH. Acid- Base Equilibria. Calculation of. pH. Sorensen's. Acids and Bases Equilibria—Analytical Applications. Front Matter. Pages PDF · Definitions of Acids and Bases: Strength of Acids and Bases. Jean-Louis.
Chemistry for Advanced Level. John Murray. Martin S. Silberberg McGraw Hill. An acid must thus contain H in its formula. Acidbase reactions do not only occur in aqueous solutions. They can also occur between gases, in nonaqueous solutions and in heterogeneous mixtures. What can you say about the acidbase nature of water?
H2O is amphoteric can act as an acid or a base, depending on the other substance present. Other definitions of acids and bases: Arrhenius acidbase definition. A base is a substance that dissociates in water to produce OH aq. Acids are species that accept an electron pair, e. AlCl3 Bases are species that donate an electron pair e.
NH3 2. Conjugate AcidBase Pair Consider the following acidbase reaction: CH3CO2H acid. In the forward reaction: In the reverse reaction: In general, conjugate acidbase pair acid 1. An acid after donating a proton forms its conjugate base.
A base after accepting a proton forms its conjugate acid. A conjugate acidbase pair constitutes two species which differ from each other by a proton.
In any acidbase reaction, there are two conjugate acidbase pairs. The conjugate base of a strong acid has a very low tendency to accept a proton. So the reverse reaction is negligible singleheaded arrow. Other examples of strong acids: The conjugate base of a weak acid has a tendency to accept a proton. So the reverse reaction occurs to some extent doubleheaded arrow. Other examples of weak acids: Other examples of strong bases: The strength of an acid or base is different from the concentration of the acid or base.
It is not a measure of the strength of the acid unless the two acids being compared have the same initial concentration see Section 2. Similarly, a convenient indication of the concentration of OH aq is to express it in terms of its negative logarithm to base Ionic product of water, Kw Pure water conducts electricity slightly.
This indicates the presence of trace concentrations of ions which arise from the very slight selfionisation or autoionisation of water.
H2O l. Since the extent of ionisation is very small, [H2O] is almost constant at Taking lg on both sides of equation: The value of Kw depends only on temperature. It is always a constant at constant temperature. Experimentally, at 25oC,. Thus, an acidic solution contains a few hydroxide ions, and an alkaline solution contains a few hydrogen ions. The association of water molecules by hydrogen bonding increases as temperature increases. We will need to include it in the calculation.
An aqueous solution contains 1. An aqueous solution contains 0. Calculate the concentration OH in mol dm3, and hence the pH of this solution at 25 oC. Calculate the pH of this solution at 25 oC.
Solution X contains HCl, it has pH 2. The two solutions are mixed. Calculate the pH of the acid mixture. An aqueous solution contains 1 x mol dm3 of HNO3.
Calculate the pH of this solution. In dilute aqueous solution, [H2O] is almost constant at The equilibrium constant may be written as:. For strong acids e.
HCl, Ka is very large and not used. Like other equilibrium constants, the value of Ka depends only on temperature. The value of Ka indicates the extent to which the weak acid dissociates in water at the specified. To compare the strength of two weak acids, compare their Ka values. At the same temperature, larger Ka smaller pKa stronger acid.
Base dissociation constant, Kb A weak base dissociates partially in water to give OH. The equilibrium constant may be written as: For strong bases e. NaOH, Kb is very large and not used.
The value of Kb depends only on temperature. The value of Kb indicates the extent to which the weak base dissociates in water at the specified. To compare the strength of two weak bases, compare their Kb values.
At the same temperature, larger Kb smaller pKb stronger base. Hence, we can determine the value of Ka, given the Kb value of its conjugate base and vice versa. Hence CH3CO2 has a tendency to accept a proton from water.
Degree of Dissociation The fraction of molecules which is ionised into ions in water is called the degree of dissociation,. For an acid,. Generally, for a weak monobasic acid,. Calculate the pKa of the acid. Write the Kb expression of caffeine and calculate its value at 25 oC. Calculate the degree of dissociation, for caffeine at 25 oC. Strength of Acids and Bases The strength of an acid or base: Now, find a the pH and b the degree of dissociation of ethanoic acid of concentration 0. How have the pH and been affected by the dilution?
Since pH and degree of dissociation of an acid or base will change with change in concentration, they are not reliable indicators of acid or base strength. Hence, the values of pH, and Ka for a weak acid at a constant temperature will vary with dilution i.
Ka does not vary with concentration.
It is a constant at constant temperature. Therefore Ka is the best indicator of the strength of a weak acid. Suggest with brief reasoning which is the stronger acid and which is the stronger base. Kc is very small, this suggests that the position of equilibrium lies very much to the left.
Kc is very large, this suggests that the position of equilibrium lies very much to the right. Recall from Section 2. From part b , we see that the stronger acid is able to protonate the conjugate base of the weaker acid.
Strong acids e. HCl are often used to liberate the weak acids e. CH3CO2H from their salt e. Salt Hydrolysis and pH of Salt Solutions An ionic salt may be prepared from a reaction between an aqueous acid and an aqueous base.
Salts derived from a strong acid and a strong base e. NaCl, KNO3 do not undergo salt hydrolysis. Salt derived from a strong acid and a weak base E. The anion Cl does not hydrolyse.
Hence a reversible arrow is used in the salt hydrolysis equation. The anion Cl derived from strong acid HCl is a weaker base than water so does not hydrolyse. Salt derived from a weak acid and a strong base E. CH3CO2 aq. Salt derived from a weak acid and a weak base E. Salt containing an aqueous metal cation with high charge density E. The hydrated cations are coordinated to water molecules through dative coordinate bonds to form complexes.
When one OH bond breaks, a proton is released: The Ka values of hydrated metal cations are given in the table: In the presence of a base stronger than water e. OH, further abstraction of protons can occur. Write an equation, including state symbols, for any hydrolysis reaction. Na2CO3 aq , given: Both ions undergo hydrolysis: H2S is a weak diprotic acid Ka1: The salt solution is alkaline. Ka value for ethanoic acid is 1. What are Buffer Solutions? A buffer solution is one that is able to resist pH changes upon addition of a small amount of acid or base.
Types of buffer solution: Together, they form a conjugate acidbase pair. Acidic buffer E. Adding sodium ethanoate completely soluble in water to this solution adds a lot of extra ethanoate ions. Concept Check: What happens to the pH of the solution when sodium ethanoate is added to ethanoic acid?
The solution now contains the following: Action of Acidic Buffer i. Buffer action equations are written with single nonreversible arrows. Alkaline buffer E. Adding ammonium chloride completely soluble to this solution adds a lot of extra ammonium ions. What happens to the pH of the solution when ammonium chloride is added to ammonia? Action of Alkaline Buffer i. Acidic buffer In a buffer consisting of HA and its salt A, the following equilibrium exists: Hence pH can be found easily.
Alternatively, to find pH of a buffer, take lg on both sides of boxed equation,. It is important to use the correct concentrations in the numerator and denominator of the last term in the equation. Alkaline buffer In a similar manner as above, the pOH and pH of an alkaline buffer can be found using: The more concentrated the components of a buffer, the greater the buffer capacity.
The process is known as salt hydrolysis. Aqueous solution of salt of strong acid and strong base is neutral Aqueous solution of salt of a weak acid and a strong base is alkaline due to anionic hydrolysis, and aqueous solution of salt of strong acid and a weak base is acidic due to cationic hydrolysis with dilution degree of hydrolysis increases.
Hydrolysis is a reverse process of neutralisation. Common Ion Effect It is. According to Le-Chatelier principle, because of the presence of common ion. Colnmon ion effect is used in Purification of common salt Salting out of soap Qualitative analysis, II group radicals are precipitated out in the presence of HCI which suppress the S2- ion concentration, which is just sufficient to precipitate only II group radicals.
Isohydric Solutions If the concentration of the common ions in the solution of two alectrolytes, e. Such solution are called isohydric solutions. Solubility Product It is defined as the product of the concentrations of the ions of the salt in its saturated solution at a given temperature raised to the power of the ions produced by the dissociation of one mole of the salt. It is denoted by Ksp. Consider the dissociation of an electrolyte AxBy Application of Solubility 1.
The concept Product of Ksp helps in predicting the formation of precipitate.
In predicting the solubility of a sparingly soluble salt knowing the values of Ksp, x and y, the solubility of the salt can be Computed.