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.
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 more similar the concentrations of the buffer components, the greater the buffer capacity. This buffer is said to have the maximum buffer capacity and it can most effectively resist a change in pH in either direction i.
Effective Buffer Range The further the buffercomponent concentration ratio is from 1, the less effective the buffering action i. For your choice, calculate the ratio of the salt conjugate base to the acid concentrations required to attain the desired pH. The best choice is methanoic acid because its pKa lies closest and within 1 unit of the desired pH. Buffer System in Human Blood independent learning subtopic The pH of our blood must be kept constant at about 7.
If blood pH rises above about 7. This can arise from hyperventilation or oxygen deficiency at high altitude. It can lead to over excitability of the central nervous system, muscle spasms and death. One way to treat alkalosis is to breath into a paper bag. The CO2 exhaled is recycled into the body. Can you explain how that helps?
Importance of Buffer Solutions in Everyday Life independent learning subtopic Injections and drips into a patients body must be carefully buffered so that the pH of body fluids does not change much. In bacteriological research, pH of culture medium is maintained to control the growth of bacteria.
In agriculture, the pH of soil is maintained to optimize plant growth. The pH of many industrial processes must also be carefully controlled, e. Baking soda, which is sodium hydrogencarbonate, NaHCO3, is sometimes added to swimming pools to control the pH of the water.
On addition of small amount of acid,.
AcidBase Indicators independent learning subtopic An indicator is a weak acid, whose acid form, HIn, is a different colour from its ionised form, In. At different pH values, the proportion of HIn to In is different, giving rise to different colours.
For litmus: The position of equilibrium lies on the left. Predominant form of litmus is HIn so the solution will appear red. The position of equilibrium lies on the right. Predominant form of litmus is In hence the solution will appear blue.
Working pH Range of Indicators independent learning subtopic For most indicators, the predominant form must be at least 10 times more concentrated than the other form for its colour to be distinguished from the other coloured form. Given that the dissociation constant of an indicator, termed KIn, is. 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.
Thermodynamics and Equilibrium. Activities and Activity Coefficients.
Definitions of Acids and Bases: Strength of Acids and Bases. Calculations of pH Values in Aqueous Solutions.
Buffer Solutions. Some General Points Concerning Titrations. Neutralization or Acid-Base Indicators. Acid-Base Titration Curves. Acid—Base Titrations: Further Theoretical Studies. Acid—Base Reactions and Chemical Analysis. Generalities on Oxidation-Reduction. Redox Reactions and Electrochemical Cells. Predicting Redox Reactions. Predicting Redox Reactions by Graphical Means.