Ionic Equilibrium

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It is the presentation describes the ionic equilibrium in chemistry created by : CHIRAG PADARIYA

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Presentation Transcript

Acid and Bases : 

Acid and Bases

Acid and Bases : 

Acid and Bases

The pH scale is a way of expressing the strength of acids and bases. Instead of using very small numbers, we just use the NEGATIVE power of 10 on the Molarity of the H+ (or OH-) ion.Under 7 = acid 7 = neutralOver 7 = base : 

The pH scale is a way of expressing the strength of acids and bases. Instead of using very small numbers, we just use the NEGATIVE power of 10 on the Molarity of the H+ (or OH-) ion.Under 7 = acid 7 = neutralOver 7 = base

pH of Common Substances : 

pH of Common Substances

Slide 6: 

H2O can function as both an ACID and a BASE. In pure water there can be AUTOIONIZATION Equilibrium constant for water = Kw Kw = [H3O+] [OH-] = 1.00 x 10-14 at 25 oC

pOH : 

pOH Since acids and bases are opposites, pH and pOH are opposites! pOH does not really exist, but it is useful for changing bases to pH. pOH looks at the perspective of a base pOH = - log [OH-] Since pH and pOH are on opposite ends, pH + pOH = 14

Slide 9: 

pH [H+] [OH-] pOH

Slide 10: 

[OH-] [H+] pOH pH 10-pOH 10-pH -Log[H+] -Log[OH-] 14 - pOH 14 - pH 1.0 x 10-14 [OH-] 1.0 x 10-14 [H+]

Slide 11: 

HNO3, HCl, H2SO4 and HClO4 are among the only known strong acids. Strong and Weak Acids/Bases The strength of an acid (or base) is determined by the amount of IONIZATION.

Slide 12: 

Weak acids are much less than 100% ionized in water. One of the best known is acetic acid = CH3CO2H Strong and Weak Acids/Bases

Slide 13: 

Strong Base: 100% dissociated in water. NaOH (aq) ---> Na+ (aq) + OH- (aq) Strong and Weak Acids/Bases Other common strong bases include KOH and Ca(OH)2. CaO (lime) + H2O --> Ca(OH)2 (slaked lime)

Slide 14: 

Weak base: less than 100% ionized in water One of the best known weak bases is ammonia NH3 (aq) + H2O (l)  NH4+ (aq) + OH- (aq) Strong and Weak Acids/Bases

Weak Bases : 

Weak Bases

Equilibria Involving Weak Acids and Bases : 

Equilibria Involving Weak Acids and Bases Consider acetic acid, HC2H3O2 (HOAc) HC2H3O2 + H2O  H3O+ + C2H3O2 - Acid Conj. base (K is designated Ka for ACID) K gives the ratio of ions (split up) to molecules (don’t split up)

Ionization Constants for Acids/Bases : 

Ionization Constants for Acids/Bases Acids Conjugate Bases Increase strength Increase strength

Equilibrium Constants for Weak Acids : 

Equilibrium Constants for Weak Acids Weak acid has Ka < 1 Leads to small [H3O+] and a pH of 2 - 7

Equilibrium Constants for Weak Bases : 

Equilibrium Constants for Weak Bases Weak base has Kb < 1 Leads to small [OH-] and a pH of 12 - 7

pH testing : 

pH testing There are several ways to test pH Blue litmus paper (red = acid) Red litmus paper (blue = basic) pH paper (multi-colored) pH meter (7 is neutral, <7 acid, >7 base) Universal indicator (multi-colored) Indicators like phenolphthalein Natural indicators like red cabbage, radishes, morning glory flowers, turmeric.

Relation of Ka, Kb, [H3O+] and pH : 

Relation of Ka, Kb, [H3O+] and pH

Paper testing : 

Paper testing Paper tests like litmus paper and pH paper Put a stirring rod into the solution and stir. Take the stirring rod out, and place a drop of the solution from the end of the stirring rod onto a piece of the paper Read and record the color change. Note what the color indicates. You should only use a small portion of the paper. You can use one piece of paper for several tests.

Slide 23: 

pH paper

pH indicators : 

pH indicators Indicators are dyes that can be added that will change color in the presence of an acid or base. Some indicators only work in a specific range of pH Once the drops are added, the sample is ruined Some dyes are natural, like radish skin or red cabbage

Slide 25: 

An acid-base indicator is itself a weak acid (or its conjugate base). An acid-base indicator is a weak acid having a different colour in aqueous solution from its conjugate base. Consider methyl orange, if the acid form of the indicator is represented by HIn and its conjugate base form by In-, the following equilibrium exists in aqueous solution:

Slide 26: 

How simple indicators work Indicators as weak acids Litmus Litmus is a weak acid. It has a seriously complicated molecule which we will simplify to HLit. The "H" is the proton which can be given away to something else. The "Lit" is the rest of the weak acid molecule. There will be an equilibrium established when this acid dissolves in water. Taking the simplified version of this equilibrium: The unionised litmus is red, whereas the ion is blue.

Slide 27: 

Adding hydroxide ions: Adding hydroxide ions:

Slide 28: 

Adding hydrogen ions:

Slide 29: 

Methyl orange Methyl orange is one of the indicators commonly used in titrations. In an alkaline solution, methyl orange is yellow and the structure is:

Slide 30: 

when you add an acid: the hydrogen ion attaches to one of the nitrogens in the nitrogen-nitrogen double bond to give a structure which might be drawn like as follows:

Slide 31: 

nitrogen double bond give a structure which might be drawn like this:

Slide 32: 

The colours are different.

Slide 33: 

In the methyl orange case, the half-way stage where the mixture of red and yellow produces an orange colour happens at pH 3.7 – nowhere near traditional neutral i.e.7 pH.

Slide 34: 

Phenolphthalein Phenolphthalein is another commonly used indicator for titrations, and is another weak acid.

Slide 36: 

Phenolphthalein In this case, the weak acid is colourless and its ion is bright pink. Adding extra hydrogen ions shifts the position of equilibrium to the left, and turns the indicator colourless. Adding hydroxide ions removes the hydrogen ions from the equilibrium which tips to the right to replace them - turning the indicator pink.

Slide 37: 

The pH range of indicators The importance of pKind HInd - where "Ind" is all the rest of the indicator apart from the hydrogen ion which is given away: : :

Slide 38: 

what happens half-way through the colour change. At this point the concentrations of the acid and its ion are equal. In that case, they will cancel out of the Kind expression .

Slide 39: 

convert to pH and pKind, you get:

Slide 40: 

That means that the end point for the indicator depends on what its pKind value is.

Slide 41: 

The pH range of indicators Indicators don't change colour sharply at one particular pH (given by their pKind). Instead, they change over a narrow range of pH. the equilibrium is firmly to one side, but now you add OH- to shift it. As the equilibrium shifts, you will start to get more and more of the second colour formed, and at some point the eye will start to detect it.

Slide 42: 

For example, suppose you had methyl orange in an alkaline solution so that the dominant colour was yellow. Now start to add acid so that the equilibrium begins to shift.

Slide 43: 

At some point there will be enough of the red form of the methyl orange present that the solution will begin to take on an orange tint. As you go on adding more acid, the red will eventually become so dominant that you can no longer see any yellow.

Slide 44: 

There is a gradual smooth change from one colour to the other, taking place over a range of pH. As a rough "rule of thumb", the visible change takes place about 1 pH unit either side of the pKind value.

Slide 45: 


Slide 46: 

The litmus colour change happens over an unusually wide range, but it is useful for detecting acids and alkalis in the lab because it changes colour around pH 7.

Slide 47: 

This is more easily seen diagramatically.

Slide 48: 

methyl orange would be yellow in any solution with a pH greater than 4.4. It couldn't distinguish between a weak acid with a pH of 5 or a strong alkali with a pH of 14.

Slide 49: 

Choosing indicators for titrations Remember that the equivalence point of a titration is where you have mixed the two substances in exactly equation proportions. You obviously need to choose an indicator which changes colour as close as possible to that equivalence point. That varies from titration to titration.

Slide 50: 

Strong acid v strong base The next diagram shows the pH curve for adding a strong acid to a strong base. Superimposed on it are the pH ranges for methyl orange and phenolphthalein. . .

Slide 51: 

You can see that neither indicator changes colour at the equivalence point. However, the graph is so steep at that point that there will be virtually no difference in the volume of acid added whichever indicator you choose. However, it would make sense to titrate to the best possible colour with each indicator. If you use phenolphthalein, you would titrate until it just becomes colourless (at pH 8.3) because that is as close as you can get to the equivalence point. On the other hand, using methyl orange, you would titrate until there is the very first trace of orange in the solution. If the solution becomes red, you are getting further from the equivalence point.

Slide 52: 

Strong acid v weak base

Slide 53: 

This time it is obvious that phenolphthalein would be completely useless. However, methyl orange starts to change from yellow towards orange very close to the equivalence point. You have to choose an indicator which changes colour on the steep bit of the curve.

Slide 54: 

Weak acid v strong base

Slide 55: 

This time, the methyl orange is hopeless! However, the phenolphthalein changes colour exactly where you want it to.

Slide 56: 

Weak acid v weak base The curve is for a case where the acid and base are both equally weak - for example, ethanoic acid and ammonia solution. In other cases, the equivalence point will be at some other pH.

Slide 58: 

You can see that neither indicator is any use. Phenolphthalein will have finished changing well before the equivalence point, and methyl orange falls off the graph altogether.