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Salt Solutions

Pearson
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In this demonstration, we're going to look at the pH effects of salt solutions by analyzing the effect that a cation and an anion have on pH. In front of me here, I have equal molar concentrations of chloride, fluoride, and acetate anions. As we can see, the pH of our chloride solution is roughly 7. When I take my pH probe and go from the chloride solution to the fluoride solution, what will happen to the pH? As we can see from the pH meter, the pH went from 7 to around 7.5 here for the fluoride solution, indicating that the fluoride ion is a little bit more basic than we had from the chloride ion. Now what will happen when we take this same pH probe and stick it into a solution containing the acetate ion? As we can see from the pH meter, the acetate ion in solution is even more basic than the fluoride ion. So how can we explain this trend that we just observed experimentally? If we think about comparing strong and weak acids, a strong acid fully dissociates. So, for example, HCl is a strong acid. It completely dissociates into H+ plus Cl- anions. The chloride ion then does not interact with the water. We have no reaction, and the pH remains 7. In the case of the fluoride ion reacting with the water, it's going to be in equilibrium with HF plus OH-, indicating that this particular solution is basic. Another anion from a weak acid, such as the acetate ion, also interacts with water, forming in equilibrium with HC2H3O2 plus OH-. The presence of these OH- ions indicate that our solution will be basic, which is what we saw here. The question is why is the acetate ion in solution more basic than the HF ion in solution. In this case, we have to compare the Ka values of HF and HC2H3O2. We see that HF is a stronger acid. So its conjugate base, F-, is going to be weaker, okay? For acetic acid, our acetate ion, or the conjugate base of acetic acid, is going to be stronger. This is why we see that change in pH. If we want to look at a reference in order to see and be able to predict these, we can look at a Ka table. The Ka table tells us the relative strength of our acids. We don't see a Ka for HCl because it's a strong acid. For HF and acetic acid, we see 6.8 times 10 to the minus 4 and 1.9 times 10 to the minus 5. If we have a larger Ka, we have a stronger acid. Okay, that means HF will be a stronger acid than acetic acid. That means when we look at the conjugate bases, we have to come up with a relationship. The stronger the acid, the weaker the conjugate base. So when we're comparing the conjugate bases of these two weak acids, we would protect that F- would be a weaker base than C2H3O2-. We saw this when we observed it in our particular experiment. Now that we looked at the trends at anions, let's see if we can apply the same principles to cations. In the beaker on my left, we have a solution of Na+ ions in solution. That gives the pH of around seven. What will happen when I take this pH probe and insert it into a solution containing NH4+ cations? From the pH meter, we can see that the pH went down, or the solution became more acidic. Let's see if we can apply the same principles involving weak acids to week bases. If we have a Na+ ion in solution, it's going to be the conjugate of a strong base, or sodium hydroxide. When we add water to that, we're going to see no reaction, and the pH is going to stay at 7. When we have NH4+, which will be the conjugate base of ammonia, which ammonia, being a weak base, means that this cation right here will be acidic. We see that. But we can then put forth as an explanation is the fact that the NH4+ interacts with water to give us NH3 and H3O+ ions in solution. These ions produced in the solution make it acidic, and this is why we are seeing the pH drop from 7 down to 5.
In this demonstration, we're going to look at the pH effects of salt solutions by analyzing the effect that a cation and an anion have on pH. In front of me here, I have equal molar concentrations of chloride, fluoride, and acetate anions. As we can see, the pH of our chloride solution is roughly 7. When I take my pH probe and go from the chloride solution to the fluoride solution, what will happen to the pH? As we can see from the pH meter, the pH went from 7 to around 7.5 here for the fluoride solution, indicating that the fluoride ion is a little bit more basic than we had from the chloride ion. Now what will happen when we take this same pH probe and stick it into a solution containing the acetate ion? As we can see from the pH meter, the acetate ion in solution is even more basic than the fluoride ion. So how can we explain this trend that we just observed experimentally? If we think about comparing strong and weak acids, a strong acid fully dissociates. So, for example, HCl is a strong acid. It completely dissociates into H+ plus Cl- anions. The chloride ion then does not interact with the water. We have no reaction, and the pH remains 7. In the case of the fluoride ion reacting with the water, it's going to be in equilibrium with HF plus OH-, indicating that this particular solution is basic. Another anion from a weak acid, such as the acetate ion, also interacts with water, forming in equilibrium with HC2H3O2 plus OH-. The presence of these OH- ions indicate that our solution will be basic, which is what we saw here. The question is why is the acetate ion in solution more basic than the HF ion in solution. In this case, we have to compare the Ka values of HF and HC2H3O2. We see that HF is a stronger acid. So its conjugate base, F-, is going to be weaker, okay? For acetic acid, our acetate ion, or the conjugate base of acetic acid, is going to be stronger. This is why we see that change in pH. If we want to look at a reference in order to see and be able to predict these, we can look at a Ka table. The Ka table tells us the relative strength of our acids. We don't see a Ka for HCl because it's a strong acid. For HF and acetic acid, we see 6.8 times 10 to the minus 4 and 1.9 times 10 to the minus 5. If we have a larger Ka, we have a stronger acid. Okay, that means HF will be a stronger acid than acetic acid. That means when we look at the conjugate bases, we have to come up with a relationship. The stronger the acid, the weaker the conjugate base. So when we're comparing the conjugate bases of these two weak acids, we would protect that F- would be a weaker base than C2H3O2-. We saw this when we observed it in our particular experiment. Now that we looked at the trends at anions, let's see if we can apply the same principles to cations. In the beaker on my left, we have a solution of Na+ ions in solution. That gives the pH of around seven. What will happen when I take this pH probe and insert it into a solution containing NH4+ cations? From the pH meter, we can see that the pH went down, or the solution became more acidic. Let's see if we can apply the same principles involving weak acids to week bases. If we have a Na+ ion in solution, it's going to be the conjugate of a strong base, or sodium hydroxide. When we add water to that, we're going to see no reaction, and the pH is going to stay at 7. When we have NH4+, which will be the conjugate base of ammonia, which ammonia, being a weak base, means that this cation right here will be acidic. We see that. But we can then put forth as an explanation is the fact that the NH4+ interacts with water to give us NH3 and H3O+ ions in solution. These ions produced in the solution make it acidic, and this is why we are seeing the pH drop from 7 down to 5.