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18.1 The Danger of Antifreeze -- Part 1
The Henderson-Hasselbalch equation can be used to calculate the pH of a buffer solution.
Section 18.1 states that a buffer can be destroyed by adding too much acid or base.
When the concentrations of acid and conjugate base are equal, a buffer is most effective.
Both solutions have the same amount of acid and conjugate base, but solution II has more acid than solution I.
The Henderson-Hasselbalch equation can be used to calculate the initial pH values.
The more effective buffer is the buffer with equal amounts of acid and conjugate base.
As the relative amounts of acid and conjugate base increase, a buffer becomes less effective.
An effective buffer must have a base>acid ratio in the range of 0.10 to 10.
The acid and base in solution I are ten times more concentrated than the acid and base in solution II.
The buffer with greater amounts of acid and conjugate base is more effective than the one with less acid and conjugate base.
It can add more acid to the base.
The relative amounts of acid and conjugate base can be adjusted.
The buffer would be the most effective because it would be exactly the same as the pH.
The ratio of the conjugate base to the acid is the best choice.
In Section 18.1, we discussed the ionized equilibrium of blood.
The most important part of the lac buffering systems is tic acid, which is composed of carbonic acid and the bicarbonate ion.
The con carbonic acid at the body temperature is 6.1.
Carbon dioxide is removed from our blood when we breathe.
The blood's pH is 7.4.
The higher the ion concentration in the blood, the greater the buffer capacity of the blood is for acid.
Lactic acid is produced when we exercise.
Normal blood has a pH of 7.4.
The higher the buffer capacity, the more concentrated the weak acid and conjugate base is.
A 1.0-L buffer solution has 0.10 M in HF and 0.050 M in NaF.
Section 5.7 contains acid-base titrations.
The acid and base combine to destroy each other.
The number of moles of acid and base are complete and so the titration is named.
When the equivalence point is reached, neither reactant is in excess and the same number of moles of the reactants are related by the reaction stoichiometry.
In this section, we look at acid-base titrations more closely, focusing on the pH changes that occur during the titration.
The pH is low before any base is added to the solution.
As OH- is added in a titration, it creates a difference between the H+ and the water.
The titration is complete at the equivalence point.
The ionized equilibrium is acidic because of the NaOH.
The equivalence point is in the middle of the curve.
The solution is basic because the HCl has been completely neutralized and excess base is being added to it.
The strength of the acid is one of the factors that affect the shape of the pH curve.
Let's take a look at several combinations.
Consider the amount of 25.0 mL of 0.
100 M HCl and 0.
100 M NaOH.
When the number of moles of base added equals the number of moles of acid in solution, the equivalency point is reached.
The amount of NaOH must be added.
The equivalency point is reached when NaOH is added.
The volume of NaOH solution required to reach the equivalence point is the same as the volume of the HCl solution that is being titrated.
The initial pH of the solution is 0.100 M HCl.
The number of moles of H3O+ is just calculated.
The equivalence point of a strong acid-strong base titration will always be 7.00.
Water is the only source of hydronium ion.
The water's pH is 7.00 and it's 3O+ at 25 degrees.
The excess reagent is created when NaOH is added beyond the equivalence point.
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