Aqueous+Equilibria

Aqueous equilibria is the combination of acids and bases to produce a water and salt. The point of the titration is to find the equivalence point -- the end point is just a very close approximation to it. This is because the pH of the solution changes very rapidly close to the equivalence point. Therefore, the indicator will change color very close to the equivalence point because of the steepness of the pH change.
 * Aqueous Equilibria **
 * (Chapter 17, pages: 726-765) **
 * A. Be able to define the following concepts: **
 * 1. Aqueous Equilibria: ﻿﻿ ** Aqueous equilibria plays a crucial role in many biological and environmental processes. The pH of human blood, for example, is carefully controlled at a value of 7 by equilibria involving, primarily, the conjugate acid-base pair H2CO3 and HCO3-. The pH of many lakes and streams must remain near 5.5 for plant and aquatic life to flourish.
 * 2. Titration: The technique used to measure the concentration of a compound in a solution. **

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Titration is a common laboratory method of quantitative chemical analysis that is used to determine the unknown concentration of a known reactant. Because volume measurements play a key role in titration, it is also known as //volumetric analysis//. For acid base titrations a pH indicator or a pH meter is used to indicate when the titration is completed. An important tool used in titration is a buret.

"An instrument that determines the concentration of a substance in a solution by slowly adding a standard reagent of known concentration until a reaction is completed as shown by a color change or electrical measurement." pH to buffer cannot be more than +/- 1 pH unit away from the pKa indicates the pH of of an (aq) solution It is simple and speedy device to measure the acidity and alkalinity of a fluid. A pH meter acts as a volt meter that measures the electrical potential difference between a pH electrode and a reference electrode and displays the result in terms of the pH value of the solution in which they are immersed. It needs to be calibrated and/or rinsed between each test. When using a pH meter note that the value on the screen will sometimes continue to decrease and you have to know when to take your reading
 * 3. Analyte: is a substance or chemical constitute that is determined in an analytical procedure such as titration. What is being titrated, this can be either an acid or a base and it's titrant will be it's opposite. **
 * 4. Titrant: volume or concentration of a solution. It is the solution that is added from the burette. In experiments you will know the concentration of the titrant and how much added to find the pH/concentration of the analyte. Usually a base(NaOH) in this class. **
 * 5. Buffer: ﻿A solution that contains significant concentrations of both members of a conjugate pair. Buffers resist changes in pH upon addition of small amounts of either an acid or a base. **
 * 6. Neutralization: to be neutral. to counteract or counterbalance effects. A neutralization reaction takes place between an acid and a base, in which the products are water and a salt made up of the cation from the base and the anion from the acid. **
 * 7. pH meters **

"A pH meter is an electronic instrument measuring the [|pH] ( [|acidity] or [|alkalinity] ) of a liquid (though special probes are sometimes used to measure the pH of semi-solid substances). A typical pH meter consists of a special measuring probe (a [|glass electrode] ) connected to an electronic meter that measures and displays the pH reading." Usually a weak organic acid or base for which the ionized and un-ionized forms are different colors. They exude different colors at different pH levels and different colors for different chemicals. "A pH indicator is a [|halochromic] chemical [|compound] that is added in small amounts to a [|solution] so that the [|pH] ( [|acidity] or [|basicity] ) of the solution can be determined visually"
 * 8. pH indicators **

Here is an example of a pH chart and colors that various chemicals turn at specific pH levels. Ksp=[products] Ksp is just a specially subscripted Keq. It is the equilibrium constant that indicates to what extent a slightly soluble ionic compound dissociates in water.Ksp is measured as the product of the concentrations of aqueous products.
 * 9. Solubility product constant (Ksp) **
 * 10. Molar solubility: ﻿The numer of moles of solute in one liter of saturated solution. " **Molar solubility is the number of moles of a substance (the solute) that can be dissolved per liter of solution before the solution becomes saturated. It can be calculated from a substance's Solubility Product constant (Ksp) and Stoichiometry. "
 * 11. Chelating agents - **a substance whose molecules can form several bonds to a single metal ion. In other words, a chelating agent is a multidentate ligand.
 * 12. Monodentate ligands :(meaning "one-toothed" ligand) a ligand that bonds to a metal atom through one atom of the ligand. **
 * 13. Polydentate ligands: ("two-toothed" ligand) a ligand that bonds to a metal atom through two atoms of the liqand. **

You can do this by finding the pKas of the given Kas for indicators by taking the -log(Ka) and find the BR point on the graph which will give the pKa of the solution and find the best fit a buffer must be made up of a weak acid and its conjugate weak base or a weak base and its conjugate weak acid. If it is a strong base and its conjugate acid or a strong acid and its conjugate base, it can not be a buffer. An example of a buffer would be HCO3(-) and CO3 (2-) or CH3COOH and CH2COO- or Citric Acid and Sodium Citrate. The best buffers are going to be the acid-base conjugate pair with the pKa value that is closest to whatever the value of the pH you want. The reasoning can be explained by the Henderson-Hasselback equation. According to the equation, the pH of a buffer is going to be optimal at its pKa value. For example, if you wanted a solution that would keep the pH at around 4.75, a buffer of equal concentrations of acetic acid and acetate would be an ideal choice because the Ka of acetic acid is 1.8*10^(-5) and so the pKa is 4.74. This can be used to find how much of the acid disociated. This also indicates the pH of the Solution. You can use the equivilance point to find the initial amount of moles of acid or the initial concentration. The equivalence point is the point where the number of moles of base equal the number of moles of acid. The end point is the point where the indicator being used changes color (also 'indication point)'.
 * B. **** Be able to identify the most appropriate indicator depending on the titration reaction. **
 * C. Explain the relationship between buffer capacity and concentration **
 * ** The greater the concentration of the conjugate acid-base pair, then the greater the capacity of the buffer to resist changes in its pH. This is because the buffer will be able to neutralize larger amounts of strong acid/strong base if the concentration of the conjugate acid-base pair is also large. For example, say you made a buffer with some weak acid and its conjugate base and the the relative concentration were .1 molar and .1 molar respectively. If you were going to add some strong acid or strong base which at equilibrium will cause the ratio of the concentration of the conjugate acid-base pair to not differ by more then 10, then you could only add (.9/11)molar of strong acid or base. Now say you made a new buffer of the same weak acid and conjugate base, but now with different amounts of weak acid and conjugate base, say relative concentrations of 10 molar and 10 molar respectively. Here, the concentration of strong acid or strong base you add can't exceed (90/11)molar before the ratio of the equilibrium concentrations of the conjugate acid-base pair exceeded a value of 10. As you can see, the buffer's capacity to resist pH (stay within that ratio of 10) is much more flexible when the relative concentration of acid/conjugate base were 10 molar rather then .1 molar.
 * D. Be able to identify and predict possible buffer combinations **
 * E. Be able to identify the best buffers depending on the pH to buffer **
 * F. Be able to identify the following areas and values using a titration graph **
 * 1. Initial pH- ** T he very left of the graph, when the x-axis is at zero.
 * 2. Buffer region- ** the buffer region is the area on a titration graph right before the pH changed from acid to base, or vice versa. This region is usually indicated by a near 0 slope.
 * 3. pKa- ** at the halfway point of sudden increase or decrease in pH, after the buffer region. The pKa = pH.
 * 4. Equivalence point- ** the point where the number of moles of the base is equal to the number of moles of the acid.

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Here is an image of a graph
 * 5. Initial concentration of analyte **

[|Titration Graph]

The Henderson-Hasselback equation allows us to calculate the pH or pOH of a conjugate acid-base pair from the pKa or pKB and the relative equilibrium concentrations of the conjugate acid-base pair. So, for an acid, the pH is equal to the pKa plus the log of the equilibrium concentration of conjugate base divided by the equilibrium concentration of the conjugate acid. For a base, the pOH is equal to the pKb plus the log of the equilibrium concentration of the conjugate acid divided by the equilibrium concentration of the base. When the relative equilibrium concentrations of the conjugate acid/base pair are equal, then the pH is equal to the pKa (or the pOH is equal to the pKb). This is the optimal pH of the buffer, which means that the pH at this pKa/pKb value will have the largest amount of spread within the pH range whereby a buffer will be effective in its resilience to changes in pH of the solution when a strong acid/strong base are added to solution. Moreover, The //Henderson//-//Hasselbalch equation// can be used to prepare buffer solutions and to estimate charges on ionizable species in solution The pH = pKa at the "halfway point" - halfway to the equivalence point, because at that point the concentration of the initial weak acid and the resulting conjugate base are equal. Ph value less than the Pka required more acid than base A strong acid and a strong base will equate to a product that will be neutral at the equivalence point (H2O + salt)(pH=7).
 * G. Be able to predict the pH of a solution using: **
 * 1. Henderson-Hasselback equation **
 * 2. Ka, pKa and pH values - ** pH = -log[H+] => [H+] = sqrt(initial concentration x Ka) => Ka = 10^(-pKa)
 * 3. After combining a strong acid and a strong base **
 * 4. After combining a weak acid and a strong base **


 * A weak acid and a strong base will equate to a solution that will be a weak base (pH>7). **

At the equivalence point, the only thing left in solution will be the conjugate base in the products, which is basic. This results in a pH greater than 7.
 * 5. After combining a strong acid and a weak base **

A strong acid and a weak base will equate to a solution that will be a weak acid (pH<7). ﻿pH will be neutral. Approximately 7. Ex: titrating HCl and NaOH pH will be basic. pH>7 Ex: titrating NaOH and CH3COOH weak acid HPr and strong base NaOH weak acid HCN and strong base NaOH weak acid HF and strong base KOH pH will be acidic pH<7 Ex: titrating HCl and NH3 Ksp = [Am+]^n[Bn-]^m //K////sp// expression for a salt is the product of the concentrations of the ions, with each concentration raised to a power equal to the coefficient of that ion in the balanced equation for the solubility equilibrium. AgBr(s) --> Ag+ (aq) + Br- (aq) Ksp=7.7*10^(-13) Ksp=[Ag+][Br-] 7.7*10^(-13)=[x^2] x=√(7.7*10^(-13)) x=8.8*10^(-7)
 * H. Be able to determine the volume of titrant necessary to neutralize the following solutions: **
 * 1. Acidic solution **
 * To fully neutralize the acid, the acid and the base should have the same amount of moles. Hence knowing the molarity and the voluma of the acidic solution, we can find the amount of moles of the acid. Then use that number to find out the volume of the necessary basic solution. **
 * 2. Basic solution **
 * To fully neutralize the base, the acid and the base should have the same amount of moles. Hence knowing the molarity and the voluma of the basic solution, we can find the amount of moles of the base. Then use that number to find out the volume of the necessary acidic solution. **
 * I. Be able to calculate the pH and build a titration graph (pH versus volume of titrant added) for the following problems: **
 * 1. Titration of a strong acid with a strong base **
 * 2. Titration of a weak acid with a strong base **
 * 3. Titration of a strong base with a strong acid **
 * pH will be neutral Approximately 7 **
 * 4. Titration of a weak base with a strong acid **
 * J. Be able to calculate the product solubility constant (Ksp) for a solid, given the concentration in equilibrium of the dissolved ions. **
 * K. Be able to calculate the concentration in equilibrium of soluble ions, given its product solubility constant of a solid. **
 * L. Explain how to prepare a buffer 1) from a conjugate pair, and 2) from a titration of WA with SB or WB with SA. **
 * M. Be able to determine the concentration in a solution of a compound based on its Ksp value. **