Final Solution

Examples of Final Solution in the following topics:

• Preparing a Buffer Solution with a Specific pH

  • It is used to prevent any change in the pH of a solution, regardless of solute.
  • In the first method, prepare a solution with an acid and its conjugate base by dissolving the acid form of the buffer in about 60% of the volume of water required to obtain the final solution volume.
  • Once the pH is correct, dilute the solution to the final desired volume.
  • Both solutions must contain the same buffer concentration as the concentration of the buffer in the final solution.
  • To get the final buffer, add one solution to the other while monitoring the pH.

• Dilutions of Solutions

  • This process keeps the amount of solute constant, but increases the total amount of solution, thereby decreasing its final concentration.
  • Dilution can also be achieved by mixing a solution of higher concentration with an identical solution of lesser concentration.
  • M1 denotes the concentration of the original solution, and V1 denotes the volume of the original solution; M2 represents the concentration of the diluted solution, and V2 represents the final volume of the diluted solution.
  • 175 mL of a 1.6 M aqueous solution of LiCl is diluted with water to a final volume of 1.0 L.
  • What is the final concentration of the diluted solution?

• Prediction

  • We found the equation of the best fit line for the final exam grade as a function of the grade on the third exam.
  • Suppose you want to estimate, or predict, the final exam score of statistics students who received 73 on the third exam.
  • We predict that statistic students who earn a grade of 73 on the third exam will earn a grade of 179.08 on the final exam, on average.
  • What would you predict the final exam score to be for a student who scored a 66 on the third exam?
  • What would you predict the final exam score to be for a student who scored a 90 on the third exam?

• Calculating Changes in a Buffer Solution

  • These solutions are known as buffers.
  • It is possible to calculate how the pH of the solution will change in response to the addition of an acid or a base to a buffer solution.
  • What is the pH of the solution?
  • Solving for the pH of a 0.0020 M solution of NaOH:
  • Calculate the final pH of a solution generated by the addition of a strong acid or base to a buffer.

• Molarity

  • Molarity is defined as the moles of a solute per volume of total solution.
  • In chemistry, concentration of a solution is often measured in molarity (M), which is the number of moles of solute per liter of solution.
  • A solution that contains 1 mole of solute per 1 liter of solution (1 mol/L) is called "one Molar" or 1 M.
  • To calculate the molarity of a solution, the number of moles of solute must be divided by the total liters of solution produced.
  • This relationship is represented by the equation c1V1 = c2V2 , where c1 and c2 are the initial and final concentrations, and V1 and V2 are the initial and final volumes of the solution.

• Solutions to Exercises in Chapter 13

  • The x values in the data are between 65 and 75. 90 is outside of the domain of the observed x values in the data (independent variable), so you cannot reliably predict the final exam score for this student.
  • The final exam score is predicted to be 261.19.
  • The largest the final exam score can be is 200.

• Precipitation Reactions

  • Precipitation reactions transform ions into an insoluble salt in aqueous solution.
  • A precipitation reaction can occur when two solutions containing different salts are mixed, and a cation/anion pair in the resulting combined solution forms an insoluble salt; this salt then precipitates out of solution.
  • This reaction can be also be written in terms of the individual dissociated ions in the combined solution.
  • A final way to represent a precipitation reaction is known as the net ionic equation.
  • For instance, if silver nitrate is added to a solution of an unknown salt and a precipitate is observed, the unknown solution might contain chloride (Cl-).

• Acid-Base Titrations

  • Rinse the burette with the standard solution, the pipette with the unknown solution, and the conical flask with distilled water.
  • At this stage, we want a rough estimate of the amount of known solution necessary to neutralize the unknown solution.
  • Record the initial and final readings on the burette, prior to starting the titration and at the end point, respectively.
  • (Subtracting the initial volume from the final volume will yield the amount of titrant used to reach the endpoint.)
  • The pH of a weak acid solution being titrated with a strong base solution can be found at each indicated point.

• The Henderson-Hasselbalch Equation

  • The Henderson–Hasselbalch equation connects the measurable value of the pH of a solution with the theoretical value pKa.
  • The equation is also useful for estimating the pH of a buffer solution and finding the equilibrium pH in an acid-base reaction.
  • Distributing the negative sign gives the final version of the Henderson-Hasselbalch equation:
  • An example of how to use the Henderson-Hasselbalch equation to solve for the pH of a buffer solution is as follows:
  • What is the pH of a buffer solution consisting of 0.0350 M NH3 and 0.0500 M NH4+ (Ka for NH4+ is 5.6 x 10-10)?

• Strong Acid-Weak Base Titrations

  • A strong acid will react with a weak base to form an acidic (pH < 7) solution.
  • A small amount of the acid solution of known concentration is placed in the burette (this solution is called the titrant).
  • As the equivalence point is approached, the pH will change more gradually, until finally one drop will cause a rapid pH transition through the equivalence point.
  • In the example of the titration of HCl into ammonia solution, the conjugate acid formed (NH4+) reacts as follows:
  • A depiction of the pH change during a titration of HCl solution into an ammonia solution.