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UNC-Chapel Hill BIOC 107 - _4B_Buffers-1

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Important buffer systemsLAB 4B: BUFFERS AND TITRATIONObjectives: By the end of this lab session, you should be able to:1. Identify a buffer.2. Calculate the pH of a buffer solution upon the addition of either a strong acid or base.3. Calculate the amount of strong base that is required to make a weak acid solution.4. Predict the pH at which a buffer will be effective based on the pKa of the buffer.4.Understand balanced neutralization reactions and titration graphs in order to determine the total amount of acid or base in a solution.5.Understand why buffers are important to living organisms.NOTE: Lab #4 is found in two different documents on the “Recitation Labs” Page for “Lab #4”. The first document you should look at is “#4A_Molly_Case_Study”. This is a Case Study provided by the National Center for Case Study Teaching in Science. The first 5 pages [Pre-Class Assignment: The Case of the Mortified Mom: Acids, pH and Buffers] should be read before the Lab on Thursday. The first section (3 pages) is mainly a fun review of topics we covered in last week’s lab, and the second section (2 pages) is the introduction to the Case Study about Molly. The third section [In-Class Handout: The Case of the Mortified Mom: Acids, pH and Buffers] (2 pages) is the section of the Case Study to be covered during class time. Please come to class having read the Pre-Class Assignment section of the Case Study.The second document you should look at is this document.If you feel as though you need more practice with questions about buffers, following is a link to a website where you can check your comprehension of your buffers. Please note that this is not a requirement and you will not be graded on this. This is purely for your benefit.http://tll.mit.edu/help/buffershttps://www.khanacademy.org/science/chemistry/acid-base-equilibriumhttp://employees.oneonta.edu/viningwj/sims/titrations_t.htmlhttp://www.biology.arizona.edu/biochemistry/problem_sets/medph/intro.htmlhttp://www.biology.arizona.edu/biochemistry/problem_sets/medph/02q.htmlhttp://group.chem.iastate.edu/Greenbowe/sections/projectfolder/flashfiles/acidbasepH/pHbuffer20.htmlhttp://employees.oneonta.edu/viningwj/sims/buffer_solutions_s.htmlhttp://employees.oneonta.edu/viningwj/sims/NEUTRALIZATION OF ACIDS AND BASESNeutralization reactions are reactions in which a base is added to a solution of acid, or an acid is added to a solution of a base. Acids and bases react with each other in aqueous solution to produce water and a salt. The neutralization reaction for HCl and NaOH is given below:Biochem 107L4-2 HCl H+ + Cl–NaOH Na+ + OH–SUM: H+ + Cl– + Na+ + OH– Na+ + Cl– + H2OThe overall effect of the two combined reactions is removal of the H+ and OH– ions from the solution, forming neutral water and a salt (sodium chloride). The pH of this solution should be near neutrality (pH = 7). Neutralization reactions require equivalent amounts of acid (H+) and base (OH-). If either acidor base is present in excess, the solution will not be neutralized and will have a pH lower than 7 if excess acid is present, or a pH greater than 7 if excess base is present. Optional Link:http://employees.oneonta.edu/viningwj/sims/titrations_t.htmlHENDERSON-HASSELBALCH EQUATIONThe relationship between changes in pH and changes in the concentrations of the acid and its corresponding salt ([HA] and [A–], respectively) is conveniently expressed by the Henderson-Hasselbach equation. This equation is helpful in understanding buffers and how they work. It is derived by rearranging the equation for the dissociation of an acid as shown below: [H +]= K a[ HA ][ A – ]Then taking the negative log of the entire equation yields the Henderson-Hasselbach equation:pH = p K a + log[A – ][HA]You don’t need to know how to derive the equation, but you should see that it is equivalent to a mathematical derivation of the original dissociation (equilibrium) equation. pH = p K a + log[salt][acid]MEMORIZE THE HENDERSON-HASSELBACH EQUATION! The equation relates the pH of a solution to the relative concentrations of a weak acid and its salt (conjugate base), and it will be used repeatedly as we discuss buffers and how they work. The pH of a solution can be calculated from the H/H equation if the molar ratio of A– to HA and the pKa are known. Consider a solution of 0.1 M acetic acid and 0.2 M acetate ion. The pKa of acetic acid is about 4.8; hence, the pH of the solution is given by pH = 4.76 + log 0.2 = 4.76 + log 2 0.1pH = 4.76 + 0.30 = 5.06Conversely, the pKa of an acid can be calculated if the molar ratio of A– to HA and the pH of the solutionare known. Note that the pKa is the pH at which the concentrations of the salt (A–) and the acid (HA)Biochem 107L4-3are equal. Try to mathematically verify this relationship using the Henderson-Hasselbach equation (Hint: What is the log of 1.0?).Optional Link: http://www.manuelsweb.com/pka.htmBUFFERSBuffers are chemical systems that resist changes in pH when acids or bases are added to the system. Buffers consist of solutions of a weak acid and its conjugate base. Only weak acids or weak bases can act as buffers. Buffers are important in laboratory science, and all body fluids and cells contain buffering systems that keep the hydrogen ion concentration (pH) constant. Alterations in acid-base balance are common in various abnormal or diseased states, and you will encounter these alterations and their consequences repeatedly in clinical settings; thus, it is important that you have a good understanding of buffers and the physiological mechanisms for maintaining normal acid-base balance. The acetic acid/acetate buffer system is a typical example. A solution containing 0.1 moles/liter of both acetic acid and sodium acetate has a pH of 4.76. When moderate amounts of either acid or base are added to this buffer system, the pH changes only very slightly. For example, the addition of 10 ml of 0.1 M sodium hydroxide (NaOH) to 1 liter of this buffer increases the pH by only 0.01 units, from 4.76to 4.77. In contrast, addition of 10 ml of 0.1 M NaOH to 1 liter of pure water increases the pH by 4.0 pHunits, from 7.0 to 11.0, a 10,000-fold decrease in [H+]. The buffer solution is able to resist changes in pH because the acid and its anion salt act as reservoirs of neutralizing power for protons and hydroxide ions, respectively, added to the solution. For example, the


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