1.) You will work with 0.10 M acetic acid and 17 M acetic acid in this experiment. What is the relationship between concentration and ionization? Explain the reason for this relationship
2.) Explain hydrolysis, i.e, what types of molecules undergo hydrolysis (be specific) and show equations for reactions of acid, base, and salt hydrolysis not used as examples in the introduction to this experiment
3.) In Part C: Hydrolysis of Salts, you will calibrate the pH probe prior to testing the pH of the various salts. In a few sentences, summarize and explain the necessity of the calibration process. Why is it necessary to work with three buffers?
4.) You prepare a buffer by adding 10.0 mL of 0.10 M acetic acid and 10.0 mL of 0.10 M sodium acetate. Calculate the pH that you expect for this buffer. Write the chemical equation for the equilibrium. Show your work using an ICE table
5.) To the solution from #4, you add 0.50 mL of 0.10 M HCl. Calculate the new pH of the buffer solution. Write the chemical equations for the reaction and the equilibrium. Show your work using an ICE table.
6.) To the solution from #4, you add 0.50 mL of 0.10 M NaOH. Calculate the new pH of the buffer solution. Write the chemical equation for the equilibrium. Show your work using an ICE table.
7.) Beginning on a new page, create data tables in your lab notebook using a straightedge to. Do not submit the tables with your pre-lab questions, as you will record your data in the tables during lab.
Acids and Bases Aqueous solutions of acids and bases are recognized as “acidic” or “basic” because they contain appreciable concentrations of either hydronium (H3O+ ) or hydroxide (OH– ) ions.
Hydronium ions are produced from the reaction of covalent molecules like HCl with water. HCl (g) + H2O (l) H3O+ (aq) + Cl– (aq) Some bases are ionic compounds that dissolve in water or react with it to produce aqueous hydroxide ions, for example: NaOH (s) Na+ (aq) + OH– (aq) BaO (s) + H2O (l) Ba2+ (aq) + 2OH– (aq) Other bases such as ammonia, NH3, and related amines produce hydroxide ions when dissolved in water by accepting a proton from the solvent H2O.
NH3 (g) + H2O (l) NH4 + (aq) + OH– (aq) Acids and bases are classified as either strong or weak. Strong acids or bases are those that are completely or almost completely ionized in dilute aqueous solution. We can easily calculate the concentration of either H3O+ or OH– ions in these solutions by assuming that the acid or base is completely dissociated into its constituent ions when it is dissolved in water. For example, 0.10 M hydrochloric acid is around 100% dissociated so the H3O+ concentration is 0.10 M.
In aqueous solution of weak acids and bases, the undissociated species predominate and the concentration of H3O+ or OH– ions is small. For example, 0.10 M acetic acid is almost 98.7% undissociated so that the H3O+ concentration is 0.0013 M, or 1.3 × 10-3 M. CH3COOH (aq) + H2O (l) H3O+ (aq) + CH3COO– (aq) 98.7% unreacted Small amount Principles of Chemistry Lab II
The pH Scale The pH scale was developed to express low concentrations of H3O+ without the inconvenience of using decimal numbers and negative powers. We define pH in terms of the molar concentration of H3O+ using Equation 1. pH = – log [H3O+ ] (1) The H3O+ concentration and the pH have an inverse logarithmic relationship; the higher the H3O+ concentration, the lower the pH. A change in pH of 1 unit results in a 10× change in [H3O+ ].
For example, a solution with pH of 5.00 is 100 times more acidic than a solution with a pH of 7.00. This means that a 0.1 M solution of a strong acid (completely ionized) such as HCl has a lower pH value (pH = 1.0) than a 0.01 M solution of the same acid (pH = 2.0). In aqueous solutions, there is also a relationship between the H3O+ concentration and the OH– concentration. To understand this relationship we must first understand something about the nature of water.
Water “self-ionizes” to a very slight extent, as shown in the following equation. H2O (l) + H2O (l) H3O+ (aq) + OH– (aq) In this reaction one molecule of water acts as a Bronsted-Lowry acid (a proton donor) while the other acts as a Bronsted-Lowry base (a proton acceptor). In absolutely pure water the concentration of H3O+ and OH– are exactly the same. For every water molecule that dissociates, one hydronium and one hydroxide ion are formed.
In pure water at 25 °C the concentration of each ion is 1.0 × 10–7 M and the pH of pure water is 7.0. Since the concentrations of H3O+ and OH– are equal, pure water is said to have a neutral pH. The product of these two concentrations is known as Kw, which is calculated in Equation 2. Kw = [H3O+ ][OH– ] = 1.0 × 10-14 (2) Equation 2 holds for most any aqueous solution. For aqueous solutions, if we know the concentration of H3O+ , we can calculate the concentration of OH– , and vice versa. We can now define acidic and basic solutions in a more quantitative manner. A solution is acidic if its H3O+ concentration is greater than 1.0 × 10-7 M or it is has a pH which is less than 7. In acidic solution, the H3O+ concentration is greater than OH– concentration. Conversely, a solution is basic if its H3O+ concentration is less than 1.0 × 10-7 M or it is has a pH which is greater than 7. In basic solutions, the OH– concentration is greater than the H3O+ concentration.
Indicators The pH of a solution can be measured precisely using a pH meter. Frequently, however, we need only an approximate pH value, which can be determined using a suitable acid-base indicator. Indicators are complex organic molecules which change color as they lose or gain a hydrogen ion. If we represent the acid form of the indicator as H:In, and the base form of the indicator as :In– , we can write its reaction with a base :B– or acid H:B as the following reversible equation. H:In (aq) + :B– (aq) :In– (aq) + H:B (aq) Acid form of indicator Base Base form of indicator acid Each indicator has its characteristic tendency to ionize; therefore, its color change will occur over a specific range of pH values.
Table I shows some common indicators with their color changes and the pH ranges at which they occur. Table I. Common Acid-Base Indicators Indicator pH range of color change Color in acid Color in base Methyl violet 0 – 2.0 Yellow Violet Congo red 3.0 – 5.0 Blue Red Litmus 5.0 – 8.2 Red Blue Phenolphthalein 8.3 – 10.0 Colorless Red/pink Alizarin Yellow R 10.1 – 12.0 Yellow Red Hydrolysis Weak acids are poorly ionized in aqueous solution; that is, they are poor proton donors or BronstedLowry acids. Conversely, the bases formed when weak acids dissociate (their conjugate bases) have relatively strong affinity for protons and are comparatively strong Bronsted-Lowry bases. Consider the dissociation of H2CO3 below to produce the conjugate base HCO3 – . Correspondingly, the conjugate acids formed when weak bases react with H2O are relatively strong.
For example, consider dissolving HCl in water. Since HCl is a strong acid, Cl– , its conjugate base, is weak. Further, since we know HCl is strong and does dissociate completely, it is a stronger acid than H3O+ . Likewise, since NH4 + , the conjugate acid of NH3, does in fact react with OH– to produce NH3, NH4 + is a stronger acid than H2O. Similarly, H3O+ is a stronger acid than H2CO3. Table II lists some common acids and bases in order of their relative acid and base strengths. This table will aid you in your interpretation of the observations you will make in the hydrolysis experiments. Table II. Relative strengths of acids and their conjugate bases Acid Conjugate Base Strongest H2SO4 HSO4 – Weakest HCl Cl– H3O+ H2O HSO4 – SO4 2– H3PO4 H2PO4 – HC2H3O2 C2H3O2 – H2CO3 HCO3 – H2PO4 – HPO4 2– NH4 + NH3 H2O OH– HCO3 – CO3 2– Weakest HPO4 2– PO4 3– Strongest Principles of Chemistry Lab II Montgomery College, Rockville 5 Buffers The proper function of biochemical systems requires that the pH be maintained within a few tenths of a pH unit, yet many cellular reactions produce hydronium ions which could radically alter the pH. Blood normally has a pH of about 7.4; any significant variation can result in death. Many chemical reactions occurring in the laboratory or in industry also require careful control of pH.
The pH of a system can be controlled by the use of a buffer. A buffer is a substance or combination of substances capable of consuming limited amounts of added H+ or OH– , thereby preventing significant changes in the pH of the system. Consider one of the major buffers in blood, the carbonic acid/bicarbonate ion system. Carbonic acid (aqueous CO2) is a weak acid that ionizes as follows. H2CO3 (aq) + H2O (l) H3O+ (aq) + HCO3 – (aq) If we prepare a solution containing both carbonic acid and bicarbonate ion (from NaHCO3) at equal concentrations, as shown below, we see that added OH– simply reacts with H2CO3. H2CO3 (aq) + OH– (l) H2O (aq) + HCO3 – (aq) Similarly, added H+ reacts with HCO3 – .
HCO3 – (aq) + H+ (l) H2O (aq) + H2CO3 (aq) Acid or base added to the system is removed by reacting with one component of the buffer system, converting it into the other buffer component. Thus, as long as we do not exceed the buffering capacity (which depends on the concentration of the buffer components), addition of H+ or OH– to the buffer solution results in only very slight pH changes.
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