Activity Titrations & ph

Transcription

1 Activity Titrations & ph Directions: This Guided Learning Activity (GLA) focuses on chemical calculations related to acids, bases and ph. Part A gives basic information about acids and bases, and describes the ionization of water. Part B discusses K W and the ph scale. Part C uses stoichiometry to solve titration problems. The worksheet is accompanied by instructional videos. See for additional materials. Part A The Basics of Acids & Bases Early in the class, you probably learned that an acid is identifiable because the chemical formula is written with hydrogen as the first element. Acids are a class of chemicals that are able to ionize in aqueous solutions to produce hydronium, H 3O +, ions. For example, hydrochloric acid, or HCl, is a covalent compound. Yet when it is placed in an aqueous solution, the hydrogen from the HCl binds to a water molecule to form a hydronium ion, leaving the chloride ion in solution. HHHHHH (gg) + HH 22 OO(ll) HH 33 OO + (aaaa) + CCll (aaaa) All acids will ionize to some degree to form H 3O + ions in solution. Because of this, all acids share some characteristics. For example acids are typically sour, can dissolve metals, and turn litmus paper red. Bases, on the other hand, are compounds that increase the hydroxide, or OH -, concentration in solution. In Chemistry 151, the bases you will encounter are soluble or slightly soluble hydroxide salts. Bases have a bitter taste, a slippery soap-like feel, and will turn litmus blue. The reason acids and bases are chemically important is because the presence of an acid or base will affect the concentration of H 3O + and OH - in solution. In pure water, a very small fraction of water molecules reacts to form both of these two ions: HH 22 OO(ll) + HH 22 OO(ll) HH 33 OO + (aaaa) + OOHH (aaaa) Adding an outside source of H 3O + of OH - will affect the fraction of water molecules that are ionized. Whenever the amount of H 3O + is increased, the amount of OH - in the solution will decrease. When the amount of OH - is increased, the amount of H 3O + decreases. This occurs because when the ions encounter each other, they react to form water. Often, the hydronium ion is written as simply H +. This communicates the reactive portion of the ion (the ionized hydrogen from the acid). In general, H + and H 3O + can be used interchangeably. But you should know that a hydrogen ion (H + ) cannot exist in isolation, is always carried by another substance typically another water molecule. In this GLA, the hydronium ion is written only as H 3O +. Activity Page 1 of 5

2 Practice: Classify each of the following as an acid, a base, or neither; then determine whether the compound will release H 3O +, OH - or neither into solution. HCl is an acid and produces H 3O + in solution. SrCl 2 neither and produces n/a in solution. HC 2H 3O 2 and produces in solution. KOH and produces in solution. NaF and produces in solution. Ba(OH) 2 and produces in solution. Part B K W and the ph Scale As mentioned above, adding an acid or a base to water will change both the amount of H 3O + and the amount of OH - in the solution. Usually, the amount of H 3O + and OH - is expressed in molarity because acids and bases exist in aqueous solutions. (See GLA for a review of molarity and concentration units). The two concentrations are related by the equation: [HH 33 OO + ] [OOHH ] = Where [H 3O + ] is the molar concentration of hydronium ions, [OH - ] is the molar concentration of hydroxide ions, and 1x10-14 is a constant called K W, or the ionization constant for water. Example #1. A basic solution contains M hydroxide ions. What is the concentration of H 3O + in this solution? Because only a very small fraction of water ionizes, typical concentrations of H 3O + and OH - are quite low (~1x10-7 M). This characteristic has led to the development of the ph scale. The ph scale is defined based on the molar concentration of H 3O + in solution: This can also be expressed as: pppp = llllll [HH 33 OO + ] [HH 33 OO + ] = 1111 pppp Activity Page 2 of 5

3 According to this scale, acidic solutions (which have a high [H 3O + ]) will have a low ph (below 7) and basic solutions (which have a high [OH - ]) will have a high ph (above 7). Solutions that have a ph equal to 7 are considered neutral and have [H 3O + ] = [OH - ]. There is no physical limit to the ph scale, but it is generally drawn with boundaries around 0 and 14. Keep in mind that if we are given ph, we can determine [OH - ] by using the K w mentioned earlier. Example #2. What is the ph of a solution with a H 3O + concentration of 1x10-4 M? Is the solution acidic, basic, or neutral? Part C Titrations Titrations are a specific application of the stoichiometry concepts already discussed in GLA Specifically, an acid-base titration is a double replacement reaction between an acid and a base. These reactions will always form a salt and water. (Refer to GLA-6 Predicting Products in Chemical Reactions for additional guidance.) Because titrations are performed in aqueous solutions, generally the amount of reactant is given in terms of molarity, a concentration. Remember, a 1.0 molar solution contains 1.0 mole of solute in 1.0 liter of solution. A molar concentration can easily be written as a conversion factor. (Refer to GLA Units of Concentration for additional guidance.) 0.3 MM NNNNNNNN = 0.3 mmmmmm NNNNNNNN 1 LL ssssssssssssssss = 0.3 mmmmmm NNNNNNNN 1000 mmmm ssssssssssssssss During a titration, a solution with a known concentration (a titrant) is added slowly to another solution with an unknown concentration (an analyte). Titrations utilize indicators that undergo some change, typically a color change, when the amount of titrant added is enough to react completely with the analyte. This is termed the equivalence point. The equivalence point is where you see the most dramatic ph change with even small additions of titrant. Because both the concentration of the titrant and the amount of titrant added are known, the amount of analyte present can be found. (Refer to GLA Introduction to Stoichiometry for additional guidance.) Activity Page 3 of 5

4 Example #3. Label the analyte, titrant, indicator, and equivalence point in the following diagram ph ml Titrant Added Example # ml of a hydrobromic acid solution is titrated with ml of M NaOH. How many moles of hydrobromic acid were present in the 50.0 ml solution? Solution: HHHHHH(aaaa) + NNNNNNNN(aaaa) NNNNNNNN(aaaa) + HH 22 OO(ll) 11 LL NNNNNNNN mmmmmm NNNNNNNN 11 mmmmmm HHHHHH ( mmmm NNNNNNNN) mmmm NNNNNNNN 11 LL NNNNNNNN 11 mmmmmm NNNNNNNN = mmmmmm HHHHHH Example #4b. What was the concentration of the initial hydrobromic acid solution? Example #5. How many moles of hydrochloric acid are needed to react completely with a 75.0 ml solution of M barium hydroxide? Activity Page 4 of 5

5 Part D Extra Practice 1. What is the [OH - ] in a solution with a ph of 9.4? Is the solution acidic, basic, or neutral? 2. What is the ph of 12 M hydrochloric acid? Is the solution acidic, basic or neutral? (*Hint: HCl is a strong acid, so it ionizes completely, producing 12 M H 3O + ions.) 3. Complete the following table: [H 3O + ] [OH - ] ph Acidic, basic or neutral? 1.4 x x Neutral 4. Determine the ph, poh, [H 3O + ], or [OH - ] for the following scenarios. a. If [H 3O + ] = 0.10 M, ph = b. If [H 3O + ] = 0.10 mm, [OH - ] = c. If [OH - ] = 10 μm, ph = d. If [OH - ] = M, [H 3O + ] = e. If [OH - ] = M, ph = f. If [OH - ] = M, poh = ml of H 3PO 4 is titrated with M NaOH ml of the NaOH solution is needed to reach the equivalence point. What is the concentration of H 3PO 4 in the original solution? ml sample of acetic acid (unknown concentration) was titrated with ml of M KOH. What is the concentration of the acetic acid? 7. How much M Ba(OH) 2 is required to reach the equivalence point when titrating 16.7 ml of a solution of 2.00% (m/v) HC 2H 3O 2? Activity Page 5 of 5