Lesson Five: Acids, Bases, ph, and Buffers

Transcription

1 Lesson Five: Acids, Bases, ph, and Buffers Arrhenius Acids and Bases Acids and bases can be defined a number of ways. One of the oldest and most common ways is the definition according to Arrhenius, named in honour of the Swedish scientist Svante Arrhenius ( ). Acids According to Arrhenius, an acid is a chemical compound which dissolves in water to produce hydrogen ions (H + ) with the following properties: can be liquid or solid corrosive sour taste (vinegar (acetic acid), lemon juice (citric acid)) good conductor of electricity reacts with many metals (Mg, Zn, K, Na...) to produce hydrogen gas reacts with carbonates to produce CO 2 gas makes Litmus paper become red makes bromthymol blue turn yellow If the acid dissociates (or ionizes) very much (almost completely), the acid is considered a strong acid. Some examples are H 2 SO 4 and HCl. If the acid dissociates very little, the acid is considered as a weak acid. Some examples are CH 3 COOH and organic acids. The most common acids in the lab are: hydrochloric acid (HCl) sulphuric acid (H 2 SO 4 ) nitric acid (HNO 3 ) acetic acid (CH 3 COOH) Some common household acids are: acetylsalicylic acid (ASA in aspirin) ascorbic acid (vitamin C) acetic acid (vinegar) citric acid (lemon juice) carbonic acid (soft drinks) boric acid (antiseptic) sulphuric acid (car batteries)

2 Bases/Alkali Bases are considered the chemical opposites of acids. According to Arrhenius, a base is a chemical compound which dissolves in water to produce hydroxide ions (OH - ) with the following properties: can be liquid or solid has a soapy/slimy feeling (slippery) very corrosive, or caustic (they can burn) has a bitter taste good conductor of electricity reacts with a acids does not react with metals to form hydrogen makes red litmus paper turn blue makes phenolphthalein become pink makes bromthymol blue become blue If the base is soluble in water, it is considered as a strong base. Some examples are NaOH and KOH. If the base dissolves very little in water, it is considered as a weak base. Some examples are Al(OH) 3 and Fe(OH) 3. The most common bases in the lab are: sodium hydroxide (NaOH) potassium hydroxide (KOH) calcium hydroxide (Ca(OH) 2 ) sodium carbonate (Na 2 CO 3 ) sodium bicarbonate (NaHCO 3 ) ammonium hydroxide (NH 4 OH) Some common household bases are: household ammonia (Windex) lye sodium hydroxide (used in drain cleaners(drano)) toothpaste is a mild base. It neutralizes the acids which are harmful to your teeth, formed when bacteria in saliva act on the sugars in food.

3 Calculate the Concentration of Strong Acids/Bases What is the concentration of hydrogen ions in a solution made by dissolving 165 g of pure H 2 SO 4 in 750 ml of water? By dividing mass by molar mass, you can find moles. You need a periodic table to find molar mass. But since each H 2 SO 4 molecule contains 2 hydrogen ions, [H + ] = 2 x [H 2 SO 4 ] [H +] = 2 x [H 2 SO 4 ] = 2 x 2.24M = 4.48 M So the concentration of hydrogen ions in this solution is 4.48 M. The same method is used for all other strong acids and bases. Interaction of Water with Acids and Bases Bases can neutralize the chemical properties of acids. If you mix an acid and a base, the H + and OH - ions react together to form water: Since water contains very little of its H + and OH - ions, the reaction tends towards the right side (H 2 O). Water is simultaneously an acid and a base, because it liberates H + and OH - ions. The expression of the constant, K, for the preceding reaction is: [H 2 O] is a constant factor because very little water dissociates into H + and OH - ions. So, the expression for K can be re-written as K(H 2 O)= [H + ][OH - ] But water is a unique substance and has a K that can be given its own specific name. This is the equilibrium constant for water and its symbol is K w.

4 K w = [H + ][OH - ] = 1.0 x at 25 C In pure water, each time that an H+ ion dissociates, an OH- ion must also dissociate. So, for pure water, [H + ] = [OH - ] and [H + ][OH - ] = 1.0 x Therefore, K w = [H + ][OH - ] = 1.0 x [H + ] = [OH - ] = (1.0 x ) 1/2 = 1.0 x 10-7 M So, the concentration of H + and OH - ions in pure water is 1.0 x 10-7 M. Note that the value of K w is a constant and always has the value of 1.0 x If the [H + ] increases in an aqueous solution, the [OH - ] decreases in order to maintain the constant value of K w. Example: Show that the addition of 0.01 mol of NaOH (s) to 1 L of water reduces the [H + ] to 1.0 x M. Solution The two reactions that are occurring in the solution So, [NaOH] dissolved = [OH - ] = 0.01 mol/l = 1.0 x 10-2 M And, [OH - ] total = [OH - ] water + [OH - ] added = 1.0 x 10-7 M x 10-2 M Finding the total [OH - ] But 10-7 is very negligible compared to 10-2, so it is not considered in the calculations. Thus, [OH - ] = 1.0 x 10-2 M K w = [OH - ][H + ] = 1.0 x Finding the new [H + ] So, the new concentration of H + ions in the solution is 1.0 x M (a lot lower than in pure water).

5 Example: Suppose that 3.65 g of HCl are dissolved in 10.0 L of water. a. What is the value of [H + ]? b. Show that [OH - ] is 1.00 x M. Solution: Part A Find number of moles: So, [HCl] = 0.01 mol/l = 0.01 M = [H+] added Therefore [H + ] total = [H + ] added + [H + ] water = 0.01 M x 10-7 (the latter term is negligible) = 0.01 M = 1.0 x 10-2 M Solution: Part B K w = [OH - ][H + ] = 1.0x10-14 Therefore, the concentration of hydroxide ions is 0.01 M. Brønsted-Lowry Acids and Bases A hydrogen atom has an atomic number of 1. Therefore, it contains 1 proton, 1 electron and no neutrons. The H + ion is a hydrogen atom which has lost its electron. What is left over is 1 proton. Therefore the H + ion can also be called a proton. Acids can be defined as being proton donors Bases are proton recipients

6 Water can be considered as an acid and a base. So, you can consider one of the water molecules as a proton donor and the other as a proton recipient. The proton donor (or H + donor) goes from H 2 O to OH -. Consider the following reaction: NH 3 is a base and HCl is an acid. If you consider the reverse reaction, you can consider NH 4 + an acid and Cl - a base because NH 4 + must donate a proton to become NH 3 and Cl - must accept a proton to become HCl. So, every acid-base reaction can be considered as a reaction which produces a new acid and a new base: These are called conjugate acid-base pairs. Note that the base paired to acid 1 is base 1. This are called a conjugate pair. Also, the base paired with a strong acid is a weak base. The base paired with a weak acid is a strong base. Example: In the following reactions, does ammonia act as an acid or a base? a. b. c. Solution a. The ammonia goes from NH 3 to NH 4 +, thus gaining H+ (a proton). So, it is a proton acceptor. By definition, a proton acceptor is a base. b. It is a base here. c. It acts as an acid here.

7 Lewis Acids and Bases In 1923, Gilbert Newton Lewis, proposed a definition of acids and bases that emphasizes the shared electron pair. A Lewis acid is defined as an electron-pair acceptor. A Lewis base is an electron-pair donor. For example, in the neutralization reaction By eliminating the spectator ions, the reaction is essentially: So the electron pair is donated by the OH-, thus it is the base, and the H+ accepts the electron pair and is the acid. The ph scale: A Measure of [H + ] Since the concentration of H + ions is usually very small for many weak acids (ex: 10-8, ), a scale has been established which corresponds to the exponent of the [H+], and measures the strength of the hydrogen ion in solution. This scale is called the ph scale (power of hydrogen). Keep in mind that the ph scale is not based on the exponent (logarithmic) and thus a drop in ph from 4 to 2 is not a 50% change in acidity but a 100 fold increase in the concentration of hydronium (or H + ) ions. Up to this point, you have defined the strength of an acid (or base) according to its concentration of H + (or OH - ) ions. You have used qualitative terms such as "concentrated", "diluted", "weak", and "strong". But, quantitative terms should be much more precise, so you will now use the ph scale, which has values from 0 to 14. For acids, the ph is below 7 For bases the ph is above 7 For neutral substances the ph is 7 Here are some sample values for ph. The ph of oven cleaners is about 14. Your blood maintains a ph between 7.35 and 7.45; if blood changes more than a few tenths of a ph unit from this normal range, the results could be fatal since too much acid or base interferes with the ability of the blood to pick up, carry and replace oxygen. Most plants grow best in soil with a ph value between 6 and 7; higher or lower values of ph prevent plants from absorbing nutrients from the soil. Most bacteria that cause food spoilage cannot grow in solutions having the low ph value of vinegar. Shampoos normally have a ph of about 8; your scalp has a ph of about 6.

8 In order to calculate the ph or the concentration of the hydrogen ions in a solution, you can use the following formulas: ph = -log[h + ] and [H + ] = 10 -ph Since the ph changes exponentially, a ph of 1 is 10 times more acidic than a ph of 2. Similarly, a ph of 14 is 10 times more basic than a ph of 13. Example: Find the ph of a solution that has 4.2 x 10-4 M hydrogen ions. Solution ph = -log[h + ] = -log(4.2 x 10-4 M) = -(-3.38) Or you can put all the information into your calculator at once and get the same answer! = 3.38 The ph of the solution is 3.38 Example: Calculate the concentration of hydrogen ions in a solution that has a ph of Solution [H + ] = 10 -ph = = 3.98 x So the concentration of H + ions is 3.98 x M. The ph of a substance can be determined by observing the change in colour of certain indicators, which change to different colours in different ph ranges. These indicators can be in the form of solutions or in the form of reactive pieces of paper (ph paper) which contain the solution. An indicator that contains a mix of indicators is called a universal indicator. The ph can also be measured using an electronic device known as a ph-meter. Another important relation is shown below. poh is similar to ph, except [OH - ] is used instead of [H + ]. ph + poh = 14 where poh = -log[oh - ] This allows you to calculate the [H + ] and ph of a basic solution.

9 Example: What is the ph of a solution containing 10-2 M NaOH? Solution poh = -log[oh - ] = -log (10-2 M) =2 ph + poh = 14 so, ph = 14 poh ph = 14 2 =12 The ph of the NaOH solution is 12. Buffers Many biological processes require specific ph levels. In order to maintain these very narrow ranges, special types of solutions are needed: buffers. Buffers play an important role in biological systems maintaining the ph levels with an acceptable range to permit chemical reactions to proceed normally. A buffer is a solution that resists changes in ph and maintains ph levels by taking up or releasing hydrogen ions or hydroxyl ions in a solution. It is made with a weak acid and a soluble salt containing the conjugate base of the weak acid or a weak base and a soluble salt containing the conjugate acid of the weak base. A solution of the acid and its conjugate base form a compound that undergoes little change in ph when small quantities of strong acid or base are added. The weak acid reacts with any extra OH - ions put into the solution and the conjugate base reacts with any extra H + ions put into the solution. Example An example is carbonic acid and sodium bicarbonate system which occur in the blood as a buffer. The ph of a healthy person's blood plasma is between 7.35 and The blood's ph must be in the range of in order to survive. If the ph is in the range the person will feel tired, have trouble breathing, and may even be disoriented. If the ph of the blood is in the range, the person will feel dizzy and rather agitated. If the ph is out of the range, oxygenated hemoglobin releases its oxygen which can eventually lead to death. The addition of hydroxide ions will cause a shift to the left because the OH - will react with the H +. If excess H + is added, the reaction will again shift left. This helps maintain our blood s ph.

10 Another example of a buffer system is the acetic acid-buffer equilibrium: CH 3 COOH(l) + H 2 O(l) H 3 O + (aq) + CH 3 COO - (aq) The addition of hydroxide ions will cause a shift of the acetic acid-buffer equilibrium to the left because the OH - will react with the H 3 O + H 3 O + (aq) + OH - (aq) 2H 2 O(l) The increase in water molecules will cause a shift back to the right in the acetic acid-buffer equilibrium. CH 3 COOH(l) + H 2 O(l) H 3 O + (aq) + CH 3 COO - (aq) The net result of adding the hydroxide ion to the system is: CH 3 COOH(l) + OH - (aq) H 2 O(l) + CH 3 COO - (aq)