battery acid the most widely used industrial chemical Hydrochloric acid, HCl muriatic acid stomach acid Nitric acid, HNO 3

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1 BRCC CHM 101 Chapter 9 Notes (Chapter 8 in older text versions) Page 1 of 9 Chapter 9: Acids and Bases Arrhenius Definitions more than 100 years old Acid a substance that produces H + in water (H + is actually the species H 3 O + called hydronium ion) Base a substance that produces OH - in water (OH - is called hydroxide ion) Important acids and bases Strong acids : Sulfuric acid, H 2 SO 4 battery acid the most widely used industrial chemical Hydrochloric acid, HCl muriatic acid stomach acid Nitric acid, HNO 3 strong oxidizing agent turns skin yellow Weak acids: Acetic acid, CH 3 COOH vinegar is a 5% solution in pure form called glacial acetic acid & is flammable Boric acid, H 3 BO 3 used as an antiseptic Phosphoric acid, H 3 PO 4 important acid in biochemistry Strong bases: Sodium hydroxide, NaOH called lye used in making soap and glass Potassium hydroxide, KOH similar to NaOH Weak bases: Ammonia, NH 3 widely used, especially in fertilizers Magnesium Hydroxide, used in medicine milk of magnesia Mg(OH) 2 Acid & Base Strength strong acid reacts completely with water to form H 3 O + (dissociates completely in water) weak acid reacts incompletely with water to form a small amount of H 3 O + example: CH 3 COOH + H 2 O CH 3 COO - + H 3 O + the equilibrium lies far to the left The strength of an acid or base is not related to its concentration. For example, HCl is strong whether it is dilute or concentrated. There are 6 strong acids. o H 2 SO 4 o HCl o HNO 3 o HI o HBr o HClO 4 sulfuric acid hydrochloric acid nitric acid hydroiodic acid hydrobromic acid perchloric acid

2 BRCC CHM 101 Chapter 9 Notes (Chapter 8 in older text versions) Page 2 of 9 strong base a base that readily accepts a proton (H + ) to form OH - weak base a base that is less willing to accept a proton to form a small amount of OH - example: NH 3 + H 2 O NH OH - There are 4 common strong bases. the equilibrium lies far to the left o NaOH sodium hydroxide o KOH potassium hydroxide o LiOH lithium hydroxide o Ba(OH) 2 barium hydroxide Brrnsted-Lowry Acids and Bases These are more general definitions needed because not all acids and bases involve water. o acid proton donor o base proton acceptor o acid-base reaction involves the transfer of a proton from the acid to the base CH 3 COOH + H 2 O CH 3 COO - + H 3 O + gives up a proton accepts a proton can take back a proton can give up a proton When an acid gives up a proton, it then becomes a base called the conjugate base of the original acid. When a base accepts a proton, it becomes an acid called the conjugate acid of the original base. Together, they are called a conjugate acid-base pair. CH 3 COOH + H 2 O CH 3 COO - + H 3 O + acid base conjugate base conjugate acid Acids and bases can be positive, negative, or neutral. Most acids give up one proton and are called monoprotic. Some acids can give up 2 or 3 protons. These are called polyprotic acids. o diprotic 2 protons ex. H 2 SO 4 o triprotic 3 protons ex. H 3 PO 4

3 BRCC CHM 101 Chapter 9 Notes (Chapter 8 in older text versions) Page 3 of 9 o Brrnsted-Lowry acids give up one proton at a time in multiple steps. loose the first H + H 2SO 4 + H 2 O HSO 4 - loose the second H + HSO H 2 O SO H 3 O + + H 3 O + Some substances appear as both an acid and a base. ex. HCO 3 - behaves as an acid HCO H 2 O CO H 3 O + behaves as a base HCO H 3 O + H 2 CO 3 + H 2 O A substance that can behave as both an acid and a base is called amphoteric. Water is the most important. In order to be a Brrnsted-Lowry acid, a substance must have a proton (H + ) that it can give up. ex. Phenol has 4 hydrogens, but can only give up one of them. Brrnsted-Lowry Acid-Base Strength acid strength the willingness to give up a proton base strength the willingness to accept a proton HCl + H 2 O Cl - + H 3 O + acid base conjugate base conjugate acid The equilibrium for this reaction is far to the right. Because HCl is more willing to give up its proton than H 3 O +, HCl is a stronger acid. H 2 O is more willing to accept a proton than Cl - so it is a stronger base. The stronger base always wins. The stronger the acid, the weaker its conjugate base is. Acid Dissociation Constants Acid-base reactions are all equilibrium reactions. The equilibrium constant can be used to tell how strong or weak any acid is.

4 BRCC CHM 101 Chapter 9 Notes (Chapter 8 in older text versions) Page 4 of 9 general reaction HA + H 2 O A - + H 3 O + example HCl + H 2 O Cl - + H 3 O + [A - ] [H 3 O + ] K = [HA ] [H 2 O ] ³ Water is the solvent, so its concentration does not change appreciably. It can be considered a constant. K [H 2 O ] = [A - ] [H 3 O + ] [HA ] ³ The concentration of water is combined with the constant K to make Ka, the acid dissociation constant. Ka = K [H 2 O ] = [A - ] [H 3 O + ] [HA ] The larger the Ka, the stronger the acid. Ka immediately tells us how strong an acid is. (See table 8.3) Properties of Acids and Bases Acids taste sour and bases taste bitter. Both are dangerous, but bases are more dangerous to eyes. The most important reaction is neutralization H 3 O + + OH - Y 2 H 2 O Acids react with metals to make hydrogen gas. Self-ionization of Water Pure water produces a small amount of H3O+ and OH - ions. H 2 O + H 2 O If you write an equilibrium constant for this reaction, you get: H 3 O + + OH - The equilibrium lies far to the left. K = [H 3 O + ][OH - ] [H 2 O ] 2 ³ Water concentration can be considered a constant. K [H 2 O ] 2 = [H 3 O + ][OH - ] = Kw ion product of water Kw = [H 3 O + ][OH - ] = 1.0 x Kw is valid for any water solution [H 3 O + ] = [OH - ] = 1.0 x 10-7

5 BRCC CHM 101 Chapter 9 Notes (Chapter 8 in older text versions) Page 5 of 9 example. If a strong base has a [OH - ] concentration of 1 x 10-2, then what is the [H 3 O + ] concentration? Kw = [H 3 O + ][OH - ] = 1.0 x [H 3 O + ] = 1.0 x x [OH - = ] 1.0 x 10-2 = 1.0 x The concept of ph was developed from Kw. ph = - log [H 3 O + ] ph is the negative log of the hydronium ion concentration. [H 3 O + ] = 1.0 x 10 -A M so that ph = A example. If a soap solution has [H 3 O + ] = 1.0 x 10-9 M, then what is the ph? 9 example. If ph of tomato juice is 4, then what is the [H 3 O + ]? 1.0 x 10-4 M ph is easier to use than [H 3 O + ] Neutral solutions have a ph of 7 Acids have a ph < 7 Bases have a ph > 7 How do you calculate the ph of a solution from [H 3 O + ] concentration? Use a calculator: o Enter [H 3 O + ] o Press LOG key o Change the sign from negative to positive How do you calculate [H 3 O + ] concentration from ph? Use a calculator: o Enter the ph o Change the sign from positive to negative o Press the inverse log key ex. [10 x ], or [LOG -1 ], or [INV][LOG] There are 2 common ways of measuring ph: o ph paper paper coated with an indicator that you read against a color key o ph meter an electrode that measures [H 3 O + ] and displays the output directly Acidosis & Alkalosis Blood ph is between 7.35 & 7.45 Acidosis is when the ph is too low Alkalosis is when the ph is too high respiratory acidosis, called hypoventilation, is caused by difficulty breathing by obstruction or

6 BRCC CHM 101 Chapter 9 Notes (Chapter 8 in older text versions) Page 6 of 9 disease you can produce mild acidosis by holding your breath. metabolic acidosis - comes from fasting, starvation, or heavy exercise the mechanism is that the body burns fats and produces acidic compounds or the exercise produces lactic acid respiratory alkalosis can be caused by fever, infection, or drugs it can be brought about by the excessive loss of CO 2 called hyperventilation acidosis can depress the nervous system causing fainting or coma alkalosis can cause overstimulation of the nervous system causing cramps and convulsions both can cause death Buffers buffer solution a solution whose ph does not change much when H 3 O + or OH - ions are added to it. How do you make a buffer solution? o One way is to add equal amounts of a weak acid and its conjugate base to water. o example: Add 1 mol of acetic acid (CH 3 COOH) and 1 mol of sodium acetate (CH 3 COO - Na + ) to water. Acetic acid is the weak acid and sodium acetate is its conjugate base. How does it work? o There is a substantial concentration of both acid and conjugate base in the solution. In the above example this would be CH 3 COOH and CH 3 COO -. o If H 3 O + is added, CH 3 COO - reacts to remove it. CH 3 COO - + H 3 O + CH 3 COOH + H 2 O o If OH - is added, CH 3 COOH reacts to remove it. CH 3 COOH + OH - CH 3 COO - + H 2 O o The equilibria for both of these reactions lie mostly to the right. The concentrations of CH 3 COOH and CH 3 COO - change only slightly when H 3 O + and OH - are added, so ph is maintained. o example: When HCl is added to pure water, the ph goes from 7 to 1. When HCl is added to a buffer, the ph could go from 7.21 to Every buffer solution has 2 characteristics: ph and capacity o buffer ph For equal concentrations of weak acid and conjugate base, the ph of the solution is always equal to the pka of the weak acid.

7 BRCC CHM 101 Chapter 9 Notes (Chapter 8 in older text versions) Page 7 of 9 You can look up the pka of common weak acids in tables, such as the table in the textbook. If the concentrations of the weak acid and conjugate base are not equal, then use the Henderson-Hasselbalch Equation. ph = pka + log [A - ] ³ conjugate base concentration [HA ] ³ weak acid concentration What is the ph of a H 2 PO 4 - / HPO 4 2- buffer system in which [HPO 4 2- ] is twice [H 2 PO 4 - ]? From table 8.3, you find the pka of H 2 PO 4 - to be Using the Henderson- Hasselbalch equation: ph = log 2 1 = = 7.51 o buffer ph What is the ph of a buffer solution in which [CH 3 COO - ] = 0.20 M and [CH 3 COOH] = 0.60 M? From table 8.3, you find the pka of CH 3 COOH to be Using the Henderson- Hasselbalch equation: ph = log = = 4.27 buffer capacity how much acid or base the buffer is able to neutralize this depends on the concentrations of weak acid and conjugate base. example: A buffer with 1 mol each of weak acid and conjugate base has ten times as much capacity as a buffer with.1 mol each of weak acid and conjugate base. o The most important buffer systems are the 3 blood buffer systems. They are carbonate, phosphate, and proteins. carbonate buffer system [HCO 3 - ]/[H 2 CO 3 ] is in a 10 to 1 ratio of conjugate base to weak acid this makes for a good buffer for acids under normal conditions, larger amounts of acidic substances enter the blood than basic this is the most important blood buffering system phosphate buffer system [HPO 4 2- ]/[H 2 PO 4 - ] is in a 1.6 to 1 ratio the second most important blood buffering system

8 BRCC CHM 101 Chapter 9 Notes (Chapter 8 in older text versions) Page 8 of 9 Titration, Equivalents, and Normality titration a technique for determining how much total acid or how much total base is present in a solution. o If the solution is acidic, base is added to it until all the acid is neutralized. o And, if the solution is basic, acid is added to it until all the base is neutralized. o You need a means of measuring exactly how much base or acid you add to the solution. o You need a means of telling when all the acid or base in the solution is neutralized. o buret glassware graduated to deliver exact amount of liquid. o indicator a substance which shows a sudden large change in ph by changing color. o endpoint the sudden large change in ph. Titration does not determine the acidity or basicity of a solution ph does. o example: 0.1 M HCl has a ph of 1 and 0.1M acetic acid has a ph of 2.9 o These 2 solutions in the example have very different acidities, but titration will tell you that both of them have the concentration 0.1 M. titrant this is the acid or base solution of known concentration that is added to neutralize the solution of unknown concentration. (You make the titrant up yourself.) We can read from the buret how many mls of titrant it took to neutralize the solution of unknown concentration. Then we must calculate the amount of acid or base in that solution. equivalent weight (EW) the formula weight (FW) divided by the number of H 3 O + or OH - ions produced. (Remember diprotic, triprotic, etc...) Look at the following table: Species Formula Weight FW / # of ions Equivalent weight HNO / 1 63 NaOH / 1 40 H 2 SO / 2 49 Mg(OH) / 2 29 equivalent (eq) the equivalent weight expressed as grams o 1 eq of HNO3 = 63g o 1 eq of H 2 SO 4 = 49 g o Equivalents are similar to moles. o example: How many equivalents are there in 18 grams of H 2 SO 4? 18 g H 2 SO 4 1 eq H 2 SO 4 X 49 g H 2 SO 4 = 0.37 eq H2SO4

9 BRCC CHM 101 Chapter 9 Notes (Chapter 8 in older text versions) Page 9 of 9 In acid-base reactions, one equivalent of any acid exactly neutralizes one equivalent of any base. normality (N) equivalents of solute per liter of solution similar to molarity (M) o Normality is the molarity multiplied by the number of H 3 O + or OH - ions produced per molecule. example: a 6 M solution of H 2 SO 4 is 12 N o Normality is always equal to or greater than molarity. o example: What is the normality of a solution made by dissolving 8.5 g of H 2 SO 4 in enough water to make 500 ml of solution? 8.5 g H 2 SO 4 1 eq H 2 SO ml X X 500 ml 49 g H 2 SO 4 1 L = 0.34 N At the end point of a titration, the equivalents of acid equal the equivalents of base. o Eq acid = Eq base eq Eq Since N = = L V ³ volume o This is the basis for the titration equation: V acid X N acid = V base X N base o example: If 50.0 ml of an unknown concentration of acid is titrated with 0.32 N base and it takes 24.6 ml of base to reach the end point, then what is the concentration of the acid? N acid V base N base 24.6 ml 0.32 N = X = X = 0.16 N 50.0 ml V acid