Chapter 17. How are acids different from bases? Acid Physical properties. Base. Explaining the difference in properties of acids and bases

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1 Chapter 17 Acids and Bases How are acids different from bases? Acid Physical properties Base Physical properties Tastes sour Tastes bitter Feels slippery or slimy Chemical properties Chemical properties Turns blue litmus red Turns red litmus blue Reacts with some metals => H 2 gas released - Denatures proteins React with carbonates => CO 2 released Explaining the difference in properties of acids and bases Arrhenius acid is a substance that produces H + (H 3 O + ) in water Explaining the difference in properties of acids and bases Brønsted-Lowry Concept At the molecular level: Acids are proton (H + ) donors in aqueous solutions Hydrogen ions, H + aqua = water Bases are proton (H + ) acceptors in aqueous solutions Arrhenius base is a substance that produces OH - in water 1

2 An example of a Brønsted-Lowry Acid: Here s an example of a Brønsted-Lowry Base: H 3 O + is called a hydronium ion, and is how acids exist in water. Water here is acting as a base, even though we don t normally think of it that way. A hydroxide ion (OH - ) is produced. This is how bases are usually expressed in water. Notice that here the water is acting as an acid. When acids and bases react together, the result is a salt, plus water. Here are some examples: HCl + NaOH NaCl + H 2 O Hydrochloric sodium sodium water acid hydroxide chloride Polyprotic Acids & Bases Some acids have more than one hydrogen to donate; some bases can accept more than one hydrogen. They are polyprotic (more than one proton). HNO 3 + KOH KNO 3 + H 2 O Nitric potassium potassium water acid hydroxide nitrate 2HCl + Ca(OH) 2 CaCl 2 + H 2 O Hydrochloric calcium calcium water acid hydroxide chloride A salt is a soluble ionic compound. Other examples include H 2 CO 3, H 3 PO 4, citric acid, etc.. 2

3 Conjugate acids and bases In most acid base reactions, there is an equilibrium; all four components are present. (But not at equal concentrations). This is what happens when acetic acid (the acid in vinegar) dissolves in water. In this case, the equilibrium lies very strongly to the left. The equilibrium constant, K, is a very small number. In an equilibrium involving a weak acid or base, we call the acid & base on the right the conjugate acid and conjugate base. You can identify these by remembering the Brønsted-Lowry definition: acids donate protons (H + ), bases accept protons. So in the previous reaction: C 2 H 3 O 2 H + H 2 O C 2 H 3 O H 3 O + acid base conj. base conj. acid The pattern is always the same: you have all four components; an acid reacts with a base; the acid s partner on the other side is a base, and the base s partner is an acid. Acid-Base Properties of Water H 2 O (l) + H 2 O (l) H O + H O H H 3 O + (aq) + OH- (aq) autoionization of water H base H O H + - H O [ ] + H conjugate acid H 2 O + H 2 O H 3 O + + OH - acid conjugate base At 25 C, [H 3 O + ] = 1.0 x 10-7 = [OH - ] The Equilibrium Expression for Water H 2 O (l) + H 2 O (l) H 3 O + (aq) + OH- (aq) K - w = [H 3 O + ][OH ] K c = [H 3O + ][OH - ] [H 2 O] 2 [H 2 O] = constant The equilibrium constant for the ionization of water (K w ) is the product of the molar concentrations of H + and OH - ions at a particular temperature. At 25 C K w = [H 3 O + ][OH - ] = 1.0 x [H 3 O + ] = [OH - ] [H 3 O + ] > [OH - ] [H 3 O + ] < [OH - ] Solution Is neutral acidic basic See Problem

4 ph and the ph scale ph and the ph scale ph is a way to put a number on how acidic or basic a solution is. Mathematically, ph is found by: ph = -log [H 3 O + ] Since the ph scale is logarithmic, a change of one ph unit is a ten-fold increase (or decrease) in acid concentration. Example: A ph of 3 is ten times more acid than a ph of 4 and 1000 times more acid than a ph of 6. Since ph is a negative log, the lower the ph, the more acid it is. The ph scale usually runs from 0 to 14. Since [H 3 O + ] of pure water is 1.0 x 10-7, its ph is 7.0. When the ph is <7, there is more H 3 O + (the acid component), so the solution is acidic. When the ph is >7, there is more OH - (the base component). So the solution is basic. If you know the ph, then [H 3 O + ] = 10 -ph ph scale (cont.) The ph scale can be used to measure how acidic or basic many common substances are: Acidic Basic Calculating ph and poh ph = -log [H 3 O + ] So [H -ph 3 O + ] = 10 Similarly, poh = -log[oh - ], and [OH - ] = 10 -poh battery acid soda urine sea water household ammonia gastric juice blood baking rain soda drain cleaners (Drano, Liquid Plumr) Since K w = [H O + ][OH - 3 ], taking negative logs gives: ph + poh = See Review & Check,

5 Acids and bases can be strong or weak. When you dissolve HCl in water, there is (for all intents and purposes) no equilibrium; all you have is chloride ion and hydronium ion. Strong Acid 100% dissociation HCl (g) H 2 O H + (aq) + Cl - (aq) No equilibrium K is huge Weak Acid not completely dissociated CH 3 COOH CH 3 COO - (aq) + H + (aq) Equilibrium present K is < 1 This makes HCl a strong acid. There are only three common strong acids: HCl, HNO 3, and H 2 SO 4. All other acids are weak. ph of Strong Acids Strong Acids are strong electrolytes HCl (aq) + H 2 O (l) H 3 O + (aq) + Cl- (aq) With strong acids, there is no equilibrium, all the acid is dissociated in water: HA H + + A - So the H + concentration is the same as the HA concentration, and ph can be found assuming all the HA is converted to H + and A - HNO 3(aq) + H 2 O (l) H 3 O + (aq) + NO 3 - (aq) H 2 SO 4(aq) + H 2 O (l) H 3 O + (aq) + HSO 4 - (aq) Weak Acids are weak electrolytes HF (aq) + H 2 O (l) H 3 O + (aq) + F- (aq) HNO 2(aq) + H 2 O (l) H 3 O + (aq) + NO 2 - (aq) HSO 4 - (aq) + H 2 O (l) H 3 O + (aq) + SO 4 2- (aq) H 2 O (l) + H 2 O (l) H 3 O + (aq) + OH- (aq) 5

6 Strong Bases Soluble metal hydroxides (e. g., NaOH, KOH) are strong bases; all other bases are weak. The ph of a strong base may be calculated in a fashion similar to that for strong acids. Except that you will need to derive the ph from K w = [H + ][OH - ] or, ph + poh = Acid - Base Equilibrium In an equilibrium we have: 1. All compounds present at once, 2. but not all in the same concentration. 3. These concentrations remain constant over time. Since the concentrations don t change, we can define an equilibrium constant. For the reaction: aa + bb cc + dd We can express an equilibrium constant as: K = [C]c [D] d [A] a [B] b The ionization of a weak acid is: HA + H 2 O H 3 O + + A - Relative strength of some weak acids Since this is an equilibrium, we can have an equilibrium expression: K = [H 3 O + ][A - ] [HA][H 2 O] Since water is a pure liquid, we can ignore it and simplify this to: K a = [H 3 O + ][A - ] [HA] Here K a is the acid dissociation constant. We can also have similar expressions for bases. Smaller K a means a weaker acid. Identical concentrations of a weaker acid will have a higher ph ([H 3 O + ] is lower). 6

7 ph of Weak Acids To find the ph of a weak acid: In weak acids, there is an equilibrium, so the equilibrium expression will need to be used: K a = [H 3 O + ][A - ] [HA] K a must be known. Solve for [H + 3 O ] using the ICE (Initial / Change / Equilibrium) method. Assume x is much smaller than [HA] (so we can avoid using the quadratic formula). What is the ph of a 0.50 M HF solution (at 25 C)? HF (aq) H + (aq) + F - (aq) K a = [H+ ][F - ] = 7.1 x 10 [HF] -4 Initial (M) Change (M) Equilibrium (M) K a = x 2 HF (aq) x +x x H + (aq) + F - (aq) x x - = 7.1 x 10-4 K a << 1 So 0.50 x x x 2 K a 0.50 = 7.1 x 10-4 x 2 = 3.55 x 10-4 x = M [H + ] = [F - ] = M ph = -log [H + ] = 1.72 x [HF] = 0.50 x = 0.48 M See Problem 17.5 When can I use the approximation? When K a is less than A 0 / 100. K a << x 0.50? K a = 7.1 x /100 = 5.0 x 10-3 Approximation OK. What is the ph of a 0.05 M HF solution (at 25 0 C)? K a = 7.1 x /100 = 5.0 x 10-4 Approximation not OK. Must solve for x exactly using quadratic equation. Solving Weak Acid Equilibrium Problems 1. List the major species in solution. 2. Choose the species which can produce H + (i. e., are acids) and write balanced equations. 3. Choose which reaction will dominate in making H + (i. e., greatest K a ). H 2 O H + + OH - can usually be ignored. 4. Write the equilibrium expression for this reaction. 5. Use ICE method list initial concentration(s). 6. Define change to reach equilibrium (i. e., x) 7. Add up to get equilibrium concentrations, in terms of x. 8. Substitute equilibrium concentrations into equilibrium expression. 9. Use approximation x is small compared to original [HA]. 10. See if approximation is valid. 11. Calculate [H + ] and ph 7

8 Solving Weak Acid Equilibrium Problems In a similar way, we can use the ph to calculate the K a. So, for the generic reaction: HA H + + A - We can get [H + ] from the ph; [A - ] must be the same as [H + ] (why?); and [HA] is the given concentration. So the only thing we don t know is K a. See Problem 17.4 NH 3(aq) + H 2 O (l) = [NH 4 + ][OH - ] K b [NH3 ] Weak Bases or, NH 4 + (aq) + OH - (aq) = [BH+ ][OH - ] K b [B] K b is the base ionization constant K b weak base strength Solve weak base problems like weak acids except solve for [OH - ] instead of [H + ], and then calculate ph. Some examples of weak bases. Notice that they are all organic compounds (amines) except NH 3 Conjugate Acids & Bases To form the conjugate acid of a compound or ion, you add a hydrogen; to form the conjugate base you remove a hydrogen. So, for example, the conjugate acid of the bicarbonate ion (HCO 3- ) is carbonic acid (H 2 CO 3 ), and its conjugate base is carbonate ion (CO -2 3 ). A relatively strong acid will have a relatively weak conjugate base, and vice versa. A relatively strong base will have a relatively weak conjugate acid, and vice versa. 8

9 Conjugate Acids & Bases Conjugate acid & base K s are mathematically related: For carbonic acid: H 2 CO 3 H + + HCO 3 - K a = 4.2 x 10-7 HCO H 2 O H 2 CO 3 + OH - K b = 2.4 x 10-8 Polyprotic Acids Polyprotic acids can lose more than one proton (H + ). Thus there are several K a s: K a1, K a2, etc. Usually the K a s for the loss of each proton are quite different, making calculations easier. If you add these two reactions up, you get: H 2 O H + + OH - K w = 1.0 x So, for a conjugate acid base pair: K a K b = K w See Problem 17.9 Acid-Base Properties of Salts Salts can be acidic, basic, or neutral. A salt of a strong acid and a strong base will be neutral (ph=7). Ex.: NaCl will be neutral because NaOH is a strong base and HCl is a strong acid. A salt of a weak acid and a strong base will be basic (ph>7). Ex.: KCN is basic because KOH is a strong base and HCN is a weak acid. A salt of a strong acid and a weak base will be acidic (ph<7). Ex.: NH 4 Cl is acidic because NH 4 OH is a weak base and HCl is a strong acid. See Problem 17.2 Acid-Base Properties of Salts The ph of a salt of a weak acid and a weak base depends on the relative strength of the acid and base (K a and K b ). Ex.: The ph of NH 4 C 2 H 3 O 2 is around 7 because the K a of C 2 H 3 O 2 H is approximately equal to the K b of NH 4 OH. 9

10 Acid-Base Properties of Salts You can calculate the ph of (for example) the salt of a weak acid and a strong base by using the same ICE method as before, but with a different equilibrium: u A - + H 2 O HA + OH - This equilibrium is used because you start off with the anion A - totally dissociated, from (e. g.) Na + A -, plus the solution must be basic (OH - > H 3 O + ). See Problems 17.7, 17.8 Acid-Base Properties of Oxides Covalent oxides (oxides of non-metals) form acids: SO 3 + H 2 O H 2 SO 4 CO 2 + H 2 O H 2 CO 3 Ionic oxides (oxides of metals) form bases: CaO + H 2 O Ca(OH) 2 Fe 2 O 3 + 3H 2 O 2Fe(OH) 3 Molecular Structure and Acid Strength - + Z O H Z O - + H + The O-H bond will be more polar and easier to break if: Z is very electronegative or Z is in a high oxidation state (more oxygens attached) Lewis Acids and Bases The Lewis Acid-Base Model is yet more general (more inclusive) than the Brønsted-Lowry model. A Lewis acid is an electron-pair acceptor. A Lewis base is an electron-pair donor. Weak acid weaker than weaker than Strong acid 10

11 Here s an example: Lewis Acids & Bases Lewis Acids & Bases Many reactions of organic compounds involve Lewis acids & bases. Here s a simple example: Here Cu +2 accepts the electron pair, so it s a Lewis Acid. The ammonia donates the electrons, so it s a Lewis Base. Notice no hydrogens involved! This is why CO 2 is considered an acidic oxide it s a Lewis acid. This is a complex ion more about them in the next chapter. Lewis Acids & Bases Almost all the bases in organic chemistry contain nitrogen. It has a free electron pair which it can donate thus it s a Lewis Base. Note all these can accept protons, and so can also be Brønsted bases. 11

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