Acids and Bases

What are acids and bases? Lemons, grapefruit, vinegar, etc. taste sour because they contain acids. Acid in our stomach helps food digestion Acid from bacteria turn milk sour and are used to make yogurt and cheese. oBases neutralize acids (pH = 7) oAntacids are taken to offset the effect of too much acid in the stomach. Other examples are drain cleaners and oven cleaners.

Naming Acids Acids are substances that dissolve in water to produce hydrogen ions (H + ) and a simple nonmetal anion. HCl (g)  H + (aq) + Cl - (aq) Hydro…… is used before the name of the nonmetal, and its …..ide ending is changed to ………ic acid. When the acid contains a polyatomic ion the name comes from the polyatomic ion… If it ends in …..ate, replace with …….ic acid If it ends in …..ite, replace with …..ous acid

Naming Bases Bases are ionic compounds that dissociate into a metal ion and hydroxide ions (OH - ) when they dissolve in water. NaOH (s)  Na + (aq) + OH - (aq) Bases are named as hydroxides.

Bronsted-Lowry acids and bases Bronsted-Lowry Acid - donates a proton (hydrogen ion, H + ) to another substance Bronsted-Lowry Base - donates accepts a proton. The proton does NOT actually exist in water………a hydronium ion (H 3 O + ) is formed when the proton is attracted to the polar water molecule. H - O + H +  H - O - H + l l H H Water proton hydronium ion

Bronsted-Lowry acids and bases Examples: HCl + H 2 O  H Cl - When ammonia (NH 3 ) reacts with water, the nitrogen has a stronger attraction for the proton than water. NH 3 + H 2 O  NH OH -

Water can be an acid or a base! Identify the reactant that is an acid ( H + donor) and the reactant that is a base (H + acceptor). HBr + H 2 0  H Br - H 2 O + CN -  HCN + OH - Thus, H 2 O can be an acid or a base!

Conjugate acid-base pairs When molecules are related by the loss or gain of one H+ (proton) they make a conjugate acid-base pair. The protons are transferred both forward and reverse! Conjugate acid-base pair HA + B  A - + BH + Acid 1 Base 2 Base 1 Acid 2 H+ donor H+ acceptor H+ acceptor H+ donor

Problem Time! Identify the conjugate acid-base pairs in the following reactions; HF + H 2 O  F - + H NH 3 + H 2 0  NH OH -

Bronsted-Lowry acids and bases

Classifying acids and bases Acids -classified according to their ability to donate protons Strong acids………give up protons easily Weak acids………..give up only a few protons and most molecules keep their protons. Bases - classified in terms of their ability to accept protons. Strong bases……..have a strong attraction for protons Weak bases……….have little attraction for protons

Strong and Weak Acids Strong acids dissociate almost completely to give H ions and anions. HCl + H 2 0  H Cl - Weak acids dissociate only slightly and produce small concentrations of H30+ ions. CH 3 COOH + H 2 O  H CH 3 COO -

Strong and Weak Bases Strong bases are ionic compounds that dissociate in water to give an aqueous solution of a metal ion and hydroxide ion. KOH (s)  K + (aq) + OH - (aq) Weak bases are poor acceptors of protons. NH 3 (g) + H 2 O (l)  NH 4 + (aq) + OH - (aq)

Ionization of Water Remember water can act as BOTH an acid AND a base…………….how? One water molecule can donate a proton to another to produce H 3 O + and OH - Conjugate acid-base pair  H 2 O (l) + H 2 0 (l)  H 3 O + (aq) + OH - (aq)  Conjugate acid-base pair

Ion-product constant In pure water the transfer of proton between two water molecules produces EQUAL numbers of hydronium and hydroxide ions. [H ] = [OH - ] = 1.0 x M Ion-product constant for water (K w ) = [H 3 O + ] x [OH - ] = 1.0 x (no units) The K w values applies to any aqueous solution. ALL aqueous solutions have have H 3 O + and OH - ions. [H 3 O + ] = [OH - ] ……solution is neutral [H 3 O + ] > [OH - ]……………………. ……..acid [H 3 O + ] < [OH - ] ……………………..…..basic Remember the sum of both ions is always = 1.0 x at 25 o C.

pH Scale On the pH scale a number between 0 and 14 represents the H 3 O + concentration. Acidic solutionpH 1.0 x M Neutral solutionpH = 7[H 3 O + ] = 1.0 x M Basic solutionpH > 7[H 3 O + ] < 1.0 x M

Calculating the pH of solutions pH scale is a log scale. ………………… pH = -log[H 3 O + ] Example…………. lemon juice solution has [H 3 O + ] = 1.0 x M pH = - log[ 1.0 x 10-2] = - (- 2.00) = 2.00 Remember the number of significant figures……….answer should have the same number of sig. figs. as the [H 3 O + ] [H 3 O + ] = 1.0 x pH = 3.00   2 significant figs.2 decimal places

Measuring pH

Reactions of Acids and Bases Acids will react with - metals - carbonates and bicarbonates - hydroxides

Acids + Metals Acid + “active metals”  “salt” + hydrogen gas These are single replacement reactions. Active metals are ; potassium, sodium, calcium, magnesium, aluminium, zinc, iron, tin. Example: Mg (s) + 2HCl (aq)  MgCl 2(aq) + H 2(g) Metalacid salthydrogen

Acids + carbonates & bicarbonates Strong acid + carbonate/bicarbonate  CO 2 + H 2 O + salt Example: HCl (aq) + NaHCO 3(aq)  CO 2(g) + H 2 O + NaCl (aq)

Acids + Hydroxides Neutralization occurs between an acid and a base to form a salt and water. HCl + NaOH  NaCl + H 2 O ( H + Cl - Na + OH - Na + Cl - H 2 O ) One can write the net ionic equation ( omitting the metal ions and chloride ions that are not reacting) as follows: H + + OH -  H 2 O

Buffers BUFFER SOLUTIONS resist changes in pH when small amounts of acids or base are added. Buffers contain an acid to react with any OH - that is added and a base to react with any H 3 O + added………BUT that acid and base must not neutralize each other. So, a combination of an acid-base pair conjugate is used in buffers ie. a weak acid and its salt and a weak base and its salt. CH 3 COOH (aq) + H 2 O (l)  H 3 O + (aq) + CH 3 COO - (aq)Large amount

Buffers 1. If a small amount of acid is added then: CH 3 COOH (aq) + H 2 O (l)  H 3 O + (aq) + CH 3 COO - (aq) Acid will combine with the acetate ion and shift the equilibrium…  [CH 3 COO - ]  [CH 3 COOH]no change in H 3 O + 2. If a small amount of base is added then: CH 3 COOH (aq) + H 2 O (l)  H 3 O + (aq) + CH 3 COO - (aq) Base is neutralized by the acetic acid to produce water and shift the equilibrium.  [CH 3 COO - ]  [CH 3 COOH]no change in H 3 O +

Buffers in the Blood Cells can only function properly when the pH is between 6.8 – 8.0. The normal pH of arterial blood is 7.35 – The bicarbonate/carbonic acid system is an important buffer system in the blood. H 2 O CO 2 + H 2 O  H 2 CO 3  H 3 O + + HCO 3 - Excess H 3 O + entering the body fluids reacts with the HCO 3 - and excess OH - reacts with carbonic acid.

Buffers in the Blood H 2 O CO 2 + H 2 O  H 2 CO 3  H 3 O + + HCO 3 -  CO2   H3O+   pH……………………. Acidosis …….emphysema, difficulty with breathing, medulla affected by depressive drugs or accident trauma.  CO2   H3O+   pH……………………. Alkalosis …….hyperventilation causes expiration of a lot of carbon dioxide eg. excitement, trauma, high temperatures