CHAPTER 14 ACIDS AND BASES Properties of Acids and Bases 1. Acidic Compounds 1. react with metals that are more reactive than hydrogen. 2. react with carbonates such as marble, seashells and limestone. 3. Change litmus paper to red 4. are electrolytes- they conduct electricity when dissolved in water 5. pH is less than 7 6. they taste sour

1. are electrolytes 2. feel slippery 3. taste bitter 2. Basic Compounds 1. are electrolytes 2. feel slippery 3. taste bitter 4. pH is greater than 7 5. Turn litmus paper blue

B. Arrhenius Acids  and Bases 1. Acids a. monoprotic b. diprotic c. triprotic hydrogen-containing compounds that ionize to yield hydrogen ions (H+) Acids that contain one ionizable hydrogen Ex. HNO3 Acids that contain two ionizable hydrogens Ex. H2SO4 Acids that contain three ionizable hydrogens Ex. H3PO4

2. Bases Ionize to yield hydroxide ions (OH-) in solution Ex. NaOH C. Brønsted-Lowry Acids and Bases  1. Acids  2. Bases Hydrogen ion donors Hydrogen ion acceptors The Brønsted-Lowry theory of acids and bases includes all of the acids as described by Arrhenius and also some bases that were not included.

Example: NH3(aq) + H2O (l) → NH4+ (aq) + OH- (aq) B-L base B-L acid ammonium hydroxide    ion ion 3. conjugate acids and  bases  a. Conjugate acid  b. conjugate base the product formed when a base gains a hydrogen ion The product that remains when an acid has donated a hydrogen

 c. conjugate  acid-base pair Consists of two substances related by the loss or gain of a single hydrogen ion. NH3(aq) + H2O (l) → NH4+ (aq) + OH- (aq) In this example the NH3 and NH4+ are a conjugate acid-base pair. The H2O and the OH- are also a conjugate acid-base pair.

HCl (aq) + H2O (l) → H3O+ + Cl- In this example HCl is an acid and Cl- is the conjugate base. H3O+ is the conjugate acid of the base H2O

E. Summary of both  theories  1. acids  2. bases A substance that produces hydrogen ions when dissolved in water. Acids are known as proton (H+) donors. a substance that produces hydroxide ions (OH-) when place in water. Some bases are proton acceptors.

F. Naming rules  1. acids To name acids focus on the name of the anion of the acid. The anion may be monatomic such as HCl or polyatomic such as HNO3 1. When the name of the anion ends in “-ide”, the acid name begins with “hydro-“. The stem of the anion has the suffix “-ic” and the name is followed by “acid” Ex. HCl Hydrogen Chloride When dissolved in water it becomes Hydrochloric acid

2. When the anion ends in “-ite”, the acid name is the stem of the anion with the suffix “ous” followed by the word “acid” Ex. HNO2 (NO2 = nitrite) Name: nitrous acid 3. When the anion ends in “-ate”, the acid name is the stem of the anion with the suffix of “-ic” followed by the word “acid” Ex. H2CO3 Hydrogen Carbonate Name: Carbonic Acid

Name the following compounds and then name them as acids H2SO4  HNO3 H2SO3  HNO2 HCl  H3PO4 HC2H3O2  2. Bases Naming of bases is just the same as naming ionic compounds Ex. Ca2+ + OH1- ® Ca(OH)2 Calcium Hydroxide NaOH  KOH  NH3

G. Hydrogen Ions from Water When water molecules move around and collide they occasionally react and produce Hydronium ions H3O+ (acid) and hydroxide ions OH- (base)

The reaction between two water molecules that produces ions (OH- and H3O+ or just H+) The concentration of each ion, H+ and OH- in pure water is 1 x 10-7 mol/L Their concentrations are the same so the solution is said to be neutral. Pure water is a neutral solution.  1. self ionization of  water

a. Kw Ion-product constant for water, the product of the concentrations of the hydrogen ions and hydroxide ions Kw = [H+] X [OH-] [ ] brackets mean concentration Kw= [1 x 10-7 M] X [1 x 10-7 M] Where M stands for (mol/L) So: Kw = 1 x 10-14 M2 If additional H+ ions are added to the solution the equilibrium shifts. The concentrations of hydroxide ions will decrease. This causes the overall Kw to always be 1 x 10-14 M.

b. acidic and basic  solutions  1. acidic solution  2. basic solution Not all solutions are neutral. To determine if a solution is acidic, basic or neutral, use the following equation. Kw = [H+] X [OH-] Acids always have a H+ concentration greater than 1 x 10-7 and the [H+] is greater than [OH-] Bases have a H+ concentration less than 1 x 10-7 and the [H+] is less than [OH-] Also known as an alkaline solution.

Example: Colas have an [H+] = 1 x 10-5 M Is the solution acidic or basic? The solution is acidic because the [H+] is greater than 1 x 10-7 M What is the [OH-]? Kw = [1 x 10-5 M] X [OH-] = 1 x 10-14 M [OH-]= 1 x 10-9 M

 H. pH A measure of the hydronium ion (H3O+) concentration in solution 1. Proposed by Søren Sørensen in 1909 while he was brewing beer. 1. calculation pH It is the negative logarithm of the hydrogen ion concentration. pH = -log[H+] If the [H+] = 1 x 10-7 pH = -log(1 x 10-7) = 7 What is the pH of a solution with an [H+] = 4.5 x 10-5 M? pH = -log(4.5 x 10-5) = 4.35

2. pH scale The scale ranges from 0-14. Which comes from the concentration of H3O+, which ranges from 100 to 10-14M. 100 is a high concentration 10-14 is a low concentration pH means power of hydrogen pH of 7 is neutral pH of 0 is strongly acidic pH of 14 is strongly basic pH is measured using  a. Indicators such as litmus  paper and phenolphthalein.  Which are just dyes.  b. Instruments such as a pH  meter, which measure the  conductivity of the solution.  

a. Strong acid b. Weak Acid c. Strong Base d. Weak Base Dissociates completely in solution HCl + H2O à H3O+ + Cl- (no HCl remains) Does not completely dissociates, some of the original acid is left. Example: Acetic acid , only 0.5% reacts in water dissociates completely in solution to a metal ion and hydroxide ion Example: NaOH Does not completely dissociate. React to form hydroxide ion and conjugate acid of the base Example: ammonia, only 0.5% tends to react in water