Chapter 14 Properties of Acids and Bases

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1 Section 14.1 Defining Acids and Bases Properties of acids and bases Chapter 14 Properties of Acids and Bases taste sour Acids taste bitter Bases conduct electricity no characteristic feel react with metals to produce hydrogen gas turns indicators color conduct electricity feel slippery no reaction with metals turns indicators color If red litmus paper turns blue = basic If blue litmus paper turns red = acidic If red stays red and blue stays blue = neutral Definitions of Acids and Bases: 1. Arrhenius ACID - a substance that dissociates in water to produce hydrogen ions (H+) eg HCl > H + + Cl - BASE - a substance that dissociates in water to produce hydroxide ions (OH-) eg. NaOH > Na + + OH - limitations - hydrogen ions NEVER exist alone; they attach to water to become a HYDRONIUM ION - H3O+} - cannot explain that some compounds act as bases even though they have no OH- 2. Modern Arrhenius theory - acknowledges the role of water ACID - any substance that dissociates to form H3O+ in aqueous solution eg. HCl + H2O > H3O + + Cl - BASE - any substance that dissociates to form OH- in aqueous solution eg. NH3 + H2O > NH4 + + OH - limitations - assumes that all acid-base reactions occur in water, but they do occur in other solutions - some compounds beginning with H are not always acids eg HPO4 2-

2 3. Bronsted - Lowry (1923) ACID - a substance that DONATES a proton (hydrogen ion) BASE - a substance that ACCEPTS a proton (hydrogen ion) - acid definition is basically the same, but bases do not rely on the presence of OH- - so, acid-base reactions involve the transfer of a proton (hydrogen ion) CONJUGATE BASE - the particle that remains when a H+ ion is removed from the acid CONJUGATE ACID - the particle that results when a base receives the proton from the acid CONJUGATE ACID-BASE PAIR - two particles related by the transfer of a H+ AMPHOTERIC - a substance that can act as both an acid and a base eg. Water examples: HCl + H2O < > H3O + + Cl -

3 Section 14.2 A Strong and Weak Acids and Bases p Acid and base strength are determined by two factors: dissociation (ionization) and concentration acid/base strength extent of dissociation extent of concentration strong weak dilute concentrated *** The terms strong and weak are ONLY used to describe dissociation (ionization) STRONG = dissociates completely into ions in water WEAK = dissociates slightly into ions in water eg vinegar is a weak, dilute acid There are three ways to classify acids: 1. MONOPROTIC - contain 1 ionizable hydrogen eg HCl 2. DIPROTIC - contain 2 ionizable hydrogens eg H2SO4 3. TRIPROTIC - contain 3 ionizable hydrogens eg H3PO4 **The acid that is formed by the first ionization is stronger than the acid that is formed by the second ionization, etc. Strong Acids Weak Acids Strong Bases Weak Bases H2SO4 H3BO3 LiOH NH3 HCl HC2H3O2 NaOH NaCN HNO3 H2CO3 KOH (many compounds that contain nitrogen) HClO4 RbOH HBr CsOH HI Ca(OH)2 Sr(OH)2 Ba(OH)2

4 Section 14.2 B ph Concept *** Always note that chemists use [H+] as a shorthand for [H3O+] {background: self-ionization of water; only 2 water molecules in 1 billion dissociate - that s why pure water is a poor conductor of electricity} Occasionally, the collisions between water molecules are energetic enough that a hydrogen is transferred from one water molecule to another in a process called the SELF IONIZATION OF WATER. + O O O + < > + O H - H H H H H H H water water hydronium ion hydroxide ion In pure water the [H+] and [OH-] must be equal. Each has a concentration of 1.0 x 10-7 M. This is considered NEUTRAL. In aqueous solution the [H+] and the [OH-] are interdependent. IE if the [H+] increases then the [OH-] decreases and vice versa. Kw = [H+] x [OH-] = 1.0 x M 2 where Kw = the ion-product constant for water In ACIDIC solutions the [H+] is greater than the [OH-] ie. [H+] > 1.0 x 10-7 M In BASIC or ALKALINE solutions the [OH-] is greater than the [H+] ie. [H+] < 1.0 x 10-7 M Expressing [H+] in M is difficult and somewhat confusing. A more widely used system is the ph scale. The ph of a solution is defined as the negative logarithm of the hydrogen ion concentration. ph = - log [H+] In pure water the [H+] = 1.0 x 10-7 M ph = - log [H+] ph = -log[1.0 x 10-7 ] ph = 7.0 Therefore a ph of 0 is highly acidic and a ph of 14 is highly basic. In a similar way the poh of a solution is defined as the negative logarithm of the hydroxide ion concentration. poh = - log [OH-] Therefore, a poh of 0 is highly basic and a poh of 14 is highly acidic (note: this is exactly backwards to ph)

5 Formulas: 1. ph + poh = [H+] x [OH-] = 1.0 x M 2 3. ph = - log [H+] 4. poh = - log [OH-] 5. [H+] = antilog (- ph) *** antilog means 10 x 6. [OH-] = antilog (- poh) ph Scale poh Scale 14 Very basic 14 Very acid 10.5 Mod basic 10.5 Mod acidic 8 Slightly basic 8 Slightly acidic 7 Neutral 7 Neutral 6 Slightly acidic 6 Slightly basic 3.5 Mod acidic 3.5 Mod basic 0 Very acidic 0 Very basic {practice problems and assignment}

6 Section 14.2 C Determining the ph of a Strong Acid/Base 1. What is the ph of M HCl? 2. What is the ph of M H2SO4? 3. What is the ph of M NaOH? 4. What is the ph of M Ca(OH)2? Section 14.2 D Effect of Dilution We only do effect of dilution problems for a strong acid/base. 1. If 230.-ml HCl with a ph of 2.59 is diluted to 2.30 L, what is the ph of the dilution? 2. If 125-ml of HNO3 with a ph of 1.56 is diluted to 2.50 L, what is the ph of the dilution? 3. If 422-ml of KOH with a ph of is diluted to 844 ml, what is the ph of the dilution? 4. If 300.-ml of HBr with a [H+] = 1.32 x 10-2 is diluted to 3.00 L, what is the ph of the dilution?