Presentation on theme: "Acids and Bases. Properties of Acids Sour taste React w/ metals to form H 2 Most contain hydrogen Are electrolytes Change color in the presence."— Presentation transcript:
Properties of Acids Sour taste React w/ metals to form H 2 Most contain hydrogen Are electrolytes Change color in the presence of indicators (turns litmus red) Has a pH lower than 7
Two Types of Acids Strong acids –Any acid that dissociates completely in aqueous sol’n Weak acids –Any acid that partially dissociates in aqueous sol’n
Properties of Bases Bitter taste Slippery feel Are electrolytes Change color in the presence of indicators (turns litmus blue) Has a pH higher than 7
Types of Bases Strong Base –Any base that dissociates completely in aqueous sol’n Weak Base –Any base that partially dissociates in aqueous sol’n
Neutralization Neutralization rxn: a rxn of an acid and a base in aqueous sol’n to produce a salt and water Salt: compound formed from the positive ion of a base and a negative ion of an acid Properties of the acid and base cancel each other
Arrhenius Model of Acids and Bases Proposed the model in 1887 Acid: any compound that produces H + ions in aqueous (water) sol’n Base: any compound that produces OH - (hydroxide) ion in aqueous sol’n Offers an explanation of why acids and bases neutralize each other (H + + OH - = H 2 O)
Problems with Model Restricts acids and bases to water sol’ns (similar reactions occur in the gas phase) Does not include certain compounds that have characteristics of bases (e.g., ammonia)
Brønsted-Lowry Model of Acids and Bases Brønsted acid: a hydrogen ion donor (H +, or proton) Brønsted base: a hydrogen ion acceptor Defines acids and bases independently of how they behave in water Amphiprotic: having the property of behaving as an acid and a base –Also called amphoteric, e.g., water
Conjugate Acid-Base Pairs The rxn between Brønsted-Lowry acids and bases can proceed in the reverse direction (reversible reactions) HX (aq) + H 2 O (l) H 3 O + (aq) + X - (aq) The water molecule becomes a hydronium ion (H 3 O + ), and is an acid because it has an extra H + to donate The acid HX, after donating the H +, becomes a base X -
Conjugate Acids and Bases HX (aq) + H 2 O (l) H 3 O + (aq) + X - (aq) Acid BaseConjugate Acid Conjugate Base Forward reaction: Acid and base Reverse reaction: Conjugate acid and conjugate base
Conjugate Acid: species produced when a base accepts a hydrogen ion from an acid Conjugate Base: species produced when an acid donates a hydrogen ion to a base Conjugate Acid-Base Pair: two substances related to each other by the donating and accepting of a single hydrogen ion
Types of Acids Monoprotic acids: acids that contain only 1 hydrogen; e.g., HCl Diprotic acids: acids that contain 2 hydrogens; e.g. H 2 CO 3 Triprotic acids: acids that contain 3 hydrogens; e.g. H 3 PO 4
More Types of Acids Binary acids: acids that contain only 2 elements; e.g. HF Polyatomic acids: acids that contain more than 2 elements; e.g. H 2 SO 4 – These acids contain polyatomic ions – Also called ternary or oxy- acids
Naming Binary Acids Start with the prefix hydro- Put it in front of the root word of the anion (- charged ion) Add –ic to the end Examples – Hydrobromic (HBr) – Hydrofluoric (HF) – Hydroiodic (HI) – Hydrochloric (HCl)
Naming Polyatomic Acids Start with the root word of the name of the polyatomic ion Add –ous if name ends in –ite Add -ic if name ends in –ate Examples: – Chlorous (from chlorite, ClO 2 - ) – Nitric (from nitrate, NO 3 - ) – Sulfurous (from sulfite, SO 3 -2 )
pH and [H 3 O + ] pH: number that is derived from the concentration of hydronium ions ([H 3 O + ]) in sol’n – pH = -log [H 3 O + ] – As pH increases, [H 3 O + ] decreases Scale ranges from 0 – 14 – pH = 7 is neutral – pH < 7 is acidic – pH > 7 is basic
Dissociation Constants Acid dissociation constant: (K a ): the equilibrium constant for the rxn of an aqueous weak acid and water Base dissociation constant: (K b ): the equilibrium constant for the rxn of an aqueous weak base w/ water Both are derived from the ratio of the concentration of the products and reactants at equilibrium
Acid Dissociation Constant K a = [H 3 O + ] [A - ] [HA] K a is a measure of the strength of an acid K a values for weak acids are always less than one Used mostly w/ weak acids because the K a values for strong acids approach infinity
Examples HMnO 4 (aq) + H 2 O (l) H 2 S (aq) + H 2 O (l)
Base Dissociation Constant K b = [HB + ] [OH - ] [B] K b is a measure of the strength of a base K b values for weak bases are always less than 1 K b values for strong bases approach infinity
Examples H 2 NOH (aq) + H 2 O (l) NH 3 (aq) + H 2 O (l)
Water Water can dissociate into its component ions, H + and OH - – 2H 2 O (l) H 3 O + (aq) + OH - (aq) One water molecule acts as a weak acid, and the other acts as a weak base The ions are present in such small amounts they can’t be detected by a conductivity apparatus In pure water, [H 3 O + ] =1.0 x 10 –7 M and [OH - ] = 1.0 x 10 -7 M
Dissociation Constant for Water It is defined as K w : the ion product constant for water K w = [H 3 O + ] [OH - ] K w = (1.0 x 10 -7 )(1.0 x 10 -7 ) K w = 1.0 x 10 -14 The value of K w can always be used to find the concentration of either H 3 O + or OH - given the concentration of the other
Examples What is the pH of a 0.001 M sol’n of HCl, a strong acid?
Examples What is the pH of a sol’n if [H 3 O + ] = 3.4 x 10 -5 M?
Examples The pH of a sol’n is measured with a pH meter and determined to be 9.00. What is the [H 3 O + ]?
Examples The pH o f a sol’n is measured with a pH meter and determined to be 7.52. What is [H 3 O + ]?
Calculating K a In these problems, remember that the concentration of the [H 3 O + ] ions will equal the concentration of the conjugate base ions. –This is because for every molecule of weak acid that dissociates, there will be an equal number of H 3 O + ions and base ions
Example Assume that enough lactic acid is dissolved in sour milk to give a solution concentration of 0.100 M lactic acid. A pH meter shows that the pH of the sour milk is 2.43. Calculate K a for the lactic acid equilibrium system.
Titrations An analytical procedure used to determine the concentration of a sample by reacting it with a standard sol’n In a titration, an indicator is used to determine the end point Standard sol’n: a sol’n of precisely known concentration Indicator: any substance in sol’n that changes color as it reacts with either an acid or a base
Titrations Each indicator changes its color over a particular range of pH values (transition interval) An unknown acid sol’n will be titrated with a standard sol’n that is a strong base An unknown base sol’n will be titrated with a standard sol’n that is a strong acid
Titrations Equivalence point: point at which the concentration of H 3 O + ions is the same as the concentration of OH - ions; [H 3 O + ] = [OH - ] Endpoint: the point at which the indicator changes color Titration curve: graph that shows how pH changes in a titration
Titrations The equivalence point is at the center of the steep, vertical region of the titration curve At the equivalence point, pH increases greatly w/ only a few drops