Acids, Bases, and Buffers I PDF

PreMed1 Chemistry for Medicine Acids, Bases, and Buffers I Dr. Alya A. Arabi College of Medicine and Health Sciences UAEU Sections to be covered: • Acids and Bases: Arrhenius definition • Brnsted-Lowry Acids and Bases • The Auto-ionization of Water Acids and Bases: A Brief Review Acids: taste...

PreMed1 Chemistry for Medicine Acids, Bases, and Buffers I Dr. Alya A. Arabi College of Medicine and Health Sciences UAEU Sections to be covered: • Acids and Bases: Arrhenius definition • Brnsted-Lowry Acids and Bases • The Auto-ionization of Water Acids and Bases: A Brief Review Acids: taste sour and cause dyes to change color. Example: The purple liquid of cabbage will change to a bright pink in acidic environments (and to blue under alkaline conditions.) Bases: taste bitter and feel soapy. 1. Arrhenius definitions of acids and bases: [H+] H2O Acids are donors: HCl(g) Bases are [OH-] donors: NaOH(s) H+(aq) + Cl-(aq) H2O Na+(aq) + OH-(aq) H+(aq) and H3O+(aq) are interchangeably used. H3O+(aq) is H+ clustered with H2O Brønsted-Lowry Acids and Bases 2. Brønsted-Lowry definitions of acids and bases: acid donates H+ and base accepts H+. i.e. Brønsted-Lowry base does not need to contain OH-. Example: HCl(aq) + H2O(l) → H3O+(aq) + Cl-(aq) HCl donates a proton to water. Therefore, HCl is an acid. H2O accepts a proton from HCl. Therefore, H2O is a base. Brønsted-Lowry Acids and Bases Example: HCl + NH3 → NH4+ + Cl- HCl donates a proton to NH3. Therefore, HCl is an acid. NH3 accepts a proton from HCl. Therefore, NH3 is a base. Example: NH3 + H2O → NH4+ + OH- H2O donates a proton to NH3. Therefore, H2O is an acid. NH3 accepts a proton from H2O. Therefore, NH3 is a base. Amphoteric substances can behave as acids and bases, e.g. water is amphoteric. H2O acting as base: HCl + H2O → H3O+ + Cl- H2O acting as acid: NH3 + H2O → NH4+ + OH- Conjugate Acid-Base Pairs A conjugate base is what is left of the acid after the proton is donated. A conjugate acid is what is what the base becomes after it accepts a proton. Example: HA(aq) + H2O(l) H3O+(aq) + A-(aq) 1. After HA (acid) loses its proton it is converted into A- (base). A- is the conjugate base of the HA acid. 2. After H2O (base) gains a proton it is converted into H3O+ (acid). H3O+ is the conjugate acid of the base H2O. Conjugate acid-base pairs differ by only one proton. Practice: Label the acid, base, conjugate acid and conjugate base in each of the following equations HNO2 (aq) + H2O (l) NH3 (aq) + H2O (l) NO2- (aq)+ H3O+ (aq) NH4+ (aq)+ OH- (aq) Relative Strengths of Acids and Bases Strength of an acid is measured by its ability to give a proton. Strong Acids completely transfer their protons to water, leaving no undissociated molecules in solution. Their conjugate bases are weak, they have a weak tendency to accept the proton (or to reverse the reaction). The stronger the acid, the weaker the conjugate base. Irreversible, 100% dissociation Example: HCl H+ + ClCl- is the conjugate base of HCl: HCl is a strong acid, so Cl- is a weak conjugate base Strong acids are: Relative Strengths of Acids and Bases 2. Weak Acids dissociate only partially, they exist in the solution as a mixture of acid molecules (HA) and their conjugate ions (A-). Their conjugate bases have a slight ability to accept protons (from water). Example: H2CO3 H+ + HCO3- Reversible, only partial dissociation (i.e. not all 100% protons dissociate) https://www.slideserve.com/hector/weak-acids-weak-bases pH = -log [H+] Henderson-Hasselbalch equation https://wps.prenhall.com/wps/media/objects/602/616516/Media_Assets/Chapter15/Text_Images/FG15_01.JPG The Auto(self)ionization of Water The Ion Product of Water In pure water the following equilibrium is established at 25 C H H H O + : : 2 :O : H3O+(aq) + OH-(aq) : H2O(l) + H2O(l) H + :O H- H The above is called the auto-ionization of water. Two water molecules react together, one water molecule acts as an acid, and the other acts as a base. The Autoionization of Water The Ion Product of Water H2O(l) + H2O(l) H3O+(aq) + OH-(aq) 𝐾𝑤 = H + OH − = 1.0 × 10−14 pKw=− log H + OH − = − log(1.0 × 10−14 ) − log[ H + ] − log[ OH − ] = 14 pH + pOH = 14 Therefore, knowing any of the concentrations can simply lead to the knowledge of the other using this law. If [H+]> [OH-], then the solution is acidic If [H+] = [OH-], then the solution is neutral If [H+] < [OH-], then the solution is basic Practice Example: Calculate the concentration of OH - (aq) in a solution in which: a)[H+] = 2 x 10-6 M b) [H+] = [OH-] c) [H+] = 100 [OH-] 𝐾𝑤 = H + OH − = 1.0 × 10−14 Practice Example: Indicate whether a solution with each of the following ion concentrations is neutral, acidic, or basic: a) [H+] = 4 x 10-9 M b) [OH-] = 1 x 10-7 M c) [OH-] = 7 x 10-13 M