Acid and Base Chemistry

Transcription

1 Acid and Base Chemistry

2 Arrhenius Theory of Acids An acid is a substance that gives off H+ in water. This is a hydronium ion or proton. H2SO4 is sulfuric acid. H2SO4 2H+ + SO4 2- has one proton and one electron. H+ has lost the electron and so we call it a proton. H+ + H2O H3O+ Hydronium Ion H2SO4 + 2H2O 2H3O+ + SO4 2-

3 Arrhenius Bases According to the Arrhenius Theory, a base is a substance that dissolves in water to make OH-. NaOH(s) Na+(aq) + OH-(aq) Ca(OH)2(s) Ca+2(aq) + 2OH-(aq)

4 Overall Arrhenius Theory Acid + Base Salt + Water + Heat Salt is an ionic compound that results from the reaction between an acid and a base. A solution is acidic if there are more moles of H+ than OH-. A solution is basic if there are more moles of OH - than H+. A solution is neutral if the number of moles of H+ and OH- are the same. They then neutralize, or cancel out, each other.

5 Examples of Neutralization Equations HCl + NaOH NaCl + H2O 2HBr + Ca(OH)2 CaBr2 + 2H2O HNO3 + KOH KNO3 + H2O 3H2SO4 + 2Al(OH)3 Al2(SO4)3 + 6H2O

6 Acid Nomenclature Binary Acids = H and one nonmetal element. Name always begins hydro-, root is based on the nonmetal name with the suffix -ic, followed by the word acid. HCl = Hydrochloric Acid HBr = Hydrobromic Acid H2Se = Hydroselenic Acid

7 Oxyacids These are ternary acids, and contain 3 elements: H, O and a nonmetal. Example: H2SO4 Ternary or oxyacids do not have hydro as a prefix. Name starts with root. The suffix depends on the anion from which the acid is derived. -ate anion = -ic acid. -ite anion = -ous acid.

8 Examples HClO3 = H+ + ClO3 - (Chlorate ion) Chloric Acid HNO2 = Nitrous Acid H2CO3 = Carbonic Acid

9 Titration Titration is a process in which the amount of acid and base are carefully measured to exactly produce a neutral solution. Technically speaking, the substance being added (either acid or base) is slowly added until the solution is neutral, and the volume that is added is carefully measured and recorded.

10 Indicator Substance that changes color as the relative concentration of acid and base changes. The best indicator will change color from acid to neutral or from base to neutral. Example--phenolthalein. It is clear in an acidic solution or neutral solution, and pink in a basic solution.

11 Math on Titration Problems In a titration problem, the moles of H+ and the moles of OH- must be equal. Molarity Volume = moles, so MaVa = MbVb Ma = molarity of H+ Va = volume of acid Mb = molarity of OH- Vb = volume of base

12 Practice Problem 23.1 ml of a.100 M HCl solution is titrated with 15.7 ml of a NaOH base solution. Find the concentration of the base. MaVa = MbVb (.100 M) (23.1mL)=Mb (15.7mL) Mb=.147 M

13 Logarithms Logarithm Defined X is the logarithm (or log) of A if A can be written as 10X. A = 10X X = log A What is the log of 1,000,000? Can 1,000,000 be written as a power of 10? Yes, 1,000,000 = 10e6. So log 1,000,000 = 6.

14 ph Definition: ph is the - log of the [H+] ph = - log [H+] ph is a measure of the acidity or basicity of a solution. For a neutral solution, ph = 7. For an acidic solution, ph < 7. For a basic solution, ph > 7.

15 ph and Logarithms 104 = 10,000 4 is the logarithm of 10,000 10e3 = 1,000 3 is the logarithm of 1,000 10e2 = is the logarithm of e1 = 10 1 is the logarithm of 10 10e0 = 1 0 is the logarithm of 1 10e-1 =.1-1 is the logarithm of.1 10e-2 =.01 10e-3 =.001

16 What is the log of 4,258? 4,258 = 10X What is X? Use calc.---- log(4258)= e3.63 = 4,258

17 Other important equations ph = - log [H+] poh = - log [OH-] Kw = [H+][OH-] = 1 x ph + poh = 14 [H+] = 10e-pH [OH-] = 10e-pOH

18 Dissociation Equilibrium for Water H2O(l) H+(aq) + OH-(aq) Kw = [H+][OH-] = 1 x So as water dissociates, [H+]=[OH-]. Then, [H+]2 = 1 x 10-14, and [H+] = [OH-] = 1 x This is true for a neutral solution

19 Practice Problem A solution is.20 M in HCl. Find the ph, poh and the [OH-]. HCl H+ + Cl-.20 M HCl =.20 M H+ =.20 M Cl- HCl H+ + Cl- ph = -log [.20 M] =.70 [OH-] = 1x10e-14/.20 = 5.0x10e-14 poh = -log [5.0x10e-14] = or ph + poh = 14 poh= = 13.30

20 A solution has a ph of Find the concentration of H+ and OH-. Also find the poh. [H+] = 10e-pH = 10e-3.20 or antilog of [H+] = 6.31 x 10e-4 ph + poh = 14 poh = [OH-] = 10e-pOH or [H+][OH-]=1x10e-14 [OH-] = 1.58 x 10-11

21 Write an equation for dissolving and dissociating the following acids in water: HCl, HNO3, H2SO4, H3PO4 HCl H+(aq) + Cl-(aq) or HCl + H2O H3O+(aq) + Cl-(aq) HNO3 H+(aq) + NO3-(aq) or HNO3 + H2O H3O+(aq) + NO3-(aq) H2SO4 2H+(aq) + SO4 2-(aq) orh3po4 3H+(aq) + PO4 3-(aq)

22 Mono-, Di- and Triprotic Acids A monoprotic acid gives off one H+ for each acid molecule. HCl and HNO3 are examples of monoprotic acids. A diprotic acid gives off two H+ ions for each acid molecule. H2SO4 is an example of a diprotic acid. A triprotic acid gives off three H+ ions for each acid molecule. H3PO4 is an example of a triprotic acid.

23 Concentration Calculations with Di- and Triprotic Acids When you have di- and triprotic acids, the concentration of H+ is greater than the concentration of the original acid. Find the molarity H+ and the ph of a.200 M H2SO4 solution. H2SO4 2H+(aq) + SO4 2-(aq) (.200 moles H2SO4) (2 moles H+) =.400 M H+ ( 1 liter soln.) (1 mole H2SO4) ph = - log (.400) =.398

24 Dibasic and Tribasic Substances Bases that have more than one OH- per molecule follow a similar pattern to the acids. Ca(OH)2 Ca OH- Al(OH)3 Al3+ + 3OH-

25 Bronsted-Lowry Theory of Acids and Bases A Bronsted-Lowry Acid is a substance the donates a proton or H+ ion to another molecule or ion. A Bronsted-Lowry Base is a substance that accepts a proton or H+ ion from another molecule or ion. An Acid-Base reaction must include both a proton donor and a proton acceptor.

26 Examples of Bronsted-Lowry Acid Base Reactions Acetic Acid reacts with Ammonia to produce Ammonium Acetate. CH3CO2H + NH3 NH4CH3COo CH3CO2H + NH3 NH4+ + CH3COo -

27 Lewis Acid and Base Lewis Acid-an atom, ion, or molecule that accepts an electron pair to form a covalent bond H+ + NH3 -> NH4+ Lewis Base-an atom, ion, or molecule that donates an electron pair to form a covalent bond BF3 + F- -> BF4+ Lewis Acid-Base Reaction-the formation of one or more covalent bonds between an electron-pair donor and an electron-pair acceptor

28 Do page 482

29 Conjugate Acids and Bases A conjugate acid is an acid (as a product) that results from a Bronsted acid-base reaction. A conjugate base is a base (as a product) that results from a Bronsted acid-base reaction. A Bronsted Base accepts a proton to produce its conjugate acid. A Bronsted Acid donates a proton to produce its conjugate base.

30 Example CH3CO2H + NH3 NH4+ + CH3COo- The Bronsted acid is CH3COoH. The Bronsted base is NH3. The conjugate acid is NH4+. The conjugate base is CH3COo-.

31 Weak Acids and Bases Weak Acids do not fully dissociate when dissolved in water. This means that the concentration of the H+ ion is less than expected due to the stoichiometry. Examples of Weak Acids: Acetic Acid (CH3CO2H) Phosphoric Acid (H3PO4) Weak Bases do not fully accept one H+ per molecule as expected.

32 Strong and Weak Acids and Bases Strong Acids fully dissociate in water, and donate all possible H+ ions readily. Strong Bases fully accept all donated H+ ions readily. Weak Acids do not readily donate all possible H+ ions. Weak Bases do not readily accept all possible H+ ions.

33 Buffers A buffer is a solution that resists changes in the ph. A buffer is made from a weak acid and its conjugate base or a weak base and its conjugate acid. Since it is made of both a weak acid and a weak base, there is both acid and base in solution to react with any strong acid or base that is added to it. This minimizes the effect of adding acid or base.

34 A.001 M acetic acid solution is 5% dissociated. What is the H+ concentration? % Dis. = M of dissociated acid(100) M of acid total CH3CO2H H+ + CH3CO2- % Dis. = [CH3CO2-](100) [CH3CO2H]

35 Ionization Constant HA H+ + A- Ka= [H+] [A-] [HA] Ka = acid dissociation constant

36 A- + H2O HA + OH Kb= [HA] [OH-] [A-] Kb = base dissociation constant