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Acids & Bases

Acids & Bases. Properties of Acids. Sour taste Change color of acid-base indicators (red in pH paper) Some react with active metals to produce hydrogen gas Ba (s) + H 2 SO 4(aq) BaSO 4(s) + H 2(g) Some react with bases to neutralize and form salt and water

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Acids & Bases

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  1. Acids & Bases

  2. Properties of Acids • Sour taste • Change color of acid-base indicators (red in pH paper) • Some react with active metals to produce hydrogen gas Ba(s) + H2SO4(aq) BaSO4(s) + H2(g) • Some react with bases to neutralize and form salt and water H2SO4 (aq) + 2NaOH(aq) Na2SO4 (aq) + 2H2O(l) • Some are electrolytes

  3. Examples of Acids • Lemons and oranges - citric acid • Vinegar - 5% by mass acetic acid • Pop and fertilizer - phosphoric acid

  4. Properties of Bases • Bitter taste • Change color of acid-base indicators (blue in pH paper) • Dilute aqueous solutions feel slippery Ex. Soap • Some react with acids to neutralize and form salt and water • Some are electrolytes

  5. Examples of Bases • Soap - NaOH • Household cleaners - NH3 • Antacids - Ca(OH)2, Mg(OH)2

  6. Arrhenius Acids • Acids that increase the concentration of hydronium (H3O+) in aqueous solutions HNO3(aq) + H2O(l) H3O+(aq) + NO3-(aq) H+ + NO3- + H2O acid

  7. Why do acids produce H3O+? • H+ is extremely attracted to the unshared pair of electrons on the water molecule so it donates itself to this molecule where it becomes covalently bonded. The ion formed is known as the hydronium ion (H3O+) H+

  8. Arrenius Bases • Bases that increase the concentration of hydroxide ions (OH-) in aqueous solutions NaOH(s) Na+(aq) + OH-(aq) H2O

  9. Strength of Acids & Bases • Strong acids & bases completely ionize in aqueous solutions H2SO4 + H2O H3O+ + HSO4- NaOH Na+ + OH- • Strong acids & bases are strong electrolytes • A list of strong acids & bases can be found on pg. 460-461

  10. Weak acids & bases only partially break down into ions when in aqueous solutions HCN + H2O H3O+ + CN- NH3 + H2O NH4+ + OH- • Weak acids & bases are weak electrolytes • A list of weak acids & bases can be found on pg. 460-461

  11. Why can we drink H2O? • Water self ionizes to form equal concentrations of H3O+ and OH- H2O(l) + H2O(l) H3O+(aq) + OH-(aq) • A substance is considered “neutral” when [H3O+] = [OH-] • [H3O+] concentration = 1.0 x 10-7M • [OH-] concentration = 1.0 x 10-7 M

  12. When [H3O+] = [OH-] • If [H3O+] > 1.0 x 10-7 M, the solution is acidic • If [OH-] > 1.0 x 10-7 M, the solution is basic • To find the concentration of [H3O+] or [OH-] in acidic or basic solutions, the following equation can be used: 1.0 x 10-14 M2 = [H3O+] [OH-] 1.0 x 10-14 M2 = ionization constant for H2O (Kw)

  13. Sample Problem • A 1.0 x 10-4 M solution on HNO3 has been prepared for laboratory use. a. Calculate the [H3O+] of this solution b. Calculate the [OH-] of this solution c. Is this solution acidic or basic? Why? d. Substitute H2SO4 as the acid. How would the calculations change?

  14. Sample Problem • An aqueous 3.8 x 10-3 M NaOH solution has been prepared for laboratory use. a. Calculate the [H3O+] of this solution b. Calculate the [OH-] of this solution c. Is this solution acidic or basic? Why? d. Substitute Ca(OH)2 as the base. How would the calculations change?

  15. Practice Problems • Complete practice problems on pg. 484 #1-4

  16. The pH scale • The pH scale measures the power of the hydronium ion [H3O+] in a solution • The scale typically goes from 1-14 (although it can extend below or above it under extreme conditions) • The following equations can be used to determine the pH or [H3O+] of a solution: pH = -log [H3O+] [H3O+] = antilog (-pH) [H3O+] = 1 x 10-pH

  17. pH > 7 basic pH = 7 neutral pH < 7 acidic

  18. The pOH scale • The pOH scale measures the power of the hydroxide ion [OH-] in a solution • The scale typically goes from 1-14 (although it can extend below or above it under extreme conditions) • The following equations can be used to determine the pOH or [OH-] of a solution: pOH = -log [OH-] [OH-] = antilog (-pOH) [OH-] = 1 x 10-pOH

  19. pH + pOH = 14

  20. Sample Problems • Calculate the pH of each of the following. Classify as acidic or basic. • 1.3 x 10-5 M NaOH • 1.0 x 10-4 M HCl

  21. Sample Problems • What is the [H3O+] for each of the following? Classify as acidic or basic. • pH = 5.8 b. pOH = 8.9

  22. Sample Problems • What is the [OH-] for each of the following? Classify as acidic or basic. • [H3O+] = 9.5 x 10-10 M • pOH = 1.3

  23. Practice Problems • Complete practice problems on pg. 487 #1 pg. 488 #1-4 pg. 490 #1-4

  24. Expansion of the Acid-Base Theory • Substances can still act as an acid or base if they are not dissolved in water to make a solution

  25. Bronsted-Lowry Acids • A molecule or ion that is a proton (H+) donor HCl(g) + NH3(g) NH4(g)+ + Cl-(g) H+ donor

  26. Bronsted-Lowry Bases • A molecule or ion that is a proton (H+) acceptor HCl(g) + NH3(g) NH4+(g) + Cl-(g) • In a Bronsted-Lowry acid-base reaction, protons (H+) are transferred from one reactant (the acid) another (the base) H+ acceptor

  27. Monoprotic versus Polyprotic Acids • Monoprotic acids can only donate 1 proton per molecule HCl(g) + H2O(l) H3O+(aq) + Cl-(aq) Monoprotic

  28. Polyprotic acids can donate more than one proton per molecule H2SO4(aq) + H2O(l) H3O+(aq) + HSO4-(aq) Polyprotic HSO4-(aq) + H2O(l) H3O+(aq) + SO4-2(aq) One additional proton can still be donated

  29. Conjugate acids & bases • A conjugate acid is the species that is formed when a Bronsted-Lowry base gains a proton • A conjugate base is the species that remains after a Bronsted-Lowry acid has given up a proton HF(aq) + H2O(l) F-(aq) + H3O+(aq) acid base Conjugate acid Conjugate base

  30. More examples CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO-(aq) HCl(aq) + H2O(l) H3O+(aq) + Cl-(aq) CB CA acid base acid base CA CB Proton transfer reactions favor the production of the weaker acid and base. Use table 15-6 on pg. 471 in your text to compare the relative strengths of acids and bases

  31. Is H2O an acid or a base? • H2O is amphoteric, it can react as either an acid or a base • If H2O reacts with a compound that is a stronger acid than itself, it acts as a base • If H2O reacts with a weaker acid, it will act as the acid H2SO4(aq) + H2O(l) H3O+(aq) + HSO4-(aq) Base H+ acceptor NH3(aq) + H2O(l) NH4+(aq) + OH-(aq) Acid H+ donor

  32. OH- in a molecule • When an OH- group is covalently bonded in a molecule, it is referred to as a hydroxyl group • Hydroxyl groups are present in many organic compounds Ex. Acetic acid (HC2H3O2) or CH3COOH Hydroxyl group

  33. How does the OH- make something acidic? • In order for a compound with an OH- group to be acidic, H2O must be able to attract the H atom from the OH- group and act as a proton donor CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO-(aq) The more O atoms bonded to the OH- group, the more acidic the compound is likely to be. Oxygen is highly electronegative and will attract electrons closer to it, making the OH- bond more polar. This will allow H2O to “steal” the H atoms more easily.

  34. Why are substances with OH- covalently bonded to it sometimes not acidic? • Ex. Acetic acid (CH3COOH) versus ethanol (C2H5OH) Ethanol Acetic acid Acetic acid- the 2 O atom on the C atom draws electron density away from the OH- group, making the bond more polar. This allows the H+ to be donated more easily Ethanol- this compound is essentially neutral. It does not have a second O atom to make the bond as polar. It would be classified as a very weak acid because it is harder to donate H+.

  35. Further expansion of acid-base theory • Substances can still act like an acid or base if they do not contain hydrogen at all

  36. Lewis acids & bases • A Lewis acid is an atom, ion, or molecule that accepts an electron pair to form a covalent bond Ag+(aq) + 2NH3(aq) [H3N-- Ag--NH3]+ • A Lewis base is an atom, ion, or molecule that donates an electron pair to form a covalent bond e- pair acceptor e- pair donator

  37. Sample Lewis acid-base problem • For the following equation, which reactant is the Lewis acid? Lewis base? BF3(aq)+ F-(aq) BF4-(aq)

  38. BF3 is the Lewis acid because it is the e- pair acceptor • F- is the Lewis base because it is the e- pair donor

  39. Review of acid-base categorization

  40. Strong Acid-Base Neutralization • When equal parts of acid and base are present, neutralization occurs where a salt and water are formed HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

  41. Sample Problems • H2CO3 + Sr(OH)2 • HClO4 + NaOH • HBr + Ba(OH)2 • NaHCO3 + H2SO4

  42. Titrations • When you have a solution with an unknown concentration, you can find it by reacting it completely with a solution of known concentration • This process is known as titrating • To perform a titration, an instrument called a buret can be used to precisely measure amounts of solution, drop by drop

  43. Titration Termonology • Equivalence point - the point at which the known and unknown concentration solutions are present in chemically equivalent amounts moles of acid = moles of base Indicator - a weak acid or base that is added to the solution with the unknown concentration before a titration so that it will change color or “indicate” when in a certain pH range (table 16-6 on pg. 495 in your text will show various indicators and their color ranges)

  44. End point - the point during a titration where an indicator changes color • The 2 most common indicators we will use in our chemistry class will be: • Phenolphthalein - turns very pale pink at a pH of 8-10 • Bromothymol blue - turns pale green at a pH of 6.2-7.6 Phenolpthalein is clear at pH<8, pale pink at pH 8-10 and magenta at pH >10 Bromothymol blue

  45. Practice Titration for an unknown acid • 1. Titrate 5.0 of mL of unknown HCl into a 250 mL erlenmeyer flask - *remember to document the starting amount and ending amount of acid on the buret to prevent error • 2. Add 2 drops of indicator (phenolphthalein) to the flask - the color of the solution should be clear • 3. Titrate with .5M NaOH, continuously swirling the flask, until the solution turns very pale pink for 30 seconds - *remember to document the starting amount and ending amount of base on the buret • 4. Mathematically determine the concentration of the unknown HCl solution by using the following equation:

  46. Titration Equation MAVA = MBVB MA = molarity (mol/L) of acid VA = volume in L of acid MB = molarity (mol/L) of base VB = volume in L of base molesA = molesB 5. After calculating the molarity of the unknown acid experimentally, get the theoretical molarity and calculate % error

  47. Practice titration for an unknown base • 1. Titrate 5.0 of mL of unknown NaOH into a 250 mL erlenmeyer flask - *remember to document the starting amount and ending amount of base on the buret to prevent error • 2. Add 2 drops of indicator (phenolphthalein) to the flask - the color of the solution should be magenta • 3. Titrate with .5M HCl, continuously swirling the flask, until the solution turns very pale pink for 30 seconds - *remember to document the starting amount and ending amount of acid on the buret • 4. Mathematically determine the concentration of the unknown NaOH solution by using MAVA = MBVB • 5. After calculating the molarity of the unknown base experimentally, get the theoretical molarity and calculate % error

  48. How do pH indicators work? • Acid-base indicators are usually weak acids or bases that are in equilibrium and show color changes when a stress is applied HIn H+ + In- In acidic solutions, the H+ concentration increases. The stress will cause a shift to the left (red color). In basic solutions, the OH- concentration increases. These ions will combine with H+ which will cause a shift to the right (blue color) red blue

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