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Acids and Bases

Acids and Bases. Some Properties of Acids. Produce H + (H 3 O + ) ions in water. The Hydronium Ion (H 3 O + ) is an H + (proton) attached to a water molecule. Taste Sour. React with certain metals to produce H 2 gas and a salt.

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Acids and Bases

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  1. Acids and Bases

  2. Some Properties of Acids • Produce H+ (H3O+) ions in water. • The Hydronium Ion (H3O+) is an H+ (proton) attached to a water molecule. • Taste Sour. • React with certain metals to produce H2 gas and a salt. • Salt – ionic-metal or a positive polyatomic ion bonded with a negative ion other than OH- • Example: MgCl2, NH4Cl • Aqueous solutions of acids conduct electricity. • Electrolytes – the greater the concentration of ions in solution, the greater the electrical conductivity. • Strong Acids & weak Acids

  3. Some Properties of Acids • React with carbonates and bicarbonates to produce carbon dioxide gas. • HCl(aq) + NaHCO3(aq) NaCl(aq) + H2O(l) + CO2(g) • React with bases to form a salt and water. • Neutralization Reaction (Double Replacement) • HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l) • pH is less than 7 • pH scale expresses the amount of H+ as a number from 0 to 14 • pH of 0 is strongly acidic and has the highest amount of H+ ions, pH of 7 is neutral, pH of 14 is strongly basic and has the fewest H+ ions. • “ABC easy as 123” • Cause acid-base indicators to change color. • Acids turn Blue litmus Red

  4. Acid Nomenclature

  5. Acid Nomenclature • Use the flowchart to name the following acids. • HBrHydrobromic Acid • H2CO3 Carbonic Acid • H2SO3 Sulfurous Acid

  6. Going Backwards… • Write H first. • Write the 2nd ion. (you may have to check table E) • Assign charges. • Criss Cross, if necessary. • Examples • Sulfuric Acid _____ H2SO4__________ • Nitrous Acid _______HNO2 __________ • Oxalic Acid ______ H2C2O4_________

  7. Name ‘Em • HI(aq) • HCl(aq) • H2SO3 • HNO3 • HClO4

  8. Some Properties of Bases • Produce OH- ions in water. • Taste Bitter, chalky. • Aqueous solutions of bases conduct electricity. • Electrolytes – the greater the concentration of ions in solution, the greater the electrical conductivity. • Strong bases & weak bases • Feel soapy, slippery. • This is because they break down the normal body fat in your hands or whatever part of your body they come into contact with. • NaOH Before After

  9. Some Properties of Bases • React with acids to form a salt and water. • Neutralization Reaction (Double Replacement) • HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l) • pH is greater than 7 • pH scale expresses the amount of H+ as a number from 0 to 14 • pH of 0 is strongly acidic and has the fewest OH- ions, pH of 7 is neutral, pH of 14 is strongly basic and has the greatest amount of OH- ions. • “ABC is easy as 123” • Cause acid-base indicators to change color. • Bases turn Red litmus Blue

  10. Naming Bases • Name the Metal. • If the metal has only one possible charge, just write it’s name. • If the metal has more than one possible charge, use Roman Numerals to indicate the charge. • Follow with Hydroxide • Examples: • LiOHLithium Hydroxide • Fe(OH)3Iron(III) Hydroxide

  11. Going Backwards… • Write the symbol of the metal. • Write OH- • Assign Charges. • Criss Cross, if necessary • Examples • Cesium Hydroxide CsOH • Chromium(III) Hydroxide Cr(OH)3 • Strontium Hydroxide SrOH

  12. Practice Name each of the following… • HBr • H2SO3 • H2C2O4 • HClO • Ca(OH)2 • AgOH • HgOH • HF • HI • HClO4 • HCl • LiOH • Sn(OH)2 • Ti(OH)3

  13. Practice Now go backwards…. • nitric acid • carbonic acid • dichromic acid • acetic acid • nitrous acid • potassium hydroxide • cesium hydroxide • barium(II) hydroxide • aluminum(III) hydroxide • strontium(III) hydroxide

  14. Explaining Acids and Bases • There have been several attempts to explain the properties of acids and bases. • These explanations define how acids and bases behave. • There are three such definitions. • Arrhenius Theory • Brønsted – Lowry • Lewis Acids & Bases

  15. Arrhenius Theory • Acids – produce H+ ions (or hydronium ions H3O+) as the onlypositive ion. HCl(l) Cl- + H+ • A substance with a carboxyl group(COOH) looks like a base when you look at the chemical formula but it is an acid. (Acetic Acid = HC2H3O2 = CH3COOH) CH3COOH + H2O  CH3COO- + H+

  16. Arrhenius Theory • Bases – produce OH- ions (or hydroxide ions). • Some bases DO NOT have hydroxide ions attached. • Amines – organic compounds containing C and N. Amines are bases even though they do not have an hydroxide ion. Instead they react with water to produce the OH-ion. NH3 + H2O  NH4+ + OH- ~Caution~ • Alcohols – contain an –OH but ARE NOTbases. • Example: CH3OH (hydroxyl group on a carbon chain)

  17. Types of Acids and Bases • Acids • Monoprotic: produce one H+ ion. • HCl • Diprotic: produce two H+ ions. • H2SO4 • Triprotic: produce three H+ ions. • H3PO4 • Bases • Monohydroxy: produce one OH- ion. • NaOH • Dihydroxy: produce two OH- ions. • Ba(OH)2 • Trihydroxy: produce three OH- ions. • Al(OH)3

  18. Strength of Acids and Bases Determined by the amount of Ionization. • Strong Acids • 100% dissociation in water. • Great conductors of electricity. HNO3(aq) + H2O(l) H3O+(aq) + NO3-(aq) • HI, HCl, HBr, H2SO4, and HClO4 are strong acids. • Weak Acids • Much less than 100% dissociation. • Poor conductors of electricity. CH3COOH(aq) + H2O(l) CH3COO-(aq)+ H3O+(aq) • Acetic Acid(CH3COOH)

  19. Strength of Acids and Bases Determined by the amount of Ionization. • Strong Bases • 100% dissociation (ionization) in water. • Great conductors of electricity. NaOH(aq) Na+(aq) + OH-(aq) • KOH, Ca(OH)2, Group 1 or 2 metals with hydroxide!! • Weak Bases • Much less than 100% dissociation (ionization). • Poor conductors of electricity. NH3(aq) + H2O(l) NH4+(aq)+ OH-(aq) • Ammonia (NH3)

  20. Brønsted-Lowry Acids and Bases • Acids – Proton Donors • According to the Brønsted-Lowy concept, an acid is the chemical species that donates the proton in a proton transfer reaction. • Bases – Proton Acceptors • According to the Brønsted-Lowy concept, a base is the chemical species that accepts the proton in a proton transfer reaction. A “proton” is really just a hydrogen that has lost its electron…H+

  21. Conjugate Pairs • The Brønsted-Lowry concept defines a species as an acid or a base according to its function in the proton-transfer reaction. • Consider the Reaction of NH3 + H2O • In the forward reaction, NH3 accepts a proton donated by H2O. Thus, NH3 is a base and H2O is an acid.

  22. Conjugate Pairs • In the reverse reaction, NH4+ donates a proton to OH- which accepts it. Thus, NH4+ is acid and OH- is the base.

  23. Conjugate Pairs • The species NH4+ and NH3 are a conjugate acid-base pair. • A conjugate acid-base pair consists of two species in an acid-base reaction, one acid and one base, that differ by the loss or gain of a proton. • NH4+ is the conjugate acid of NH3 • NH3 is the conjugate base of NH4+ • The species OH- and H2O are a conjugate-acid base pair as well. • OH- is the conjugate base of H2O • H2Ois the conjugate acid of OH-

  24. Conjugate Pairs

  25. Conjugate Pairs…Practice Problems • Label the Acid, Base, Conjugate Acid, Conjugate Base in each reaction.

  26. Strength of Acid-Base Conjugate Pairs • Strong Acids (Proton Donors) have weak conjugate bases. • Strong Bases (Proton Acceptors) have weak conjugate acids. • Strong acids and strong bases are always on the same side of an equation. • An acid can donate it H+ to any base EXCEPT it’s conjugate base. • Example: H3PO4 can donate to F-, but not to PO43-

  27. Practice Problems • Write the conjugate base for each. • HCl ______________ • H2CrO4 ___________ • NH4+ _____________ • NH3 ______________ • Write the conjugate acid for each. • F- _________________ • H2PO4- ___________ • NH3 ______________ • HSO4- ____________

  28. Practice Problems • CH3COO- + H30+ CH3COOH + H2O • HCl + H2O  H3O+ + Cl- • NH2- + H2O  NH3 + OH- • H3O + OH-  H2O +H2O • CN- + H2O  HCN + OH- • HClO4 + CH3COOH  ClO4- + CH3COOH2+ • HCN + H2O  H3O+ + CN-

  29. Practice Problems • HSO4- + HCl H2SO4 +Cl- • SO42- + HNO3  HSO4- + NO3- • NH4+ + HSO4-  NH3 + H2SO4 • HCl + Al(H2O)5(OH)2+  Cl- + Al(H2O)63+ • NH3 + NH3  NH4+ + NH2-

  30. Amphoteric (Amphiprotic) Species • Substances that act as both an acid or a base. • Depends on chemical environment. • Examples: H2O, HSO4-, HS- • In the reaction between NH3 and H2O, water is an acid. • In the reaction between HNO2 and H2O, water is a base. • Water (H2O) is an amphoteric substance.

  31. Lewis Acids and Bases • Lewis Acid – a substance that ACCEPTS an electron(e-) pair. • Lewis Base – a substance that DONATES an electron(e-) pair. Formation of the Hydronium Ion is an excellent example.

  32. Lewis Acid/Base Reaction

  33. Reactions Involving Acids • Steps… • Check the metal on Table J. If it is above H2 proceed. • Write H2 as a product. • Combine the metal with the negative (-) ion to form an ionic salt. (write the metal first, followed by the negative ion.) • Assign charges • Criss Cross, if necessary • Balance the equation.

  34. Reactions Involving Acids, Examples… • HCl + Sr __________ + __________ • H3PO4 + Zn  __________ + __________ • HNO3 + Au  __________ + __________ • HC2H3O2 + K  __________ + __________ • HF + Cu  __________ + __________

  35. Neutralization Acid + Base  Salt + H2O • Steps… • Form water, H2O. • Get rid of all H+ on the acid, and all OH- on the base. • Write the metal 1st and the negative(-) ion 2nd. • Assign Charges. • CrissCross, if necessary. • Balance the equation.

  36. Neutralization…Examples • CH3COOH + NaOH __________ + __________ • KOH + HCl  __________ + __________ • HCl + NaOH  __________ + __________ • H2SO4 + NaOH  __________ + __________ • HCl + Ba(OH)2  __________ + __________ • HNO3 + LiOH  __________ + __________

  37. Spectator Ions • Ions found on both sides of the equation that are not involved in making water. • Not part of the Net Arrhenius Equation. • Spectator Ions for the previous example… • CH3COO-, Na+

  38. Net Arrhenius Equation • Does NOT include spectator ions!! • Net Arrhenius Equation is always… H+ + OH- H2O

  39. The pH Scale • The pH Scale expresses the strength of acids and bases. • Logarithmic Scale – one jump on the scale represents a tenfold change in [H+] • [ ] = concentration (usually Molarity) • pH > 7 is a Base • pH = 7 is Neutral • pH < 7 is an Acid • As pH ↑ [H+] ↓ • The stronger the acid the more H+ ions it produces. • The stronger the base the more OH- ions it produces.

  40. pH Examples • A solution with a pH of 1 has how many times the amount of H+ compared to a solution with a ph of 6? • A solution with a pH of 2 has how many times the amount of H+ compared to a solution with a pH of 5? • If the [H+] increases, the [OH-] decreases by the same amount. • If the pH changes from 8 to 13, the [H+] decreases _______ times and the [OH-] increases _______ times. • If the pH changes from 6 to 2, the [H+] increases _______ times and the [OH-] decreases _______ times.

  41. Calculating the pH of a substance pH = -log [H+] Recall that the [ ] mean concentration (usually Molarity) • Example: If [H+] = 1.0 x 10-10 what is the pH? pH = -log [H+] pH = -log (1.0 x 10-10) pH = 10 • Example: If [H+] = 1.8 x 10-5 what is the pH? pH = -log [H+] pH = -log (1.0 x 10-5) pH = 4.74

  42. Calculating the pH of a substance • Find the pH of these: • A 0.15M solution of HCl. • A 3.00 x 10-7M solution of HNO3

  43. More About water • H2O can function as both an ACID and a BASE. • Amphotericsubstance • In pure water there can be Autoionization. • Equilibrium Constant for water = Kw • Kw = [H3O+][OH-] = 1.00 x 10-14at 25oC • In a neutral solution [H3O+] = [OH-] • So… [H3O+] = [OH-] = 1.00 x 10-7

  44. pOH • Strong Acids and Strong Bases are opposites. • pH and pOH are opposites as well. • pOH scale does not really exist, but it is useful for determining the pH of a base. pOH = -log [OH-] • Since pH and pOH are on opposite ends, pH + pOH = 14

  45. [H30+], [OH-], pH and pOH…Problems • The pH of rainwater collected in a certain region of the northeastern United States on a particular day was 4.82; What is the H+ ion concentration of the rainwater? • The OH- ion concentration of a blood sample is 2.5 x10-7 M. What is the pH of the blood?

  46. General Summary of Formulas • pH = -log [H+] • pOH = -log [OH-] • [H+][OH-] = 1.0 x10-14 • [H+] = 10-pH • [OH-] = 10-pH • pH + pOH = 14

  47. Calculating [H30+], [OH-], pH and pOH Problem 1: A chemist dilutes concentrated hydrochloric acid to make two solutions: (a) 3.0M and (b) 0.0024M. Calculate the [H3O+], pH, [OH-] and pOH of the two solutions at 25oC.

  48. Calculating [H30+], [OH-], pH and pOH Problem 2: What is the [H30+], [OH-] and pOH of a solution with pH = 3.67? Is this an acid, base, or neutral?

  49. Calculating [H30+], [OH-], pH and pOH Problem 3: What is the [H30+], [OH-] and pOH of a solution with pH = 8.05? Is this an acid, base, or neutral?

  50. Using Table M

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