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Unit 10: Acids and Bases

Unit 10: Acids and Bases. Flowers – the sign of love, affection, and the acidity of the soil!. Common Occurrences . We encounter acids and bases daily Acids Foods and Drinks Bases Household products (cleaners) The hydrangeas on the previous slide Blue – produced by a flower in acidic soil

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Unit 10: Acids and Bases

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  1. Unit 10: Acids and Bases Flowers – the sign of love, affection, and the acidity of the soil!

  2. Common Occurrences • We encounter acids and bases daily • Acids • Foods and Drinks • Bases • Household products (cleaners) • The hydrangeas on the previous slide • Blue – produced by a flower in acidic soil • Pink – produced by a flower in basic soil • These colors are opposite of the common “litmus” test.

  3. Acid Properties • Regardless of the strength of the acid, they exhibit the same properties • 1) Acids have a sour taste. • 2) Acids change the color of an acid-base indicator. • 3) Some acids react with metals to produce hydrogen gas. • Ba (s) + H2SO4 (aq)  BaSO4 (s) + H2 (g) • 4) Acids react with bases to produce salts and water. • Neutralization • 5) Acids conduct electric current.

  4. Industrial Acids • Sulfuric Acid, H2SO4 • 37 million metric tons produced a year in the US! • Used in petroleum refining to paper production. • Nitric Acid, HNO3 • Explosives to pharmaceuticals • Phosphoric Acid, H3PO4 • Beverages to fertilizers • Hydrochloric Acid, HCl • Digestion, food processing, cleaning agent (muriatic) • Acetic Acid, CH3COOH • Vinegar, food supplements (lysine – amino acid)

  5. Base Properties • The properties of bases differ slightly from those of acids • 1) Bases have a bitter taste • Having one’s mouth washed out with soap….. • 2) Bases change the color of acid-base indicators. • 3) Bases feel slippery. • 4) Bases react with acids to produce salts and water. • Neutralization • 5) Bases conduct an electric current.

  6. Naming Acids • Naming acids follows a system • Binary Acids • Composed of 2 different elements. • 1) HCl • Hydrogen • Chlorine • Hydro + chlor + ic + acid • Hydrochloric acid

  7. Naming Acids II • Binary Acids • There are only a few, common binary acids. • 1) HCl • Hydrochloric acid • 2) HF • Hydrofluoric acid • 3) HI • Hydroiodic acid • 4) HBr • Hydrobromic acid • 5) H2S • Hydrosulfuric acid

  8. Naming Acids III • Oxyacids • These acids contain oxygen • PO4-3 • Phosphate ion • H3PO4 is phosphoric acid • NO3-1 • Nitrate ion • HNO3 is nitric acid • Ions that end with “ate” produce “ic” acids • “You ate something ic-y”

  9. Naming Acids IV • Oxyacids Part Deux • PO3-3 • Phosphite ion • H3PO3 is phoshorous acid • NO2-1 • Nitrite ion • HNO2 is nitrous acid • Ions that end with “ite” produce “ous” acids • Use the vowels: • A(t)E – ic • ite – OU(s)

  10. Naming Acids V • 1) H2CO3 • Carbonic acid • 2) HClO2 • Chlorous acid • 3) HIO3 • Iodic acid • 4) H2SO3 • Sulfurous acid • 5) HClO4 • Perchloric acid • 6) H2C2O4 • Oxalic acid

  11. Naming Bases • The naming system for bases is more basic than acids. • Follow the same system has if naming a compound. • You do not add “base” to the end of the name. • NaOH • Na – sodium • OH – hydroxide • Sodium hydroxide • NH3 • ammonia

  12. Naming Bases II • 1) Ba(OH)2 • Barium hydroxide • 2) Ca(OH)2 • Calcium hydroxide • 3) Cu(OH)2 • Copper (ii) hydroxide • 4) KOH • Potassium hydroxide • 5) RbOH • Rubidium hydroxide • 6) CsOH • Cesium hydroxide

  13. Strengths of Acids and Bases • Strong Acids • “Complete ionization” (break into ions) • HCl, HBr, HI, HNO3, HClO4, H2SO4, HClO3 • Weak Acids • “Incomplete ionization” (break into few ions) • HF, H2CO3, H2S, H3PO4 • Strong Bases • “Complete dissociation” • Ca(OH)2, Sr(OH)2, Ba(OH)2, NaOH, KOH, RbOH, CsOH • Weak Bases • “Incomplete dissociation” • NH3

  14. Types of Acids and Bases • Acids and bases are common substances that have been studied for hundreds of years. • It is no surprise that there are competing theories about how acids and bases behave at a molecular/atomic level. • Arrhenius • Bronsted-Lowry • Lewis*

  15. Arrhenius Definition • Arrhenius knew that acids and bases conducted electric current in solution. • Therefore, they must produce ions in solution. • Acids • A compound that increases the [H+] in solution. • H+ = H3O+ • HNO3, HCl • Bases • A compound that increases the [OH-] in solution. • NaOH, Ba(OH)2

  16. Arrhenius Acid • Arrhenius acids all produce H3O+ (hydronium ions) • HNO3 (l) + H2O (l)  H3O+ (aq) + NO3- (aq) • HCl (g) + H2O (l)  H3O+ (aq) + Cl- (aq)

  17. Arrhenius Base • Arrhenius bases all produce OH- (hydroxide ions) • NaOH (s) + (H2O (l))  Na+ (aq) + OH- (aq) • Sodium hydroxide dissociates (splits) • NH3 (aq) + H2O (l)  NH4+ (aq) + OH-(aq) • Ammonia is a weak base(more on that later) and is considered an Arrhenius base since OH- is produced.

  18. Bronsted Definition • Bronsted aimed to expand the Arrhenius definition of acids and bases. • Now the role of the H+ (proton) is key. • Acids • A compound that is a H+ donor. • HNO3, HCl • These compounds become negative as they lose positive ions. • Bases • A compound that is a H+acceptor. • NH3 • These compounds become positive as they accept positive ions!

  19. Bronsted Acid and Base • To identify a Bronsted Acid or Base, we must look at the movement of H • HCldonates the H, H2O accepts the H • H2O donates the H, NH3accepts the H

  20. Lewis Definition • To identify a Lewis Acid or Base, we must look at the movement of a pair of electrons. • Lewis acids accepts the electrons • Lewis bases donate the electrons BF3 accepts a pair (L.A.) of electrons from F- (L.B.) forming BF4-.

  21. Identify the Type(s) of Acid • An example may fit more than one description • Also identify as strong or weak • 1) HNO3 (aq) + H2O (l)  H3O+(aq) + NO3-(aq) • Arrhenius and Bronsted • Strong • 2) H2SO4 (aq) + H2O (l)  H3O+(aq) + HSO4-(aq) • Arrhenius and Bronsted • Strong • 3) HSO4-(aq) + H2O(l) H3O+(aq) + SO4-2(aq) • Arrhenius and Bronsted • Weak • 4) HCN (aq) + H2O (l)  H3O+(aq) + CN-(aq) • Arrhenius and Bronsted • Weak

  22. Identify the Type(s) of Base • An example may fit more than one description • Also identify as strong or weak • 1) KOH (s)  K+(aq) + OH-(aq) • Arrhenius • Strong • 2) NH3 (aq) + H2O (l)  NH4+(aq) + OH-(aq) • Arrhenius, Bronsted • Weak • 3) H2O (l) + HCl (aq)  H3O+(aq) + Cl-(aq) • Bronsted • Weak • 4) HPO4-2 (aq) + H3O+ (aq)  H2O (l) + H2PO4- (aq) • Bronsted • Weak

  23. Neutralization • Antacid tablets are bases which try to reduce the excess acid in the stomach, making you feel better. • The process forms water and a salt. • HCl (aq) + NaOH (aq) form ions in solution • H+(aq) + Cl- (aq) + Na+ (aq) + OH- (aq) • HOH(l) + NaCl (aq)

  24. Acid Rain • Rainwater is slightly acidic • Acid rain is very acidic rainwater. • Industrial processes produce gasses that, when dissolved in water, produce an acid. • NO, NO2, CO2, SO2, SO3 • SO3 (g) + H2O (l)  H2SO4 (aq) • Thanks to the Clean Air Act, the occurrence of acid rain has dropped in the US since 1990.

  25. Go GO Power HYDROGEN • The true pH of a solution is dependent upon concentration of H+ (H3O+) ions. • The higher the concentration, the lower the pH.

  26. pH and Concentration • [H3O+]> [OH-] = acid • [H3O+] < [OH-] = base • [H3O+] = [OH-] = neutral (pH = 7) • When the concentration (Molarity) of H3O+ is known, the pH can be easily calculated. • pH = -log[H3O+] • A neutral solution has a [H3O+] = 1 x 10-7 M. • pH = -log [1 x 10-7] = -(-7.0) • pH = 7.0

  27. pH and Concentration II • When you use your calculator: • - log ( ### x 10 ^ -# ) = • Example: What is the pH of a solution if the [H3O+] = 3.4 x 10-5 M? • Answer • pH = - log (3.4 x 10 ^ -5) = • pH = -log(3.4 x 10-5) • pH = 4.46

  28. pH and Concentration III • pH = -log[H3O+] • 1) What is the pH of a solution if the [H3O+] = 6.7 x 10-4 M? • pH = 3.17 • 2) What is the pH of a solution with a hydronium ion concentration of 2.5 x 10-2 M? • pH = 1.60 • 3) Determine the pH of a 2.5 x 10-6 HNO3 solution. • pH = 5.60

  29. pOH • When the concentration (Molarity) of OH- is known, the pOH can be easily calculated. • pOH = -log[OH-] • A neutral solution has a [OH-] = 1 x 10-7 M. • pOH = -log [1 x 10-7] = -(-7.0) • pOH = 7.0 • Therefore • pH + pOH = 14.0*

  30. pOH and Concentration • pOH = -log[OH-] • Determine the pOH for the following solutions: • 1) 1 x 10-3 M NaOH • pOH = -log(1 x 10-3) = 3.0 • 2) 1.0 x 10-5 M Ca(OH)2 = • pOH = -log(1.0 x 10-5) = 5.0 • 3) 6.29 x 10-6 M RbOH • pOH = -log(6.29 x 10-6) = 5.20 • 4) 7.4 x 10-4 M KOH • pOH = -log(7.4 x 10-4) = 3.13

  31. pH from pOH Concentration • pOH = -log[OH-] • pH + pOH = 14 • Determine the pH for the following solutions: • 1) 3.8 x 10-3 M NaOH • pOH = -log(3.8 x 10-3) = 2.42 • pH + 2.42 = 14.0 • pH = 14.0 – 2.42 = 11.58 • 2) 4.19 x 10-4 M LiOH • pOH = -log(4.19 x 10-4) = 3.38 • pH + 3.38 = 14.0 • pH = 14.0 – 3.38 = 10.62

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