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Acids and Bases Chapter 15

Acids and Bases Chapter 15. Properties of Acids. Sour taste Change color of vegetable dyes React with “active” metals Like Al, Zn, Fe, but not Cu, Ag or Au Zn + 2 HCl  ZnCl 2 + H 2 Corrosive React with carbonates, producing CO 2 Marble, baking soda, chalk

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Acids and Bases Chapter 15

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  1. Acids and BasesChapter 15

  2. Properties of Acids • Sour taste • Change color of vegetable dyes • React with “active” metals • Like Al, Zn, Fe, but not Cu, Ag or Au Zn + 2 HCl ZnCl2 + H2 • Corrosive • React with carbonates, producing CO2 • Marble, baking soda, chalk CaCO3 + 2 HCl CaCl2 + CO2 + H2O • React with bases to form ionic salts • And often water

  3. Properties of Bases • Also Known As Alkalis • Taste bitter • Feel slippery • Change color of vegetable dyes • Different color than acid • Litmus = blue • React with acids to form ionic salts • And often water • Neutralization

  4. Arrhenius Theory • Acids ionize in water to H+1 ions and anions • Bases ionize in water to OH-1 ions and cations • Neutralization reaction involves H+1 combining with OH-1 to make water • H+ ions are protons • Definition only good in water solution • Definition does not explain why ammonia solutions turn litmus blue • Basic without OH- ions

  5. Brønsted-Lowery Theory • H+1 transfer reaction • Since H+1 is a proton, also known as proton transfer reactions • Acid is H+ donor; Base is H+ acceptor • Base must contain an unshared pair of electrons • In the reaction, a proton from the acid molecule is transferred to the base molecule • H forms a bond to lone pair electrons on the base molecule • We consider only 1 H transferred in each reaction • Products are called the Conjugate Acid and Conjugate Base • After reaction, the original acid is the conjugate base and the original base is changed to what is now called the conjugate acid

  6. Brønsted-Lowery Theory H-A + :B  A-1 + H-B+1 A-1 is the conjugate base, H-B+1 is the conjugate acid • Conjugate Acid-Base Pair is either the original acid and its conjugate base or the original base and its conjugate acid • H-A and A-1 are a conjugate acid-base pair • :B and H-B+1 are a conjugate acid-base pair • The conjugate base is always more negative than the original acid; and the conjugate acid is always more positive than the original base

  7. Example #1 Write the conjugate base for the acid H3PO4 • Determine what species you will get if you remove 1 H+1 from the acid • The Conjugate Base will have one more negative charge than the original acid H3PO4  H+1 + H2PO4-1

  8. Brønsted-Lowery Theory • In this theory, instead of the acid, HA, dissociating into H+1(aq) and A-1(aq); The acid donates its H to a water molecule HA + H2O  A-1 + H3O+1 A-1 is the conjugate base, H3O+1 is the conjugate acid • H3O+1is called hydronium ion • In this theory, substances that do not have OH-1 ions can act as a base if they can accept a H+1 from water H2O + :B  OH-1 + H-B+1

  9. Strength of Acids & Bases • The stronger the acid, the more willing it is to donate H • Strong acids donate practically all their H’s HCl + H2O  H3O+1 + Cl-1 • Strong bases will react completely with water to form hydroxides CO3-2 + H2OHCO3-1 + OH-1 • Weak acids donate a small fraction of their H’s • The process is reversible, the conjugate acid and conjugate base can react to form the original acid and base HC2H3O2 + H2O  H3O+1 + C2H3O2-1 • Only small fraction of weak base molecules pull H off water HCO3-1 + H2OH2CO3 + OH-1

  10. Figure 15.1: Graphical representation of the behavior of acids in aqueous solution

  11. Figure 15.2: The relationship of acid strength and conjugate base strength

  12. Multiprotic Acids • Monoprotic acids have 1 acid H, diprotic 2, etc. • In oxyacids only the H on the O is acidic • In strong multiprotic acids, like H2SO4, only the first H is strong; transferring the second H is usually weak H2SO4 + H2O  H3O+1 + HSO4-1 HSO4-1 + H2O  H3O+1 + SO4-2

  13. Water as an Acid and a Base • Amphoteric substances can act as either an acid or a base • Water as an acid, NH3 + H2O NH4+1 + OH-1 • Water as a base, HCl + H2O  H3O+1 + Cl-1 • Water can even react with itself H2O + H2O H3O+1 + OH-1

  14. Autoionization of Water • Water is an extremely weak electrolyte • therefore there must be a few ions present H2O + H2O  H3O+1 + OH-1 • all water solutions contain both H3O+1 and OH-1 • the concentration of H3O+1 and OH-1 are equal • [H3O+1] = [OH-1] = 10-7M @ 25°C • Kw = [H3O+1] x [OH-1] = 1 x 10-14 @ 25°C • Kw is called the ion product constant for water • as [H3O+1] increases, [OH-] decreases

  15. 1 x 10-14 [OH-1] 1 x 10-14 [H+1] [H+1] = [OH-1] = Acidic and Basic Solutions • acidic solutions have a larger [H+1] than [OH-1] • basic solutions have a larger [OH-1] than [H+1] • neutral solutions have [H+1]=[OH-1]= 1 x 10-7 M

  16. Example #2 Determine the [H+1] and [OH-1] in a 10.0 M H+1 solution • Determine the given information and the information you need to find Given [H+1] = 10.0 M Find [OH-1] • Solve the Equation for the Unknown Amount

  17. Example #2 Determine the [H+1] and [OH-1] in a 10.0 M H+1 solution • Convert all the information to Scientific Notation and Plug the given information into the equation. Given [H+1] = 10.0 M = 1.00 x 101 M Kw = 1.0 x 10-14

  18. pH & pOH • The acidity/basicity of a solution is often expressed as pH or pOH • pH = -log[H3O+1] pOH = -log[OH-1] • pHwater = -log[10-7] = 7 = pOHwater • [H+1] = 10-pH [OH-1] = 10-pOH • pH < 7 is acidic; pH > 7 is basic, pH = 7 is neutral • The lower the pH, the more acidic the solution; The higher the pH, the more basic the solution • 1 pH unit corresponds to a factor of 10 difference in acidity • pOH = 14 - pH

  19. Figure 15.3: The pH scale and pH values of some common substances

  20. Figure 15.4: A pH meter

  21. Figure 15.5: Indicator paper being used to measure the pH of a solution

  22. Example #3 Calculate the pH of a solution with a [OH-1] = 1.0 x 10-6 M • Find the concentration of [H+1]

  23. Example #3 Calculate the pH of a solution with a [OH-1] = 1.0 x 10-6 M • Enter the [H+1] concentration into your calculator and press the log key log(1.0 x 10-8) = -8.0 • Change the sign to get the pH pH = -(-8.0) = 8.0

  24. Example #4 Calculate the pH and pOH of a solution with a [OH-1] = 1.0 x 10-3 M • Enter the [H+1] or [OH-1]concentration into your calculator and press the log key log(1.0 x 10-3) = -3.0 • Change the sign to get the pH or pOH pOH = -(-3) = 3.0 • Subtract the calculated pH or pOH from 14.00 to get the other value pH = 14.00 – 3.0 = 11.0

  25. Example #5 Calculate the [OH-1] of a solution with a pH of 7.41 • If you want to calculate [OH-1] use pOH, if you want [H+1] use pH. It may be necessary to convert one to the other using 14 = pH + pOH pOH = 14.00 – 7.41 = 6.59 • Enter the pH or pOH concentration into your calculator • Change the sign of the pH or pOH -pOH = -(6.59) • Press the button(s) on you calculator to take the inverse log or 10x [OH-1] = 10-6.59 = 2.6 x 10-7

  26. Calculating the pH of a Strong, Monoprotic Acid • A strong acid will dissociate 100% HA  H+1 + A-1 • Therefore the molarity of H+1 ions will be the same as the molarity of the acid • Once the H+1 molarity is determined, the pH can be determined pH = -log[H+1]

  27. Example #6 Calculate the pH of a 0.10 M HNO3 solution • Determine the [H+1] from the acid concentration HNO3 H+1 + NO3-1 0.10 M HNO3 = 0.10 M H+1 • Enter the [H+1] concentration into your calculator and press the log key log(0.10) = -1.00 • Change the sign to get the pH pH = -(-1.00) = 1.00

  28. Buffered Solutions • Buffered Solutions resist change in pH when an acid or base is added to it. • Used when need to maintain a certain pH in the system • Blood • A buffer solution contains a weak acid and its conjugate base • Buffers work by reacting with added H+1 or OH-1 ions so they do not accumulate and change the pH • Buffers will only work as long as there is sufficient weak acid and conjugate base molecules present

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