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Acid Base Rxns

Acid Base Rxns. Reaction with an “active metal” and an acid. Mg + 2HCl H 2 + MgCl 2. Mg + HCl. Zn + 2HCl H 2 + ZnCl 2. Zn + HCl. Al + HCl. 2Al + 6HCl 3H 2 + 2AlCl 3. Write the balanced equation for the reaction that occurs between. What type of Reaction?.

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Acid Base Rxns

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  1. Acid Base Rxns

  2. Reaction with an “active metal” and an acid

  3. Mg + 2HCl H2 + MgCl2 Mg + HCl Zn + 2HCl H2 + ZnCl2 Zn + HCl Al + HCl 2Al + 6HCl 3H2 + 2AlCl3 Write the balanced equation for the reaction that occurs between

  4. What type of Reaction? Single Replacement Reactions

  5. Neutralization Reactions Despite heavy fire, McAllister’s aim was true, and his carefully measured hydrochloric acid found its mark in the enemy’s reservoir of sodium hydroxide McAllister grinned, one of the enemy’s strongest bases had finally been neutralized

  6. Hydrochloric acid + sodium hydroxide ? ACID + BASE SALT + WATER HCl + NaOH NaCl + HOH Neutralization: An Acid Base Rxn Hydroxide ion Complete neutralization occurs when The acid H+ and the base OH- are 1:1

  7. Magnesium hydroxide + nitric acid Neutralization Reactions Phosphoric acid + ammonium hydroxide

  8. Net ionic equation: H+ + OH- HOH Complete neutralization . . . • Occurs when H+ and OH- are added together in a 1:1 molar ratio • MaVa = MbVb • BE CAREFUL • If the acid or base has a subscript you have to account for it in the formula. • saMaVa = sbMbVb • Ex. If you have 1L of 1M HCL, what M is necessary to neutralize 1L of Mg(OH)2?

  9. Titrations • Lab procedure used to determine a concentration (M) of a solution. • Titrations use an indicator to signal when the “endpoint” is reached • Endpoint is when the solution is neutralized • Indicator changes color

  10. Titration / Neutralization practice • 55mL of an 8M HCl solution is completely neutralized by 26mL of Ca(OH)2 • Write the balanced neutralization reaction. • What is the concentration of base that was used? • What is the net ionic equation? • What are the spectator ions?

  11. Salt Hydrolysis Reaction of a salt with water • Salts are ionic compounds • contain a + cation and a – anion • Derived from acids and bases • So when dissolved in water some salts produce: • A neutral solution • A basic solution • An acidic solution

  12. Cation of NaOH Anion of CH3COOH Salts that hydrolize water & produce alkaline solns • Consists of a cation of a strong base and a anion of a weak acid NaCH3COO

  13. NaCH3COO  Na+ + CH3COO- • When this salt dissolves in water dissociation occurs. • Water also ionizes H2O  H+ + OH- • There is a force of attraction between the • H+ of the water and the acetate anion • & a force of attraction between the Na+ and • OH-

  14. YES I KNOW . . . • Acetic acid CH3COOH is formed • BUT it has a low Ka value and doesn’t ionize completely • Most of the H+ stays bonded • NaOH is formed BUT it stays completely ionized as Na+ and OH- • If there are more OH- than H+ in solution the solution is considered BASIC!!!!!!!!!

  15. Weak base bonds Strong acid stays ionized Some salts hydrolyze water & produce acidic solns • The salt of a strong acid and a weak base NH4Cl dissolved in water produces: NH4+ Cl- H+ and OH-

  16. Salts of a weak base and a weak acid may hydrolyze water, but pH needs to be tested. • Salts of strong bases and strong acids do not hydrolyzewater and form neutral solutions • NaCl

  17. Buffers • A buffer is a mixture of chemicals that make a solution resist a change of pH • pH remains relatively constant when adding an acid or base • A buffer is either a solution of a weak acid and one of its salts or a weak base and one of its salts • The salt cation is only a spectator ion and is not involved in the reaction

  18. Buffer Capacity • Buffers can’t stop the change in pH they can only resist it • There comes a point when the buffer is “used up” • It can no longer take H+ ions out of solution and the pH begins to change

  19. This is important when the system is fragile • Bloodstream pH 7.3 > 7.5 • If < 6.9 or > 7.7 the person dies • Buffer aids in maintaining homeostasis of the individual • The buffer species in the blood is carbonic acid / hydrogen carbonate ion H2CO3 / HCO3-

  20. If the blood is too alkaline, the reaction decreases the amount of OH- in the bloodstream H2CO3 + OH- HOH + HCO3- • If excess H+ is in the blood, the reaction • decreases it • H+ + HCO3- H2CO3

  21. Practice • Write an equation to show the addition of an acid and a base to the following buffer systems NH4+ / NH3 CH3COOH / CH3COO- H2PO4- / HPO4-

  22. THE END

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