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Acids and Bases

Acids and Bases. Chapter 15. What are acids and bases Arrhenius definition. Arrhenius suggested that acids are compounds that contain hydrogen and can dissolve in water to release hydrogen ions into solution. For example, hydrochloric acid (HCl) dissolves in water as follows:

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Acids and Bases

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  1. Acids and Bases Chapter 15

  2. What are acids and bases Arrhenius definition • Arrhenius suggested that acids are compounds that contain hydrogen and can dissolve in water to release hydrogen ions into solution. For example, hydrochloric acid (HCl) dissolves in water as follows: HCl +H2OH+(aq) + Cl-(aq)

  3. Arrhenius defined bases as substances that dissolve in water to release hydroxide ions (OH-) into solution. For example, a typical base according to the Arrhenius definition is sodium hydroxide (NaOH): • NaOH +H2ONa+(aq)  +  OH-(aq)

  4. Acids release H+ into solution and bases release OH-. If we were to mix an acid and base together, the H+ ion would combine with the OH- ion to make the molecule H2O, or plain water: • H+(aq) +   OH-(aq)   H2O

  5. What are acids and BasesBronsted Lowry theory • The Bronsted -Lowry definition of acids is very similar to the Arrhenius definition, any substance that can donate a hydrogen ion is an acid (under the Bronsted definition, acids are often referred to as proton donors because an H+ ion, hydrogen minus its electron, is simply a proton).

  6. The Bronsted definition of bases is, however, quite different from the Arrhenius definition.  The Bronsted base is defined as any substance that can accept a hydrogen ion. A base is the opposite of an acid.  NaOH and KOH, would still be considered bases because they can accept an H+ from an acid to form water. 

  7. The Bronsted-Lowry definition also explains why substances that do not contain OH- can act like bases.  Baking soda (NaHCO3), for example, acts like a base by accepting a hydrogen ion from an acid as illustrated below: • Acid Base Salt HCl+ NaHCO₃  H2CO3  +  NaCl

  8. Neutralization reaction • We have seen from the equations, acids release H+ into solution and bases release OH- If we were to mix an acid and base together, the H+ ion would combine with the OH- ion to make the molecule H2O, or plain water: • H+(aq) +   OH-(aq)   H2O. The neutralization reaction of an acid with a base will always produce water and a salt: • AcidBaseWaterSalt • HCl  +  NaOH  H2O  +  NaCl • HBr  +  KOH   H2O  +  KBr

  9. Self ionization of water • Water molecules collide with one another to cause the self-ionization reaction represented by this equation: 2H2O H3O+ + OH- • It is a reversible reaction so the equation is usually written with the arrows going in both directions: 2H2O(l) ↔H3O+ (aq)+ OH- (aq) • The self ionization of pure water produces equal amounts of H₃O+ ions and OH- ions. Therefore the concentrations of these ions in pure water must be equal. [OH- ]=[H3O+ ]

  10. It has been found that at 25⁰C the concentration of these two ions are equal to 1.00x10-7 mol/L • Recall the equation for equilibrium constant Keq which can be calculated from the equation. Since self ionization is an equilibrium reaction, the equilibrium constant called Kw, can be calculated as follows. Kw=[H3O+ ] [OH- ]=(1.00x10-7 )(1.00x10-7)= 1.00x10-14 Kw= 1.00x10-14 Kw is called as autoionisation constant.

  11. Class Practice • What is [OH- ] in a 3.00x10-5 M solution of HCl? • Section review on page 549

  12. Acids donate protons • Any molecule or ion that can transfer a proton; a hydrogen atom nucleus to another species. • Let us take the example of sulfuric acid • H2SO4(l) + H2O(l)- H 3O+ (aq) +HSO 4- (aq) • The hydrogen sulfate ion, is also a Bronsted Lowry acid because it too can donate a proton to a water molecule. • HSO 4- (aq) +H2O(l)↔ H3O+ (aq) +SO42- (aq)

  13. Monoprotic , diprotic and triproticacids • Monoprotic acids are acids that can donate one hydrogen ion per molecule. • Example nitric acid HNO₃ or hydrochloric acid HCl. • Diprotic acids are acids that can donate two hydrogen ions per molecule. • Examples sulfuric acids H₂SO₄. • Triprotic acids are acids that can donate three hydrogen ions per molecule. • Example phosphoric acid H₃PO₄

  14. Bases accepts protons • Any atom, or molecule that receives a proton from another species. • Ammonia is a typical base • NH₃(aq)+H₂O(l) ↔NH₄+ (aq)+ OH- (aq) • In the gas phase ammonia accepts a proton from HCl and forms ammonium chloride which is composed of ammonium and chloride ion.

  15. Species that are both acids and bases. • Water can act as an acid by donating a proton, and can act as a base by accepting a proton. • H₂O(l)+H₂O(l)↔OH- (aq) +H₃O+ (aq) • A species that can act as either an acid or a base is called as amphoteric. • Hydrogen carbonate ion is also an example of amphoteric species. • HCO₃- (aq) + OH- (aq) ↔CO₃2- (aq)+H₂O(l) • It behaves as abase in the presence of an acid such as formic acid • HCOOH (aq) + HCO₃- (aq) ↔HCOO- (aq) + H₂CO₃(aq)

  16. Conjugate acids and bases • Conjugate acids are formed when a base accepts a proton. • Conjugate base are formed when an acid donates a proton. • NH₃(aq)+H₂O(l)↔NH₄+ (aq) + OH- (aq) • Ammonium ion is the conjugate acid and OH ion is the conjugate base.

  17. Class Practice • Page 554 concept check

  18. Weak acids and bases • Weak acids and bases are partially ionized in their solutions, whereas strong acids and bases are completely ionized when dissolve in water. Common Weak Acids • Formic HCOOH • Acetic CH3COOH • Trichloro acetic CCl3COOH • Hydrofluoric HF • Hydrocyanic HCN • Hydrogen sulfide H2S • Water H2O

  19. Common Weak Bases • ammonia NH3 • Trimethyl ammonia N(CH3)3 • pyridine C5H5N • Ammonium hydroxide NH4OH • water H2O

  20. Acid ionization constant • The equilibrium constant of the ionization of a weak acid in water is known as the acid ionization constant Ka. • CH₃COOH(aq)+H₂O(l)↔CH₃COO- (aq) +H₃O+ (aq) • The equilibrium expression for this reaction is written as follows • Ka=[H₃O+ ][CH₃COO- ]/[CH₃COOH] =1.75x10-5

  21. Class practice • A vinegar sample is found to be 0.837M acetic acid. Its hydronium ion concentration is measured as 3.86x10-3 mol/L. Calculate Ka for CH₃COOH.

  22. Home work • Page 558 • Total recall

  23. pH • What is pH? • The negative logarithm of the hydronium ion concentration in a solution. • pH of a solution ranging from 1-7 are usually acidic. • pH of a solution ranging from 7-14 are usually basic. • pH of a solution which is 7 are neutral.

  24. Class Practice • Page 562 • Practice problems all

  25. pH and Kw • Kw=[H₃O+][OH-]=1.00x10-14 • What is the pH of a 0.0136M solution of Ba(OH)₂ a strong base?

  26. Class Practice • Page 563 • Practice problems

  27. Buffers • A buffer solution is one which resists changes in pH when small quantities of an acid or an alkali are added to it. • An acidic buffer solution is the one which has a pH less than 7. Acidic buffer solutions are commonly made from a weak acid and one of its salts often a sodium salt.

  28. An alkaline buffer solution has a pH greater than 7. Alkaline buffer solutions are commonly made from a weak base and one of its salts.

  29. Titration • The operation of gradually adding one solution to another to reach an equivalence point. • Equivalence point is the point in a titration when the amount of added base or acid exactly equals the amount of acid or base originally in solution.

  30. Titrant: The solution added to another solution in a titration.

  31. Home work • Term Review all Page 580 • 15,17,29,41 ,47,66 • Test prep all

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