Acids and Bases

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Acids and Bases

• Chapter 15

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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)

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• 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)

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• 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

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What are acids and Bases Bronsted
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).

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• 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. 

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• 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

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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
• Acid Base Water Salt
• HCl    NaOH ?  H2O    NaCl
• HBr    KOH  ? H2O    KBr

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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

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• It has been found that at 25C 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.
• KwH3O OH- (1.00x10-7 )(1.00x10-7)
• 1.00x10-14
• Kw 1.00x10-14 Kw is called as
  autoionisation constant.

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Class Practice

• What is OH- in a 3.00x10-5 M solution of HCl?
• Section review on page 549

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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)

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Monoprotic , diprotic and triprotic acids

• 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

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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.

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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)

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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.

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Class Practice

• Page 554 concept check

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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

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

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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
• KaH3O CH3COO- /CH3COOH 1.75x10-5

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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.

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Home work

• Page 558
• Total recall

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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.

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Class Practice

• Page 562
• Practice problems all

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pH and Kw

• KwH3OOH-1.00x10-14
• What is the pH of a 0.0136M solution of Ba(OH)2 a
  strong base?

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Class Practice

• Page 563
• Practice problems

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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.

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• 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.

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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.

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• Titrant The solution added to another solution
  in a titration.

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Home work

• Term Review all Page 580
• 15,17,29,41 ,47,66
• Test prep all