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FUNDAMENTALS OF ACID – BASE (pH and Buffers). Herbert M. San Pedro Lecturer. Acid - Base. An important factor in the proper behavior of many biochemical phenomena (processes)

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acid base
Acid - Base
  • An important factor in the proper behavior of many biochemical phenomena (processes)
  • Any deviation from the expected acid-base balance may lead to some disorders, most common are the alkalosis and acidosis.
  • Can be measured by knowing the Hydronium ion concentration [ ] (established by Sorensen) - pH
dissociation principle
DISSOCIATION PRINCIPLE
  • Ability to form ions in solution.
  • Water, plays a vital role in the determination of the degree of acidity (alkalinity) of many inorganic substances, likewise in the biochemical processes.
  • Water being a naïve substance, can act as both an acid or base ( amphoterism )
slide4
Cont.
  • Water molecules themselves may react with each other and create ions (auto-ionization of water – autoprotolysis)
  • Eqn: H2O + H2O ↔ H3O+ + - OH
  • Or Simplified by:

H2O ↔ H+ + - OH

slide5
Cont.
  • The presence of such ions contribute significantly to the over-all properties of water.
  • Going further analytically: the auto-ionization of water is expressed in terms of equilibrium expression:

K = [H+ ][- OH] / [H2O}

K = dissociation constant

slide6
Cont.
  • To further simplify the equation: it is considered that the concentration of pure water being experimentally to high (55.56 mol/L) is doubtedly affected by dissociation (making it constant).
  • Eqn becomes: K[H2O] = Kw (water const.)

Kw = [H+ ][- OH ] = ion product of water

  • Simplifying things out, it is from the above eqn.

wherein the concept pH was introduced and also applied to acidic and basic solutions.

slide7
Cont.
  • pH – degree of acidity, value is affected by the amount of hydronium ion present in solution.
  • Low pH – High concentration of hydronium ion. (below 7)
  • High pH – Low concentration of hydronium ion. (above 7)
  • The formula: pH = - log [H+ ] : UE
sample calculations
Sample Calculations:

What is the pH of a solution whose hydrogen ion concentration is 3.2 x 10-4 mol/L.

Solution:

pH = - log [H+ ]

= - log (3,2 x 10-4 )

= - (-3.5)

pH = 3.5

slide9
Cont:

2. What is the pH of a solution whose hydroxide ion concentration is 4,0 x 10-4 mol/L: (in similar fashion pOH = - log [- OH])

Considering analytically the ion product of water Kw = [H+ ][- OH] = 1 x 10-14 taking the log of both sides this becomes:

pH + pOH = 14 : UE

slide10
Cont.

To solve the problem:

[- OH] = 4.0 x 10-4

pOH = - log [- OH]

= - log (4.0 x 10-4 )

= - (-3.4)

pOH = 3.4

pH + pOH = 14

pH = 14 – pOH

= 14 – 3.4

pH = 10.4

slide11
Cont.

Considering the vice-versa:

[H+ ] = 10-pH : UE

Given the pH of a substance to be equal to 5. Calculate its hydrogen ion concentration.

[H+ ] = antilog (-pH)

= antilog ( - 5)

[H+ ] = 1 x 10-5 mol/L

the weak acid weak base
The Weak Acid / Weak Base
  • Many biochemicals possess functional groups that are weak acids or bases. Thus knowledge of the dissociation of WA/ WB is basic to understanding the influence of intracellular pH on structure and biologic activity.
  • And these therefore require a different approach in calculating for their acidity and basicity (acidity / basicity dissociation constant; Ka or Kb) :

HA ↔ H+ + - A (weak acid)

Ka = [H+ ][- A] / [HA]

henderson hasselbach equation
HENDERSON-HASSELBACH EQUATION
  • A very useful equation (UE) in calculating for pH involving weak acids/weak base and their conjugates.
  • Derivation:

HA ↔ H+ + - A

Ka = [H+ ][- A]/[HA]

Ka[HA] = [H+ ][- A]

[H+ ] = Ka[HA]/[- A]

slide14
Cont.

Log [H+ ] = Log (Ka[HA]/[- A])

= Log Ka + Log [HA]/[- A]

Multiplying both sides by -1

-Log [H+ ] = - Log Ka – Log [HA]/[- A]

pH = pKa – Log [HA]/[- A]

pH = pKa + log [- A]/[HA] : UE

buffers
BUFFERS
  • Solution of Weak Acid/Base and their conjugate. (salts)
  • Has the ability to resist change in pH after the addition of strong acids/bases.
  • To measure the pH changed that occurred after the addition of either strong acid/base the HHE is being used.
slide16
Cont.
  • There are so many buffers in the human body and all of these are responsible to maintain the proper pH the body requires (physiologic buffers)
  • a. Bicarbonate, Phosphate, Protein Buffers
  • Any deviation from the desired pH may lead to disorders like acidosis or alkalosis.
acid base balance
Acid-Base Balance
  • Blood place a crucial role In the over-all acid-base balance, It must retain (through the help of buffers) its normal pH (even after the addition of strong acid/bases) so as not jeopardize the biochemical processes.
  • The Kidney and the Lungs are the two major organs that help the blood to regulate its pH
slide18
Cont.
  • The exhaling of CO2 by the lungs and regulation of HCO3- by the kidney, are the responsible processes how the blood is capable of maintaining its pH.

> Any trouble acquired by either of the organs or conditions/situations again acquired by any persons may suggest possible imbalance to the equilibrium

slide19
Cont.

CO2 + H2O ↔ H2CO3 ↔ H+ + HCO3-

thank you
THANK YOU!!!

Long Examination on Thursday (July 15, 2010)

from the Beginning –up to Last Topic

--

No Meeting on Saturday (July 10, 2010)