The concept of pH and pKa - PowerPoint PPT Presentation

Loading...

PPT – The concept of pH and pKa PowerPoint presentation | free to download - id: 569a93-NDJiN



Loading


The Adobe Flash plugin is needed to view this content

Get the plugin now

View by Category
About This Presentation
Title:

The concept of pH and pKa

Description:

The concept of pH and pKa Lecture 3 Handout continued base dissociation constant (or Kb) a measure of basicity pKb is the negative log of Kb and related to the pKa by ... – PowerPoint PPT presentation

Number of Views:596
Avg rating:3.0/5.0
Slides: 57
Provided by: spa956
Category:
Tags: concept | pka

less

Write a Comment
User Comments (0)
Transcript and Presenter's Notes

Title: The concept of pH and pKa


1
The concept of pH and pKa
  • Lecture 3
  • Handout

2
Introduction
  • Why is pH so important for maintaining
    homeostasis?
  • pH of blood
  • pH and diseases

3
Introduction
  • pH the measure of the acidity or alkalinity of
    a solution (pH stands for "power of hydrogen)
  • a measure of the activity of dissolved hydrogen
    ions (H)
  • for very dilute solutions ? the molarity (molar
    concentration) of H may be used as a substitute
    with little loss of accuracy
  • In solution ? hydrogen ions occur as a number of
    cations including hydronium ions (H3O)

4
continued
  • pure water at 25 C ? the concentration of H
    equals the concentration of hydroxide ions (OH-)
  • "neutral" and corresponds to a pH level of 7.0
  • Solutions ? the concentration of H exceeds that
    of OH- have a pH value lower than 7.0 acids
  • Solutions ? OH- exceeds H have a pH value
    greater than 7.0 bases
  • pH is dependent on ionic activity

5
(No Transcript)
6
Definition
  • pH a measurement of the concentration of
    hydrogen ions in a solution
  • low pH values ? associated with solutions with
    high concentrations of hydrogen ions
  • high pH values ? solutions with low
    concentrations of hydrogen ions
  • Pure water ? a pH of 7.0, and other solutions are
    usually described with reference to this value
  • Acids ? solutions that have a pH less than 7
    (i.e. more hydrogen ions than water)
  • Bases ? a pH greater than 7 (i.e. less hydrogen
    ions than water)

7
continued
  • definitions of weak and strong acids, and weak
    and strong bases do not refer to pH
  • It describe whether an acid or base ionizes in
    solution

8
Explanation of pH
  • the number (pH) arises from a measure of the
    activity of hydrogen ions or their equivalent in
    the solution
  • pH scale an inverse logarithmic representation
    of hydrogen proton (H) concentration
  • pH unit is a factor of 10 different than the next
    higher or lower unit
  • a change in pH from 2 to 3 represents a 10-fold
    decrease in H concentration, and a shift from 2
    to 4 represents a one-hundred (10 10)-fold
    decrease in H concentration

9
. The formula for calculating pH
  • aH denotes the activity of H ions, is
    dimensionless
  • Activity a measure of the effective
    concentration of hydrogen ions (rather than the
    actual concentration)
  • other ions surrounding hydrogen ions will shield
    them and affect their ability to participate in
    chemical reactions

10
dilute solutions (tap water) ? activity is
approximately equal to the numeric value of the
concentration of the H ion denoted as H
(H3O) measured in moles per litre (also known
as molarity) often convenient to define pH as
11
continued
  • log10 denotes the base-10 logarithm
  • therefore pH defines a logarithmic scale of
    acidity

12
continued
  • . For example, if one makes a lemonade with a H
    concentration of 0.0050 moles per litre, its pH
    would be

13
continued
  • A solution of pH 8.2
  • have an H concentration of 10-8.2 mol/L, or
    about 6.31 10-9 mol/L
  • its hydrogen activity aH is around 6.31 10-9
  • solution at 25 C, a pH of 7 indicates neutrality
    (i.e. the pH of pure water)
  • because water naturally dissociates into H and
    OH- ions with equal concentrations of 110-7 mol/L

14
continued
  • lower pH value (for example pH 3) indicates
    increasing strength of acidity
  • higher pH value (for example pH 11) indicates
    increasing strength of basicity
  • (pure water, when exposed to the atmosphere, will
    take in carbon dioxide, some of which reacts with
    water to form carbonic acid and H, thereby
    lowering the pH to about 5.7)

15
Calculation of pH for weak and strong acids
  • stronger or weaker acids are a relative concept
  • a strong acid a species which is a much
    stronger acid than the hydronium (H3O) ion
  • the dissociation reaction (strictly
    HXH2O?H3OX- but simplified as HX?HX-) goes
    to completion, i.e. no unreacted acid remains in
    solution
  • Dissolving the strong acid HCl (hydrochloric
    acid) in water
  • HCl(aq) ? H Cl-

16
continued
  • in a 0.01 mol/L solution of HCl it is
    approximated that there is a concentration of
    0.01 mol/L dissolved hydrogen ions
  • the pH is pH -log10 H
  • pH -log (0.01)
  • It equals 2

17
continued
  • weak acids
  • dissociation reaction does not go to completion
  • equilibrium is reached between the hydrogen ions
    and the conjugate base
  • equilibrium reaction between methanoic acid and
    its ions
  • HCOOH(aq) ? H HCOO-
  • We must know ? the value of the equilibrium
    constant of the reaction for each acid in order
    to calculate its pH
  • In the context of pH ? this is termed the acidity
    constant (Ka) of the acid
  • Ka hydrogen ionsacid ions / acid

18
continued
  • For HCOOH Ka 1.6 10-4
  • When calculating the pH of a weak acid, it is
    usually assumed that the water does not provide
    any hydrogen ions
  • it simplifies the calculation, and the
    concentration provided by water, 110-7 mol/L, is
    usually insignificant

19
continued
  • With a 0.1 mol/L solution of methanoic acid
    (HCOOH), the acidity constant is equal to
  • Ka HHCOO- / HCOOH
  • Given that an unknown amount of the acid has
    dissociated, HCOOH will be reduced by this
    amount, while H and HCOO- will each be
    increased by this amount

20
continued
  • HCOOH may be replaced by 0.1 - x, and H and
    HCOO- may each be replaced by x, giving us the
    following equation
  • Solving this for x yields 3.910-3 the
    concentration of hydrogen ions after dissociation
  • the pH is -log(3.910-3) or about 2.4

21
pH can be measured
  • by addition of a pH indicator into the solution
    under study
  • by using a pH meter together with pH-selective
    electrodes
  • by using pH paper, indicator paper that turns
    colour corresponding to a pH on a colour key

22
Fluid pH
gastric acid 0.7
lysosome 5.5
granule of chromaffin cell 5.5
Neutral H2O at 37C 6.81
cytosol 7.2
CSF 7.3
arterial blood plasma 7.4
mitochondrial matrix 7.5
exocrine secretions of pancreas 8.1
pH in body fluids
23
Acids
  • An acid (often represented by the generic formula
    HA HA-) ? any chemical compound that, when
    dissolved in water, gives a solution with a
    hydrogen ion activity greater than in pure water
    (a pH less than 7.0)
  • an acid as a compound which donates a hydrogen
    ion (H) to another compound (called a base)

24
continued
  • In water the following equilibrium occurs between
    a weak acid (HA) and water, which acts as a base
  • HA(aq) H2O ? H3O(aq) A-(aq)
  • acidity constant (or acid dissociation constant)
    is the equilibrium constant for the reaction of
    HA with water

25
  • Strong acids have large Ka values (the reaction
    equilibrium lies far to the right the acid is
    almost completely dissociated to H3O and A-)
  • Strong acids include the heavier hydrohalic
    acids hydrochloric acid (HCl), hydrobromic acid
    (HBr), and hydroiodic acid (HI)

26
continued
  • Weak acids ? have small Ka values (i.e. at
    equilibrium significant amounts of HA and A-
    exist together in solution modest levels of H3O
    are present the acid is only partially
    dissociated)
  • Most organic acids ? weak acids
  • nitrous acid, sulfurous acid and hypochlorous
    acid are all weak acids

27
Neutralization
  • the reaction between an acid and a base,
    producing a salt and neutralized base
  • hydrochloric acid and sodium hydroxide form
    sodium chloride and water
  • HCl(aq) NaOH(aq) ? H2O(l) NaCl(aq)

28
continued
  • Neutralization ? the basis of titration, where a
    pH indicator shows equivalence point when the
    equivalent number of moles of a base have been
    added to an acid
  • It is often wrongly assumed that neutralization
    should result in a solution with pH 7.0 (is only
    the case with similar acid and base strengths
    during a reaction)

29
continued
  • Neutralization with a base weaker than the acid ?
    weakly acidic salt
  • E.g. weakly acidic ammonium chloride (produced
    from the strong acid hydrogen chloride and the
    weak base ammonia)
  • neutralizing a weak acid with a strong base gives
    a weakly basic salt, e.g. sodium fluoride from
    hydrogen fluoride and sodium hydroxide

30
Biological occurrence of acids
  • In humans ? hydrochloric acid is a part of the
    gastric acid secreted within the stomach
  • hydrolyze proteins and polysaccharides
  • converting the inactive pro-enzyme, pepsinogen
    into the enzyme, pepsin

31
Bases
  • A strong base ? a base which hydrolyzes
    completely, raising the pH of the solution
    towards 14
  • weak bases (ammonia)
  • Arrhenius bases ? water-soluble and these
    solutions always have a pH greater than 7
  • alkali is a special example of a base, where in
    an aqueous environment, hydroxide ions (also
    viewed as OH-) are donated

32
Bases and pH
  • pure water ? molecules dissociate into hydronium
    ions (H3O) and hydroxide ions (OH-), according
    to the following equation
  • 2H2O(l) ? H3O(aq) OH-(aq)
  • concentration, measured in molarity (M or moles
    per dm³), of the ions ? indicated as H3O and
    OH-

33
continued
  • their product is the dissociation constant of
    water has the value 10-7 M
  • A base accepts (removes) hydronium ions (H3O)
    from the solution, or donates hydroxide ions
    (OH-) to the solution
  • Both actions will lower the concentration of
    hydronium ions, and thus raise pH
  • an acid donates H3O ions to the solution or
    accepts OH-, thus lowering pH

34
continued
  • base dissociation constant (or Kb) ? a measure of
    basicity
  • pKb is the negative log of Kb and related to the
    pKa by the simple relationship pKa pKb 14
  • Alkalinity is a measure of the ability of a
    solution to neutralize acids to the equivalence
    points of carbonates or bicarbonates

35
Neutralization of acids
  • When dissolved in water, the strong base sodium
    hydroxide decomposes into hydroxide and sodium
    ions
  • NaOH ? Na OH-
  • in water hydrogen chloride forms hydronium and
    chloride ions
  • HCl H2O ? H3O Cl-
  • When the two solutions are mixed, the H3O and
    OH- ions combine to form water molecules

36
continued
  • H3O OH- ? 2 H2O
  • If equal quantities of NaOH and HCl are dissolved
    ? the base and the acid exactly neutralize,
    leaving only NaCl (table salt) in solution

37
Confusion between alkali and base
  • The terms "base" and "alkali" are often used
    interchangeably, since most common bases are
    alkalis
  • . It is common to speak of "measuring the
    alkalinity of soil" when what is actually meant
    is the measurement of the pH (base property). In
    a similar manner, bases that are not alkalis,
    such as ammonia, are sometimes erroneously
    referred to as alkaline
  • not all or even most salts formed by alkali
    metals are alkaline this designation applies
    only to those salts that are basic

38
continued
  • most electropositive metal oxides are basic only
    the soluble alkali metal and alkaline earth metal
    oxides can be correctly called alkalis
  • This definition of an alkali as a basic salt of
    an alkali metal or alkaline earth metal does
    appear to be the most common, based on dictionary
    definitions (however conflicting definitions of
    the term alkali do exist)

39
Weak acid/weak base equilibria
  • In order to lose a proton, it is necessary that
    the pH of the system rise above the pKa of the
    protonated acid
  • decreased concentration of H in that basic
    solution shifts the equilibrium towards the
    conjugate base form (the deprotonated form of the
    acid)

40
continued
  • In lower-pH (more acidic) solutions, there is a
    high enough H concentration in the solution to
    cause the acid to remain in its protonated form,
    or to protonate its conjugate base (the
    deprotonated form)
  • Solutions of weak acids and salts of their
    conjugate bases form buffer solutions

41
The HendersonHasselbalch equation
  • describes the derivation of pH as a measure of
    acidity (using pKa, the acid dissociation
    constant) in biological and chemical systems.
  • also useful for estimating the pH of a buffer
    solution and finding the equilibrium pH in
    acid-base reactions
  • Two equivalent forms of the equation

42
                            and
                                  
43
continued
  • pKa is - log(Ka)
  • where Ka is the acid dissociation constant that
    is

44
continued
  • In these equations
  • A - the ionic form of the relevant acid
  • Bracketed quantities such as base and acid
    denote the molar concentration of the quantity
    enclosed
  • In analogy to the above equations, the following
    equation is valid

45
continued
  • B denotes the salt of the corresponding base B

46
Inorganic buffer
  • A buffer solution an aqueous solution
    consisting of a mixture of a weak acid and its
    conjugate base or a weak base and its conjugate
    acid
  • has the property that the pH of the solution
    changes very little when a small amount of acid
    or base is added to it
  • Buffer solutions are used as a means of keeping
    pH at a nearly constant value in a wide variety
    of chemical applications

47
In a simple buffer solution ? an equilibrium
between a weak acid, HA, and its conjugate base,
A-
  • HA H2O ?? H3O A-

48
continued
  • hydrogen ions are added to the solution ? the
    equilibrium moves to the left (as there are
    hydrogen ions on the right-hand side of the
    equilibrium expression)
  • hydroxide ions are added ? the equilibrium moves
    to the right (as hydrogen ions are removed in the
    reaction H OH- ? H2O)
  • some of the added reagent is consumed in shifting
    the equilibrium and the pH changes by less than
    it would do if the solution were not buffered

49
The acid dissociation constant for a weak acid,
HA, is defined as
50
Simple manipulation with logarithms gives the
Henderson-Hasselbalch equation, which describes
pH in terms of pKa
51
continued
  • A- is the concentration of the conjugate base
  • HA is the concentration of the acid
  • Applies ? when the concentrations of acid and
    conjugate base are equal
  • often described as half-neutralization, pHpKa

52
The same considerations apply to a mixture of a
weak base, B and its conjugate acid BH
  • B H2O ? ? BH OH-

53
continued
  • In general ? a buffer solution may be made up of
    more than one weak acid and its conjugate base
  • if the individual buffer regions overlap a wider
    buffer region is created by mixing the two
    buffering agents

54
Applications
  • resistance to changes in pH ? makes buffer
    solutions very useful for chemical manufacturing
    and essential for many biochemical processes
  • ideal buffer for a particular pH has a pKa equal
    to that pH, since such a solution has maximum
    buffer capacity
  • Buffer solutions are necessary to keep the
    correct pH for enzymes in organisms to work

55
continued
  • Many enzymes work only under very precise
    conditions if the pH strays too far out of the
    margin, the enzymes slow or stop working and can
    denature, thus permanently disabling its
    catalytic activity
  • A buffer of carbonic acid (H2CO3) and bicarbonate
    (HCO3-) is present in blood plasma ? to maintain
    a pH between 7.35 and 7.45

56
Textbook
  • In your text (by Kier and Dowd)
  • Pg 56-65
About PowerShow.com