pH Change

12 Key excerpts on "pH Change"

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  eBook - ePub

  Science in Nursing and Health Care

  • Tony Farine, Mark A. Foss(Authors)
  • 2013(Publication Date)
  • Routledge

    (Publisher)

  When using the pH scale, a number of points must be considered. First, since it is a logarithmic scale, every change of one unit in pH represents a tenfold change in hydrogen ion concentration, a change of two units in pH a 100-fold change in hydrogen ion concentration, and so on. For this reason, the normal range of blood pH (7.35–7.45) is not as narrow as it first appears, and apparently small changes in blood pH represent large changes in hydrogen ion concentration, which you may need to report. Second, the pH scale is a negative scale – that is, a falling pH represents a rise in hydrogen ion concentration and a rising pH represents a falling hydrogen ion concentration.

  Pure water has a pH of 7 and an identical concentration of hydrogen ions and hydroxide ions, and therefore is referred to as neutral . If hydrogen ions are added, then [H+ ] rises and pH falls – that is, acids have a pH of less than 7. In contrast, if hydrogen ions are removed, then [H+ ] falls and pH rises – that is, bases have a pH of greater than 7.

  Salts as acids and bases

  When acids and bases react together, the salt (ionic compound) formed may be neutral, acidic or basic, depending on the strengths of the acid and base used in the reaction. If a strong acid is added to a strong base, or a weak acid is added to a weak base, then the resultant salt is neutral. In contrast, the reaction between a strong acid and a weak base results in the formation of an acidic salt, while the reaction between a weak acid and a strong base produces a basic salt.

  In-text review

    Acids are substances that donate hydrogen ions during a chemical reaction.

    Bases are substances that accept hydrogen ions during a chemical reaction.

    Acids and bases are described as weak or strong, depending upon the extent of their dissociation.

    A solution of an acid or base can be concentrated or dilute, irrespective of whether the acid or base is strong or weak.

    Acids and bases react together to produce a salt and water.

    The concentration of hydrogen ions is described in terms of pH.

  Acid–base balance

  We have already looked at the concept of homeostasis in Chapter 2

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  The Science For Conservators Series

  Volume 2: Cleaning

  • Matthew Cushman, Conservation Unit Museums and Galleries Commission(Authors)
  • 2005(Publication Date)
  • Routledge

    (Publisher)

  tiny concentration it does mean that even the purest of pure water is not, chemically, a single molecular species. Moreover, because the ions are chemically more reactive than their parent molecules, their presence strongly influences the chemical interaction of water with other substances.

  A chemical equilibrium, like other forms of equilibrium or stability, can be upset by suitable external influences. The conditions of acidity and alkalinity are just this. The equilibrium is disturbed so that the concentrations of H3 O+ ions or OH– ions are no longer one ten-millionth of a mole per litre. In acidic solutions the concentration of H3 O+ is increased by hundreds, thousands or millions of times. Alkaline solutions, conversely, have the concentrations of OH– ions dramatically increased. Thus, the chemical behaviour of the solution becomes controlled by the behaviour of these ions. The compounds called acids and alkalies can bring about these remarkable changes in water when they go into solution.

  acidity and alkalinity

  A2  The pH Scale for Hydrogen Ion Concentrations

  The concentration of H3 O+ and OH– ions in pure water is one ten-millionth of a mole per litre. Written as a fraction this is which can be written more compactly as 10–7 , to be read as “ten to the minus seven”.

  The convention for describing numbers like this is simply to count how many noughts there are in the number. Numbers bigger than 1 are given a plus index; thus 1000 is 10+3 , ie “ten to the plus three” (normally just 103 or “ten to the power of three”). Fractions are indicated with a minus index. The fraction is 10–3 , “ten to the minus three”.

  It is long-winded to refer to concentrations in moles per litre when the numbers become awkward mouthfuls like “one ten-millionth” so a shorthand convention based on the “ten-to-the-something” system has been adopted. When used for describing acids and alkalies this is known as the pH scale and describes the concentration of hydrogen ions (more strictly, of H3 O+

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  Analytical Chemistry

  • Clyde Frank(Author)
  • 2012(Publication Date)
  • Academic Press

    (Publisher)

  − ions. Although acidity and basicity of a solution can be expressed in terms of pOH it is customary to use the pH scale.

  Table 8-3 pH-pOH Chart

  It is possible to have negative pH values which means that the H3 O+ concentration is greater than 1 M. However, this is not common since very concentrated solutions of strong acids are not fully dissociated. Second, the approximation that the activity of H3 O+ and concentration of H3 O+ are equal, that is the activity coefficient is 1 (see Chapter 3 , page 27), is no longer valid in concentrated solutions.

  pH OF ACID AND BASE SOLUTIONS

  Strong Acids and Bases

  Strong acids and bases are classified as strong electrolytes, which means that they are virtually 100% dissociated in water. Thus, for a strong acid, such as hydrochloric acid, the concentration of hydronium ion in water is equal to the original analytical concentration of the acid. A strong base solution can be described in the same way. Therefore, the pH for solutions of strong acids and bases are easily calculated since the hydronium or hydroxide ion concentrations are arrived at directly from the analytical concentrations of the strong acid or base.

  Example 8-1

  Calculate the H3 O+ , OH− , pH, and pOH for 100 ml of 0.0250 F HCl solution.

  Example 8-2

  Calculate the pH and pOH for the solution made by mixing 400 ml of water and 200 ml of 0.0500 F NaOH. Assume that the volumes are additive.

  For very dilute solutions, < 10−6 M , of strong acids or bases the ionization of water and its contribution to the equilibrium concentration of [H3 O+ ] and [OH−

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  General Chemistry

  • Linus Pauling(Author)
  • 2014(Publication Date)
  • Dover Publications

    (Publisher)

  A slightly acidic solution, containing 10 times as many hydrogen ions (concentration 10 −6, p H 6), also contains some hydroxide ions, one-tenth as many as a neutral solution. A solution containing 100 times as much hydrogen ion as a neutral solution (concentration 10 −5, p H 5) contains a smaller amount of hydroxide ion, one one-hundredth as much as a neutral solution; and so on. A solution containing 1 mole of strong acid per liter has hydrogen-ion concentration 1, and p H 0; such a strongly acidic solution also contains some hydroxide ion, the concentration of hydroxide ion being 1 × 10 −14. Although this is a very small number, it still represents a large number of actual ions in a macroscopic volume. Avogadro’s number is 0.602 × 10 24, and accordingly a concentration of 10 −14 moles per liter corresponds to 0.602 × 10 10 ions per liter, or 0.602 × 10 7 ions per milliliter. 14-3. Indicators Indicators such as litmus may be used to tell whether a solution is acidic, neutral, or basic. The change in color of an indicator as the p H of the solution changes is not sharp, but extends over a range of one or two p H units. This is the result of the existence of chemical equilibrium between the two differently colored forms of the indicator, and the dependence of the color on the hydrogen-ion concentration is due to the participation of hydrogen ion in the equilibrium. Thus the red form of litmus may be represented by the formula HIn and the blue form by In −, resulting from the dissociation reaction In alkaline solutions, with [H + ] very small, the equilibrium is shifted to the right, and the indicator is converted almost entirely into the basic form (blue for litmus). In acidic solutions, with [H + ] large, the equilibrium is shifted to the left, and the indicator assumes the acidic form. Let us calculate the relative amount of the two forms as a function of [H + ]

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  Soil Science

  Methods & Applications

  • David L. Rowell(Author)
  • 2014(Publication Date)
  • Routledge

    (Publisher)

  Section 8.1 explains the terms involved. It should be noted that when applied to soils, ‘neutral’ is given a slightly different meaning, being a range from about pH 6.5 to 7.

  Soil acidity involves more than just the pH of the soil solution. This is still the main principle and the measurement of soil pH (Section 8.1 ) is normally made in a suspension of soil in water such that the value obtained is primarily related to the solution pH. However, hydrogen ions are also present on cation exchange sites and have an effect on the measurement. Also as soils become more acidic (pH 7 → 3), there are associated changes in the following properties:

  •  The amounts of exchangeable Ca2+ and Mg2+ decrease. These together with exchangeable K+ , Na2+ and are known as the basic cations: their total amount is often expressed as a percentage of the CEC which is termed the percentage base saturation (Section 8.2 ).

  •  The amount of exchangeable Al3+ increases and is often expressed as the percentage aluminium saturation of the ECEC (Section 8.2 ).

  •  The negative charge on humus decreases and the positive charge on sesquioxides increases (Sections 7.1 and 7.5 ).

  •  The availability of plant nutrients is changed. For example, phosphate solubility is reduced (Ch. 10 ).

  •  The availability of toxic elements is changed. For example, aluminium and manganese become more soluble in acid soils (Section 8.3 ).

  •  The activity of many soil organisms is reduced resulting in an accumulation of organic matter, reduced mineralization and a lower availability of N, P and S. THE DEVELOPMENT OF SOIL ACIDITY

  In pure water the concentration of H+ ions is 10−7 mol 1−1 and the pH is 7. When in contact with the atmospheric concentration of CO2 a dilute carbonic acid solution is formed with a pH of 5.6. Distilled or deionized water in the laboratory therefore has a pH of about 5.6. For the pH to differ from this value some other acid or base must be present. Thus ‘acid rain’ contains nitric and sulphuric acid dissolved from the atmosphere (or ammonia and oxides of N and S which can form these acids). Its pH is below 5.6; the average pH of rain over eastern Britain is about 4.4 (DOE, 1990). Even in unpolluted air rain picks up small amounts of naturally occurring acid and has a pH of about 5. Ammonia and oxides of N and S are also deposited dry on vegetation and soil and are washed into the soil by rain where they produce acidity. Thus the atmosphere is an external source of acidity (Fig. 8.1

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  Basic Physical Chemistry for the Atmospheric Sciences

  • Peter V. Hobbs(Author)
  • 2000(Publication Date)
  • Cambridge University Press

    (Publisher)

  We see from definition (5.14) that (1) the greater the hydrogen ion concentration (i.e., the more acidic the solution) the smaller is the pH value of the solution, and (2) a change in the hydrogen ion concentration by a factor of ten (e.g., from 10 1 to 10~ 2 M) changes the pH value by unity. At the beginning of this section we defined a solution as being neutral if [H + (aq)] = [OH(aq)]. Pure water is neutral; therefore, from Eqs. (5.12) and (5.13) [H 3 O + (aq)][OH-(aq)] = l(r 14 or, [H 3 O + (aq)] 2 =10-Therefore, for pure water [H 3 O + (aq)] = [H + (aq)] = 10 7 M Hence, the pH of pure water is -log(10~ 7 ) = 7. It follows that acidic solu-tions have pH < 7 and basic solutions have pH > 7. Observed pH values in nature are generally between 4 and 9. Seawater normally has a pH between 8.1 and 8.3. Streams in wet climates generally have a pH between 5 and 6.5 and in dry climates between 7 and 8. Soil water in the presence of abundant decaying vegetation may have a pH of 4 or lower. The pH of rainwater can range from quite acidic (around 4.0) in industrial regions to about 5.6 in very clean regions. We will discuss the acidity of rainwater in some detail at the end of this chapter, but the following exercise illustrates why even clean rainwater does not have a pH of 7. Exercise 5.2. The pH of natural rainwater is about 5.6. Assum-ing that all of this acidity is due to the absorption of CO 2 by the rain, determine how many moles of CO 2 would have to be absorbed in 1L of rainwater. Solution. Since the pH of rainwater is 5.6, the concentration of H 3 O + (aq) in natural rainwater is given by pH = 5.6 = -log[H 3 O + (aq)] Therefore, [H 3 O + (aq)] = 0.25xl0-5 M 90 Acids and bases The main source of H 3 O + (aq) when CO 2 dissolves in water is CO 2 (g) + H 2 O(l)+±H 2 CO 3 (aq) (5.15a) H 2 CO 3 (aq) + H 2 O(1) ?± HCO 3 -(aq) + H 3 O + (aq) (5.15b) We see from Reactions (5.15) that for every mole of CO 2 that is absorbed in water, one mole of H 3 O + (aq) is produced.

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  Visualizing Everyday Chemistry

  • Douglas P. Heller, Carl H. Snyder(Authors)
  • 2015(Publication Date)
  • Wiley

    (Publisher)

  [H 3 O + ] [H 3 O + ] A basic solution has a pH greater than 7. 10 –14 10 –13 10 –12 10 –11 10 –10 10 –9 10 –8 10 –7 10 –6 10 –5 10 –4 10 –3 10 –2 10 –1 10 0 10 –14 10 –13 10 –12 10 –11 10 –10 10 –9 10 –8 10 –7 10 –6 10 –5 10 –4 10 –3 10 –2 10 –1 10 0 Low concentration High concentration Low concentration High concentration pH = 3 Here, [H 3 O + ] = 10 –3 , so pH = 3. [H 3 O + ] [H 3 O + ] [OH – ] An acidic solution has a pH less than 7. 10 –14 10 –13 10 –12 10 –11 10 –10 10 –9 10 –8 10 –7 10 –6 10 –5 10 –4 10 –3 10 –2 10 –1 10 0 10 –14 10 –13 10 –12 10 –11 10 –10 10 –9 10 –8 10 –7 10 –6 10 –5 10 –4 10 –3 10 –2 10 –1 10 0 Low concentration High concentration Low concentration High concentration [OH – ] Any increase in the hydronium ion concentration above 10 −7 (with an associated decrease in the hydroxide ion concentration) produces an acidic solution: Any decrease in the hydronium ion concentration below 10 −7 (with an associated increase in the hydroxide ion concentration) produces a basic solution: The pH Scale 245 increases [OH − ] and lowers [H 3 O + ]. We can see this see- saw effect and its relation to pH in Figure 8.11 In Words, Math, and Pictures. To summarize, at 25°C: • The pH of a neutral solution equals 7. • The pH of an acidic solution is less than 7. • The pH of a basic solution is greater than 7. We can see the relationship between hydronium ion con- centration and pH in Figure 8.12. We can increase the H 3 O + concentration of a solution by adding acid, such as HCl. Conversely, we can increase the OH − concentration by adding base, such as NaOH. Furthermore, the value obtained by multiplying the hy- dronium ion concentration of a solution, [H 3 O + ], by its hydroxide ion concentration, [OH − ], is always constant, regardless of the addition of acid or base to the water. Since the value of this product remains fixed at any spe- cific temperature, adding acid not only increases [H 3 O + ] but lowers [OH − ] as well.

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  General, Organic, and Biological Chemistry

  An Integrated Approach

  • Kenneth W. Raymond(Author)
  • 2012(Publication Date)
  • Wiley

    (Publisher)

  Increasing the concentration of a reactant will, for example, cause a net forward reaction to occur until equilibrium is reestablished. 4. Use H 3 O + concentration In a neutral solution, [H 3 O + ] = 1 * 10 -7 M and 7.5, 7.6–7.8 7.43–7.70 and pH to identify a pH = 7, in an acidic solution [H 3 O + ] 7 1 * 10 -7 M 7.6 solution as being acidic, and pH 6 7, and in a basic solution basic, or neutral. [H 3 O + ] 6 1 * 10 -7 M and pH 6 7. 5. Use K a and pK a values to The larger the value of K a or the smaller the value 7.7 7.9 7.71–7.80 assess the relative strength of pK a , the stronger an acid. A strong acid has a of acids and state the weak conjugate base and a weak acid has a strong relationship between the conjugate base. strength of an acid and the strength of its conjugate base. 6. Describe the processes of Neutralization involves reacting an acid with a base 7.8 7.10 7.81–7.90 neutralization and titration. to produce water, a salt, and a neutral solution. The concentration of an unknown acidic solution can be determined through titration, in which just enough of a basic solution of known concentration is added to neutralize the acid. 7. Explain how the pH of a In a solution containing an acid (HA) and its conjugate 7.9, 7.11, 7.12 7.91–7.104 solution can affect the relative base (A - ), when pH = pK a , [HA] = [A - ]. When 7.10 concentrations of an acid and pH 6 pK a , [HA] 7 [A - ] and when pH 7 pK a , its conjugate base and [HA] 6 [A - ]. Buffers, solutions that resist changes in describe buffers. pH when small amounts of acid or base are added, can be made by combining a weak acid and its conjugate base. Buffers are most effective when the pH is within one unit of the pK a for the weak acid. 8. Describe the role of buffers, Carbonic acid and hydrogen carbonate ion make up the 7.11 7.13 7.105–7.112 respiration, and the kidneys most important of the blood serum buffers. in maintaining a stable H 2 CO 3 + H 2 O N HCO 3 - + H 3 O + blood serum pH.

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  Analytical Chemistry

  • Gary D. Christian, Purnendu K. Dasgupta, Kevin A. Schug(Authors)
  • 2013(Publication Date)
  • Wiley

    (Publisher)

  ranges, from 1 M or greater to 10 −14 M or less. To construct a plot of H + concentration against some variable would be very difficult if the concentration changed from, say, 10 −1 M to 10 −13 M. This range is common in a titration. It is more convenient to 228 CHAPTER 7 ACID–BASE EQUILIBRIA compress the acidity scale by placing it on a logarithm basis. The pH of a solution was defined by Sørenson as pH = −log[H + ] (7.15) The minus sign is used because most of the concentrations encountered are less than pH is really –log a H + . This is what a pH meter (glass electrode) measures—see Chapter 13. 1 M, and so this designation gives a positive number. (More strictly, pH is now defined as − log a H + , but we will use the simpler definition of Equation 7.15.) In general, pAnything = − log Anything, and this method of notation will be used later for other numbers that can vary by large amounts, or are very large or small (e.g., equilibrium Carlsberg Laboratory archives In 1909, Søren Sørenson, head of the chemistry department at Carlsberg Laboratory (Carlsberg Brewery) invented the term pH to describe this effect and defined it as −log[H + ]. The term pH refers simply to “the power of hydrogen.” In 1924, he realized that the pH of a solution is a function of the “activity” of the H + ion, and published a second paper on the subject, defining it as pH = − log a H + . constants). Example 7.2 Calculate the pH of a 2.0 × 10 −3 M solution of HCl. Solution HCl is completely ionized, so [H + ] = 2.0 × 10 −3 M pH = − log(2.0 × 10 −3 ) = 3 − log 2.0 = 3 − 0.30 = 2.70  A similar definition is made for the hydroxide ion concentration: pOH = −log[OH − ] (7.16) Equation 7.13 can be used to calculate the hydroxyl ion concentration if the A 1 M HCl solution has a pH of 0 and pOH of 14. A 1 M NaOH solution has a pH of 14 and a pOH of 0. hydrogen ion concentration is known, and vice versa.

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  Basic Chemistry Concepts and Exercises

  • John Kenkel(Author)
  • 2010(Publication Date)
  • CRC Press

    (Publisher)

  The pH electrode is a small glass or plastic enclosure—the size of a small test tube—that is electrically sensitive to the hydrogen ions in a solution . When this electrode is immersed into the solution to be tested, an electrical potential or voltage is produced across the glass membrane at the tip that separates the interior of the electrode from the solution . This voltage can be measured with a simple voltage measuring device, a voltmeter . The concentration of the hydrogen ions in the solution determines the voltage level that develops, and thus this voltage is proportional to [H + ] . Since pH is also proportional to [H + ], it is possible to relate this voltage directly to the pH of the solution . The pH meter is a voltmeter that utilizes an electrical circuit that automatically con-verts the voltage level to pH and displays the pH on its readout . Thus, simply stated, the pH of a solution can be measured or monitored by immersing this pH probe, connected to the pH meter, into the solution of interest and reading the pH on the readout of the meter . Since the meter is a sort of modified voltmeter, the measurement is a rela-tive measurement . This means that the voltage developed at the tip of the probe is measured relative to some other voltage . This other voltage is the 319 Acids, Bases, and pH voltage of a reference electrode . Modern pH probes have a reference elec-trode built in and are called combination pH electrodes . The interior design of these is beyond our scope . 12.6 Theories of Acids and Bases 12.6.1 Arrhenius Theory The view of acids and bases as we have been discussing them thus far is know as the Arrhenius theory , named after Svante Arrhenius, a Swedish scientist who introduced this theory in 1884 . According to this theory, an acid is defined as a substance that releases hydrogen ions when dissolved in water . A base is defined as a substance that releases hydroxide ions when dissolved in water .

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  Chemical Processes for Pollution Prevention and Control

  • Paul Mac Berthouex, Linfield C. Brown(Authors)
  • 2017(Publication Date)
  • CRC Press

    (Publisher)

  − ], and vice versa. For example,

  [H+ ] = 10−2 and [OH− ] = 10−12

  [H+ ] = 10−3 and [OH− ] = 10−11

  [H+ ] = 10−9 and [OH− ] = 10−5

  where [H+ ] = hydrogen ion concentration (mol/L)

        [OH− ] = hydroxyl ion concentration (mol/L)

  KW

  is temperature dependent, as shown in Table 7.1 .

  TABLE 7.1The Ionization of Water (

  KW

  ) Is a Function of Temperature

  7.4   pH

  pH is a fundamental characteristic of chemical solutions. It determines the extent of ionization of soluble compounds and the formation of solids by ions that tend to precipitate. Low pH indicates acidic conditions.

  The definition of pH is

  pH = −log10 [H+ ]

  This convenient “p” notation saves us having to write the concentrations as exponentials. The “p” comes from “potential,” and the formula −log10 [X ] is used in other cases where logarithmic values are easier to understand.

  Taking negative logarithms of the equilibrium expression for water,

  KW = [H+ ][OH− ] = 10−14
  gives	−log10 KW = −log10 [H+ ]−log10 [OH− ] = 14
  and	p KW = pH + pOH = 14

  The pH scale for aqueous solutions runs from 0 to 14. A change of one pH unit represents a 10-fold change in [H+].

  At	pH = 2	[H+] = 10−2 and [OH− ] = 10−12
  pH = 3	[H+ ] = 10−3 and [OH− ] = 10−11
  pH = 9	[H+ ] = 10−9 and [OH− ] = 10−5

  Table 7.2 shows how

  KW

  and the neutral pH depend on temperature. Even though the pH of pure water changes with temperature, it is still neutral. In pure water, there will always be the same number of hydrogen ions and hydroxide ions. That means that the pure water remains neutral even if its pH Changes.

  We are strongly conditioned to the idea of 7 being the pH of pure water, and anything else feels strange. Remember that if the value of

  KW

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  Foundations of Chemistry

  An Introductory Course for Science Students

  • Philippa B. Cranwell, Elizabeth M. Page(Authors)
  • 2021(Publication Date)
  • Wiley

    (Publisher)

  • A reversible chemical reaction reaches equilibrium when the rates of the forward and backward reactions become equal. • A chemical equilibrium is a dynamic equilibrium because reaction con-tinues to occur in the forward and backward directions. • The equilibrium constant for a reaction, K c , relates the concentrations of products at equilibrium to the concentrations of reactants. • Le Châtelier ’ s principle states that when a reaction at equilibrium is sub-jected to a change, the equilibrium shifts to minimise the change. • Changes in temperature, pressure, and concentration affect the position of equilibrium, but only a change in temperature affects the equilibrium constant. • The equilibrium constant, K c , can be written in terms of concentrations of reactants and products and K p can be written for gaseous reactions in terms of partial pressures of reactants and products. • A homogeneous equilibrium is one in which the reactants and products are in the same physical state whereas a heterogeneous reaction has reactants and products in more than one physical state. • The Brønsted – Lowry theory of acids and bases defines an acid as a proton donor and a base as a proton acceptor. • A strong acid is totally ionised (or dissociated) in aqueous solution, whereas a weak acid is only partially ionised (or dissociated). • The pH of a solution is a measure of the acidity of the solution and is expressed on a scale of 0 to 14. The expression for pH = − log 10 [H + ]. • The acid dissociation constant, K a , represents the equilibrium constant for the dissociation of a weak acid. • Acid strengths are measured by p K a values, where p K a = − log 10 [ K a ]. 250 Chemical equilibrium and acid-base equilibrium • The base dissociation constant, K b , represents the equilibrium constant for the dissociation of a weak base. • Base strengths are measured by p K b values, where p K b = − log 10 [ K b ]. • A conjugate acid and base are linked by the gain and loss of a proton.

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