pH and Solubility

11 Key excerpts on "pH and Solubility"

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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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  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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  Elementary Medical Biochemistry

  • J. M. M. Brown, G. G. Járos(Authors)
  • 2013(Publication Date)
  • Butterworth-Heinemann

    (Publisher)

  If the expression (2) is to hold, then the concentration of hydroxyl ions in these solu-tions of acids must decrease correspondingly below the value of 10~ 7 mole/litre. A solution in which the concentration of hydrogen ions is in excess of 1 0 -7 mole/litre is known as an acidic solution. Conversely if the concentration of hydroxyl ions is greater than 10~ 7 mole/litre (which means that [H + ] will be less than this value) we speak of an alkaline solution. The numerical values used above, e.g. 1 0 -7 mole/litre, are cumber-some when it comes to expressing the hydrogen ion concentration, for instance the expression [H + ] = 10~ 7 mole/litre, reads, the hy-drogen ion concentration equals ten to the power of minus seven mole/litre. The pH system is used to simplify expression of state-ments like the above one to a remarkable degree. We merely replace the previous statement by the simple expression pH = 7 where pH is taken to mean the hydrogen ion concentration and seven is the negative power often in the term 10~ 7 . It is this power which is the most meaningful in expressions of this nature, so we use it as a whole number, whilst bearing in mind what we mean by it. By the same token, [H + ] = 10~ 3 or [H + ] = 1 0 ~ 12 can be written as pH = 3 or pH = 12 respectively. We must bear in mind that 10~ 3 (or 1/1 000) is a larger number than 10~ 4 (or 1/10 000). In other words the smaller the value of the negative power, the larger is the number, the smaller the pH value, the greater is the hydrogen ion concentration and vice versa. A pH of less than 7 indicates an acidic solution, while a pH of greater than 7 indicates an alkaline solution. A pH of 7 indicates neutrality in the absolute sense. 5: 12 Elementary Medical Biochemistry When we consider the hydrogen ion concentration of body fluids, neutrality of these does not coincide with the absolute value of 7, but rather with the pH values 7,3 - 7,4 i.e.

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  Chemistry and the Environment

  • Sven E. Harnung, Matthew S. Johnson(Authors)
  • 2012(Publication Date)
  • Cambridge University Press

    (Publisher)

  However, since then the usage has changed. Nowadays, the pH of a solution is called its acidity or degree of acidity , and, with the appearance of environmental chemistry, the word alkalinity has been given its present meaning, namely, the definition in Equation 5.24 . Buffer value Definition . The buffer value of an aqueous solution is the derivative of the alkalinity with respect to pH. 178 Chemistry of the hydrosphere From Equation 5.24 , one obtains β def = d A d pH (5.25a) = c α 0 α 1 + [H + ] + [HO − ] × ln 10 (5.25b) Further differentiation shows that d β/ dpH = 0 for α 1 = α 0 , implying that the buffer value has an extremum (more precisely, it is a maximum) for α 1 = α 0 = 1/2 with β = ( c / 4) ln 10 ≈ 0.58 c . This happens at pH = p K a . In general, at pH values where | pH − p K a < 1, the buffer equation pH ≈ p K a + lg c KA c HA (5.26) gives a good approximation of pH (see Equation 5.18 ). In medicinal and biochemical literature, this equation is sometimes called the Henderson-Hasselbalch equation (Henderson, 1908, and Hasselbalch, 1927). Note that the buffer value is also large at low and high pH values, because it is proportional to the hydron concentration as well as the hydroxide concentration. Example 1. A conjugate acid-base pair The basic properties of a conjugate acid-base pair in aqueous solution are shown in Figure 5.2 with acetic acid–acetate as the example. The experimental data c , p K a , and p K w are given in the legend. A titration of 0.10 m acetate with strong acid starts at pH = 8.88, A (8.88) = 0.10 m , and the end point occurs at pH = 2.88, A (2.88) = 0.00 m . Negative values of A mean that the end point has been exceeded. Example 2. Several conjugate acid-base pairs Figure 5.3 shows an acid-base titration. The titrand is a basic solution of ammonia and acetate and the titrant (titrator) is a strong acid. Consider some characteristic values of the alkalinity as a function of pH, A (pH).

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  Chemistry

  The Molecular Nature of Matter

  • James E. Brady, Neil D. Jespersen, Alison Hyslop(Authors)
  • 2014(Publication Date)
  • Wiley

    (Publisher)

  According to the color code, the pH of the solution is closer to 3 than to the color for pH 5. Andy Washnik 806 Chapter 16 | Acid–Base Equilibria in Aqueous Solutions | Summary Organized by Learning Objective Define pH and explain the use of “p” notation Water reacts with itself to produce small amounts of H 3 O + (often abbreviated H + ) and OH - ions. The concentrations of these ions, both in pure water and dilute aqueous solutions, are related by the expression 3 H + 4 3 OH - 4 = K w = 1.0 Ž 10 -14 (at 25 °C) K w is the ion product constant of water. In pure water 3 H + 4 = 3 OH - 4 = 1.0 Ž 10 -7 The pH of a solution is defined by the equation, pH = -log[H + ]. As the pH decreases, the acidity, or [H + ], increases. The compa- rable term, pOH (= -log[OH - ]), is used to describe a solution that is basic. A solution is acidic if the hydrogen ion concentration exceeds 1.0 Ž 10 -7 or the pH is less than 7.00. Similarly, a solution is basic if the hydroxide ion concentration exceeds 1.0 Ž 10 -7 or if the pH is greater than 7.00. Explain how to determine the pH of strong acids or bases in aqueous solution When calculating the pH of strong acids or strong bases, we assume that they are 100% ionized. Write expressions for the acid ionization constant, K a , and base ionization constant, K b , and explain how they are related to each other A weak acid H A ionizes according to the general equation H A + H 2 O m H 3 O + + A - or more simply, H A mH + + A - The equilibrium constant is called the acid ionization con- stant, K a : K a = 3 H 3 O + 4 3 A - 4 3 H A 4 A weak base B ionizes by the general equation B + H 2 O mBH + + OH - The equilibrium constant is called the base ionization con- stant, K b : K b = 3 BH + 4 3 OH - 4 3 B 4 The smaller the values of K a (or K b ), the weaker are the sub- stances as Brønsted acids (or bases).

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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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  Introduction to the Chemistry of Food

  • Michael Zeece(Author)
  • 2020(Publication Date)
  • Academic Press

    (Publisher)

  Writing numbers this way is tedious and leads to mistakes (e.g., adding or leaving out a zero) that are avoided by using exponential (scientific) notation as shown below. 0.000000 1 M H 3 O + = 1.0 × 10 − 7 M H 3 O + pH and measuring acidity pH is the term used to express the measurement of hydronium ion concentration in solution. The chemical definition of pH is given as “the negative log of the hydronium concentration” and corresponds to the following equation. pH = − Log [ H 3 O + ] This equation of acidity is very useful and often employed to find the pH of materials (in solution) including water, soil, and food. The hydronium ion concentration of pure water at 25 °C is 1.0 × 10 − 7 M. Using the equation for pH, the value for water under these conditions is 7. pH = − Log [ 1.0 × 1 0 − 7 M ] pH = − (− 7) pH = 7 This equation can be used to find the pH of soda pop and egg white, as shown below. Soda pop contains phosphoric acid and is an acidic food. The hydronium ion concentration of a typical pop is 1.0 × 10 − 4 M, therefore its pH is: pH = − Log (1.0 × 1 0 − 4 M) pH = − (− 4) pH = 4 Egg white, in contrast to soda pop, represents one of the few alkaline foods. This means it contains a enough base to raise its pH above 7. The hydronium ion concentration of egg white is 1.0 × 10 − 8 M. Therefore, its pH is: pH = − Log (1.0 × 1 0 − 8 M) pH = − (− 8) pH = 8 The pH scale ranges from 1 to 14, with pH 1 being very acidic and pH 14 being strongly basic (alkaline). One of the only places to find a substance with a pH of 1 is inside the stomach. Its digestive fluid contains HCl and the pepsin enzyme essential to digestion. The highly acidic environment denatures proteins and makes their breakdown by pepsin more effective

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  Aquaculture Engineering

  • Odd-Ivar Lekang(Author)
  • 2019(Publication Date)
  • Wiley-Blackwell

    (Publisher)

  5 Controlling pH, Alkalinity and Hardness

  5.1 Introduction

  At some sites the freshwater pH is too low or too high to achieve optimal growth for fish or shellfish. At other sites, the buffering capacity of the water is low, and it is difficult to avoid pH fluctuations in the water. This again results in suboptimal growth. Sites where acid rain is a problem are particularly exposed to low pH. Further, in reuse systems (Chapter 14 ), biological filters (nitrification) are used to remove ammonia. This causes a drop in pH that must be corrected to maintain optimal growing conditions for fish. Of this reason technology and chemicals must be used for pH regulation.

  The terms pH, alkalinity and hardness of water are related to each other and often affect each other. However the chemistry behind these terms can be quite difficult to understand because a number of chemical interactions are involved and it also include some not basic mathematics. Typically are adapted computer software used when calculations are necessary within this area, but rough estimates can be achieved by using some basic calculations and diagrams described later in this chapter. If an acid, or actually H+ , is added to water, both pH and alkalinity will be affected, and if the alkalinity of the water is not known, it is impossible to know what is going to happen with the pH. Will there be a drop in pH, a drop in alkalinity or both? The aim of the chapter is to give a basic understanding of the concepts including some basic chemistry. To understand pH and how to regulate and control pH and alkalinity in the water, it is of major importance to understand the carbonate systems. This is related to the CO2 gas in the water. In air atmosphere, CO2 can be measured directly, and it is not mixed with other gases. In water it can also be measured, but here CO2 will react with substances in the water. How the reactions occur and which reactions to occur depend on the pH. It can also be said that it is the carbonate system that controls the pH in natural water. A lot of literature is available about these topics including general water chemistry (for example Refs. 1 –4 , those are also base for general values in the chapter) and specifically for aquaculture.

  5 ,6

  pH regulation of water is also performed in water/wastewater treatment and literature from here can also be of good help.7

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  Water Chemistry

  Green Science and Technology of Nature's Most Renewable Resource

  • Stanley E. Manahan(Author)
  • 2010(Publication Date)
  • CRC Press

    (Publisher)

  56 Water Chemistry 3.4 ALKALINITY The.capacity.of.water.to.accept.H + .ions.(protons).is.called. alkalinity . .Alkalinity.is. important.in.water.treatment.and.in.the.chemistry.and.biology.of.natural.waters . . Frequently,. the. alkalinity. of. water. must. be. known. to. calculate. the. quantities. of. chemicals.to.be.added.in.treating.the.water . .Highly.alkaline.water.often.has.a.high. pH.and.generally.contains.elevated.levels.of.dissolved.solids . .These.characteristics. may.be.detrimental.for.water.to.be.used.in.boilers,.food.processing,.and.municipal. water.systems . .Alkalinity.serves.as.a.pH.buffer.and.reservoir.for.inorganic.carbon,. thus. helping. to. determine. the. ability. of. water. to. support. algal. growth. and. other. aquatic.life,.so.alkalinity.can.be.used.as.a.measure.of.water.fertility . .Generally,.the. basic.species.responsible.for.alkalinity.in.water.are.bicarbonate.ion,.carbonate.ion,. and.hydroxide.ion: . HCO CO 3 2 -+ + H H O + arrowpairrightleft 2 . (3.21) . CO CO 3 3 2 --+ H H + arrowpairrightleft . (3.22) . OH O -+ H H 2 + arrowpairrightleft . (3.23) Other,.usually.minor,.contributors.to.alkalinity.are.ammonia.and.the.conjugate. bases.of.phosphoric,.silicic,.boric,.and.organic.acids . At.pH.values.below.7,.[H + ].in.water.detracts.significantly.from.alkalinity,.and.its. concentration.must.be.subtracted.in.computing.the.total.alkalinity . .Therefore,.the. following.equation.is.the.complete.equation.for.alkalinity.in.a.medium.where.the. only.contributors.to.it.are. HCO 3 − ,. CO 3 2 − ,.and.OH − : . [alk] CO = + + ----+ [ ] [ ] [ ] [ ] H CO OH H 3 3 2 2 . (3 .24) Alkalinity.generally.is.expressed.as. phenolphthalein alkalinity ,.corresponding.to. titration.with.acid.to.the.pH.at.which. HCO 3 − .is.the.predominant.carbonate.species. (pH.8 .3), .or. total alkalinity ,.corresponding.to.titration.with.acid.to.the.methyl.orange. endpoint. (pH. 4 .3), . where. both. bicarbonate. and. carbonate. species. have. been. con-verted.into.CO 2 .

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  Fundamentals of Sustainable Chemical Science

  • Stanley E. Manahan(Author)
  • 2009(Publication Date)
  • CRC Press

    (Publisher)

  the desired solution containing 20 mg of calcium per liter. Molar Concentration of H Molar Concentration of H Ion and pH Ion and pH Concentrations are important in expressing the degree to which solutions are Concentrations are important in expressing the degree to which solutions are acidic or basic. Recall that acidic or basic. Recall that acids acids , such as HCl and H , such as HCl and H 2 SO SO 4 , produce H , produce H ion, whereas ion, whereas bases bases , such as sodium hydroxide and calcium hydroxide (NaOH and Ca(OH) , such as sodium hydroxide and calcium hydroxide (NaOH and Ca(OH) 2 , , respectively), produce hydroxide ion, OH respectively), produce hydroxide ion, OH . Molar concentrations of hydrogen ion, . Molar concentrations of hydrogen ion, [H [H ], range over many orders of magnitude and are conveniently expressed by pH ], range over many orders of magnitude and are conveniently expressed by pH defined as defined as pH pH log[H log[H ] (7.5.11) ] (7.5.11) 256 Fundamentals of Sustainable Chemical Science 256 Fundamentals of Sustainable Chemical Science In absolutely pure water the value of [H In absolutely pure water the value of [H ] is exactly 1 ] is exactly 1 10 10 7 mol/L, the pH is 7.00, mol/L, the pH is 7.00, and the solution is and the solution is neutral neutral (neither acidic nor basic). (neither acidic nor basic). Acidic Acidic solutions have pH values solutions have pH values of less than 7 and of less than 7 and basic basic solutions have pH values of greater than 7. solutions have pH values of greater than 7. Solubility Solubility If a chemist were to attempt to make a solution containing calcium ion by dis-If a chemist were to attempt to make a solution containing calcium ion by dis-solving calcium carbonate, CaCO solving calcium carbonate, CaCO 3 in water, not much of anything would happen.

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  Chemistry

  • John A. Olmsted, Gregory M. Williams, Robert C. Burk(Authors)
  • 2020(Publication Date)
  • Wiley

    (Publisher)

  The figure shows its colour behaviour. Does the Result Make Sense? We see that x is small relative to 10 −2 , and an acidic pH at the stoichio- metric point is what we expect for this titration of a weak base by a strong acid. Practice Exercise 16.15 A chemist needs to titrate a solution of hydrofluoric acid that is approximately 0.15 M. What indi- cator would be appropriate for the titration? NEED MORE PRACTICE? Try Problems 16.27 and 16.28 16.4 Solubility Equilibria The limestone deposits that decorate many caverns are the result of the solubility equilibrium of calcium carbonate in groundwater, as described in Chapter 14: CaCO 3 (s) ⟶ ⟵ Ca 2+ (aq) + CO 3 2− (aq) K sp = [Ca 2+ ] eq [CO 3 2− ] eq Solubility equilibria are important in the natural environment and in our daily lives. Coral reefs are built up over many years through deposits of calcium carbonate from tiny sea organ- isms. Thermal vents on the ocean’s floor spew out insoluble metal sulphides and other miner- als. Deposits from hard water can wreak havoc with plumbing. 16.4 Solubility Equilibria 805 R. Mclntyre/ShutterStock.com Ralph White/Getty Images Martyn F. Chillmaid/Science Source Galyna Andrushko/123 RF In an aqueous solubility reaction, a salt dissolves to yield ions in solution. The amount of a salt that dissolves in water varies over a large range, as the following examples show: AgI(s) ⟵ ⟶ Ag + (aq) + I − (aq) K sp = 8.5 × 10 −17 Cu(OH) 2 (s) ⟵ ⟶ Cu 2+ (aq) + 2 OH − (aq) K sp = 1.1 × 10 −15 CaCO 3 (s) ⟵ ⟶ Ca 2+ (aq) + CO 3 2− (aq) K sp = 3.36 × 10 −9 Ag 2 SO 4 (s) ⟵ ⟶ 2Ag + (aq) + SO 4 2− (aq) K sp = 1.2 × 10 −5 NaCl(s) ⟵ ⟶ Na + (aq) + Cl − (aq) K sp = 38.4 Substances that have K sp < 10 −5 , such as calcium carbonate, copper(II) hydroxide, and silver iodide, are commonly classified as insoluble (even though they are not totally insoluble).

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