Chemistry
pH Scale
The pH scale is a measure of the acidity or alkalinity of a solution, ranging from 0 to 14. A pH of 7 is considered neutral, while lower values indicate acidity and higher values indicate alkalinity. The scale is logarithmic, meaning each unit change represents a tenfold difference in acidity or alkalinity.
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10 Key excerpts on "pH Scale"
eBook - PDF- Saeed Farrokhpay(Author)
- 2023(Publication Date)
- Arcler Press(Publisher)
The pH of a solution is a measure of how basic or acidic it is. H+ activity in a solution is measured using this method, which is represented as a negative logarithm in the equation. The pH readings are presented on a scale ranging from 0.0 to 14.0 (neutral). Pure water contains a pH of 7.0 and is considered neutral; water having a pH < 7.0 is considered acidic, while water having a pH >7.0 is considered basic or alkaline. Conditions Monitoring pH and Alkalinity of Water 89 containing pH levels ranging from around 6.5 to 8.5 are preferred by most estuary species (Bae et al., 2020; Boyd, 2020). pH values are expressed on a logarithmic scale, which means that for every one-unit change in pH, acidity or alkalinity increases or decreases by a factor of ten; for example, a pH of 5.0 is 10 times more acidic than a pH of 6.0 and 100 times more acidic than a pH of 7.0. The pH of a solution is 7 when the hydroxyl and hydrogen ions are mixed in sufficient amounts (this is known as the neutral point). 3.3. THE ROLE OF PH IN THE ESTUARINE ECOSYSTEM It is possible for plants and animals to change the pH of water by photosynthesis and respiration, as well as by the minerals dispersed in the dust, aerosols, and water from the human-made pollutants and air. Human activities that result in substantial, long-term acidification of a waterbody or short-term variations in pH are extremely detrimental to the environment. For example, algal blooms, which are frequently triggered by an excess of nutrients, may affect pH levels to change substantially over a short period of time, putting a significant amount of stress on nearby species. It is possible that acid precipitation in the higher freshwater goes of an estuary will reduce the existing rate of eggs laid by breeding fish in that area (Figure 3.1) (Jarvis et al., 2006; Boyd et al., 2011). Figure 3.1. Estuarine ecosystem. Source: https://unacademy.com/lesson/estuarine-ecosystem/Y1FL6LFP.
eBook - ePubSoil 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 ACIDITYIn 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
eBook - ePub- Jeffrey Gaffney, Nancy Marley(Authors)
- 2017(Publication Date)
- Elsevier(Publisher)
+ ]. Although theoretically the pH Scale is open-ended, most pH values fall in the range from 0 to 14.Table 5.3 Molar Hydronium and Hydroxide Ion Concentrations for pH Values From 0 to 14pH [H+ ] [OH− ] 0 1.0 0.00000000000001 1 0.1 0.0000000000001 2 0.01 0.000000000001 3 0.001 0.00000000001 4 0.0001 0.0000000001 5 0.00001 0.000000001 6 0.000001 0.00000001 7 0.0000001 0.0000001 8 0.00000001 0.000001 9 0.000000001 0.00001 10 0.0000000001 0.0001 11 0.00000000001 0.001 12 0.000000000001 0.01 13 0.0000000000001 0.1 14 0.00000000000001 1.0 The pH of pure water at 25°C can be calculated from the concentration of [H3 O+ ] in pure water, which is equal to 1.0 × 10− 7 M.pH = − log1 ×= −10− 7log 1.0 + log= −10− 70 += 7.00− 7Acidic solutions will have a higher [H3 O+ ] than pure water (values > 1 × 10− 7 ). So, the pH values of acidic solutions will be less than 7.0. Since basic solutions will have a lower concentration of [H3 O+ ] than pure water (values < 1 × 10− 7 ), the pH values of basic solutions will be greater than 7.0. The lower the pH value, the higher the hydronium ion concentration and the stronger the acid solution. Similarly, the higher the pH value, the lower the hydronium ion concentration and the stronger the basic solution (Fig. 5.6 ).Fig. 5.6
eBook - PDF- Douglas P. Heller, Carl H. Snyder(Authors)
- 2015(Publication Date)
- Wiley(Publisher)
The pH Scale 243 pH: The Measure of Acidity As we just saw, we can make things simpler for ourselves by writing the value of [H 3 O + ] as an exponent of 10. For example, a hydronium ion concentration of 0.0000001 M is the same as 1 × 10 −7 M, which is simply 10 −7 M. We can carry this a step further by dispensing with both the 10 and the negative sign. This is exactly what the Danish biochemist Søren Sørensen did in 1909 when he proposed that concentrations of H + (or, as we now know, H 3 O + ) be treated as exponential values. (Mathematically, this process of using the exponent or power of a number as a value is called “taking the logarithm” of that num- ber.) Following Sørensen’s recommendation, we now consider the concentration of hydro- nium ion, [H 3 O + ], in terms of pH (Figure 8.10). The letters pH represent the “power of the Hydrogen (or Hydronium) ion.” As a symbol for acidity, pH reflects nicely the international character of chemistry. The let- ter p begins the English word power as well as its French and German equivalents, puissance and Potenz. At the time of Sørensen’s suggestion, English, French, and German were the world’s dominant scientific languages. In chemically pure water, the concentrations of all the transient hydronium (H 3 O + ) and hydroxide (OH − ) ions that exist at equilibrium are not only fixed at any given temperature but are also equal to each other. They must always equal each other in pure water be- cause the ionization of each water molecule produces an equal number of ions—one hydronium ion and one hydroxide ion. Experimental measurements show that each of these ions is present in pure, neutral water at a concentration of 0.0000001 moles per li- ter at 25°C. For brevity, we write [H 3 O + ] for “the molar concentration of H 3 O + ” and we express the value of the molar concentration in exponential notation, using the italicized capital M as the symbol for moles/liter.
eBook - ePub- 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 basesWhen 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.Acid–base balanceIn-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.We have already looked at the concept of homeostasis in Chapter 2
eBook - ePubChemistry
Concepts and Problems, A Self-Teaching Guide
- Richard Post, Chad Snyder, Clifford C. Houk(Authors)
- 2020(Publication Date)
- Jossey-Bass(Publisher)
The pH, sometimes called the “hydrogen ion exponent,” is defined mathematically as the negative logarithm to the base 10 of the hydrogen ion concentration. For the following H + ion concentrations, compare the power of 10 exponent with the corresponding value of pH. Answer: 6 If the hydrogen ion concentration is 1 × 10 –(exponent), then the pH is equal to the exponent. Or, expressed another way: If a solution has a pH of 3, what is the hydrogen ion concentration? _____ Answer: [H + ] = 1 × 10 −3 M What is the pH of pure neutral water? __________ Answer: A solution is acidic if [H + ] is greater than 1 × 10 −7 M. A solution is alkaline if [H + ] is less than 1 × 10 −7 M. Is a solution with a pH of 6 acidic or alkaline? __________ Answer: acidic (The [H + ] = 1 × 10 −6, which is greater than 1 × 10 −7.) To determine the pH of a solution if the [H + ] is something other than 1 × 10 –(exponent) requires use of common logarithms (log to the base 10). Some instructors may arrange problems so that all [H + ] will be 1 × 10 –(exponent) to simplify calculations. More often than not, [H + ] is something other than 1 × 10 –(exponent). We have chosen such examples and included a log table for your use in the Appendix. You may also use your scientific calculator or the handy calculators available on the Internet by way of a web browser and search engines. In this book, the answers to problems involving logarithms will assume you understand and know how to use logarithms. To determine the pH of a hydrogen ion concentration of n × 10 –(exponent), where n is a number other than 1, we must determine the logarithm of the number n. Determine the pH of a solution if [H + ] = 3 × 10 −5 (to the nearest hundredth)
eBook - ePubThe 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 alkalinityA2 The pH Scale for Hydrogen Ion ConcentrationsThe 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+
eBook - ePub- Odd-Ivar Lekang(Author)
- 2019(Publication Date)
- Wiley-Blackwell(Publisher)
This is actually not exactly correct due to the fact that the total concentration is not exactly 14 as can be seen from above. It is temperature dependent, and at low temperatures (0 °C), pH 7 will be slightly acidic: Example The pH in pure water not affected by the atmosphere or any buffering system (explained later) shall be found. The water temperature is set to room temperature: In pure water the concentration of H + and OH − must be equal, meaning Figure 5.1 The pH Scale. Example It is 10 −14 mol/L H + ions in the water. What is the pH? pH = −log 10 (H +) pH = 14 It is 10 −8 mol/L H + ions in the water. What is the pH? pH = 8 It is 10 −6 mol/L H + ions in the water. What is the pH? pH = 6 It is 10 −2 mol/L H + ions in the water. What is the pH? pH = 2 This shows that for 1 pH unit the concentration of H + ions is changing by 10. Example The water has a pH of 6. Find the number of OH − ions in the water with a temperature of 20 °C. Even at pH 2, there is still a very limited amount of free H + ions in the solution, 0.01 mol/L, compared to the amount of water molecules, 55.5 mol/L. 5.2.3 The carbonate system Natural water will however not be completely pure because gases in the atmosphere, the earth/rock/soil where water is draining through and biological processes in the water interact with the water and affect the composition (add substances). Looking at the gases in the water first, there is an important equilibrium between the CO 2 gas in the atmosphere and the CO 2 concentration in the water. This is described by Henry's law (see chapter 12) and states that if there is CO 2 in the atmosphere above water surface, some of this gas will be transferred and dissolved in the water, assuming available surface for gas transfer and enough time for this gas transfer to take place
eBook - PDFGeneral, Organic, and Biological Chemistry
An Integrated Approach
- Kenneth W. Raymond(Author)
- 2012(Publication Date)
- Wiley(Publisher)
7.6 The pH Scale 253 SOLUTION a. [H 3 O + ] = 10 - pH = 10 - 6 = 1 * 10 - 6 M (acidic) b. [H 3 O + ] = 10 - pH = 10 - 6.5 = 3 * 10 - 7 M (acidic) c. [H 3 O + ] = 10 - pH = 10 - 1.2 = 6 * 10 - 2 M (acidic) Solving part a does not require a calculator. All three answers are reported with one signifi- cant figure (Math Support—Logs and Antilogs). PRACTICE PROBLEM 7.8 What is [OH - ] in the following solutions? Indicate whether each is acidic, neutral, or basic. a. pH = 7.2 b. pH = 9.1 c. pH = 3.3 MATH SUPPORT—LOGS AND ANTILOGS On your scientific calculator you should be able to calculate logarithms (logs). If you try a few calculations, you will see that log 100 = 2 log 1000 = 3 log 0.0001 = -4 Converting these three numbers (100, 1000, and 0.0001) into scientific notation (Section 1.4) should help explain what the log or logarithm of a number is. log 100 = log 10 2 = 2 log 1000 = log 10 3 = 3 log 0.0001 = log 10 -4 = -4 The log of a number is the power to which ten must be raised to equal the number (log 10 n = n). The log of 100 is 2, because 10 2 equals 100 and the log of 0.0001 is -4 because 10 -4 = 0.0001. When a number has the value 1 * 10 n , where n is an integer (-2, 5, etc.), its log can be determined without using a calculator. In these cases, the log is equal to the value of n. log 1 * 10 -2 = -2 log 1 * 10 5 = 5 In all other cases (7.9 * 10 2 , 2.2 * 10 -5 , etc.), a calculator will be required. log 7.9 * 10 2 = 2.90 log 2.2 * 10 -5 = -4.66 Reversing this process gives the antilog or antilogarithm of a number (antilog n = 10 n ). log of 10 2 = 2, so antilog 2 = 10 2 log of 10 -8 = -8, so antilog -8 = 10 -8 When n is an integer (2, -8, etc.), its antilog can be determined without using a calculator, because the antilog is equal to 10 n (see the two examples directly above). At other times (3.5, -1.3, etc.) a calculator must be used.
eBook - PDF- Gary D. Christian, Purnendu K. Dasgupta, Kevin A. Schug(Authors)
- 2013(Publication Date)
- Wiley(Publisher)
272 CHAPTER 7 ACID–BASE EQUILIBRIA ● buffer capacity, ionic strength, fractional distribution, activities and apparent dissociation constants of all species at equilibrium. It has a higher learning curve, but with practices, it provides a wealth of information. You might try doing the same calculations using both programs. You should get the same result. We can also solve the pH of mixtures of acids and bases using Goal Seek. See the video illustrating this for a mixture of NaOH and H 2 CO 3 . Professor’s Favorite Problems Contributed by Professor Michael De Grandpre, University of Montana Example 7.25 pH of Seawater Did you know that ocean pH is decreasing (see the figures on the next page)? A portion of the CO 2 from fossil fuel combustion is absorbed by the oceans, forming carbonic acid. This process, named “ocean acidification,” has decreased the average pH in the surface oceans by 0.1 units over the past ∼100 years. Chemical oceanographers have long realized the importance of tracking and studying CO 2 in the oceans and in the 1980s began developing improved analytical tools to do this, including new methods for measuring pH. Glass pH electrodes, the workhorse for most pH applications, are not accurate enough to document these small pH changes over time. Oceanographers revitalized an old but rarely used method, the spectrophotometric measurement of pH using indicators. As you may have guessed, the function for deriving the pH is simply the Henderson – Hasselbalch equation: pH = pK a + log [A − ] [HA] where pH is defined on the total hydrogen ion scale (for a definition and description see References 17 and 19.) . The pK a is the apparent dissociation constant, and [A − ] and [HA] are the unprotonated and protonated forms of the pH indicator. The improvement came in the determination of pK a on a pH Scale consistent with the CO 2 equilibria in seawater (see References 19 and 20).
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