Chapter 5: Acids, Bases, and Salts
- Page ID
- 38898
\( \newcommand{\vecs}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)
\( \newcommand{\vecd}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash {#1}}} \)
\( \newcommand{\id}{\mathrm{id}}\) \( \newcommand{\Span}{\mathrm{span}}\)
( \newcommand{\kernel}{\mathrm{null}\,}\) \( \newcommand{\range}{\mathrm{range}\,}\)
\( \newcommand{\RealPart}{\mathrm{Re}}\) \( \newcommand{\ImaginaryPart}{\mathrm{Im}}\)
\( \newcommand{\Argument}{\mathrm{Arg}}\) \( \newcommand{\norm}[1]{\| #1 \|}\)
\( \newcommand{\inner}[2]{\langle #1, #2 \rangle}\)
\( \newcommand{\Span}{\mathrm{span}}\)
\( \newcommand{\id}{\mathrm{id}}\)
\( \newcommand{\Span}{\mathrm{span}}\)
\( \newcommand{\kernel}{\mathrm{null}\,}\)
\( \newcommand{\range}{\mathrm{range}\,}\)
\( \newcommand{\RealPart}{\mathrm{Re}}\)
\( \newcommand{\ImaginaryPart}{\mathrm{Im}}\)
\( \newcommand{\Argument}{\mathrm{Arg}}\)
\( \newcommand{\norm}[1]{\| #1 \|}\)
\( \newcommand{\inner}[2]{\langle #1, #2 \rangle}\)
\( \newcommand{\Span}{\mathrm{span}}\) \( \newcommand{\AA}{\unicode[.8,0]{x212B}}\)
\( \newcommand{\vectorA}[1]{\vec{#1}} % arrow\)
\( \newcommand{\vectorAt}[1]{\vec{\text{#1}}} % arrow\)
\( \newcommand{\vectorB}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)
\( \newcommand{\vectorC}[1]{\textbf{#1}} \)
\( \newcommand{\vectorD}[1]{\overrightarrow{#1}} \)
\( \newcommand{\vectorDt}[1]{\overrightarrow{\text{#1}}} \)
\( \newcommand{\vectE}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash{\mathbf {#1}}}} \)
\( \newcommand{\vecs}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)
\( \newcommand{\vecd}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash {#1}}} \)
\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)After reading this section1, you should be able to:
- Explain the properties of water
- Describe inorganic and organic acids
- Describe inorganic bases
- Explain pH, acids, and bases
- Describe buffers
Water is the most abundant and important inorganic compound on earth. It makes up 60 to 89-percent of the volume of most living cells, and it possesses several properties that make it vital to life. The capability for a molecule to absorb heat energy is called heat capacity. Water has the highest heat capacity of any liquid. It would require calorie of heat energy to change the temperature of water by 1. As a result, it takes water a long time to heat and a long time to cool. This property prevents sudden changes in temperature caused by external factors, such as sun or wind exposure, or by internal conditions that release heat rapidly, such as vigorous muscle activity. As part of blood system or the environment, water redistributes heat among adjacent structures ensuing the temperature remains homeostatic.2
When water evaporates, or vaporizes, water changes from a liquid to a gas, water vapor. The transformation requires that large amounts of heat be absorbed to break hydrogen bonds that hold water molecules together. This property is beneficial because as water evaporates from an object or organism large amounts of heat are removed providing efficient cooling. This property is referred to as high heat of vaporization.
Water is the best solvent in nature. It is called the universal solvent. Biological molecules do not react chemically unless they are in solution, and virtually all chemical reactions that occur in the living cells depend on water’s solvent properties. Water molecules are referred to as being polar. They orient with their slightly negative ends toward the positive ends of the solutes. . This characteristic is called polarity, and it explains the reason that ionic compounds and other small reactive molecules, such as acids and bases, dissociate in water, where their ions separating from each other and become evenly scattered in the water forming a true solution.
Water also forms layers of water molecules, called hydration layers, around large, charged molecules such as protein, shielding them from the effects of other charged substances in the areas and preventing them from settling out of solution. Such protein water mixtures are biological colloids. Water is also the major transport medium because it is an excellent polar solvent. Nutrients, gases, and metabolic wastes are carried dissolved in water based fluids. Wastes are excreted form living organisms in watery fluids. Specialized molecules that lubricate organisms also use water as the dissolving medium.
As a reactant, water is involved in many important chemical reactions. Nutrients are decomposed by adding a water molecule to each chemical bond that is broken. Decomposition reactions are more specifically referred to as hydrolysis reactions. When large carbohydrates or protein molecules are synthesized from smaller molecules, a water molecule is removed for every bond formed, a reaction that is called dehydration synthesis.
Water forms resilient cushions, cushioning around certain biological structures providing protection from physical trauma.
Other than binary molecular compounds, most inorganic compounds can be classified as acids, bases, or salts. Acids are characterized by the H+ion, the presence of which in water, makes the water acidic. An acid either contains this ion or produces it when it dissolves in water. A base either contains hydroxide ion, OH-, or reacts with water to produce hydroxide. Most bases that contain hydroxide consist of metal cations and hydroxide; examples are sodium hydroxide, NaOH, and calcium hydroxide, Ca(OH)2. Acids and bases react to form a salt, an ionic compound that has a cation other than H+ and an anion other than OH-. This kind of reaction always produces water.
Acids
Salts, acids, and bases are electrolytes. They ionize and dissociate in water and can conduct an electrical current.
Acids have a sour taste, and they can react with or dissolve metals. The definition of an acid is a substance that releases hydrogen ions (H+) in measurable amounts. Acids are also characterized as being proton donors because a hydrogen ion is a hydrogen nucleus, or a single proton.
When acids dissolve in water, they release hydrogen ions (protons) and anions, negative charged particles. The concentration of the protons determines the acidity of the solution. The anions have little or no effect on the acidity of the solution. Hydrochloric acid (HCl) dissociates into a proton and a chloride ion:
HCl → H+, proton + Cl-, anion
Living organisms require a balance of acids and bases to function. We can measure this balance by determining the number by the amount of H+, and OH- dissociated in a solution. Simply stated, more H+ hydrogen ions in a solution means more acidic. More OH- hydroxide ions means more basic (alkaline).
Biochemical reactions are sensitive to small changes in the acidity or alkalinity of the environment in which they occur. H+ and OH- are involved in almost all biochemical processes, and any deviation from a narrow band of the normal H+ and OH- concentration dramatically modifies the systems’ functioning. Acids and bases that are formed in living systems must be kept in balance.
It is convenient to express the amount of H+ in a solution by a logarithmic pH scale that ranges from 0 to 14. The term pH means potential of hydrogen ion concentration. On a log scale, a change of one whole number has 100 times more hydrogen ions than a solution of pH 2, and a pH of 2 has 100 times more hydrogen ions than a solution of pH 3.
Acidic solutions contain more H+ than OH- and have a pH lower than 7. If a solution has more OH- than H+, it is a basic or alkaline solution. In pure water, a small percentage of molecules are dissociated into H+ and OH- ions so that it has a pH of 7. When the ion concentrations of a solution are equal, the solution has a pH of 7 and is considered to be neutral.
The pH of a solution can be managed with buffers. Buffers resist a change in pH and will be discussed later in the chapter.
Bases
Bases have a bitter taste, feel slippery, and are proton acceptors. They take up hydrogen ions in measurable amounts. Common inorganic bases include the hydroxides like magnesium hydroxide and sodium hydroxide. Lye acids, hydroxides dissociate when dissolved in water. In this case, hydroxyl ions (OH-) and cations are released. Ionization of sodium hydroxide produces a hydroxyl ion and a sodium ion. The hydroxyl ion binds to a proton present in the solution. This reaction produces water and simultaneously reduces the acidity of the solution by taking up free H+ ions:
NaOH → Na+ + OH-
OH- + H+ → H2O
The term pH is used to express the intensity of an acid or alkaline solution.
Salts and Neutralization
Acids and bases react to form salts, an ionic compound that has a cation other than H+ and an anion other than OH-. This kind of reaction always produces water and is known as a neutralization reaction. The most well-known salt is sodium chloride, NaCl.
HCl + NaOH → NaCl + H2O
Acid Base Salt Water
The salt produced is dissolved in the aqueous solution and disassociated into Na+ and Cl- ions. Although sodium chloride is commonly what one means in referring to “salt,” there are many other salts as well. For example, Sodium sulfate, Na2SO4, dissociates into two Na+ ions and one SO4-2. The salt dissociates because the ions are formed. The water overcomes the attraction between the oppositely charged ions, and they disassociate.
All ions are electrolytes which are substances that conduct an electrical current in solution. Salts dissociate in aqueous solutions into ions, and the most common salts are sodium salts. In their ionized form, salts play a vital role in nature and aqueous solutions.
Buffers3
How can organisms whose bodies require a near-neutral pH ingest acidic and basic substances (a human drinking orange juice, for example) and survive? Buffers are the key. Buffers usually consist of a weak acid and its conjugate base; this enables them to readily absorb excess H+ or OH–, keeping the system’s pH within a narrow range. Organisms are extremely sensitive to slight changes in the pH of the environment. In high concentrations, acids and bases are damaging to cells. Homeostasis of acid-base balance is regulated by chemical systems called buffers.
Buffers resist abrupt and large changes in the pH of a solution by releasing hydrogen ions when the pH begins to rise and by binding hydrogen ions when the pH drops.
Chemical buffer systems react by binding hydrogen ions or by releasing hydrogen ions. The acidity of a solution reflects the free hydrogen ions and not the hydrogen ions bound to anions. Acids that dissociate completely and irreversibly in water are called strong acids. They can dramatically change the pH of a solution. Acids that do not dissociate completely, like carbonic acid (H2CO3) and acetic acid, are weak acids. Undissociated acids do not affect pH, so that acetic acid solutions are much less acidic than hydrochloric acid solution. Weak acids disassociate in predictable ways, and molecules of the intact acid are in dynamic equilibrium with the dissociated ions.
For this reason, when a strong acid is added to a solution of a weak acid the equilibrium will shift to the left and some H+ will recombine to form acetic acid. On the other hand, if a strong base is added and the pH begins to rise, the equilibrium shifts to the right and more acetic molecules disassociate to release H+ ions. This characteristic of weak acids allows them to play a role in the chemical buffer systems found in nature.
H+ acetic acid ↔ H+ + Acetic-
Bases are proton acceptors. Strong bases are those bases, like hydroxides, which dissociate easily in water and quickly bind H+ ions. Sodium bicarbonate ionizes incompletely and reversibly. Since it accepts relatively few protons, its released bicarbonate ion is considered to be a weak base.
One of the buffer system that helps to maintain the pH in aqueous solutions is the carbonic acid-bicarbonate system. Carbonic acid dissociates reversibly in aqueous solutions releasing bicarbonate ions and protons, H+. The chemical equilibrium between carbonic acid, a weak acid, and bicarbonate ion, a weak base, resists changes in pH by shifting to the right or the left as H+ ions are added to or removed from the solution:
Rise in pH
H2CO3 ↔ H+ + HCO3-
Drop in pH
As the pH rises and becomes more alkaline, the equilibrium shifts to the right, forcing more carbonic acid to dissociate. Similarly, as the pH begins to drop, the equilibrium shifts to the left as more bicarbonate ions begin to bind with protons. Strong bases are replaced by a weak base, bicarbonate ion. Protons are released by strong acids and are tied up in weak acids, carbonic acid. As a result, the pH changes are much less than they would be in the absence of the buffering system.
Key Terms
- acid – a substance that dissociates into one or more hydrogen ions (H+); proton donors; acids release hydrogen ions in measurable amounts; sour taste
- base – a substance that dissociates into one or more positive ions (cations) that can accept or combine with protons; proton acceptors; bases take up hydrogen ions in measurable amounts; bitter taste
- buffer – buffers prevent the pH of a solution from changing drastically
- pH – the potential of hydrogen ion concentration
[1] This page titled 4.10: Acids, Bases and Salts is shared under a CC BY-NC-SA 4.0 license and was authored, remixed, and/or curated by Stanley E. Manahan.
[2] This page titled 2.14: Water - High Heat Capacity is shared under a CC BY-SA 4.0 license and was authored, remixed, and/or curated by Boundless.
[3] This page titled 2.4.2: pH, Buffers, Acids, and Bases is shared under a CC BY-SA 4.0 license and was authored, remixed, and/or curated by Boundless.