CHEM 101 Ch.4
Flow Diagram for Acid Base Titration
For the balanced equation: NaOH + HCl --> NaCl + H2O Given volume of NaOH use the molarity of NaOH as a conversion factor to find moles of NaOH, then use the coefficients in the balanced equation to find mole ratios to find moles of HCl, and then divide by the volume of HCl to find the molarity of HCl.
Using molarity as a conversion factor between moles and volume in stoichiometry calculations.
For the balanced equation: aA + bB --> cC + dD Given Volume of solution A Use molarity as a conversion factor to find moles of A Use coefficients in the balanced equation to find the A: B mole ratio to find moles of B needed Use molarity as a conversion factor to find the volume of solution pf B
Common Acids and Bases Strong to Weak
HClO4 H2SO4 HBr HCl HNO3 H3PO4 HF HNO2 CH3COOH KOH NaOH Ba(OH)2 Ca(OH)2 NH3
To do that, you must know the solubility of each potential product—how much of each compound will dissolve in a given amount of solvent at a given temperature. If a substance has a high solubility in water, no precipitate will form. If a substance has a low solubility in water, it's likely to precipitate from an aqueous solution.
Solubility is a continuum with some substances having a high solubility and others having a low solubility. In this section, we will define a substance as soluble if it dissolves to give a concentration of 0.01 M or greater. Solubility can be predicted by looking at the cations and anions that make up the compound.
Note that pure water is a nonelectrolyte because it does not dissociate appreciably into H+ and OH− ions.
Strong Electrolytes HCl, HBr, HI HClO4 HNO3 H2SO4 KBr NaCl NaOH, KOH Other soluble ionic compounds Weak Electrolytes CH3CO2H HF HCN Nonelectrolytes H2O CH3OH C2H5OH C12H22O11 Most compounds of carbon *organic*
The volume to be diluted is withdrawn using a calibrated tube called a pipet, placed in an empty volumetric flask of the chosen volume, and diluted to the calibration mark on the flask.
The one common exception to this order of steps is when diluting a strong acid such as H2SO4 where a large amount of heat is released. In such instances, it is much safer to add the acid slowly to the water rather than adding water to the acid.
Note that a forward-and-backward double arrow (⇌) is used in the dissociation equation to indicate that the reaction takes place simultaneously in both directions. That is, dissociation is a dynamic process in which an equilibrium is established between the forward and reverse reactions. Dissociation of acetic acid takes place in the forward direction, while recombination of H+ and CH3CO2− ions takes place in the reverse direction. The size of the equilibrium arrow indicates whether the equilibrium reaction forms mostly products or mostly reactants. Ultimately, the concentrations of the reactants and products reach constant values and no longer change with time. We'll learn much more about chemical equilibria in Chapter 15.
A chemical equilibrium, as we'll see in Chapter 15, is the state in which a reaction takes place in both forward and backward directions so that the concentrations of products and reactants remain constant over time.
A compound is soluble if it meets either (or both) of the following criteria:
A compound is soluble if it contains one of the following cations: Li+, Na+, K+, Rb+, Cs+ (group 1A cations) NH+4 (ammonium ion) That is, essentially all ionic compounds containing an alkali metal cation or ammonium cation are soluble in water and will not precipitate, regardless of the anions present. A compound is soluble if it contains one of the following anions: Cl−, Br−, I− (halide) except: Ag+, Hg2+2 and Pb2+ halides NO−3 (nitrate), ClO−4 (perchlorate), CH3CO−2 (acetate), and SO2−4 (sulfate) except: Sr2+, Ba2+, Hg2+2 and Pb2+ sulfates That is, most ionic compounds containing a halide, nitrate, perchlorate, acetate, or sulfate anion are soluble in water and will not precipitate regardless of the cations present. The exceptions that will precipitate are silver(I), mercury(I) and lead(II) halides and strontium, barium, mercury(I), and lead(II) sulfates. On the other hand, a compound that does not contain one of the cations or anions listed above is not soluble. Thus, carbonates (CO32−) sulfides (S2−) phosphates (PO43−) and hydroxides (OH−) are generally not soluble unless they contain an alkali metal or ammonium cation. The main exceptions are the sulfides and hydroxides of Ca2+, Sr2+ and Ba2+.
We all know that both sugar (sucrose) and table salt (NaCl) dissolve in water. The solutions that result, though, are quite different. When sucrose, a molecular substance, dissolves in water, the resulting solution contains neutral sucrose molecules surrounded by water. When NaCl, an ionic substance, dissolves in water, the solution contains separate Na+ and Cl− ions surrounded by water. Because of the presence of the charged ions, the NaCl solution conducts an electric current, but the sucrose solution does not.
A molecule is a unit of matter that results when two or more nonmetal atoms are joined by covalent bonds in which electrons are shared. An ionic substance is formed when a metal and nonmetal atom form an ionic bond in which electrons are transferred from the metal to the nonmetal to form ions. The electrical conductivity of an aqueous NaCl solution is easy to demonstrate using a battery, a light bulb, and several pieces of wire, connected as shown in Figure 4.3. When the wires are dipped into an aqueous NaCl solution, the positively charged Na+ ions move through the solution toward the wire connected to the negatively charged terminal of the battery, and the negatively charged Cl− ions move toward the wire connected to the positively charged terminal of the battery. The resulting movement of electrical charges allows a current to flow, so the bulb lights. When the wires are dipped into an aqueous sucrose solution, however, there are no ions to carry the current, so the bulb remains dark.
We saw in Section 4.9 that the concentration of an acid or base solution can be determined by titration. A measured volume of the acid or base solution of unknown concentration is placed in a flask, and a base or acid solution of known concentration is slowly added from a buret. By measuring the volume of the added solution necessary for a complete reaction, as signaled by the color change of an indicator, the unknown concentration can be calculated. Remember . . . The reaction used for a titration must go to completion and have a yield of 100%
A similar procedure can be carried out to determine the concentration of many oxidizing or reducing agents using a redox titration. All that's necessary is that the substance whose concentration you want to determine undergo an oxidation or reduction reaction in 100% yield and that there be some means, such as a color change, to indicate when the reaction is complete. The color change might be due to one of the substances undergoing reaction or to some added indicator. Let's imagine that we have a potassium permanganate solution whose concentration we want to find. Aqueous KMnO4 reacts with oxalic acid, H2C2O4, in acidic solution according to the following net ionic equation (K+ is a spectator ion): 5 H2C2O4(aq)+2 MnO4−(aq)+6 H+(aq)→10 CO2(g)+2 Mn2+(aq)+8 H2O(l) The reaction goes to completion with 100% yield and is accompanied by a sharp color change when the intense purple color of the MnO4− ion disappears. The strategy used is outlined in Figure 4.9. As with acid—base titrations, the general idea is to measure a known amount of one substance—in this case, H2C2O4—and use mole ratios from the balanced equation to find the number of moles of the second substance—in this case, KMnO4—necessary for complete reaction. With the molar amount of KMnO4 thus known, titration gives the volume of solution containing that amount. Dividing the number of moles by the volume gives the concentration.
The rules for assigning oxidation numbers are as follows:
An atom in its elemental state has an oxidation number of 0. An atom in a monatomic ion has an oxidation number identical to its charge. Review Section 2.13 to see the charges on some common ions. An atom in a polyatomic ion or in a molecular compound usually has the same oxidation number it would have if it were a monatomic ion. - In general, the farther left an element is in the periodic table, the more probable that it will be cationlike. Metals, therefore, usually have positive oxidation numbers. The farther right an element is in the periodic table, the more probable that it will be anionlike. Nonmetals, such as O, N, and the halogens, usually have negative oxidation numbers. ----Hydrogen can be either +1 or −1. When bonded to a metal, such as Na or Ca, hydrogen has an oxidation number of −1 When bonded to a nonmetal, such as C, N, O, or Cl, hydrogen has an oxidation number of +1. ----Oxygen usually has an oxidation number of −2 The major exception is in compounds called peroxides, which contain either the O22− ion or an O—O covalent bond in a molecule. Both oxygen atoms in a peroxide have an oxidation number of −1. ----Halogens usually have an oxidation number of −1. The major exception is in compounds of chlorine, bromine, or iodine in which the halogen atom is bonded to oxygen. In such cases, the oxygen has an oxidation number of −2, and the halogen has a positive oxidation number. In Cl2O, for instance, the O atom has an oxidation number of −2, and each Cl atom has an oxidation number of +1. The sum of the oxidation numbers is 0 for a neutral compound and is equal to the net charge for a polyatomic ion. This rule is particularly useful for finding the oxidation number of an atom in difficult cases. The general idea is to assign oxidation numbers to the "easy" atoms first and then find the oxidation number of the "difficult" atom by subtraction.
When an acid and a base are mixed in the right stoichiometric proportions, both acidic and basic properties disappear because of a neutralization reaction that produces water and an ionic salt. The anion of the salt (A−) comes from the acid, and the cation of the salt (M+) comes from the base: HA(aq) acid + MOH(aq) base --> H2O(l) water + MA(aq) salt
Because salts are generally strong electrolytes in aqueous solution, we can write the neutralization reaction of a strong acid with a strong base as an ionic equation: H+(aq)+A−(aq)+M+(aq)+OH−(aq)→H2O(l)+M+(aq)+A−(aq) Canceling the ions that appear on both sides of the ionic equation, A− and M+ gives the net ionic equation, which describes the reaction of any strong acid with any strong base in water. Net Ionic Equation orH+(aq)+OH−(aq)→H2O(l)H3O+(aq)+OH−(aq)→2 H2O(l) For the reaction of a weak acid with a strong base, a similar neutralization occurs, but we must write the molecular formula of the acid rather than simply H+(aq) because the dissociation of the acid in water is incomplete. Instead, the acid exists primarily as the neutral molecule. In the reaction of the weak acid HF with the strong base KOH, for example, we write the net ionic equation as HF(aq)+OH−(aq)→H2O(l)+F−(aq)
Redox reactions take place with every element in the periodic table except helium and neon and occur in a vast number of processes throughout nature, biology, and industry. Here are just a few examples:
Combustion. Combustion is the burning of a fuel by oxidation with oxygen in air. Gasoline, fuel oil, natural gas, wood, paper, and other organic substances of carbon and hydrogen are the most common fuels. Even some metals, such as magnesium and calcium, will burn in air. CH4(g)Methane(Natural gas)+2 O2(g)→CO2(g)+2 H2O(l) Bleaching. Bleaching uses redox reactions to decolorize or lighten colored materials. Dark hair is bleached to turn it blond, clothes are bleached to remove stains, wood pulp is bleached to make white paper, and so on. The exact oxidizing agent used depends on the situation—hydrogen peroxide (H2O2) is used for hair, sodium hypochlorite (NaOCl) is used for clothes, and ozone or chlorine dioxide is used for wood pulp—but the principle is always the same. In all cases, colored impurities are destroyed by reaction with a strong oxidizing agent. Batteries. Although they come in many types and sizes, all types of batteries are powered by redox reactions. In a typical redox reaction carried out in the laboratory— say, the reaction of zinc metal with Ag+ to yield Zn2+ and silver metal—the reactants are simply mixed in a flask and electrons are transferred by direct contact between them. In a battery, however, the two reactants are kept in separate compartments and the electrons are transferred through a wire running between them. The inexpensive alkaline battery commonly used in flashlights and other small household items uses a thin steel can containing zinc powder and a paste of potassium hydroxide as one reactant, separated by paper from a paste of powdered carbon and manganese dioxide as the other reactant. A graphite rod with a metal cap sticks into the MnO2 to provide electrical contact. When the can and the graphite rod are connected by a wire, zinc sends electrons flowing through the wire toward the MnO2 in a redox reaction. The resultant electrical current can be used to light a bulb or power a small electronic device. The reaction is Zn(s)+2 MnO2(s)→ZnO(s)+Mn2O3(s) Metallurgy. Metallurgy, the extraction and purification of metals from their ores, makes use of numerous redox processes. Metallic zinc is prepared by reduction of ZnO with coke, a form of carbon: ZnO(s)+C(s)→Zn(s)+CO(g) Corrosion. Corrosion is the deterioration of a metal by oxidation, such as the rusting of iron in moist air. The economic consequences of rusting are enormous: It has been estimated that up to one-fourth of the iron produced in the United States is used to replace bridges, buildings, and other structures that have been destroyed by corrosion. (The raised dot in the formula Fe2O3⋅H2O for rust indicates that one water molecule is associated with each Fe2O3 in an unspecified way.) 4 Fe(s)+3 O2(g)−→H2O2 Fe2O3Rust⋅H2O(s) Respiration. The term respiration refers to the processes of breathing and using oxygen for the many biological redox reactions that provide the energy needed by living organisms. The energy is released from food molecules slowly and in complex, multi-step pathways, but the overall result of respiration is similar to that of a combustion reaction. For example, the simple sugar glucose (C6H12O6) reacts with O2 to give CO2 and H2O according to the following equation: C6H12O6Glucose(a carbohydrate)+6 O2→6 CO2+6 H2O+energy
Many common chemical reactions that take place in aqueous solution fall into one of three general categories: precipitation reactions, acid-base neutralization reactions, and oxidation-reduction reactions.
In precipitation reactions, soluble ionic reactants (strong electrolytes) yield an insoluble solid product called a precipitate, which falls out of the solution, thereby removing some of the dissolved ions. Most precipitations take place when the anions and cations of two ionic compounds change partners. In acid-base neutralization reactions, an acid reacts with a base to yield water plus an ionic compound called a salt. Acids are compounds that produce H+ ions when dissolved in water, and bases are compounds that produce OH− ions when dissolved in water. Thus, a neutralization reaction removes H+ and OH− ions from solution, just as a precipitation reaction removes metal and nonmetal ions. In oxidation-reduction reactions, or redox reactions, one or more electrons are transferred between reaction partners (atoms, molecules, or ions). As a result of this electron transfer, the charges on atoms in the various reactants change.
Approximately 71% of the Earth's surface is covered by water, and another 3% is covered by ice; 66% of the mass of an adult human body is water, and water is needed to sustain all living organisms.
It's therefore not surprising that a large amount of important chemistry, including all those reactions that happen in our bodies, takes place in water—that is, in aqueous solution.
Molarity can be used as a conversion factor to relate a solution's volume to the number of moles of solute. If we know the molarity and volume of a solution, we can calculate the number of moles of solute. If we know the number of moles of solute and the molarity of the solution, we can find the solution's volume.
Molarity = Moles of solute/Volume of solution(L) Moles of solute = Molarity x Volume of solution Volume of solution = Moles of solution/Molarity
Thus, the most useful means of expressing a solution's concentration is molarity (M), the number of moles of a substance, or solute, dissolved in enough solvent to make 1 liter of solution.
Molarity(M)=Moles of solute/Liters of solution Note that it's the final volume of the solution that's important, not the starting volume of the solvent used. The final volume of the solution might be a bit larger than the volume of the solvent because of the additional volume of the solute. In practice, a solution of known molarity is prepared by weighing an appropriate amount of solute and placing it in a container called a volumetric flask, as shown in Figure 4.1. Enough solvent is added to dissolve the solute, and further solvent is added until an accurately calibrated final volume is reached. The solution is then gently mixed to reach a uniform concentration.
The main thing to remember when diluting a concentrated solution is that the number of moles of solute is constant; only the volume of the solution is changed by adding more solvent. Because the number of moles of solute can be calculated by multiplying molarity times volume, we can set up the following equation:
Moles of solute (constant) = Molarity x Volume = M1V1 = M2V2 where Mi is the initial molarity, Vi is the initial volume, Mf is the final molarity, and Vf is the final volume after dilution. Rearranging this equation into a more useful form shows that the molar concentration after dilution (Mf) can be found by multiplying the initial concentration (Mi) by the ratio of initial and final volumes (Vi/Vf) M2=M1xV1/V2
Substances such as NaCl or KBr, which dissolve in water to produce conducting solutions of ions, are called electrolytes. Substances such as sucrose or ethyl alcohol, which do not produce ions in aqueous solution, are nonelectrolytes.
Most electrolytes are ionic compounds, but some are molecular. Hydrogen chloride, for instance, is a gaseous molecular compound when pure but dissociates, or splits apart, to give H+ and Cl− ions when it dissolves in water.
Compounds that dissociate to a large extent (70-100%) into ions when dissolved in water are said to be strong electrolytes, while compounds that dissociate to only a small extent are weak electrolytes.
Potassium chloride and most other ionic compounds, for instance, are largely dissociated in dilute solution and are thus strong electrolytes. Acetic acid (CH3CO2H), by contrast, dissociates only to the extent of about 1.3% in a 0.10 M solution and is a weak electrolyte. As a result, a 0.10 M solution of acetic acid is only weakly conducting, and the bulb in Figure 4.3 would only light dimly.
oxidation and reduction reactions, called half-reactions, always occur together. A redox reaction consists of two half-reactions; one oxidation half-reaction and one reduction half-reaction. Whenever one atom loses one or more electrons, another atom must gain those electrons. The substance that causes a reduction by giving up electrons—the iron atom in the reaction of Fe with O2 and the carbon atom in the reaction of C with Fe2O3—is called a reducing agent. The substance that causes an oxidation by accepting electrons—the oxygen atom in the reaction of Fe with O2 and the iron atom in the reaction of C with Fe2O3—is called an oxidizing agent. The reducing agent is itself oxidized when it gives up electrons, and the oxidizing agent is itself reduced when it accepts electrons.
Reducing agent - causes reduction, loses one or more electrons, undergoes oxidation, oxidation number of atoms increases Oxidizing agent - causes oxidation, gains one or more electrons, undergoes reduction, oxidation number of atom decreases redox reactions are common for almost every element in the periodic table except for the noble-gas elements of group 8A. In general, metals give up electrons and act as reducing agents, while reactive nonmetals such as O2 and the halogens accept electrons and act as oxidizing agents. Different metals can give up different numbers of electrons in redox reactions. Lithium, sodium, and the other group 1A elements give up only one electron and become monopositive ions with oxidation numbers of +1 Beryllium, magnesium, and the other group 2A elements, however, typically give up two electrons and become dipositive ions. The transition metals in the middle of the periodic table can give up a variable number of electrons to yield more than one kind of ion depending on the exact reaction. Titanium, for example, can react with chlorine to yield either TiCl3 or TiCl4. Because a chloride ion has a −1 oxidation number, the titanium atom in TiCl3 must have a +3 oxidation number, and the titanium atom in TiCl4 must be +4.
Solubility Table for Ionic Compounds in Water
Soluble Li, Na, K, Rb, Cs NH4 Cl, Br, I except Ag, Hg2, Pb NO3 ClO4 CH3CO2 SO4 except Sr, Ba, Hg2, Pb Insoluble Compounds CO3 except carbonates of group 1A cations, NH4 S except sulfides of group 1A cations, NH4, Ca, Sr, Ba PO4 except phosphates of group 1A cations, NH4 OH except hydroxides of group 1A cations, NH4, Ca, Sr, Ba
In 1777, the French chemist Antoine Lavoisier (1743-1794) proposed that all acids contain a common element: oxygen. In fact, the word oxygen is derived from a Greek phrase meaning "acid former." Lavoisier's idea had to be modified, however, when the English chemist Sir Humphrey Davy (1778-1829) showed in 1810 that muriatic acid (now called hydrochloric acid) contains only hydrogen and chlorine but no oxygen. Davy's studies thus suggested that the common element in acids is hydrogen, not oxygen.
Swedish chemist Svante Arrhenius (1859-1927) clarified the relationship between acidic behavior and the presence of hydrogen in a compound in 1887. Arrhenius proposed that an acid is a substance that dissociates in water to give hydrogen ions (H+) and a base is a substance that dissociates in water to give hydroxide ions (OH−). An acid HA(aq)→H+(aq)+A−(aq)A base MOH(aq)→M+(aq)+OH−(aq) In these equations, HA is a general formula for an acid—for example, HCl or HNO3—and MOH is a general formula for a metal hydroxide—for example, NaOH or KOH. Although convenient to use in equations, the symbol H+(aq) does not really represent the structure of the ion present in aqueous solution. As a bare hydrogen nucleus— a proton—with no electron nearby, H+ is much too reactive to exist by itself. Rather, the H+ bonds to the oxygen atom of a water molecule and forms the more stable hydronium ion, H3O+. We'll sometimes write H+(aq) for convenience, particularly when balancing equations, but will more often write H3O+(aq) to represent an aqueous acid solution.
Bases, like acids, can also be either strong or weak, depending on the extent to which they produce OH− ions in aqueous solution. Most metal hydroxides, such as NaOH and Ba(OH)2, are strong electrolytes and strong bases.
The dissociation reactions for sodium hydroxide and barium hydroxide are: NaOH(aq)−→Na+(aq)+OH−(aq)Ba(OH)2(aq)−→Ba2+(aq)+2 OH−(aq) Ammonia (NH3) is a weak electrolyte and a weak base. Ammonia is a weak base because it reacts to a small extent with water to yield NH4+ and OH− ions. In fact, aqueous solutions of ammonia are often called ammonium hydroxide, although this is really a misnomer because the concentrations of NH4+ and OH− ions are low. NH3(g)+H2O(l)⇌NH4+(aq)+OH−(aq) As with the dissociation of acetic acid, discussed in Section 4.3, the reaction of ammonia with water takes place only to a small extent (about 1%). Most of the ammonia remains unreacted, and we therefore write the reaction with a double arrow to show that a dynamic equilibrium exists between the forward and reverse reactions.
Most acids are oxoacids; that is, they contain oxygen in addition to hydrogen and other elements. When dissolved in water, an oxoacid yields one or more H+ ions and an oxoanion like one of those listed in Table 4.4 and discussed previously in Section 2.13. Oxoanions are polyatomic anions in which an atom of a given element is combined with different numbers of oxygen atoms.
The names of oxoacids are related to the names of the corresponding oxoanions, with the -ite or -ate ending of the anion name replaced by -ous acid or -ic acid, respectively. In other words, the acid with fewer oxygens has an -ous ending, and the acid with more oxygens has an -ic ending. In addition to the oxoacids, there are a small number of other common acids, such as HCl, that do not contain oxygen. For such compounds, the prefix hydro- and the suffix -ic acid are used for the aqueous solution.
The equations we've been writing up to this point have all been molecular equations. That is, all the substances involved in the reactions have been written using their complete formulas as if they were molecules.
The physical state of a substance in a chemical reaction is often indicated with a parenthetical state abbreviation (s) for solid, (l) for liquid, (g) for gas, and (aq) for aqueous solution (Section 3.3). A Molecular Equation This equation implies that molecules are interacting. It is the case, however, that lead nitrate, potassium iodide, and potassium nitrate are strong electrolytes that dissolve in water to yield solutions of ions. Thus, it's more accurate to write the precipitation reaction as an ionic equation, in which all the ions are explicitly shown. An Ionic Equation This ionic equation shows that the NO3− and K+ ions undergo no change during the reaction. Instead, they appear on both sides of the reaction arrow and act merely as spectator ions, whose only role is to balance the charge. A Net Ionic Equation A net ionic equation gives only the species that react (the Pb2+ and I− ions in this instance) because spectator ions are canceled from both sides of the equation. Leaving the spectator ions out of a net ionic equation doesn't mean that their presence is irrelevant. If a reaction occurs by mixing a solution of Pb2+ ions with a solution of I− ions, then those solutions must also contain additional ions to balance the charge in each. That is, the Pb2+ solution must also contain an anion, and the I− solution must also contain a cation. Leaving these other ions out of the net ionic equation only implies that these ions do not undergo a chemical reaction. Any nonreactive spectator ion could serve to balance charge.
You might notice that most of the ions that impart solubility to compounds are singly charged—either singly positive (Li+, Na+, K+, Rb+, Cs+, NH4+) or singly negative (Cl−, Br−, I−, NO3−, ClO4−, CH3CO2−). Very few doubly charged ions or triply charged ions form soluble compounds.
This solubility behavior arises because of the relatively strong ionic bonds in compounds containing ions with multiple charges. The greater the strength of the ionic bonds holding ions together in a crystal, the more difficult it is to break those bonds apart during the solution process.
A technique frequently used for determining a solution's exact molarity is called a titration.
Titration is a procedure for determining the concentration of a solution by allowing a measured volume of that solution to react with a second solution of another substance (the standard solution) whose concentration is known. By finding the volume of the standard solution that reacts with the measured volume of the first solution, the concentration of the first solution can be calculated. (It's necessary, though, that the reaction go to completion and have a yield of 100%.) To see how titration works, let's imagine that we have an HCl solution (an acid) whose concentration we want to find by allowing it to react with NaOH (a base) in an acid—base neutralization reaction. The balanced equation is NaOH(aq)+HCl(aq)→NaCl(aq)+H2O(l) We'll begin the titration by measuring out a known volume of the HCl solution and adding a small amount of an indicator, a compound that undergoes a color change during the course of the reaction. The compound phenolphthalein, for instance, is colorless in acid solution but turns red in base solution. Next, we fill a calibrated glass tube called a buret with an NaOH standard solution of known concentration and slowly add the NaOH to the HCl. When the phenolphthalein just begins to turn pink, all the HCl has completely reacted, and the solution now has a tiny amount of excess NaOH. By then reading from the buret to find the volume of the NaOH standard solution that has been added to react with the known volume of HCl solution, we can calculate the concentration of the HCl.
Although these and many thousands of other reactions appear unrelated, and many don't even take place in aqueous solution, all are oxidation—reduction (redox) reactions. Historically, the word oxidation referred to the combination of an element with oxygen to yield an oxide, and the word reduction referred to the removal of oxygen from an oxide to yield the element
Today, the words oxidation and reduction have taken on a much broader meaning. Oxidation is now defined as the loss of one or more electrons by a substance, whether element, compound, or ion, and reduction is the gain of one or more electrons by a substance. Thus, an oxidation—reduction, or redox, reaction is any process in which electrons are transferred from one substance to another. How can you tell when a redox reaction takes place? The answer is that you assign to each atom in a compound a value called an oxidation number (or oxidation state), which indicates whether the atom is neutral, electron-rich, or electron-poor. By comparing the oxidation number of an atom before and after reaction, you can tell whether the atom has gained or lost electrons. Note that oxidation numbers don't necessarily imply ionic charges; they are just a convenient device to help keep track of electrons during redox reactions.
Acids that dissociate to a large extent are strong electrolytes and strong acids, whereas acids that dissociate to only a small extent are weak electrolytes and weak acids.
We've already seen in Table 4.1, for instance, that HCl, HClO4, HNO3 and H2SO4 are strong electrolytes and therefore strong acids, while CH3CO2H and HF are weak electrolytes and therefore weak acids. You might note that acetic acid actually contains four hydrogens, but only the one bonded to the oxygen atom dissociates. We will explain the effect of molecular structure on acid dissociation in Chapter 16. Different acids can have different numbers of acidic hydrogens and yield different numbers of H3O+ ions in solution. Hydrochloric acid (HCl) is said to be a monoprotic acid because it provides only one H+ ion, but sulfuric acid (H2SO4) is a diprotic acid because it can provide two H+ ions. Phosphoric acid (H3PO4) is a triprotic acid and can provide three H+ ions. With sulfuric acid, the first dissociation of an H+ is complete—all H2SO4 molecules lose one H+—but the second dissociation is incomplete, as indicated by the double arrow in the following equation: Sulfuric acid:H2SO4(aq)+H2O(l)−→HSO4−(aq)+H3O+(aq)HSO4−(aq)+H2O(l)⇌SO42−(aq)+H3O+(aq) With phosphoric acid, none of the three dissociations is complete: Phosphoric acid:H3PO4(aq)+H2O(l)⇌H2PO4−(aq)+H3O+(aq)H2PO4−(aq)+H2O(l)⇌HPO42−(aq)+H3O+(aq)HPO42−(aq)+H2O(l)⇌PO43−(aq)+H3O+(aq)
