Tatiana Roizer
Chapter 8
ACIDS, BASIS, AND IONIC OMPOUNDS
The world is complex, to understand how living and non-living systems work we must study the chemical make up of each substance.
First let’s discuss basic definitions that are the foundation to understanding this chapter.
1. Electrolyte is a "medical/scientific" term for salts, specifically ions. The term electrolyte means that this ion is electrically-charged and moves to either a negative (cathode) or positive (anode) electrode
- Ions that move to the cathode (cations) are positively charged
- Ions that move to the anode (anions) are negatively charged
- Electrolyses is the passage of electricity through a solution holding dissolved ions
- Electrolyte is a solute that enables a solution to conduct electricity
- Electrodes are the plates or wires that dip into the solution.
For example, your body fluids -- blood, plasma, interstitial fluid (fluid between cells) -- are like seawater and have a high concentration of sodium chloride (table salt, or NaCl). The electrolytes in sodium chloride are:
- sodium ion (Na+) - cation
- chloride ion (Cl-) - anion
Electrolytes are important because they are what your cells (especially nerve, heart, muscle) use to maintain voltages across their cell membranes and to carry electrical impulses (nerve impulses, muscle contractions) across themselves and to other cells
2. Ionization is the gain or loss of electrons. The loss of electrons, which is the more common process in astrophysical environments, converts an atom into a positively charged ion, while the gain of electrons converts an atom into a negatively charged ion. The ionization energy of an atom measures how strongly an atom holds its electrons. The ionization energy is the minimum energy required to remove an electron from the ground state of the isolated gaseous atom. Note that this does not mean the energy required to remove an electron from the n=1 shell (i.e the ground state orbital), the ground state here refers to the lowest energy electron configuration for the element in question
The first ionization energy, I1, is the energy needed to remove the first electron from the atom:
Na(g) -> Na+(g) + 1e-
The second ionization energy, I2, is the energy needed to remove the next (i.e. the second) electron from the atom
Na+(g) -> Na2+(g) + 1e-
The higher the value of the ionization energy, the more difficult it is to remove the electron
As electrons are removed, the positive charge from the nucleus remains unchanged, however, there is less repulsion between the remaining electrons
- Zeff increases with removal of electrons
- Greater energy is needed to remove remaining electrons (i.e. the ionization energy is higher for each subsequent electron)
Notation for Degrees of Ionization
Suffix / Ionization / Examples / Chemist's
Notation
I / Not ionized (neutral) / H I, He I / H, He
II / Singly ionized / H II, He II / H+, He+
III / Doubly ionized / He III, O III / He++, O++
- taken from:
What is the role of an ion? To understand the role of an ion, we shall ask some more detailed questions.
What is the process of solutes when dissolved in water?
Through ionization, solutes release ions in water.
Ammonia Dissolves in Water:
Polar ammonia molecules dissolve in polar water molecules. These molecules mix readily because both types of molecules engage in hydrogen bonding. Since the intermolecular attractions are roughly equal, the molecules can break away from each other and form new solute (NH3), solvent (H2O) hydrogen bonds
Alcohol Dissolves in Water: The -OH group on alcohol is polar and mixes with the polar water through the formation of hydrogen bonds. A wide variety of solutions are in this category such as sugar in water, alcohol in water, acetic and hydrochloric acids.
*Taken from:
Do ions carry electricity in water?
Yes, ions in water do carry electricity.
Redox reactions primarily involve the transfer of electrons between two chemical species. The compound that loses an electron is said to be oxidized, the one that gains an electron is said to be reduced. There are also specific terms that describe the specific chemical species. A compound that is oxidized is referred to as a reducing agent, while a compound that is reduced is referred to as the oxidizing agent.
Notes: A great number of reactions are ion-combination reactions. These reactions involve NO change to the valency (oxidation number) of the reacting chemicals. The oxidation number is a number identical with the valency but with a sign, expressing the nature of the charge of the species in question when formed from the neutral atom. Thus, the oxidation number of chlorine in hydrochloric acid is -1, while it is +1 in hypochlorous acid. Similarly we can say that the oxidation number of chlorine in chloric acid (HClO3) is +5, and in perchloric acid
(HClO4) +7. CuSO4 + 2NaOH -> Cu(OH)2 + Na2SO4
In this ion-combination reaction Copper (Cu) is in the ionic form Cu2+ on the left hand side of the equation, and as Cu2+ on the right hand side of the equation.
Cu2+ + SO42- + 2Na+ + 2OH- -> (Cu2+ + 2OH-)Solid + 2Na+ + SO42-
What we are showing here, is that the Copper (Cu) has stayed as Cu2+. It has not gained any electrons (nor lost any electrons). This is typical of ion-combinationreactions.
Redox Reactions (Oxidation – Reductions)
In these reactions, the valency (oxidation number) of the reactantschange.
For example: 2Fe3+ + Sn2+ -> 2Fe2+ + Sn4+ (8+ each side of the equation)
The iron (iii) + tin (ii) have reacted to give iron (ii) + tin (iv) of course, this reaction is carried out in the presence of HCl (Hydrochloric Acid), but the oxidation reduction reaction is only between the iron (iii) and tin (ii).
Now, a redoxreaction is the release and uptake of electrons.
So, the Fe3+ is reduced to Fe2+, and the Sn2+ is oxidized to Sn4+.
What happened in this reaction?
Sn2+ donated electrons to the Fe3+ (an electron transfer took place).
Redoxreactions then, are the transfer of electrons from one reactant to another...When there is oxidation, there is alsoreduction.
The substance which loses electrons is oxidized.
The substance which gains electrons is reduced.
For example: Fe (metal) + Cu2+ -> Fe2+ + Cu (metal)
Fe donates two electrons to the Cu2+ to form Cu (metal). The Fe lost 2 electrons, so is oxidized.
The Cu2+gained 2 electrons, so is reduced (in its valency).
Simply put:The chemical which gains electrons is reduced(reduces its valency) and is called the oxidising agent.
The chemical which loses electrons, is oxidised (increases its valency) and is called the reducingagent.
Fe / + / Cu2+ / -> / Fe2+ / + / Cu
Oxidised
Reducing
Agent / Reduced
Oxidising
Agent
The chemical which is oxidised is the reducing agent.
The chemical which is reduced is the oxidising agent.
For Example: Zn + 2HCl -> Zn2+ + H2 +2Cl-
In this reaction Zn + 2H+ -> Zn2+ + H2 (the chlorine is not changed in its ion state, so is not oxidised or reduced).
Zn is oxidised to Zn2+ (loses 2 electrons)
H+ is reduced to H2 (gains 2 electrons)
H+ has an oxidation number (valency) of +1, and is reduced to an oxidation number (valency) of 0.
So, reductiondecreases the oxidation number (valency).
Zn has an oxidation number (valency) of 0, and is oxidised to Zn2+, an oxidation number of 2+,
So oxidationincreases the oxidation number (valency).
Redox Reactions involving acid and bases solutions
Not only are there an exchange of electrons in these reactions, but also an exchange of protons (hydronium ions), as in any base system.
CuS + HNO3 -> Cu SO4 + NO (g) + H2O (equation not balanced).
3CuS + 8HNO3 -> 3 CuSO4 + 8NO(g) + 4H2O (equation balanced)
3CuS2+ + 3S2- + 8H+ + 8NO3- -> 3Cu2+ + 3SO42- + 8NO(g) + 4H2O (equation written in ionic nomenclature)
Copper Cu2+ does not change.
8H+ goes to 4H2O (exchange of hydronium ions)
Sulphur S2- goes to S6+O4 ((S6+ (O42-))2-
So has changed from 2- to 6+ (Sulphur). Sulphur has lost electrons, therefore has been oxidized.
Nitrogen has gone from N5+ (N5+ O32-) to N2+ (N2+ O2-)
So has gainedelectrons, therefore has been reduced
Sometimes it is easier to see the transfer of electrons in the system if it is split into definite steps. This will be oxidation of one substance and reduction of the other substance.
2Fe3+ + Sn2+ -> 2Fe2+ + Sn4+
Split into 2 separate steps.
2Fe3+ + 2e- -> 2Fe2+ (reduction)
(6+) + (2-) -> (4+) (balanced for charges)
Sn2+ -> Sn4+ + 2e- (oxidation)
(2+) -> (4+) + (2-)
Add the two half equations: 2Fe3+ + 2e- + Sn2+ -> 2Fe2+ + Sn4+ + 2e-
The electrons cancel each other out, so equation is: 2Fe3+ + Sn2+ -> 2F2+ + Sn4+
By breaking down the equation into half cells, the oxidation or reduction of each chemical can be determined. The atom which gains electrons reduces its valency, therefore is reduced and is called the oxidizing agent.
The atom which loses electrons, increases its oxidation number, therefore is oxidized, and is called the reducingagent.
Let’s have a final example of what the information above means. During electrolysis we have something taking electrons from one electrode and something else putting them simultaneously on the other electrode. As an example, I shall use the electrolysis of aqueous copper (II) bromide.
At the cathode: Cu2+(aq) + 2e------) Cu(s) = Reduction
At the anode: 2Br-(aq) ------) Br2(l) + 2e- = Oxidation
Sum: Cu2+(aq) + 2Br-(aq) ------) Cu(s) + Br2(l) = Electrolysis
Note: Ions need to be mobile for electrolyses to occur…(can’t occur in solid state)
Section 2: Acid and Bases as Electrolytes For thousands of years people have known that vinegar, lemon juice and many other foods taste sour. However, it was not until a few hundred years ago that it was discovered why these things taste sour - because they are all acids. The term acid, in fact, comes from the Latin term acere, which means sour. In the seventeenth century, the English writer and amateur chemist Robert Boyle first labeled substances as either acids or bases (he called bases alkalies) according to the following characteristics:
- Acids taste sour, are corrosive to metals, change litmus (a dye extracted from lichens) red, and become less acidic when mixed with bases.
- Bases feel slippery, change litmus blue, and become less basic when mixed with acids.
While Boyle and others tried to explain why acids and bases behave the way they do, the first reasonable definition of acids and bases would not be proposed until 200 years later.
In the late 1800's, the Swedish scientist Svante Arrhenius proposed that water can dissolve many compounds by separating them into their individual ions. Arrhenius suggested that acids are compounds that contain hydrogen and can dissolve in water to release hydrogen ions into solution. For example, hydrochloric acid (HCl) dissolves in water as follows:
HCl / H2O/ H+(aq) / + / Cl-(aq)
Arrhenius defined bases as substances that dissolve in water to release hydroxide ions (OH-) into solution. For example, a typical base according to the Arrhenius definition is sodium hydroxide (NaOH):
NaOH / H2O/ Na+(aq) / + / OH-(aq)
The Arrhenius definition of acids and bases explains a number of things. Arrhenius' theory explains why all acids have similar properties to each other (and, conversely, why all bases are similar), because all acids release H+ into solution (and all bases release OH-). The Arrhenius definition also explains Boyle's observation that acids and bases counteract each other. This idea, that a base can make an acid weaker, and vice versa, is called neutralization.
Neutralization: As you can see above, acids release H+ into solution and bases release OH-. If we were to mix an acid and base together, the H+ion would combine with the OH-ion to make the molecule H2O, or plain water:
H+(aq) / + / OH-(aq) / / H2OThe neutralization reaction of an acid with a base will always produce water and a salt, as shown below:
Acid / Base / Water / SaltHCl / + / NaOH / / H2O / + / NaCl
HBr / + / KOH / / H2O / + / KBr
While Arrhenius helped explain the fundamentals of acid/base chemistry, unfortunately his theories have limits. For example, the Arrhenius definition does not explain why some substances like common baking soda (NaHCO3) can act like a base even though they do not contain hydroxide ions.
In 1923, the Danish scientist Johannes Brønsted and the Englishman Thomas Lowry published independent yet similar papers that refined Arrhenius' theory. In Brønsted's words, "... acids and bases are substances that are capable of splitting off or taking up hydrogen ions, respectively." The Brønsted-Lowry definition broadened the Arrhenius concept of acids and bases.
The Brønsted-Lowry definition of acids is very similar to the Arrhenius definition, any substance that can donate a hydrogen ion is an acid (under the Brønsted definition, acids are often referred to as proton donors because an H+ion - hydrogen minus its electron - is simply a proton).
The Brønsted definition of bases is, however, quite different from the Arrhenius definition. The Brønsted base is defined as any substance that can accept a hydrogen ion. In essence, a base is the opposite of an acid. NaOH and KOH, as we saw above, would still be considered bases because they can accept an H+ from an acid to form water. However, the Brønsted-Lowry definition also explains why substances that do not contain OH- can act like bases. Baking soda (NaHCO3), for example, acts like a base by accepting a hydrogen ion from an acid as illustrated below:
Acid / Base / SaltHCl / + / NaHCO3 / / H2CO3 / + / NaCl
pH
Under the Brønsted-Lowry definition, both acids and bases are related to the concentration of hydrogen ions present.Acids increase the concentration of hydrogen ions, while bases decrease the concentration of hydrogen ions (by accepting them). The acidity or basicity of something therefore can be measured by its hydrogen ion concentration.
In 1909, the Danish biochemist Sören Sörensen invented the pH scale for measuring acidity. The pH scale is described by the formula:
pH = -log [H+] / Note: concentration is commonly abbreviated by using square brackets, thus [H+] = hydrogen ion concentration. When measuring pH, [H+] is in units of moles of H+ per liter of solution.For example, a solution with [H+] = 1 x 10-7moles/liter has a pH equal to 7 (a simpler way to think about pH is that it equals the exponent on the H+ concentration, ignoring the minus sign). The pH scale ranges from 0 to 14. Substances with a pH between 0 and less than 7 are acids (pH and [H+] are inversely related - lower pH means higher [H+] and a stronger acid). Substances with a pH greater than 7 and up to 14 are bases (the higher the pH, the stronger the base). Right in the middle, at pH = 7, are neutral substances, for example, pure water. The relationship between [H+] and pH is shown in the table below alongside some common examples of acids and bases in everyday life.
[H+] / pH / ExampleAcids / 1 X 100 / 0 / HCl
1 x 10-1 / 1 / Stomach acid
1 x 10-2 / 2 / Lemon juice
1 x 10-3 / 3 / Vinegar
1 x 10-4 / 4 / Soda
1 x 10-5 / 5 / Rainwater
1 x 10-6 / 6 / Milk
Neutral / 1 x 10-7 / 7 / Pure water
Bases / 1 x 10-8 / 8 / Egg whites
1 x 10-9 / 9 / Baking Soda
1 x 10-10 / 10 / Tums® antacid
1 x 10-11 / 11 / Ammonia
1 x 10-12 / 12 / Mineral Lime - Ca(OH)2
1 x 10-13 / 13 / Drano®
1 x 10-14 / 14 / NaOH
* taken from Anthony Carpi, Ph.D. at vision learning.com
Water as a baseCaption / When an acid reacts with water, the water behaves as a proton acceptor to form the hydronium ion.
Notes / Water behaving as a base
Keywords / conjugate acid-base, water, base
* Taken from prentice Hall.
When a base (like ammonia) reacts with water, a proton is transferred from water to the ammonia molecule to form the ammonium ion. Therefore, water is behaving as a proton donor.
Figure below Values of the H3O+and OH-concentrations at 25°C in acidic, neutral, and basic solutions.
Ionic solids (or salts) contain positive and negative ions, which are held together by the strong force of attraction between particles with opposite charges. When one of these solids dissolves in water, the ions that form the solid are released into solution, where they become associated with the polar solvent molecules.
H2ONaCl(s) / / Na+(aq) / + / Cl-(aq)
We can generally assume that salts dissociate into their ions when they dissolve in water. Ionic compounds dissolve in water if the energy given off when the ions interact with water molecules compensates for the energy needed to break the ionic bonds in the solid and the energy required to separate the water molecules so that the ions can be inserted into solution
Solubility Rules
Thanks to Professor Kenneth W. Busch from whose Web page these data were extracted.
1. Salts containing Group I elements are soluble (Li+, Na+, K+, Cs+, Rb+). Exceptions to this rule are rare. Salts containing the ammonium ion (NH4+) are also soluble.
2. Salts containing nitrate ion (NO3-) are generally soluble.
3. Salts containing Cl -, Br -, I - are generally soluble. Important exceptions to this rule are halide salts of Ag+, Pb2+, and (Hg2)2+. Thus, AgCl, PbBr2, and Hg2Cl2 are all insoluble.
4. Most silver salts are insoluble. AgNO3 and Ag(C2H3O2) are common soluble salts of silver; virtually anything else is insoluble.
5. Most sulfate salts are soluble. Important exceptions to this rule include BaSO4, PbSO4, Ag2SO4, and CaSO4.
6. Most hydroxide salts are only slightly soluble. Hydroxide salts of Group I elements are soluble. Hydroxide salts of Group II elements (Ca, Sr, and Ba) are slightly soluble. Hydroxide salts of transition metals and Al3+ are insoluble. Thus, Fe(OH)3, Al(OH)3, Co(OH)2 are not soluble.
7. Most sulfides of transition metals are highly insoluble. Thus, CdS, FeS, ZnS, Ag2S are all insoluble. Arsenic, antimony, bismuth, and lead sulfides are also insoluble.
8. Carbonates are frequently insoluble. Group II carbonates (Ca, Sr, and Ba) are insoluble. Some other insoluble carbonates include FeCO3, PbCO3. Carbonates become soluble in acid solution.
9. Chromates are frequently insoluble. Examples: PbCrO4, BaCrO4
10. Phosphates are frequently insoluble. Examples: Ca3(PO4)2, Ag2PO4
11. Fluorides are frequently insoluble. Examples: BaF2, MgF2 PbF2.
The following solubility rules predict the solubility of many ionic compound when applied in order:
1. All Na+, K+, and NH4+ salts are soluble.
2. All NO3-, C2H3O2-, ClO3-, and ClO4- salts are soluble.
3. All Ag+, Pb2+, and Hg22+ salts are insoluble.
4. All Cl-, Br-, and I- salts are soluble.
5. All CO32-, O2-, S2-, OH-, PO43-, CrO42-, Cr2O72- and SO32- salts are insoluble. / Group IIA S2- and Ba(OH)2 are soluble.
6. All SO42- salts are soluble. / CaSO4, SrSO4, and BaSO4 are insoluble.
Let's first start with a complete chemical equation and see how the net ionic equation is derived. Take for example the reaction of lead(II) nitrate with sodium chloride to form lead(II) chloride and sodium nitrate, shown below:
Pb(NO3)2(aq) + 2 NaCl(aq) PbCl2(s) + 2 NaNO3(aq)