Introduction to Inorganic Chemistry/Acid-Base Chemistry
Chapter 3: Acid-Base Chemistry
Acids and bases are important for a number reasons in inorganic chemistry.
- Many industrially useful catalytic reactions involve inorganic acids and superacids, such as zeolites, anhydrous hydrogen fluoride, and sulfated zirconia. These acids are sufficiently strong in anhydrous media that they can protonate olefins and alcohols to produce carbocations. Carbocations are key intermediates in the transformations of hydrocarbons.
- Inorganic compounds are sometimes synthesized in strongly acidic or basic media. For example, ternary metal oxides can be synthesized and crystallized in molten NaOH or KOH, which are strongly basic. Organic fluorination reactions are often done in strongly acidic media, such as anhydrous HF. Understanding the familiar chemistry of acids and bases in water helps us understand how these non-aqueous media work.
- The acidic or basic environment of metal ions affects the stability of their oxidation states. We will learn more about this in Chapter 4.
- Transition metal complexes (coordination compounds and organometallic compounds) are essentially Lewis acid-base complexes. We can understand a great deal about their stability and reactivity by considering the acid-base character of metals and ligands. We will learn about this in Chapter 5.
3.1 Brønsted and Lewis acids and bases
There are three major classifications of substances known as acids or bases. The Arrhenius definition states that an acid produces H+ in solution and a base produces OH-. This theory was developed by Svante Arrhenius in 1883. Later, two more sophisticated and general theories were proposed. These are the Brønsted-Lowry and the Lewis definitions of acids and bases. The relationship between these theories is illustrated in the figure at the left.
In water, we normally think about the Brønsted acidity or basicity of acids and bases. Brønsted-Lowry acidity is based on the transfer of a hydrogen ion in solution. The acid is a proton donor and dissociate in order to increase [H+], and the base is a proton acceptor and dissociate in order to increase [OH-]. The hyrdogen ion (H+) is often referred to as a proton in these cases. To clarify, this theory does not imply that protons do not exist independently in solution. In an aqueous solution the proton is generally designated as H3O+. This theory encompasses any type of solution, not just aqueous ones as the Arrhenius theory does. The strength of an acid or a base varies depending on the solvent. This strength can be classified by the extent of the dissociation of the acid or base. Strong acids and bases dissociate completely in a solvent. Weak acids and bases will only partially dissociate in a solvent. An acid or base may be strong in one solvent but weak another. This means that the strength of the acid/base is relative. The strength of weak acids in solution can be measured through equilibrium calculations. The equilibrium constant (also known as acid dissociation constant), ka, is the number that represents the degree of dissociation. For the chemical equation:
HA <----> A- + H+
When an acid dissociates, it gives up a proton and a Conjugate Base. When a base gains a proton, the product is called a Conjugate Acid. Strong acids produce spectator ions as conjugate bases, while weak acids produce a weak conjugate bases. This is illustrated below for the weak acid acetic acid and its conjugate base, the acetate anion.
The strength of a conjugate acid/base varies inversely with the strength or weakness of it's parent acid or base. Any acid or base is technically a conjugate acid or conjugate base also; these terms are simply used to identify species in solution (i.e Acetic acid is the conjugate acid of acetate anions,a base, while acetate anions are the conjugate base of acetic acid, an acid). Some acids, called polyprotic acids have more than one proton to donate in acid/base reactions. For example, Sulfuric acid, H2SO4, is a strong acid that has a conjugate base that actually happens to be a weak acid itself. This means that every mole of H2SO4 in aqueous solution donates more than 1 mole of protons.
Some substances can act both as an acid and as a base. An example would be water. H2O molecules may either donate a hydrogen or accept it.This is called an amphoteric substance. In the situation where you have an acid in solution, water will act as a base. Water acts as an acid in base. The strongest acid we can make in H2O is H+ (aq). In addition, the strongest base we can make in H2O is OH- (aq).
Non-aqueous acid-bas chemistry follows similar rules. For example in liquid ammonia:
- 2NH3(l) → NH4+ (NH3) + NH2- (NH3)
where NH4+ is the strongest acid and NH2- is the strongest base.
SOLVENT LEVELING Solvent leveling is an effect that occurs when a strong acid is placed in an autoprotic solvent, such as, but not limited too, H2O. Any strong acid donates it's proton to the solvent, this leveling effect means that the strongest possible acid in aqueous solution is H3O+. This means that the strength of acids such as HCl and HBr cannot be accurately determined in water as they both are leveled by H2O. Using an autoprotic solvent of a higher dissociation constant, such as anhydrous acetic acid, can allow for the determination of strength of strong acids. with the strongest possible base of this solvent (CH3COO-) being a weaker base than aqueous solutions (OH-), the strength of strong acids can be more accurately measured, as strong acids become weak acids (note: the properties of the acid/base itself do not change, but the dissociation equilibrium begins to favor the undissociated molecule as the solvent approaches the acid/base in acidity/basicity). Solvent leveling also applies to bases, though more basic autoprotic solvents than water are used, such as ammonia (NH3)
MOLTEN SALT SOLUTIONS
When a solid salt melts, it forms a solution of the cations and anions (KOH melts and dissociates into K+ + OH- ions) which can act as a solvent for chemical reactions. When water is added to a molten KOH flux, the OH- acts as the solvent, leading to the acid base equilibrium H2O <----> OH- <----> O2- this allows for the acidity of a flux to be easily altered through the addition or boiling off of water, a "wet" flux is a more acidic flux, while a "dry" flux is more basic. Molten fluxes are often used in the synthesis of crystals, such as perovskite superconductors (K1-XBaXBiO3)
LEWIS ACIDS AND BASES
The Lewis classification of acids and bases is a much broader definition, and encompasses many more substances. Whereas the Brønsted-Lowry and the Arrhenius classifications are based on proton activity, Lewis acidity is based on the transfer of an electron pair. Lewis acids can accept an electron pair, while Lewis bases can donate an electron pair. This definition clearly allows many more substances to be considered acidic or basic. It is not necessary for a compound to contain hydrogen atoms in order for it to be a Lewis acid or base. Such a substance need not even be a compound, rather only an element. For example, Aluminum chloride, AlCl3 can be an acid when combined with an ion that donates an electron pair. The reaction is shown below.
AlCl3 + Cl- <------> AlCl4-
Here, the acid is the AlCl3 and the base is the Cl-. This is because the chloride ion is has multiple electron pairs that it can potentially donate to the aluminum chloride.
Lewis acidity is the basis for coordination chemistry. This is because coordination chemistry involves metals that are Lewis acids bonding to ligands that are Lewis bases.
Certain molecules with five membered geometries happen to be excellent Lewis acids, because when accepting another pair of electrons from a base, they form octahedral geometry. The octahedral geometry is more stable and therefore favored over the five membered arrangements. This means that the five membered geometries will readily accept an electron pair from a ligand and are therefore acidic. Some examples of these acids are PF5, AsF5, and SbF5. One interesting example of manipulating lewis acidity is the chemical synthesis of Fluorine gas. Aqueous F-(aq) ions react with MnO4-(aq) to form MnF62-, which is stabilized by 2 K+ ions. Reaction with the lewis acid, SbF5(l) forms MnF4, an unstable compound. This decays to produce MnF3 and (1/2)F2(g).
Determining the strength of a metal Lewis Acid
There are three determining factors of the strength of a metal Lewis acid. They are as follows:
1. The higher positive charge on the metal, the more acidic it is.
2. The smaller the atomic radius of the metal ion, the more acidic it is.
3. The more electronegative the metal ion is, the more acidic it will be. The trend is that the closer the metal is to Au on the periodic table, the more electronegative and acidic it is.
3.2 Hard acids and bases
Lewis acids and bases can be classified by designating them as hard or soft.
- Hard Acids/Bases:
- "Hard" acids and bases have a high charge (positive for acids, negative for bases) to atomic radius ratio along with higher oxidation states. Hard acids have LUMOs (Lowest Unoccupied Molecular Orbital) of relatively high energy, while the HOMOs (Highest Occupied Molecular Orbital) of low energy. They interact through ionic bonding, and are not very polarizable. Hard acids generally have high charge densities, which include metal ions with higher positive charges and smaller ionic sizes. Basically, the term "hard" applies to the chemical species that considered to be small in size, but have high charges, and are weakly polarizable. Therefore, the electrons present are held tightly by the nucleus. This is due to the presence of d orbitals lacking in the ability to form π bonds. Some examples of hard acids and bases include: H+, O2-, OH-, F-, Fe3+, and Al3+. Specifically, F- is a hard base and metal ions including Li+ are considered hard acids.
- F- > Cl- > Br- > I-
- F- > Cl- > Br- > I-
- Soft Acids/Bases:
- "Soft" acids or bases have a low charge to atomic radius ratio, with low oxidation states. They are normally larger acids that interact through covalent bonding and are polarizable. Soft acids have LUMOs of lower energy than harder acids, and soft bases have HOMOs of higher energy than harder bases. Soft acids tend to have lower charge densities due to lower ionic charges and increase in ionic sizes. This is due to their d orbitals being able to form π bonds. Essentially, the term soft applies to those chemical species that tend to be bigger in size, low in charge, and are strongly polarizable, where there is greater distribution of electrons. For example, I- is a soft base and low charge density transition metals, such as Ag+, are considered soft acids. Soft acids often include transition metals in the second and third row of the periodic table that have a +1 or +2 charge, as well as almost filled d orbitals.
- I- > Br- > Cl- > F-
- I- > Br- > Cl- > F-
Like binds with Like
Hard acids interact faster with hard bases than they do with soft bases, and soft acids interact faster with soft bases than hard bases; the strongest bonds are Hard-Hard and Soft-Soft bonds. Acids and bases are not strictly hard or soft, with many compounds being classified as borderline, with trimethylborane and Fe2+ & Pb2+ cations being borderline acids, and pyridine and aniline being examples of borderline bases.
- Hard binds with hard
- Soft binds with soft
3.3 Discussion questions
- Discuss periodic trends in the Lewis acidity of metal ions.
- Explain what we mean by hard and soft acids and bases, using specific examples.
1. For each pair of compounds below, indicate which one is a stronger acid and explain your choice:
(a) [Fe(H2O)6]3+ or [Fe(H2O)6]2+
(b) H3BO3 or H3PO3
(c) [Ga(H2O)6]3+ or [Al(H2O)6]3+
(d) HIO3 or HIO4
(e) H3PO4 or H2SeO4
2. The Mg2+ and Cu2+ cations have similar ionic radii, but the acidity of aqueous nitrate solutions containing the two ions are different. Explain which ion is more acidic and why.
3. (a) Sulfides of As3+, Sb3+, and Sn4+ precipitate when ammonium hydrosulfide is added to aqueous solutions containing these ions. When excess sulfide is added they redissolve as anionic complexes. In contrast, solutions of Cu2+, Pb2+, Hg2+, Bi3+ and Cd2+ precipitate as sulfides but do not redissolve. We may consider the first group of ions to be amphoteric for soft acid-base reactions involving SH- in place of OH-. The second group is less acidic. Locate the amphoteric boundary in the periodic table for sulfides implied by this information. Compare this boundary with the amphoteric boundary for oxides reacting with H+ and OH-. Does this analysis agree with describing S2- as a softer base than O2-?
(b) Au+ is also precipitated by SH- ions, and the sulfide Au2S redissolves to form a soluble complex in excess SH-. Does this fit the trend you discovered in part (a)? Does it make sense in terms of the electronegativity of Au? Explain.
Atkins, P., Overton, T., Rourke, J., Weller, M., & Armstrong, F. (2010). Inorganic Chemistry (5th ed.). Oxford.