Common acid-base theories

1. Common acid-base theories By: Dr. O. Rajabi (Pharm.D.- Ph.D.) Associate Professor of Chemistry Department of Medicinal chemistry Mashad University of Medical Sciences

2. Acid and Bases

3. Acid and Bases

4. Acid and Bases

5. 1- Lavoisier definition August 26, 1743 May 8, 1794 Lavoisier's knowledge of strong acids was mainly restricted to oxyacids, which tend to contain central atoms in high oxidation states surrounded by oxygen, such as HNO3 and H2SO4

6. An oxoacid is an acid which contains oxygen. More specifically, it is an acid which: contains oxygen; contains at least one other element; has at least one hydrogen atom bound to oxygen; and forms an ion by the loss of one or more protons

7. The name oxyacid is sometimes used, although this is not recommended Generally, oxoacids are simply polyatomic ions with a hydrogen cation Although carboxylic acids fulfill the criteria above, they are not generally considered as oxoacids

9. 2- Liebig definition This definition was proposed by Justus von Liebig circa 1838, based on his extensive works on the chemical composition of organic acids. (May 12, 1803 – April 18, 1873)

10. This finished the doctrinal shift from oxygen-based acids to hydrogen-based acids, started by Davy According to Liebig, an acid is a hydrogen-containing substance in which the hydrogen could be replaced by a metal. Liebig's definition, while completely empirical, remained in use for almost 50 years until the adoption of the Arrhenius definition.

11. 3- Arrhenius definition The Arrhenius definition of acid-base reactions is a more simplified acid-base concept devised by Svante Arrhenius, which was used to provide a modern definition of bases that followed from his work with Friedrich Wilhelm Ostwald in establishing the presence of ions in aqueous solution in 1884 Nobel prize in chemistry in 1903

12. As defined at the time of discovery, acid-base reactions are characterized by Arrhenius acids, which dissociate in aqueous solution form hydrogen or the later-termed oxonium (H3O+) ions, and Arrhenius bases which form hydroxide (OH-) ions. More recent IUPAC recommendations now suggest the newer term "hydronium“ be used in favor of the older accepted term "oxonium“.

13. Arrhenius acids form hydrogen ions in aqueous solution with Arrhenius bases forming hydroxide ions. acid+ + base- → salt + water 2NaOH + H2SO4 → 2 H2O + Na2SO4

14. 4- Brønsted-Lowry definition 1923 professor of inorganic and physical chemistry at Copenhagen Englishphysical chemist

15. The Brønsted-Lowry definition, formulated independently by its two proponents Johannes Nicolaus Brønsted and Martin Lowry in 1923 is based upon the idea of protonation of bases through the de-protonation of acids -- more commonly referred to as the ability of acids to "donate" hydrogen ions (H+) or protons to bases, which "accept" them. CH3COOH + H2O === CH3COO- + H3O+ NH3 + H2O === NH4+ + OH- CH3COO- + H2O === CH3COOH + OH- NH4+ + H2O === NH3 + H3O+

16. In contrast to the Arrhenius definition, the Brønsted-Lowry definition refers to the products of an acid-base reaction as conjugate acids and bases to refer to the relation of one proton, and to indicate that there has been a reaction between the two quantities, rather than a "formation" of salt and water, as explained in the Arrhenius definition • AH + B → BH+ + A- • General formula for representing Brønsted-Lowry reactions. • HCl (aq) + H2O → H3O+ (aq) + Cl- (aq) • Hydrochloric acid completely reacts with water to form the hydronium and chloride ions • CH3COOH + NH3 → NH4+ + CH3COO- • Acetic acid reacts incompletely with ammonia, no hydronium ions being produced

17. In differentiation from the Arrhenius definition, the Brønsted-Lowry definition postulates that for each acid, there is a conjugate acid and base or "conjugate acid-base pair" that is formed through a complete reaction, which also includes water, which is amphoteric For example, zinc oxide (ZnO) reacts differently depending on the pH of the solution: In acids: ZnO + 2H+ → Zn2+ + H2O In bases: ZnO + H2O + 2OH- → [Zn(OH)4]2- • Base (Proton Acceptor): H2O + HCl → H3O+ + Cl− • Acid (Proton Donor): H2O + NH3 → NH4+ + OH− • (Indeed, it can do both at once: 2H2O → H3O+ + OH−)

18. Ionization + H H O H O Cl Cl H H H • O (e.g.H2SO4) was originally thought to cause acidic properties. Later, H was implicated, but it was still not clear why CH4 was neutral. • Arrhenius made the revolutionary suggestion that some solutions contain ions & that acids produce H3O+ (hydronium) ions in solution. + + • The more recent Bronsted-Lowry concept is that acids are H+ (proton) donors and bases are proton acceptors

19. + + + H H O H O Cl Cl H H H The Bronsted-Lowry concept • In this idea, the ionization of an acid by water is just one example of an acid-base reaction. conjugate acid conjugate base acid base conjugate acid-base pairs • Acids and bases are identified based on whether they donate or accept H+. • “Conjugate” acids and bases are found on the products side of the equation. A conjugate base is the same as the starting acid minus H+.

20. Practice problems Identify the acid, base, conjugate acid, conjugate base, and conjugate acid-base pairs: HC2H3O2(aq) + H2O(l)  C2H3O2–(aq) + H3O+(aq) conjugate base conjugate acid acid base conjugate acid-base pairs OH–(aq) + HCO3–(aq)  CO32–(aq) + H2O(l) base acid conjugate base conjugate acid conjugate acid-base pairs

21. Acid – Base Strength • Strong Acid - Transfers all protons to water; completely ionizes; strong electrolyte; conjugate base is weaker and has negligible tendency to be protonated. • Weak Acid - Fraction of protons transferred to water; partly ionized; weak electrolyte; conjugate base is stronger readily accepting protons from water • As acid strength decreases, base strength increases. • The stronger the acid, the weaker its conjugate base • The weaker the acid, the stronger its conjugate base

23. Leveling effect The term leveling effect refers to a solvent's ability to level the effect of a strong acid or base dissolved in it. When a strong acid is dissolved in water, it reacts with it to form H3O+ in the following reaction (where "HA" is a generic strong acid such as "HCl"): HA + H2O -> A- + H3O+ Any acid that is stronger than H3O+ reacts with H2O to form H3O+; therefore, no acid can be stronger than H3O+ in H2O. This is true with any solvent, although usually to a lesser extent than with water. Bases may be also leveled in solvents

24. Which one is stronger? A) HCl B) HClO4 Becouse of this leveing effect, HClO4 and HCl behave as if they had the same acid strength; both are leveled to H3O+: HClO4 + H2O ---- H3O+ + ClO- HCl + H2O ---- H3O+ + Cl- In acetic acid solvent, which is less basic than H2O, HClO4 and HCl are not leveled to the same strength: HClO4 + CH3COOH === CH3COOH2+ + ClO4_ K= 1.3 × 10-5 HCl + CH3COOH === CH3COOH2+ + Cl- K= 2.8 × 10-9

25. 5-Lewis definition a famous Americanphysical chemist known for his 1902 Lewis dot structures, his paper "The Atom and the Molecule", which is the foundation of modern valence bond theory, developed in coordination with Irving Langmuir . his 1923 textbook Thermodynamics and the Free Energy of Chemical Substances, written in coordination with Merle Randall, one of the founding books in chemical thermodynamics Gilbert Newton Lewis (October 23, 1875 - March 23, 1946) In 1926, Lewis coined the term "photon" for the smallest unit of radiant energy.

26. The Lewis definition of acid base reactions, devised by Gilbert N. Lewis in 1923 is an encompassing theory to the Brønsted-Lowry and solvent-system definitions with regards to the premise of a donation mechanism, which conversely attributes the donation of electron pairs from bases and the acceptance by acids, rather than protons or other bonded substances and spans both aqueous and non-aqueous reactions Ag+ + 2 :NH3 → [H3N:Ag:NH3]+ A silver cation reacts as an acid with ammonia which acts as an electron-pair donor, forming an ammonia-silver adduct • Acid - An electron pair acceptor • Base - An electron pair donor

27. 6-Solvent-system definition This definition is based on a generalization of the earlier Arrhenius definition to all autodissociating solvents. In all such solvents there is a certain concentration of a positive species, solvonium cations and negative species, solvate anions, in equilibrium with the neutral solvent molecules. For example: • 2H2O ⇌ H3O+ (hydronium) + OH- (hydroxide) • 2NH3 ⇌ NH4+ (ammonium) + NH2− (amide) • or even some aprotic systems: • N2O4 ⇌ NO+ (nitrosonium) + NO3− (nitrate) • 2SbCl3 ⇌ SbCl2+ (dichloroantimonium) + SbCl4- (tetrachloroantimonate)

28. A solute causing an increase in the concentration of the solvonium ions and a decrease in the solvate ions is an acid and one causing the reverse is a base. Thus, in liquid ammonia, KNH2 (supplying NH2-) is a strong base, and NH4NO3 (supplying NH4+) is a strong acid.

29. Other acid-base theories Usanovich definition Lux-Flood definition Pearson definition

30. What is a strong Acid? An Acid that is 100% ionized in water. Strong Acids: 100% ionized (completely dissociated) in water. HCl + H2O  H3O+ + Cl- often written as: HCl  H+ + Cl-

31. Strong Acids: 100% ionized (completely dissociated) in water. HCl + H2O  H3O+ + Cl- Strong Acids: Perchloric HClO4 Chloric, HClO3 Hydrobromic, HBr Hydrochloric, HCl Hydroiodic, HI Nitric, HNO3 Sulfuric, H2SO4

32. What is a strong Base? A base that is completely dissociated in water (highly soluble). NaOH(s)  Na+ + OH- Strong Bases: Group 1A metal hydroxides (LiOH, NaOH, KOH, RbOH, CsOH) Heavy Group 2A metal hydroxides [Ca(OH)2, Sr(OH)2, and Ba(OH)2]

33. Strong Acids: 100% ionized (completely dissociated) in water. HCl + H2O  H3O+ + Cl- Note the “one way arrow”. Weak Acids: Only a small % (dissociated) in water. HC2H3O2 + H2O H3O+ + C2H3O2- Note the “2-way” arrow. Why are they different?

34. Strong Acids: HCl HCl HCl HCl HCl (H2O) ADD WATER to MOLECULAR ACID

35. Strong Acids: Cl- H3O+ (H2O) Cl- H3O+ H3O+ Cl- Cl- H3O+ H3O+ Cl- Note: No HCl molecules remain in solution, all have been ionized in water.

36. Weak Acid Ionization: HC2H3O2 HC2H3O2 (H2O) HC2H3O2 HC2H3O2  HC2H3O2 Add water to MOLECULES of WEAK Acid

37. Weak Acid Ionization: HC2H3O2 HC2H3O2 H30+ C2H3O2- HC2H3O2 (H2O) HC2H3O2  H30+ C2H3O2- HC2H3O2 HC2H3O2 Note: At any given time only a small portion of the acid molecules are ionized and since reactions are running in BOTH directions the mixture composition stays the same. This gives rise to an Equilbrium expression, Ka

38. Chemical equilibrium In a chemical process, chemical equilibrium is the state in which the chemical activities or concentrations of the reactants and products have no net change over time. Usually, this would be the state that results when the forward chemical process proceeds at the same rate as their reverse reaction. The reaction rates of the forward and reverse reactions are generally not zero but, being equal, there are no net changes in any of the reactant or product concentrations. This process is called as dynamic equilibrium

39. Berthollet was one of the first chemists to recognize the characteristics of a reverse reaction, and hence, chemical equilibrium. Potassium chlorate (KClO3), a strong oxidizer, is known as Berthollet's Salt. Non-stoichiometric compounds are also named berthollides in his honor. Claude Louis Berthollet. Berthollet was one of several scientists who went with Napoleon to Egypt !

40. The concept of chemical equilibrium was developed after Berthollet (1803) found that some chemical reactions are reversible. For any reaction such as: αA + βB === γC +δD to be at equilibrium the rates of the forward and backward (reverse) reactions have to be equal. In this chemical equation with harpoon arrows pointing both ways to indicate equilibrium, A and B are reactant chemical species, C and D are product species, and α, β, σ, and γ are the stoichiometric coefficients of the respective reactants and products. Forward reaction rate = K+[A]α[B]β Backward reaction rate = K-[C]γ[D]δ

41. where A, B, C and D are active masses and k+ and k− are rate constants. Since forward and backward rates are equal: K+[A]α[B]β= K-[C]γ[D]δ and the ratio of the rate constants is also a constant, now known as an equilibrium constant K= K+/K- = [C]γ[D]δ/[A]α[B]β