2. What are Acids and Bases ? There are three main theories to describe what acids and bases are In general each theory widens what chemical reactions can be considered “ Acid – Base” reactions, than the previous theory Let us now examine the three separate theories in detail

3. Theory 1 : The Arrhenius Theory of Acids and Bases Came in 1884 from his PhD work on proving that ions exist in water Work was not widely accepted at the time but he won the 1903 Nobel Prize in Chemistry for it Svante Arrhenius ( 1859 – 1927 )

4. According to the Arrhenius Theory, Acids and Bases are defined as follows, An Acid is a substance which produces hydrogen ions ( H + ) in aqueous solution HCl (aq) H + (aq) + Cl - (aq) Example: Hydrochloric Acid ( HCl ) A Base is a substance which produces hydroxide ions ( OH - ) in aqueous solution Example: Sodium Hydoxide ( NaOH ) NaOH (aq) Na + (aq) + OH - (aq)

5. Neutralisation When an Acid and a Base meet they can “cancel” each other out, forming a salt and water Acid + BaseSalt + Water HCl (aq) + NaOH (aq) NaCl + H 2 O There are 2 separate reactions combined in this equation 1. Na + + Cl - NaCl 2. H + + OH - H2OH2O According to Arrhenius`s Theory the charged hydrogen ion and the hydroxide ion bond together to form “Neutral” water

6. The Hydronium Ion ( H 3 O + ) Developments in Chemistry after Arrhenius showed that it was impossible for the hydrogen ion ( H + ) to exist freely in solution 1. Too Small - proton diameter 100,000 less than an atom 2. Too Reactive – “naked” + charge has to bind to something In water the H + ion immediately binds to a water molecule to form the Hydronium Ion as follows, H + + H 2 OH3O+H3O+ This fact does not really change our understanding of Acid-Base Chemistry too much. Just remember when you see H + in an equation it is really the hydronium ion H 3 O +

7. Evidence for the Hydronium Ion ( H 3 O + ) Conductivity: Positive ions in water move towards an anode at a rate consistent with them being the same size as a water molecule. Being 100,000 times smaller, protons would move much faster

8. Let us now re-examine the reaction between HCl and NaOH in the context of the hydronium ion ( H 3 O + ) HCl (aq) Cl - + H + H + + H 2 OH3O+H3O+ Overall equation: HCl (aq) Cl - + H 3 O + NaOH (aq) Na + + OH - HCl (aq) + NaOH (aq) NaCl +2H 2 O One H 2 O molecule was taken from solution to carry the H + charge so the original reaction on Slide 5 is still correct as only one new H 2 O molecule is actually produced

9. Limitations of the Arrhenius Definition 1.Many Acid-Base reactions happen in solutions other than water and some don`t require a solution at all 2.Some Acid- Base reactions don` t produce Hydronium ions ( H 3 O + ) or Hydroxide ions ( OH - ) Example: Ammonia (g) + Hydrogen Chloride (g) Ammonium Chloride (s ) NH 3 (g) + HCl (g) NH 4 Cl (s) Here Ammonia gas ( Base ) reacts with Hydrogen Chloride gas ( Acid ) to give Ammonium Chloride (solid) No solution +no Hydronium or Hydroxide ions =Need for new Theory!!

10. The Bronsted-Lowry definition of Acids and Bases Formulated independently in 1923 by the Danish chemist Johannes Nicolaus Bronsted and the English chemist Thomas Martin Lowry Both recognised the limitations of the Arrhenius definition and pro- posed a more general definition of Acids and Bases that addressed the problems on the previous slide BronstedLowry

11. According to the Bronsted-Lowry theory Acids and Bases are defined as following, an acid is a proton donor a base is a proton acceptor This definition solves all the problems Arrhenius had with this reaction NH 3(g) + HCl (g) NH 4 Cl (s) The Hydrogen Chloride gas (HCl ) donates a proton to the Ammonia gas (NH 3 ) leaving Cl -. This makes it an acid The Ammonia gas (NH 3 ) accepts a proton from the HCl becoming NH 4 +. This makes it a base Finally the Cl - ion binds ionically to the NH 4 + ion to form the salt ammonium chloride NH 4 Cl

12. Acid/Base Equilibria So far we have only discussed Acid-Base reactions in a “one way”, reactants just make products manner. A + B C + D Unfortunately things are not that straight forward and sometimes the products react together to form the reactants A + B C + D When we combine the above two equations we get a system which is said to be in Equilibrium. Reactants making products and products making reactants. A + B C + D We will now examine the concept of chemical equilibrium in relation to the Bronsted – Lowry theory of acids and bases

13. Conjugate Acids and Bases Let us examine the general case of an acid ( HA ) dissociating in water HA (aq) + H 2 O (l) H 3 O + (aq) + A - (aq) In the forward reaction HA acts as an acid by donating a proton to the H 2 O becoming A - In the backwards reaction A - acts as a base by accepting a proton from the H 3 O + ion becoming H 2 O We say that HA is the conjugate acid of A - and that A - is the conjugate base of HA Together they are called a conjugate acid-base pair

14. HA (aq) + H 2 O (l) H 3 O + (aq) + A - (aq) The same can be said of the water molecule involved in the reaction In the forward reaction it acts as a base by accepting a proton from HA and in the backwards reaction it acts as an acid by H 3 O + donating a proton to A - It is also a conjugate acid-base pair conjugate pair Because of the reversibility of Acid- Base reactions it is always possible to find two substances which act as an acid and two substances which act as bases

15. Example 2: A Base dissolving in Water NH 3(aq) + H 2 O (l) NH 4 + (aq) + OH - (aq) Here in the forward reaction H 2 O acts as an acid by donating a proton to NH 3 In the previous example ( an acid dissolving in water ) H 2 O acted as a base by accepting a proton from the acid Amphoteric Substance: A substance which can act as either an acid or a base conjugate pair

16. Strong Acids / Weak Acids : Strong Bases / Weak Bases Strong Acid: Eg: Hydrochloric Acid (HCl) HCl (aq) + H 2 O (l) H 3 O + (aq) + Cl - (aq) Cl - is a very weak conjugate base. It cannot even accept a proton from H 3 O + which is an excellent proton donor Strong Bronsted Acids have weak conjugate bases Although theoretically possible, in reality the backwards reaction does not happen so we can ignore it and rewrite the equation as; HCl (aq) + H 2 O (l) H 3 O + (aq) + Cl - (aq) Strong Acid: is one that virtually 100% ionizes in solution

17. Weak Acid: Eg. Ethanoic Acid (CH 3 COOH) CH 3 COOH + H 2 O CH 3 COO - + H 3 O + CH 3 COO - (Acetate) is a strong conjugate base which readily accepts a proton from H 3 O + to form H 2 O As a result the backward reaction is far more successful than the forward reaction with only around 1% of the CH 3 COOH ionised to CH 3 COO - in solution Weak Bronsted Acids have strong conjugate bases Weak Acid: is one which only partially ionizes in solution

18. Strong Base: Eg: Sodium Hydroxide NaOH Exactly the same as for a strong acid, the backwards reaction is ignored and the same definition is used Strong Base: is one that virtually 100% ionizes in water __________________________________________________________ Weak Base: Eg: Ammonia NH 3 NH 3(aq) + H 2 O (l) NH 4 + (aq) + OH - (aq) NH 4 + (Ammonium) is a strong conjugate acid which readily donates a proton to OH - to form H 2 O in the backwards reaction Weak Bronsted Bases have strong conjugate acids Weak Base: is one that only partially ionizes in solution

19. Theory 3 : The Lewis Theory of Acids and Bases Proposed by American Chemist Gilbert Newton Lewis in 1923, the same year as Bronsted-Lowry theory Looked more at the chemical bonding involved in Acids and Bases Gilbert N. Lewis

20. According to the Lewis Theory of Acids and Bases, an acid is an electron pair acceptor a base is an electron pair donor