6.1 – Introduction to Acids and Bases

1. 6.1 – Introduction to Acids and Bases Unit 6 – Acids and Bases

2. Introduction • In your years of studying chemistry, you have probably come across a few common acids and bases:

3. Introduction • Acids and bases are special substances with very distinct properties. It is good think of acids and bases as opposites.

4. Arrhenius’ Theory • Looking at our list of acids and bases, what do you see that is common between the acids? Most of the bases?

5. Arrhenius’ Theory • In the 1880’s, Svante Arrhenius determined that acids had their characteristic properties due to the presence of hydrogen ions, H+. • Likewise, he discovered the properties of bases are due to the presence of hydroxide ions, OH-. • These two observations together is known as the Arrhenius Theory of Acids and Bases.

6. Dissociation • Dissociation will be important in this unit. • Remember that this process is when an ionic compound is mixed with water. • Dissociation of ionic compounds occurs when water molecules “pull apart” the ionic crystal. • This occurs due to strong attractions between the polar ends of the water molecule and the positive and negative ions within the crystal. • Water molecules then surround the positive cations and negative anions • KOH(s) K+(aq) + OH-(aq) • Note that bases undergo dissociation.

7. Dissociation • There are two important things to notice about writing dissociation equations: • Generally DO NOT include H2O as a reactant. We know something has been dissolved in water when we see the (aq) notation. We will make some exceptions later to this rule • Ion charges MUST BE included!

8. Ionization • Ionization is the process of dissolving molecular compounds (covalently bonded) in water to produce ions. • Most molecular compounds do not undergo ionization. However, acids ALWAYS do. • In fact, all acids produce hydrogen ions in a solution. • HCl(g) H+(aq) + Cl-(aq) • H2SO4(g) 2H+(aq) + SO42-(aq)

9. Ionization • So what is actually happening? • Evidence suggests that the hydrogen ion actually bonds to a water molecule forming a hydronium ion, H3O+. • Ex: HCl(g) + H2O(l) H3O+(aq) + Cl-(aq) • Ex: H2SO4(g)+ 2H2O(l) 2H3O+(aq) + SO42-(aq) • You should be comfortable using either method of representation: one will mean the same as the other.

10. Why the Arrhenius Theory Isn’t Good Enough • Up until this point, we have said that a substance that produces H+ ions is an acid and one that produces OH- ions is a base • So… why is NH3 considered a base? • This may be a problem at first, but lets look at what happens when we add ammonia to water: • NH3(g) + H2O(l) NH4+(aq) + OH-(aq) • The Arrhenius Theory is unable to explain this occurrence. Luckily we have an alternative theory that works just fine…

11. Bronsted and Lowry Theory of Acids & Bases • 2 chemists working independently, Johannes Bronsted and Thomas Lowry, came up with what is now known as the “Bronsted-Lowry Theory of Acids and Bases.” • This theory states that acids are substances that can DONATE a hydrogen ion, • and bases are substances that can ACCEPT a hydrogen ion.

12. Proton Donation • How are acids “donors?” • HCl H+ + Cl- • This shows that HCl produces an H+, but to donate implies that something will receive the H+. So, we can see the donation with the ionization equation: • HCl + H2O H3O+ + Cl-

13. Proton Acceptance • Getting back to ammonia , we will see how a base can accept a hydrogen ion. • NH3(g) + H2O(l) NH4+(aq) + OH-(aq) • Notice that the ammonia has become an ammonium ion by accepting a H+ from the water. • The H+ that came from the water left its electrons behind with the remaining OH-, which gives us an H+ and an OH-.

14. Conjugate Acid-Base Pairs • Now, if NH3 can become NH4+ by gaining a hydrogen ion, then lets consider the reverse – that is, NH4 should be able to change back to NH3 by losing a hydrogen ion. • Since we have defined NH3 as a base because it can accept an H+, then its partner ion, NH4+ can be considered an acid since it can give up an H+ to become NH3.

15. Conjugate Acid-Base Pairs • Let’s consider water now. In the same equation, H2O gives up an H+ to ammonia – therefore, we should be able to consider H2O an acid. • However, in the reverse reaction, H2O’s partner ion, OH-, accepts the H+ from NH4+ to become water. This accepting of H+ makes it a base! • These two examples are called conjugate acid-base pairs.

16. Conjugate Acid-Base Pairs • Conjugate acid-base pairs differ from each other by the presence or absence of a single hydrogen ion (or proton). • Every acid has a conjugate base, and every base has a conjugate acid. • We can now express these equations with a double arrow, since it represents acid-base equilibrium

17. Examples • Example 1: Write the conjugate bases for the following acids: • A) HF • B) H2SO4

18. Answers • A) F- • B) HSO4-

19. Examples • Example 2: Write the conjugate acids for the following bases: • A) PO43- • B) SO42-

20. Answers • A) HPO4-2 • B) HSO4-

21. Amphoteric Substances • Notice that in the ammonia example, water acted as a acid. However, how does water react in the following reaction? • HCl(g) + H2O(l) H3O+(aq) + Cl-(aq) • Since water is accepting an H+ it is considered a base. • “Amphoteric substances” are those that act as an acid in one reaction but a base in another.

22. Example • Example 3: In the following two reactions which substance is amphoteric? When is it an acid? A base? • A) HSO4- + H3O+ H2SO4 + H2O • B) HSO4- + OH- SO42- + H2O • Answer: • Forward: HSO4-, A = base, B = acid • Reverse: H20, A = Base, B = acid