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: Acids: Hydrocholoric acid (HCl) Sulfuric acid (H 2 SO 4 ) Nitric acid (HNO 3 ) Acetic acid (HC 2 H 3 O 2 ) Bases: Sodium Hydroxide (NaOH) Potassium Hydroxide (KOH) Calcium hydroxide Ca(OH) 2 Ammonia (NH 3 )

3. Introduction  Acids and bases are special substances with very distinct properties. It is good think of acids and bases as opposites. Key characteristics of acidsKey characteristics of bases -Sour taste (eg. Lemons, grapefruit, vinegar, sour milk) -React with active metals such as zinc and magnesium to produce hydrogen gas -Form electrolytic solutions (conduct electricity) because they produce ions -Cause certain dyes to change color (litmus paper turns red) -Neutralized by bases (neutralized means that the substance no longer has acidic or basic properties) -Bitter taste -Generally no noticeable reaction with active metals -Form electrolytic solutions (conduct electricity) because they produce ions -Cause certain dyes to change color (litmus paper turns blue) -Slippery feel (eg. soapy feel) --Neutralized by acids

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? Acids: Hydrocholoric acid (HCl) Sulfuric acid (H 2 SO 4 ) Nitric acid (HNO 3 ) Acetic acid (HC 2 H 3 O 2 ) Bases: Sodium Hydroxide (NaOH) Potassium Hydroxide (KOH) Calcium hydroxide Ca(OH) 2 Ammonia (NH 3 )

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 H 2 O 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)  H 2 SO 4(g) 2H + (aq) + SO 4 2- (aq)

9. Ionization  So what is actually happening?  Evidence suggests that the hydrogen ion actually bonds to a water molecule forming a hydronium ion, H 3 O +.  Ex: HCl (g) + H 2 O(l) H 3 O + (aq) + Cl - (aq)  Ex: H 2 SO 4(g) + 2H 2 O(l) 2 H 3 O + (aq) + SO 4 2- (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 NH 3 considered a base?  This may be a problem at first, but lets look at what happens when we add ammonia to water:  NH 3(g) + H 2 O (l)  NH 4 + (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 + H 2 O H 3 O + + Cl -

13. Proton Acceptance  Getting back to ammonia, we will see how a base can accept a hydrogen ion.  NH 3(g) + H 2 O (l) NH 4 + (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 NH 3 can become NH 4 + by gaining a hydrogen ion, then lets consider the reverse – that is, NH 4 should be able to change back to NH 3 by losing a hydrogen ion.  Since we have defined NH 3 as a base because it can accept an H +, then its partner ion, NH 4 + can be considered an acid since it can give up an H + to become NH 3.

15. Conjugate Acid-Base Pairs  Let’s consider water now. In the same equation, H 2 O gives up an H + to ammonia – therefore, we should be able to consider H 2 O an acid.  However, in the reverse reaction, H 2 O’s partner ion, OH -, accepts the H + from NH 4 + 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) H 2 SO 4

18. Answers  A) F -  B) HSO 4 -

19. Examples  Example 2: Write the conjugate acids for the following bases:  A) PO 4 3-  B) SO 4 2-

20. Answers  A) HPO 4 -2  B) HSO 4 -

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) + H 2 O (l) H 3 O + (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) HSO 4 - + H 3 O + H 2 SO 4 + H 2 O  B) HSO 4 - + OH - SO 4 2- + H 2 O  Answer:  Forward: HSO 4 -, A = base, B = acid  Reverse: H 2 0, A = Base, B = acid