1. ACIDS AND BASES Properties of Acids and Bases Acid – Base Theories Strong and Weak Acids and Bases Understanding Indicators pH Scale Buffers and Antacids

2. Properties of Acids and Bases Macroscopic View: Acids Taste sour Produce painful sensation on skin React with certain metals (Mg, Zn, Fe) to produce H 2 gas React with limestone and baking soda to produce CO 2 Turn litmus paper red Bases Taste bitter Feel slippery on skin React with oils and greases Turn litmus paper blue React with acids to produce salt and water

3. Acids in Every Day Life Common Acid in the Home: Chemical NameCommon Name Hydrochloric Acid (HCL)Muratic Acid Acetic Acid (CH 3 COOH)Vinegar Sulfuric Acid (H 2 SO 4 )Auto Battery Acid Carbonic Acid (H 2 CO 3 )Carbonated Water Boric Acid (H 3 BO 4 )Antiseptic Eye Drops Acetylsalicylic Acid (C 16 H 12 O 6 )Aspirin

4. Acids

6. Acid Nomenclature

7. Bases in Every Day Life Common Bases in the Home: Chemical Name Common Name or Use Ammonia (NH 3 ) Cleaner Sodium Hydroxide (NaOH) Lye Sodium Bicarbonate (NaHCO 3 ) Baking Soda Magnesium Hydroxide (Mg(OH) 2 ) Milk of Magnesia Calcium Carbonate (CaCO 3 ) Antacid Aluminum Hydroxide (Al(OH) 3 ) Antacid

8. Bases

9. Properties of Acids and Bases Microscopic View: You may have noticed that all the acids contain hydrogen, while most of the bases contain the hydroxide ion (OH - ). Two main theories use these facts in their descriptions of acids and bases and their reactions: Arrhenius Theory Bronsted-Lowery Theory

10. Arrhenius Theory: Must have Water This was the first modern acid-base theory, and it tells us that when dissolved in water: –An acid yields H + ions HCl(aq)  H + + Cl - –A base yields OH - ions NaOH(aq)  Na + + OH -

11. Arrhenius Theory This theory also classifies the reaction between an acid and a base as a neutralization reaction, producing a neutral solution composed of a water and a salt. HCl(aq) + NaOH(aq)  H 2 O(l) + NaCl(aq) The water is formed from the combining of the H + and OH - ions Like all theories, this one has its limitations. Some bases don’t have hydroxide ions. To account for this, a new theory was developed.

12. Bronstead – Lowery Theory In this theory, an acid is classified as a proton (H + ) donor. A base is classified as a proton acceptor. The base accepts the H + by furnishing a pair of electrons for a coordinate-covalent bond.

13. Conjugate Pairs All acids have a conjugate base, which is formed when their proton has been donated; likewise, all bases have a conjugate acid, formed after they have accepted a proton.

14. One More – The Lewis Theory This theory extends well beyond the things you normally think of as acids and bases. The theory: –An acid is an electron pair acceptor. –A base is an electron pair donor.

15. Lewis Acid-Base Reactions

16. Water: H 2 O can function as both an acid and a base. In pure water, auto-ionization can occur: In a neutral solution, [H 3 O + ] = [OH - ]

17. Strong & Weak Acids It is important to remember that acid- base strength is not the same as concentration. Strength refers to the amount of ionization or breaking apart that a particular acid or base undergoes. Concentration refers to the amount of acid or base that you initially have. You can have a concentrated weak acid or a dilute strong acid…

18. Strong Acids Certain acids are considered to be strong, which means they are dissociated 100% in solution: –HClHydrochloric Acid –HNO 3 Nitric Acid –H 2 SO 4 Sulfuric Acid –HBrHydrobromic Acid –HIHydroiodic Acid –HClO 4 Perchloric Acid You ought to memorize this list, because almost every other acid is weak. The most common example is HCl.

19. Weak Acids A weak acid is one which doesn't ionize fully when it is dissolved in water. Ethanoic acid is a typical weak acid. It reacts with water to produce hydroxonium ions and ethanoate ions, but the back reaction is more successful than the forward one. The ions react very easily to reform the acid and the water. At any one time, only about 1% of the ethanoic acid molecules have converted into ions. The rest remain as simple ethanoic acid molecules.

20. Strong Bases A strong base is something like sodium hydroxide or potassium hydroxide which is fully ionic. You can think of the compound as being 100% split up into metal ions and hydroxide ions in solution.

21. Weak Bases A weak base is one which doesn't convert fully into hydroxide ions in solution. Ammonia is a typical weak base. Ammonia itself obviously doesn't contain hydroxide ions, but it reacts with water to produce ammonium ions and hydroxide ions. However, the reaction is reversible, and at any one time about 99% of the ammonia is still present as ammonia molecules. Only about 1% has actually produced hydroxide ions.

22. pH Scale The pH scale is a way of expressing the strength of acids and bases. Instead of using very small numbers, we just use the NEGATIVE power of 10 on the Molarity (concentration) of the H + or OH - ion. Acidic: pH < 7 Neutral: pH = 7 Basic: pH > 7

23. pH of Common Substances

24. Calculating pH pH = -log [H+] (Remember that the [ ] mean Molarity) Example: If [H + ] = 1 X 10 -10 pH = - log 1 X 10 -10 pH = - (- 10) pH = 10 Example: If [H + ] = 1.8 X 10 -5 pH = - log 1.8 X 10 -5 pH = - (- 4.74) pH = 4.74

25. pOH Since acids and bases are opposites, pH and pOH are opposites pOH does not really exist, but it is useful for changing bases to pH. pOH looks at the perspective of a base pOH = - log [OH-] Since pH and pOH are on opposite ends: pH + pOH = 14

26. pH[H+][OH-] pOH

27. pH Testing There are several ways to test pH Blue litmus paper (red = acid) Red litmus paper (blue = basic) pH paper (multi-colored) pH meter (7 is neutral, 7 base) Universal indicator (multi- colored) Indicators like phenolphthalein Natural indicators like red cabbage, radishes

28. pH Indicators Indicators are dyes that can be added that will change color in the presence of an acid or base. Some indicators only work in a specific range of pH Once the drops are added, the sample is ruined Some dyes are natural, like radish skin or red cabbage

29. Titrations Suppose you want to determine the molar concentration of an HCl solution. –You place a known volume in a flask and add a base indicator (like phenolphthalein) –You then add small amounts of a standardized base (like NaOH) with a buret –You keep adding base until the solution turns the faintest shade of pink –This is called the endpoint –Using the balanced equation and the amounts of acid and base used, you can calculate molar concentration. You can titrate a base with a standard acid solution in the same way.

30. Titration

31. 35.62 mL of NaOH is neutralized with 25.2 mL of 0.0998 M HCl by titration to an equivalence point. What is the concentration of the NaOH? M a V a = M b V b M a V a = M b V b (0.0998 M) (25.2 mL) = 0.0706 M (35.62 mL)

32. Preparing solutions by dilution: If you want to dilute a solution, use the following formula: M 1 V 1 = M 2 V 2 Example: You have a stock bottle of hydrochloric acid, which is 12.1 M. You need 400 mL of 0.10 M HCl. How much of the acid and how much water will you need? 3.3 mL of HCl and 396.7 mL of water

33. Buffers: Controlling pH Buffers, or buffer solutions, resist a change in pH caused by the addition of acids or bases. There are two types of buffers: 1.Mixtures of weak acids and bases – these may be conjugate acid-base pairs, or nonconjugate acid-base pairs 2.Amphoteric species – these are substances that can act either as an acid or a base, like water

34. Antacids: Good, Basic Chemistry The stomach secretes hydrochloric acid to activate enzymes that break down proteins. Sometimes the stomach produces too much acid and it can work its way back up the esophagus leading to heartburn. Antacids are compounds that neutralize the excess acid: –Bicarbonates – NaHCO 3 and KHCO 3 –Carbonates – CaCO 3 and MgCO 3 –Hydroxides – Al(OH) 3 and Mg(OH) 2