Unit 9 – Acids & Bases

1. Unit 9 – Acids & Bases

2. Acids and Bases – Part 1 Launch Activity – brainstorming Demos Properties

3. Essential Questions • What are acids and bases? • How can they be identified or classified? • What are some properties of acids and bases? • Do their strengths differ? • How can we measure their strengths? • How can they be neutralized?

4. Launch Activity • What are acids and bases? What do they look, smell, and feel like? What do you think about when we talk about acids and bases and our environment? • Perhaps food like fruit • Or cleaning supplies • Maybe acid rain • How about your lawn care • Or your pool chemicals • What about medicines and antacids like tums?

5. Launch Activity • Divide up into four groups • Work as a group to fill out the index cards with: • Examples of acids • Examples of bases • Properties and/or characteristics of acids • Properties and/or characteristics of bases • Tape your cards under the correct name Acid or Base at the front board

6. Demos • What do acids and bases do to different litmus paper? • Do acids and bases conduct electricity? • Do acids and bases react with metals? • Add this new information to the board! • At the end of the activity – students should write down the lists on their notes

7. Acids • The word acid comes from Latin word acidus meaning sour, sharp, or tart taste. • Vinegar means sour wine. Vinegar is dilute acetic acid. • Citrus fruits also have a sour or tart taste due to a different acid called citric acid. • Sweet-Tarts get their sour or tart taste from citric acid.

8. Properties of Acids (add to your box if not already there) • pH < 7 (lower value = stronger acid) • sour taste • unremarkable feel (not slippery) • electrolytes (conduct electricity) • turn litmus red • react with metals to form H2 gas • corrosive, which means they break down certain substances (many acids can corrode fabric, skin, paper) • neutralized when reacted with a base

9. Common Acids • HCl = hydrochloric acid - stomach • H2SO4 = sulfuric acid - car batteries • HNO3 = nitric acid - explosives • HC2H3O2 = acetic acid - vinegar • H2CO3 = carbonic acid - sodas • H3PO4 = phosphoric acid – flavorings • What do they all have in common?

10. Bases • What is the opposite of an acid? • The opposite would be something that can neutralize or cancel the acid. • The name for the opposite of acid is alkaline or basic. The common alkaline compounds come from the Alkali Metals or the Alkaline Earth Metals (no big surprise). These metal hydroxides are alkaline because they release "OH-" that can neutralize "H+" by turning it into water.

11. Properties of Bases • pH > 7 (higher value = stronger base) • bitter taste • slippery feel in water solution (hard to wash off skin) • electrolytes (conduct electricity) • turn litmus blue • generally do not react with metals • caustic (dissolves protein, i.e. YOU) • corrosive • dissolves fats and oils • interact strongly with dirt and grease • neutralized when reacted with an acid

12. Common Bases • NaOH = sodium hydroxide - soaps, drain cleaner • KOH = potassium hydroxide - oven cleaner • Mg(OH)2 = magnesium hydroxide - antacids • Ca(OH)2 = calcium hydroxide - antacids • Al(OH)3 = aluminum hydroxide - antacids, deodorants • NH4OH = ammonium hydroxide “ammonia” - glass cleaner • What do they all have in common?

13. Indicators • Indicators tell us if a substance is acid, base or neutral • Common indicators: Universal indicator, Litmus paper, Phenolphthalein, Methyl orange

14. pH Indicator Ranges & Colors

15. Salts Created from a neutralization reaction between an aqueous acid and an aqueous base Ionic compounds made up from a metal cation from the base and a nonmetalanion from the acid Double replacement reaction Many, but not all, are soluble Soluble salts dissociate in water and will conduct electricity and are called electrolytes Insoluble salts are non-electrolytes

16. Ever heard of Acid Rain? • Acid rain is formed when sulfur dioxide (SO2) and nitrogen oxides (NOx) emitted from thermal power stations, industry and motor vehicles react with water to form acids such as sulfuric acid and nitric acid. • Pure water has a pH of 7.0. Normal rain is slightly acidic because carbon dioxide dissolves into it, so it has a pH of about 5.5. As of the year 2000, the most acidic rain falling in the US has a pH of about 4.3.

18. Magic Kettle • Magic Kettle • https://www.youtube.com/watch?v=ujkuW-0cpNw • (12:51) • BBC Acids and Bases • http://www.bbc.co.uk/schools/gcsebitesize/science/add_aqa/acids/acidsbasesact.shtml

19. Acids and Bases – Part 2 Classifying and defining acids and bases Identifying them in reactions

20. Earliest Times In an attempt to unravel nature’s secrets, early scientists would taste stuff and then categorize it as ‘sour’, ‘bitter’, ‘salty’, ‘sweet’, etc. (What a brilliant idea – tasting stuff!) Sour-tasting stuff gave rise to the word ‘acid’, which comes from the Greek word oxein and mutated into the Latin verb acere meaning ‘to make sour’. Bitter-tasting stuff gave rise to the word ‘alkaline’, which was derived from the Arabic word alqily meaning ‘roasted in a pan’ or ‘the ashes of plants’.

21. Arrhenius Acids and Bases In the 1800’s chemical concepts were based on the reactions of aqueous solutions. Arrhenius developed his concept of acids and bases relevant to reactions in H2O. Arrhenius Acid – produces hydrogen ions in aqueous solutions (water). HCl (aq) → H+(aq) + Cl-(aq) Arrhenius Base – produces hydroxide ions in aqueous solutions (water). NaOH (aq) → Na+(aq) + OH-(aq)

22. Limitations of Arrhenius Theory • It restricts acids and bases to water solutions. • Only compounds with OH- can be classified as a base. • But what about ammonia, NH3? It has the properties of a base, but would not be recognized as an Arrhenius base …

23. Brønsted – Lowry Acids and Bases To overcome the limitations of the Arrhenius definition, chemists needed to define acids and bases in a broader, more general way. Enter Brønsted and Lowry. Brønsted - Lowry Acid – donates a hydrogen ion (or H+ or proton) in a reaction. Brønsted - Lowry Base – accepts a hydrogen ion (or H+ or proton) in a reaction.

24. Brønsted–Lowry Acids and Bases The Brønsted - Lowry system involves a reaction. Acids and bases always come in pairs. HCl + H2O↔ H3O+ + Cl- HCl donates a H+ to H2O so it is the acid H2O accepts the H+ from HCl so it is the base

25. Brønsted–Lowry Acids and Bases What about this reaction? NH3 + H2O↔ NH4+ + OH-  H2O donates a H+ to NH3 so it is the acid  NH3 accepts the H+ from H2O so it is the base

26. What About This …? Water “goes both ways” can be an acid, can be a base NH3 + H2O ↔ NH4+ + OH- baseacid c.a. c.b. HCl + H2O ↔ H3O+ + Cl- acid base c.a. c.b. Amphoteric – a substance that can act as both an acid and base

27. Brønsted – Lowry Acids and Bases How do we label the products on the right-hand side here? HCl + H2O↔ H3O+ + Cl- acidbase H3O+ is the particle formed when the original base accepts the H+ from HCl and is called the conjugate acid. Cl- is the remainder when the original acid donates the H+ from HCl and is called the conjugate base.

28. Why “Conjugate” Acid & Base? What does the double arrow mean? Let’s label these: HCl + H2O↔ H3O+ + Cl- NH3 + H2O ↔ NH4+ + OH- NH3 + H2O ↔ NH4+ + OH-

29. Can we say it another way • Crash course in Acids and Bases • http://www.youtube.com/watch?v=ANi709MYnWg

30. Identify the acid and base • HCl (aq) + H2O  H3O+ (aq) + Cl – (aq)

31. Identify the acid or base? • H2CO3 (g) + H2O (l)  H3O+ (aq) + HCO3 – (aq) • HCl (aq) + NH3  NH4+ (aq) + Cl – (aq)

32. Stop – let’s practice • Work with a partner to complete the Bronsted-Lowry Acids and Bases and Conjugate Acid-Base Pairs worksheets

33. Acids and Bases – Part 3 Strong vs. Weak Acids and Bases pH scale

34. Measuring the strength of acids and bases The strengths and weaknesses of acids and bases (Ted-ed) http://ed.ted.com/lessons/the-strengths-and-weaknesses-of-acids-and-bases-george-zaidan-and-charles-morton#review Bozeman’s Acids, Bases, and pH (9 min) http://www.youtube.com/watch?v=Xeuyc55LqiY

35. Demo • Can you predict which is a stronger acid? 0.1 M HCl or 0.1 M HC2H3O2 How about the stronger base? 0.1 M NaOH or 0.1M NH4OH

36. Strength The strength of acids and bases is determined by the degree to which they ionize in water.

37. So how can we see which is stronger? • How about testing for dissolved ions • Which is stronger? 0.1 M HCl or 0.1 M HC2H3O2 0.1 M NaOH or 0.1 M NH4OH

38. Strong Acids &Bases • Strong Acids and Strong Bases completely, or nearly completely, ionize in water making them excellent conductors of electricity • Acids – Hydrochloric, sulfuric, nitric acid • Bases – Sodium hydroxide, potassium hydroxide

39. Weak Acids & Bases • Weak Acids and Weak Bases only partially ionize in water making them poor conductors of electricity • Acids – Acetic, phosphoric, carbonic acid • Bases – Ammonia, baking soda

40. Concentration The concentration of acids and bases is determined by the quantity of acid or base dissolved in a solution Strength and Concentration are very different!

42. Strength vs Concentration • A very dilute strong acid or base may have little corrosive action because of the low concentration of H+ or OH- ions, while • A very concentration solution of a weak acid or base may be very dangerous

43. Amphoteric Water We’ve talked about how water is amphoteric or can be an acid, base or neutral. Water can react with itself to donate and accept H+ ions creating hydronium and hydroxide ions

44. Amphoteric Water • Small but equal amount of H3O+ and OH– is created. • Specifically, at 25°C, [H+] = [OH–] ≈ 1x10-7 moles/L Do you see any connection between this and a neutral pH of 7.0?

45. Kw of Water K = [H+][OH–] = (1×10-7)(1×10-7) = 1.0×10-14 This is called Kw - Ionization constant • Ion product constant Kw ALWAYS = 1.0 x 10-14 at 25 oC! This equilibrium constant is very important because it applies to all aqueous solutions — acids, bases, salts, and non-electrolytes — not just to pure water. Therefore if [H+] goes up [OH-] goes down Do you see any connection between this and a pH scale of 0–14?

46. A Sliding Scale pH 1pH 7pH 14 acid ← stronger neutral stronger →base The pH scale according to the late Dr. Hubert Alyea, Princeton University • pH values less than 7 have high [H+] and are acidic (6, 5, 4 . . .) • pH values greater than 7 have low [H+] and are basic (8, 9, 10 . . .)

47. The pH Scale: 0 to 14 pH = the power of hydrogen Formula: pH = –log[H+] where [H+] is the concentration of hydrogen ions in moles/L Most solutions have a pH between 0 and 14. Acidic solutions have a pH less than 7. As a solution becomes more acidic, the pH decreases. Basic solutions have a pH greater than 7. As a solution becomes more basic, the pH increases.

48. pH Scale • Measures the acidity and basicity of a solution • Represents the relative concentration of hydrogen and hydroxide ions • Used to simplify discussions concerning the amount of acid or base in the solution • Uses a logarithmic scale Image used courtesy of: http://en.wikipedia.org/wiki/PH

49. A Logarithm is an Exponent • The pH scale uses powers of ten to express the hydrogen ion concentration. • Because pH numbers represent log values based on the power of 10, • pH 5 is 10 times more acidic than pH 6 • pH 4 is 100 times more acidic than pH 6 • It’s a bit like the way earthquakes are measured on the Richter scale • 6.2-6.3 is different from 6.1-6.2 • pH values are not always whole numbers

50. pH Practice Problems Find the pH for solutions having the following [H+]: 1 x 10-4M 3) 1 x 10-11M 1 x 10-7M 4) 1.8 X 10-5 M 5) What is the pH of a 0.15 M solution of HCl (a strong acid)? What is the pH of these acid solutions? 6) 2 x 10-3 M HNO3 7) 3.00 X 10-7M HNO3