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Acids and Bases

Acids and Bases. Mr Field. Using this slide show. The slide show is here to provide structure to the lessons, but not to limit them….go off-piste when you need to!

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Acids and Bases

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  1. Acids and Bases Mr Field

  2. Using this slide show • The slide show is here to provide structure to the lessons, but not to limit them….go off-piste when you need to! • Slide shows should be shared with students (preferable electronic to save paper) and they should add their own notes as they go along. • A good tip for students to improve understanding of the calculations is to get them to highlight numbers in the question and through the maths in different colours so they can see where numbers are coming from and going to. • The slide show is designed for my teaching style, and contains only the bare minimum of explanation, which I will elaborate on as I present it. Please adapt it to your teaching style, and add any notes that you feel necessary.

  3. Menu: • Lesson 1 – Theories • Lesson 2 – Properties of Acids and Bases- • Lesson 3 – The pH Scale • Lesson 4 – Strong and Weak Acids and Bases • Lesson 5 – HL – pH, pOH and the Ionic Product of Water • Lesson 6 – HL – Ka/pKa and Kb/pKb • Lesson 7 – HL – Weak Acid/Base Calculations • Lesson 8 – HL – Buffers • Lesson 9 – HL – Salt Hydrolysis • Lesson 10 – HL– Acid-Base Titrations • Lesson 11 – HL– Indicators • Lesson 12 – Test • Lesson 13 – Test Debrief

  4. Lesson 1 Theories of Acids and Bases

  5. Overview • Copy this onto an A4 page. You should add to it as a regular review throughout the unit.

  6. Assessment • This unit will be assessed by a test (100%) • Approximately lesson 12 • The test will include both the Acids and Bases and the Oxidation and Reduction topics

  7. We Are Here

  8. Lesson 1: Theories of Acids/Bases • Objectives: • Reflect on prior knowledge of acids and bases • Understand the Bronsted-Lowry theory of acidity and identify Bronsted-Lowry acids and bases • Understand the Lewis theory of acidity and identify Lewis acids and bases

  9. Give as many answers as you can: • What is an acid? • What is an alkali?

  10. Solubility of Acids and Akalis • Most acids are soluble or react strongly with water • Some bases are soluble, some are insoluble • Soluble bases are called ALKALIS

  11. Bronsted-Lowry Acids and Bases • It’s all about protons (H+) • Acid: Proton donor • HCl(aq)  H+(aq) + Cl-(aq) • H2SO4(aq)  2H+(aq) + SO42-(aq) • Base: Proton acceptor • NH3(aq) + H2O(l)  NH4+(aq) + OH-(aq) • OH-(aq)* + H+(aq)  H2O(l) *From any soluble hydroxide or other alkali • If we mention acid/base without mentioning the type, we generally mean a Bronsted-Lowry one.

  12. Conjugate Acids and Bases • A conjugate acid/base pair are two species that differ by a single proton. • A conjugate base is a species with has one less proton • A conjugate acid is a species with one more proton • For example: • Hydrochloric acid, HCl • HCl is the acid, Cl- is a conjugate base • The HCl can donate a proton…it is an acid • The Cl- could accept a proton….it is a base • Ammonia, NH3 • NH3 is the base, NH4+ is its conjugate acid • The NH3 can accept a proton….it is a base • The NH4+ could donate a proton….it is an acid What is the link?

  13. Time to practice • Give the formula of the conjugate base for each of the following: • HF • H2SO4 • H3PO4 • CH3COOH • H2O • Give the formula of the conjugate acid for each of the following: • OH- • SO42- • HPO42- • (CH3)2NH • H2O

  14. Lewis Acids and Bases • Acid: electron pair acceptor • Species with an incomplete octet/outer-shell • Base: electron pair donor • Species with a lone pair • For example: Gilbert Lewis

  15. Some questions • Which of the following species would exhibit Lewis acid behaviour? • CH4, AlCl3, H2O, BH3, H+, Cu2+, NH3, NH4+ • Which of the following species would exhibit Lewis base behaviour? • H2O, OH-, NH3, CO2, NH4+, C2H5OH, Cl- • Draw a Venn diagram to summarise the similarities and differences between the two theories of acids/bases.

  16. Key Points • Bronsted-Lowry Theory • Acid is a proton donor • Base is a proton acceptor • Lewis Theory • Acid is electron pair acceptor • Base is electron pair donor

  17. Lesson 2 Properties of Acids and Bases

  18. Refresh • Define the terms acid and base according to the Brønsted-Lowry theory. Give one example of each. • Which statement explains why ammonia can act as a Lewis base? • Ammonia can donate a lone pair of electrons. • Ammonia can accept a lone pair of electrons. • Ammonia can donate a proton. • Ammonia can accept a proton.

  19. We Are Here

  20. Lesson 2: Properties of Acids and Bases • Objectives: • Understand the effect of acids/alkalis on indicators • Understand that acids neutralises bases (and vice versa) • Understand the reactions of acids with: • Metals • Carbonates • Hydrogen carbonates • Conduct experiments to confirm the above properties

  21. Indicators • Universal (actually a mixture of indicators): • Litmus: • Acid is Red / Base is Blue • Phenolpthalein: • Acid is Colourless / Alkali is Pink • This is sorcery due to some clever chemistry that we will meet in the HL part of the course.

  22. Neutralisation ACID + ALKALI  SALT + WATER hydrochloric acid + sodium hydroxide  sodium hydroxide + water sulfuric acid + ammonia  ammonium sulfate + water • This is always an exothermic reaction. • It is called neutralisation but won’t always lead to a perfectly neutral solution

  23. Acids and Metals ACID + METAL  SALT + HYDROGEN phosphoric acid + magnesium  magnesium phosphate + hydrogen ethanoic acid + sodium  sodium ethanoate + hydrogen • Very reactive metals may do this explosively

  24. Acids and Carbonates ACID + CARBONATE  SALT + CARBON DIOXIDE + WATER hydrochloric acid + calcium carbonate  calcium chloride + carbon dioxide + water nitric acid + sodium carbonate  sodium nitrate + carbon dioxide + water

  25. Acids and Hydrogencarbonates ACID + HYDROGENCARBONATE  SALT + CARBON DIOXIDE + WATER hydrochloric acid + calcium hydrogencarbonate  calcium chloride + carbon dioxide + water nitric acid + sodium hydrogencarbonate  sodium nitrate + carbon dioxide + water

  26. Now Check • Design and complete experiments to confirm all of the above for yourself. • Make sure you record all data clearly and explain what the data shows • In addition to standard laboratory glassware you will have: • Sodium hydroxide solution • Sodium carbonate solution • Calcium carbonate powder • Sodium hydrogen carbonate powder • Sulphuric acid solution • Hydrochloric acid solution • Magnesium metal • Zinc metal • Lime water • Indicators: universal / litmus / phenolphthalein

  27. Key Points • Acids react in similar ways to each other • Bases react in similar ways to each other • Reactions tend to neutralise acid/base properties

  28. Lesson 3 The pH Scale

  29. Refresh • If 20 cm3 samples of 0.1 mol dm–3 solutions of the acids below are taken, which acid would require a different volume of 0.1 mol dm–3 sodium hydroxide for complete neutralization? • Nitric acid • Sulfuric acid • Ethanoic acid • Hydrochloric acid

  30. We Are Here

  31. Lesson 3: The pH Scale • Objectives: • Understand how pH relates to acidity/bascicity • Understand how pH relates to changes in hydrogen ion concentration • Make a pH colour chart by diluting acids/alkalis

  32. The pH Scale • Runs from 1 for most acid up to 14 for most alkali? • Nonsense, can go below 0 for very strong acids and above 14 for very strong alkalis.

  33. What is pH?.....the ‘Power of Hydrogen’ • pH is determined by the concentration of H+ in a solution. • HL only: pH = -log10[H+] • Each one step increase in pH corresponds to a 10 fold decrease in the concentration of H+ • pH 0 : [H+] = 1.0x100 mol dm-3 (i.e. 1.0) • pH 1 : [H+] = 1.0x10-1 mol dm-3 (i.e. 0.1) • pH 2 : [H+] = 1.0x10-2 mol dm-3 (i.e. 0.01) • Follow this pattern to work out the H+ concentration required for: • pH 3 • pH 5 • pH 7 • pH 9 • pH 14

  34. So a really really weak solution of an acid is actually an alkali? WTF*? THIS IS ONLY NEEDED AT HL…MORE DETAIL LATER • No!.... Pure water exists in an equilibrium as follows: H2O  H+ + OH- • The position of the equilibrium is way over to the left. • However there is always a certain concentration of H+, even in pure water. • In pure water: [H+] = 1.00x10-7 mol dm-3 • You can only get a lower concentration than this by shifting the equilibrium to the left by adding OH- *Note: ‘WTF’ stands for the British phrase ‘What the flip?’….please bring a pure mind to lessons in the future!

  35. Task: Make a pH scale • You need to make a photographic pH scale comprising containers filled with solutions of each pH (from 0-14), coloured with universal indicator. • You should compile your photographs into an electronic document, stating the pH and H+ ion concentration of each. • You can use: • Standard laboratory glassware • 1.00 mol dm-3 hydrochloric acid (pH = 0) • 1.00 mol dm-3 sodium hydroxide (pH = 14) • Distilled water • Universal indicator • You may also wish to try the same with another indicator such as methyl orange, phenolphthalein, bromothymol blue.

  36. Key Points • pH is determined by the concentration of H+ ions • Every one step increase in pH corresponds to a ten-fold decrease in H+ concentration

  37. Lesson 4 Strong and Weak Acids and Bases

  38. Refresh • 100 cm3 of a NaOH solution of pH 12 is mixed with 900 cm3 of water. What is the pH of the resulting solution? • 1 • 3 • 11 • 13

  39. We Are Here

  40. Lesson 4: Strong and Weak Acids and Bases • Objectives: • Understand the difference between strong and weak acids and bases • Complete an experiment to explore the difference in properties of strong/weak acids and bases

  41. Strong and Weak Acids • Strong Acids: • HA(aq) H+(aq) + A-(aq)…..the acid fully dissociates into ions • For example: HCl(aq)  H+(aq) + Cl-(aq) • Includes: hydrochloric, sulfuric, phosphoric, nitric • Strong acids have weak conjugate bases • Weak Acids: • HA(aq)  H+(aq) + A-(aq) …..the acid only partially dissociates into ions • For example: HF(aq)  H+(aq) + F-(aq) • Includes: hydrofluoric, ethanoic, carbonic • Weak acids have strong conjugate bases

  42. Strong and Weak Bases • Strong Acids: • BOH(aq) B+(aq) + OH-(aq) …..the base fully dissociates into ions • For example: NaOH(aq)  Na+(aq) + OH-(aq) • Includes: group (I) hydroxides, barium hydroxide • Strong bases have weak conjugate acids • Weak Acids: • BOH(aq)  B+(aq) + OH-(aq)…..the base only partially dissociates into ions • For example: NH3(aq) + H2O(l)  NH4+(aq) + OH-(aq) • Includes: ammonia, amines • Weak bases have strong conjugate acids

  43. Simulation • http://phet.colorado.edu/en/simulation/acid-base-solutions

  44. So what? • The equilibrium has a profound effect on the properties of the acid/base • Compared with strong acids of the same concentration, weak acids: • Have lower electrical conductivity • React more slowly • pH is higher (less acid) • Change pH more slowly when diluted • However, they neutralise the same volume of alkali • Weak bases follow a similar pattern

  45. Experiment • Complete the experiment here, investigating strong/weak acids and bases • When you give explanations, focus on reasons in terms of the equilibrium rather than just ‘because it is strong/weak’

  46. Key Points • Strong acids/bases dissociate fully into ions • Weak acids/bases only partially dissociate, forming an equilibrium • The strong/weak character has a significant effect on the chemical properties

  47. Lesson 5HL Only pH, pOH and the Ionic Product of Water

  48. Refresh • Ethanoic acid, CH3COOH, is a weak acid. • Define the term weak acid and state the equation for the reaction of ethanoic acid with water. • Vinegar, which contains ethanoic acid, can be used to clean deposits of calcium carbonate from the elements of electric kettles. State the equation for the reaction of ethanoic acid with calcium carbonate.

  49. We Are Here

  50. Lesson 5: pH, pOH and the Ionic Product of Water • Objectives: • Understand the ionic product of water and use it calculate H+ and OH- concentrations • Calculate pH • Calculate pOH

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