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acids always have as the state and always have

Chapter 6: Acids & Bases. 6.1 Theories of Acids and Bases. A. Naming Acids and Bases. acids always have as the state and always have. (aq). hydrogen. Rules 1. hydrogen becomes acid

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acids always have as the state and always have

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  1. Chapter 6: Acids & Bases 6.1 Theories of Acids and Bases A. Naming Acids and Bases • acids always have as the state and always have (aq) hydrogen Rules 1. hydrogen becomes acid 2. hydrogen becomes acid 3. hydrogen becomes acid ____ide hydr____ic _____ate _____ic ____ite ____ous

  2. Examples: Change each of the following to the appropriate acid name and give the formula: HI(aq) 1. hydrogen iodide = hydroiodic acid H3PO4(aq) phosphoric acid 2. hydrogen phosphate = nitrous acid HNO2(aq) 3. hydrogen nitrite = sulphurous acid H2SO3(aq) 4. hydrogen sulphite =

  3. most bases are ionic compounds that are named accordingly Examples: Name each of the following bases: 1. NaOH(aq) = sodium hydroxide sodium hydrogen carbonate 2. NaHCO3(aq) = 3. Mg(OH)2(aq) = magnesium hydroxide 4. NH3(aq) = ammonia

  4. IUPAC names for acids and bases are simply the word “aqueous” followed by the ionic name Examples: Write the IUPAC name for each of the following acids and bases: 1. hydroiodic acid = aqueous hydrogen iodide 2. magnesium hydroxide = aqueous magnesium hydroxide 3. sulphurous acid = aqueous hydrogen sulphite 4. sodium hydrogen carbonate = aqueous sodium hydrogen carbonate

  5. B. Properties of Acids and Bases empirical properties observable properties • are of a substance • acids, bases and neutral substances have some properties that distinguish them and some that are the same

  6. Neutral Substances Acids Bases sour taste bitter taste electrolytes electrolytes electrolytes, non-electrolytes neutralize acids neutralize bases react with indicators do not react with indicators affect indicators the same way litmus - red litmus - blue bromothymol blue - yellow bromothymol blue - blue phenolphthalein - phenolphthalein - pink colourless react with to produce metals H2(g) pH greater than 7 pH of 7 pH less than 7 eg) HCl(aq), H2SO4(aq) eg) eg) Ba(OH)2(aq) NH3(aq) NaCl(aq), Pb(NO3)2(aq)

  7. C. Arrhenius Definition • first proposed theory on acids and bases Svante Arrhenius • his theory was that some compounds form electrically charged particles when in solution • his explanation of the properties of acids and bases is called the Arrhenius theory of acids and bases

  8. an Arrheniusis a substance that (because it is molecular) to form acid ionizes hydrogen ions, H+(aq), in water • an will in an aqueous solution acid increase the [H+(aq)] • an Arrheniusis a substance that to formin water base dissociates hydroxide ions, OH(aq), • a will in an aqueous solution base increase the [OH-(aq)]

  9. D. Modified Arrhenius Definition • the original definition of acids and bases proposed by Arrhenius is good but it has limitations • some substances that might be predicted to be are actually neutral basic eg) Na2CO3(aq), NH3(aq) • it has been found that not all bases contain the hydroxide ion as part of their chemical formula

  10. an Arrhenius is a substance that in aqueous solution base (modified) reacts with water to produce OH(aq) ions eg) NH3(aq) + H2O()  NH4+(aq) + OH(aq) • (Na2O, MgO etc) form in water metallic oxides bases eg) Step 1: Na2O(s) + H2O()  2 NaOH(aq) (make the ) metal hydroxide 2 NaOH(aq)  2 Na+(aq) + 2 OH(aq) Step 2: ( the metal hydroxide) dissociate

  11. when acids ionize, they produce H+(aq) eg) HCl(g) H+(aq) + Cl(aq) • it has been found using analytical technology like X-ray crystallography that in an aqueous solution H+(aq) ions do not exist in isolation • the hydrogen ion is extremely positive in charge and water molecules themselves are very polar so… it is that would exist in water without being attracted to the of other highly unlikely hydrogen ions negative poles water molecules

  12. this results in the formation of the hydronium ion + H3O+(aq)

  13. an Arrhenius is a substance that in aqueous solution acid (modified) reacts with water to produce H3O+(aq) ions eg)  HCl(aq) + H2O() Cl(aq) + H3O+(aq)  H2SO3(aq) + H2O() HSO3(aq) + H3O+(aq) • form in water non-metallic oxides (CO2, SO2 etc) acids eg) Step 1: CO2(g) + H2O()  H2CO3(aq) all (combineelements to make an ) acid + H2O() Step 2: H2CO3(aq)  H3O+(aq) + HCO3(aq) ( acid with water) react

  14. 6.2 Strong and Weak Acids and Bases • the of a substance depend on two things: acidic and basic properties 1. the of the solution concentration 2. the of the acid or base identity

  15. A. Strong Acids and Weak Acids ionizes almost 100% in water • an acid that is called a strong acid eg) HCl(aq) + H2O()  H3O+(aq) + Cl(aq) HCl(aq) • 100% of the becomes H3O+(aq) and Cl(aq) • the concentration of the is the as the concentration of the it came from H3O+(aq) same acid • strong acids are strong electrolytes and react vigorously with metals

  16. there are 6 strong acids: perchloric acid HClO4(aq) hydrobromic acid HBr(aq) hydroiodic acid HI(aq) hydrochloric acid HCl(aq) sulfuric acid H2SO4(aq) HNO3(aq) nitric acid ***on your periodic table

  17. weak aciddoes not ionize 100% • a and only a small percentage of the acid forms ions in solution eg) CH3COOH(aq) + H2O() ⇌ H3O+(aq) + CH3COO(aq) • we use the for weak acids equilibrium arrow • weak acids are react much less vigorously with metals weak electrolytes and

  18. B. Strong Bases and Weak Bases 100% • a base that dissociates into ions in water is called a strong base ionic hydroxides and metallic oxides • are strong bases eg) NaOH(aq)  Na+(aq) + OH(aq) • a and only a small percentage of the base forms weak basedoes not dissociate 100% ions in solution + NH4+(aq) eg) NH3(aq) + H2O() ⇌ OH(aq) • we use the for weak bases equilibrium arrow

  19. C. Monoprotic and Polyprotic Acids • acids that have onlyper molecule that canare called one hydrogen atom ionize monoprotic acids eg) HCl(aq), HF(aq), HNO3(aq), CH3COOH(aq) + eg) HNO3(aq) + H2O()  H3O+(aq) NO3(aq) • monoprotic acids can be strong or weak

  20. acids that contain that canare called two or more hydrogen atoms ionize polyprotic acids eg) H2SO4(aq), H3PO4(aq) • acids with are , with are two hydrogens diprotic three hydrogens triprotic

  21. when polyprotic acids ionize, only hydrogen is removed at a time, with each acid becoming one progressively weaker eg)  + H2O() + HSO4(aq) H2SO4(aq) H3O+(aq) + H2O()  H3O+(aq) + SO42(aq) HSO4(aq)

  22. D. Monoprotic and Polyprotic Bases react with water in only one step to form hydroxide ions • bases that are called monoprotic bases eg) NaOH(s) • bases that react with water in are called two or more steps polyprotic bases eg) CO32(aq), PO43(aq)

  23. one • as with polyprotic acids, only OH(aq) is formed at a time, and each new base formed is than the last weaker eg)  + H2O() + HCO3(aq) CO32(aq) OH(aq) + H2O()  OH(aq) + H2CO3(aq) HCO3(aq)

  24. E. Neutralization • the reaction between an acid and a base produces an ionic compound and water acid + base a salt + water → KCl(aq) + HOH() eg) HCl(aq) + KOH(aq)→ • the products of are both neutralization neutral • in a neutralization reaction or between a , the product is always acid-base reaction strong acid and a strong base water H3O+(aq) + OH(aq)  2 H2O()

  25. F. Acid and Base Spills • there are many uses for both acids and bases in our households and in industry • due to their, special care must be used when they are being reactivity and corrosiveness produced and transported

  26. the two ways to deal with acid or base spills are: concentration 1.dilution: reduce the by adding water weak acid or base 2.neutralization: you always use a for the neutralization so you aren’t left with another hazardous situation

  27. 6.3 Acids, Bases and pH A. Ion Concentration in Water • the “self-ionization” of water is very small (only 2 in 1 billion)  H3O+(aq) + OH-(aq) H2O() + H2O() hydronium ions • the concentration of and are hydroxide ions equal and constant in pure water [H3O+(aq)]= [OH-(aq)] = 1.0 x 10-7mol/L 1.0 x 10-7mol/L

  28. B. The pH Scale • in 1909, Soren Sorenson devised the pH scale • it is used because the [H3O+(aq)] is very small • at 25C (standard conditions), most solutions have a pH that falls between 0.0 and 14.0 • it is possible to have apHand a pH negative above 14 • it is a based on whole numbers that are powers of 10 logarithmic scale

  29. 0 7 14 • there is a for every change inon the pH scale 10-fold change in [H3O+(aq)] 1 eg) a solution with a pH of 11 is times more basic than a solution with a pH of 9 10  10 = 100 pH Scale more acidic more basic neutral

  30. C. Calculating pH and pOH pH =  log [H3O+(aq)] when reporting pH or pOH values, only the numbers to the count as significant ***New sig dig rule: right of the decimal place Try These: 1.     [H3O+(aq)] = 1x 10-10 mol/L pH = 2.     [H3O+(aq)] = 1.0 x 10-2 mol/L pH = 3.     [H3O+(aq)] = 6.88 x 10-3 mol/L pH = 4.     [H3O+(aq)] = 9.6 x 10-6 mol/L pH = 10.0 2.0 2.16 5.0

  31. Example 6.30 g of HNO3 is dissolved in 750 mL of water. What is the pH ?  H3O+(aq) + NO3-(aq) HNO3(aq) + H2O() m = 6.30 g M = 63.02 g/mol V = 0.750 L c = 0.133…mol/L x 1/1 = 0.133…mol/L n = m M = 6.30 g 63.02 g/mol = 0.0999…mol pH = -log[H3O+(aq)] = -log[0.133… mol/L] = 0.875 c = n V = 0.0999…mol 0.750 L = 0.133…mol/L

  32. [H3O+(aq)], • just as deals with deals with pH pOH [OH(aq)] ***p just means log • at SATP… pH + pOH = 14 pH 0 1 3 5 7 9 11 13 14 14 13 11 9 7 5 3 1 0 pOH

  33. to calculate the use the same formulas as pH but substitute the pOH, [OH(aq)] pOH =  log[OH(aq)] Try These: 1.[OH(aq)] = 1.0 10-11 mol/L pOH = 2.[OH(aq)] = 6.22 10-2 mol/L pOH = 3.[OH(aq)] = 9.411 10-6 mol/L pOH = 4.[OH(aq)] = 2 10-6 mol/L pOH = 11.00 1.206 5.0264 5.7

  34. you could also be given the pH or pOH and asked to calculate the [H3O+(aq)] or [OH-(aq)] [H3O+(aq)] = 10-pH [OH(aq)]= 10-pOH

  35. Try These: 1.pH 4.0 [H3O+(aq)] = 2.pH 6.21 [H3O+(aq)] = 3.pH 13.400 [H3O+(aq)] = 4.pH 7 [H3O+(aq)] = 5.pOH 1.0 [OH(aq)] = 6.pOH 13.2 [OH(aq)] = 7.pOH 6.9 [OH(aq)] = 8.pOH 0.786 [OH(aq)] = 1 x 10-4mol/L 6.2 x 10-7 mol/L 3.98 x 10-14 mol/L 10-7mol/L 0.1mol/L 6  10-14 mol/L 1  10-7 mol/L 0.164 mol/L

  36. 9. Complete the following table: 5.40 8.60 acid 4.0 x 10-6 mol/L 2.5 x 10-9 mol/L 9.500 4.500 base 3.16 x 10-10 mol/L 3.16 x 10-5 mol/L 3.30 10.70 acid 2.0  1011 mol/L 5.0 x 10-4 mol/L 15.00 10 mol/L -1.00 acid 1.0 x 10-15 mol/L base -1.00 15.00 1.0 x 10-15 mol/L base 1.36 0.044 mol/L 2.3 x 10-13 mol/L

  37. D. Measuring pH • pH can be measured using : 1. acid-base indicators 2. pH meter Indicators • an is any chemical that in an acidic or basic solution acid-base indicator changes colour • they can be dried onto strips of paper eg) litmus paper, pH paper

  38. they can be solutions eg) bromothymol blue, universal indicator, indigo carmine etc • they can be made from natural substances eg) tea, red cabbage juice, grape juice

  39. each indicator has a where it will specific pH range change colour • you can use to approximate the two or more indicators pH of a solution

  40. pH Meters • using a pH meter is the most way of measuring precise pH • it has an that compares the [H3O+(aq)] in the solution to a and it will give a of the pH electrode standard digital readout

  41. E. Diluting an Acid or Base add water acid or base • when you to an , you change the [H3O+(aq)] or the [OH(aq)] • diluting an acid will the until a pH of is reached [H3O+(aq)] decrease 7.0 • diluting a base will the until a pH of is reached [OH-(aq)] decrease 7.0

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