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Acid/Base Chemistry Part 3 (5.4-5.5)

Acid/Base Chemistry Part 3 (5.4-5.5). Science 10 CT05D05 Resource: Brown, Ford, Ryan, IB Chem. Topic 05 – Acids/Bases. 5.1 Solutions 5.2 Definitions of Acids and Bases 5.3 Properties of Acids and Bases 5.4 Calculating pH, pOH , H+, OH- 5.5 Neutralization equations

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Acid/Base Chemistry Part 3 (5.4-5.5)

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  1. Acid/Base ChemistryPart 3 (5.4-5.5) Science 10 CT05D05 Resource: Brown, Ford, Ryan, IB Chem

  2. Topic 05 – Acids/Bases • 5.1 Solutions • 5.2 Definitions of Acids and Bases • 5.3 Properties of Acids and Bases • 5.4 Calculating pH, pOH, H+, OH- • 5.5 Neutralization equations • 5.6 Titrations

  3. 5.4 Calculating pH, pOH, H+, OH- • 5.4.1 Calculate the concentration of ions (H+ and OH-) and acidity (pH and pOH) of strong acids and bases • 5.4.2 Calculate the above of a mixture of strong acids and bases

  4. 5.4 – pH Scale Proposed • Søren Sørenson • Denmark • Biochemist • Early 20th century • Proposed the pH scale

  5. 5.4 - Concentrations of [H+] • The more acidic a solution is, the more H+ ions are donated into solution • Actual concentrations of the hydronium ion, [H+], are often very small • Therefore, Sørensonproposed a manipulation of the concentration of H+ in a way that made the data much more simple to relate • The pH scale is based on the logarithm of the concentration of H+

  6. 5.4 - Concentration • We represent concentration by molarity, therefore the concentration of an acid and a base will give us the following information: • 0.45 M HCl = 0.45 mol [H+] • From an acid we find the [H+] • We can calculate pH directly • 0.65 M NaOH = 0.65 mol [OH-] • From a base we find the [OH-] • We can calculate pOH directly • If we know the concentration of one, we can find it of the other: [H+] [OH-] = Kw = 1 x 10-14

  7. 5.4 - Calculating pH • The pH of a solution is defined as the negative logarithm of the hydrogen ion concentration (in mol/L). • pH = -log[H3O+] • pH = -log[H+] • This calculation results in a pH scale, 0-14 • Therefore, the pH range of solutions are as follows: • Acidic Solutions, pH < 7.0 (0.0-6.9) • Neutral Solution, pH = 7.0 (7.0) • Basic Solutions, pH > 7.0 (7.1-14)

  8. 5.4 – Problems with scale • Find the pH of a 12M solution of HCl • pH = -log (12M) = -1.07 • Usually the pH of 12 Molar HClrepresented as 0.00 and not -1.07? • Any solution with a pH or pOH calculation that results in a negative number you are welcome to round to a pH or pOH of zero (0). • For our purposes in this class, the scale will be from 0-14!

  9. 5.4 - The pH scale is logarithmic • What does the logarithm scale mean? • The logarithm is base 10, so a change in one value has 10x the effect • Since pH 7 is neutral • pH 5 is 10 x more acidic than pH 6 • pH 4 is 100 x more acidic than pH 6 • pH 3 is 1000 x more acidic than pH 6 • pH 9 is 10 x more alkaline than pH 8 • pH 10 is 100 x more alkaline than pH 8 • pH 11 is 1000 x more alkaline than pH 8

  10. 5.4 - What is pOH? • pOH is the opposite of pH, and a measure of alkalinity (how basic) • If pH goes down, pOH goes up. • The pOH scale is (opposite pH): • Basic Solutions: pOH< 7.0 (0.0-6.9) • Neutral Solutions: pOH = 7.0 (7.0) • Acidic Solutions: pOH > 7.0 (7.1-14) • pH + pOH = 14 • 14 – pH = pOH • Calculated: pOH = -log [OH-]

  11. 5.4 – Common Solutions

  12. Practice #1 • What is the pOH of a 0.005 M Mg(OH)2 ? • What is the pH ? -log [OH-] = pOH -log [0.005] = 2.3 = pOH 14 – pOH = pH 14 – 2.3 = 11.7 = pH

  13. Practice #2 • If 20g of NH3 are dissolved in 2.5 L of distilled water, what would the pH of the solution be? 1 mol NH3 20g NH3 = 1.17 mol NH3 17g NH3 -log [OH-] = pOH 1.17 mol NH3 -log [0.47] = pOH = 0.32 = 0.47 M NH3 2.5 L solution 14 – pOH = pH 14 – 0.32 = 13.68

  14. 5.4 - Strong Acids and Bases Strong Bases • LiOH (lithium hydrox.) • NaOH (sodium hydrox.) • KOH (potassium hydrox.) • RbOH (rubidium hydrox.) • CsOH (cesium hydrox.) • Ca(OH)2 (calcium hydrox.) • Sr(OH)2 (strontium hydrox.) • Ba(OH)2 (barium hydrox.) Strong Acids • HClO4 (perchloric acid) • HI (hydroiodic acid) • HBr (hydrobromic acid) • HCl (hydrochloric acid) • H2SO4 (sulfuric acid) • HNO3 (nitric acid) We will calculate primarily with strong monoprotic acids and bases (BOLD) Mainly the acids of halides! Mainly the bases of alkali metals and some alkaline earths!

  15. Applicable Solution Def’s • Strong acids and bases will completely ionize in water and we can therefore use a “” yields symbol in the equation. (for weak acids and bases, an equilibrium arrow would be used ⇌) • Ionization: Process by which a neutral compound is split into charged particles by action when dissolved in liquid water • Equilibrium: When reactants and products are in a constant ratio. The forward and reverse reactions occur at the same rate when a system is in equilibrium.

  16. Ionization Reactions • Completion: For those that include a strong acid or strong base, the reaction will run to completion and can be shown as such with a generic ‘yields’ symbol () • Equilibrium: For those that include a weak acid or base or do not go to completion, the reaction can be represented by an equilibrium symbol ()

  17. 5.5 Neutralization equations • 5.5.1 Balance simple acid base equations • 5.5.2 Conjugate Acid/Base pairs

  18. 5.5 - Neutralization • When equal concentrations of H+ and OH- are added to one another, a neutral solution results • NaOH + HCl NaCl + H2O • Base + Acid  Salt + Water • In equal amounts this is always the case, whether the acid is strong or weak, as long as concentrations are taken into account • Try the following, assume complete neutralization: • Mg(OH)2 + HCl  • KOH + H2SO4 

  19. 5.4 - Aqueous Solutions of A&B • What happens when you put an acid or a base into water? • Each have the property of being electrolytes so will therefore dissociate • Water itself can act as an acid or a base • H+ + H2O  H3O+ (or H+ + OH-  H2O) • H3O+  H+ + H2O (or H2OH+ + OH-)

  20. 5.4 – Conjugate Acid/Base • When acids and bases go through the process of donating or accepting protons, they then switch roles as they can easily reverse the reaction • Acids become conjugate bases • Bases become conjugate acids • Strong become weak • Weak become strong

  21. Arrhenius acid is a substance that produces H+ (H3O+) in water Arrhenius base is a substance that produces OH- in water

  22. A Brønsted acid must contain at least one ionizable proton! A Brønstedacid is a proton donor A Brønstedbase is a proton acceptor conj. base conj. acid acid base

  23. 5.5 - Conjugate Acids and Bases • Conjugate pairs are two substances that differ by one H+ (they gain or lose one PROTON) HCl (acid)  Cl- (conjugate base) NH3 (base)  NH4+ (conjugate acid) • When an acid loses a proton it becomes its conjugate base H2O  OH- or HCl H+ + Cl- • When a base gains a proton it becomes its conjugate acid H2O  H3O+ or NH3 + H+  NH4+

  24. 5.5 - Conjugate pairs Conjugate base of the acid H2SO4 HSO4- Conjugate acid of the base HS- H2S Conjugate acid of H2O H3O+ Conjugate base of H2O OH-

  25. 5.5 - Bronsted-Lowry Acids and Bases NH4+ + OH- NH3 + H2O Acid Base C. Base C. Acid C2H3O2- + H3O+ HC2H3O2 + H2O Base Acid C. Acid C. Base

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