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Acids and Bases

Acids and Bases. Text – Chapter 8 Previous knowledge – Naming Acids and Bases (Gr. 11). Acid and Bases. Acid. Base. Sour Electrolytes Gritty feel pH – 0 - 6.9 Blue litmus – red React with bases to form a salt and water Put H+ into solution

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Acids and Bases

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  1. Acids and Bases Text – Chapter 8 Previous knowledge – Naming Acids and Bases (Gr. 11)

  2. Acid and Bases Acid Base • Sour • Electrolytes • Gritty feel • pH – 0 - 6.9 • Blue litmus – red • React with bases to form a salt and water • Put H+ into solution • Made by reaction of oxides and water and binary covalent molecules and water • Bitter • Electrolytes • Slippery feel • pH – 7.1 – 14 • Red litmus – blue • React with acids to form a salt and water • Put OH- into solution • Made by metallic oxides and water

  3. Acid and Base Names • Acids • Contains one or more hydrogen atoms • General formula • HnX • H – hydrogen atom • n – number of hydrogen atoms (subscript) • X – monoatomic or polyatomic anion HCl H2SO4

  4. Acid and Base Names • When the name of the anion ends in “ide” (X), the acid is a binary acid, and the prefix is “hydro” and the ending is “ic” • When there is a polyatomic ion, that makes up (X), the acid is a ternary acid. If the ion ends in “ite”, the ending for the acid is “ous” • When the polyatomic ion ends in “ate”, the ending for the acid is “ic”

  5. Acid and Base Names • Hint – If the name of the anion is “ate”, and the acid is “ic”, one less oxygen, the acid is “ous”, one more less oxygen, the acid is prefix “hypo” and ending is “ous” • If there is one more oxygen than the “ate” polyatomic ion, the name is, prefix “per” and ending “ic” • Some organic acids, you just have to memorize the name. Ex. Ethanoic Acid – CH3COOH

  6. Acid and Base Names • HCl • H2SO4 • H2SO3 • HCN • HClO3 • HClO4 • HClO • H3PO4 Cl- - chloride Binary – hydro – stem - ic Hydrochloric acid SO4 -2 - Sulfate Ternary – stem - ic Sulfuric Acid Sulfurous Acid SO3-2 - sulfite Ternary – stem - ous

  7. Acid and Base Names • Bases • Named the same as ionic compounds • Some you just have to memorize (ie. Ammonia) Positive ion – cation (+) Negative ion – anion (-) NaOH Sodium hydroxide Magnesium hydroxide Aluminum hydroxide

  8. Acid – Base Theories • Arrhenius Acids and Bases • Hydrogen containing compounds that ionize (ions – “wanderers” to produce H+ are acids • Hydroxide containing compounds that produce OH- ions in solution are called bases • Not all substances that contain hydrogen atoms or hydroxide will be acidic or basic. It depends on electronegativity and polarity between the acidic / basic unit and the bonded atoms.

  9. Acid – Base Theories • Bronsted – Lowry Acid and Bases • Acid – hydrogen ion donor • Base – hydrogen ion acceptor • Substance that accepts the hydrogen is the conjugate acid • Substance that donates the hydrogen is the conjugate base • Used for those exceptions that cannot be explained by Arrhenius • Truer theory as a “naked” hydrogen ion is very unlikely and unstable. Hydronium ion is most likely

  10. Acid – Base Theories • Water can behave as both a conjugate base and acid (can accept and donate) • Called Amphoteric substance (behave as an acid and a base)

  11. Strengths of Acids and Bases • Based upon structure • Greater the EN difference, the greater the ionization and dissociation, means the more “product” is formed, and more H+ or OH- goes into solution • Therefore, stronger acid and base (Keq greater than 1) • First ionization is the strongest • Second and subsequent ionizations are weaker. (p.607)

  12. Acid – Base Theories • The number of hydrogens will determine whether it is monoprotic, diprotic or triprotic acids. • Structure will determine the strength of the acid. • Rule – the greater the EN difference, the greater the polarity, the greater the dissociation (ionization) and strength of the acid or base. • Rule – For ternary acids, if the Oxygens out number the hydrogens by more than 2, the acid will be strong • Greater net pull, according to the first rule.

  13. Strengths of Acids and Bases • Keq for an acid is called Ka, and is the measure of how much of the acid ionizes (H+ or H3O+ and X) and how much stays together (HnX) • Keq for a base is called Kb, and is the measure of how much OH- and + ion is in solution and how much stays together. HX H+ + X- H2XOy H+ + H XOy- M(OH)z M+ + Z( OH-)

  14. Strengths of Acids and Bases • The amount that dissociates, or ionizes, is in equilibrium with the acid that stays as a “whole” • If the Keq is greater than 1, it favours “product” • For Acids and Bases, the Keq greater than 1, means it dissociates 100%, and all of the reactant (acid or the base) goes to ions.

  15. Strengths of Acids and Bases • For example • H2SO4 Ka = 1.00 x 103 • H2SO4 H+ + H SO4- • Therefore, if the concentration of the acid is 2.0 M, the concentrations of ions will be: Therefore, the concentration of the hydronium ion is 2.0 M.

  16. Strengths of Acids and Bases • Therefore, when a strong acid or base dissociates, 100% turns into ions. • What would be the OH- for a sodium hydroxide solution with a molarity of 0.5 M? Solution – sodium hydroxide is a strong base, therefore, it dissociates 100% NaOH Na+ + OH- 0.5 M O.5 M 0.5 M

  17. Strengths of Acids and Bases

  18. Strengths of Acids and Bases • What about weak acids and bases? • We need to solve for the concentrations of ions, using Ka or Kb values because they do not dissociate 100% • Ka or Kb values are less than 1, favouring the acid or base to stay together and little ionize. • They may only dissociate 50% or 10%, leaving the majority of the acid or base as a “whole” and very little in ion form

  19. Strengths of Acids and Bases • Example What is the [H+] concentration in a 1.0 M solution of carbonic acid? (Ka = 4.3 x 10-7) Hint – Look at the Ka value. It is less than 1. Therefore, it will not dissociate 100% and is a weak acid. We need to use equilibrium to determine the concentration of the [H+]. Note – In your text, the Keq expression is written using Bronsted /Lowry. You can use either, as long as you remember that there are some exceptions in which the ionization cannot be shown using Arrhenius. Also, the Keq expression uses water, then omits it. Keep in mind, water always has a concentration of 1 M. Concentration of water?!!!!

  20. Strengths of Acids and Bases • Why are we finding the H+ or OH-? pH power of hydrogen or potential of hydrogen

  21. Hydrogen Ions and pH • Based upon water • Highly polar • Made up of hydronium ion and hydroxide ion • Self Ionization of water (Kw)

  22. Hydrogen Ions and pH • Each ion has a value of 1 x 10 -7 M, which, when multiplied together forms the Keq or Kw, which is 1 x 10 -14 • Both are equal to each other in terms of their concentrations, and therefore form a neutral substance (pH = 7) • If the concentration of H+ is greater than 1 x 10 -7 M, the solution is acidic (ie. More H+ and less OH-), since all solutions are in water. • If the concentration of H+ is less than 1 x 10 -7 M, the solution is basic (ie. Less H+ and more OH-), since all solutions are in water.

  23. Hydrogen Ions and pH • pH = potential hydrogen • pH scale is based upon the [H+] found in a solution. • If the [H+] is greater than 1 x 10 -7 M, the solution will be acidic and the pH will be less than 7. • If the [H+] is less than 1 x 10 -7 M, the solution will be basic and the pH will be greater than 7. • pH is the negative log of the hydrogen ion concentration.

  24. pH Scale

  25. Measuring pH – Acid Base Indicators • Indicator – usually a weak acid that accepts hydrogen ions, and in doing so, changes its chemical structure, which facilitates a colour change. • Litmus – changes from red to blue at pH of 7, or 1 x 10 -7 M • Bromothymol Blue – is yellow below 1 x 10 -7 M, green at 1 x 10 -7 M, and blue over 1 x 10 -7 M • Phenopthlalein changes color at a pH of 7-9 (the hydrogen ion concentration of 1 x 10 -7 M to 1 x 10 -9 M)

  26. Measuring pH – Acid Base Indicators HIn (aq) H+(aq) + In-(aq) Acid Form Base Form Color #1 Color #2 OH- H+ The change is caused by the removal of a hydronium ion to form the “Base form” and the addition of a hydronium ion for the acid form.

  27. Measuring pH – Acid Base Indicators • Indicator papers are impregnated with the indicator solution and when exposed to the hydrogen ion, change color, depending on the concentration • Universal Indicators show all pH levels. What do you think they are made up of?

  28. Neutralization Reactions and Titration • If we have high levels of acid in our stomach, we take an antacid (base) to control it. • In our small intestine, the acidic chyme, is neutralized by the bile (basic) to make sure we do not ulcerate our intestine • Why does it neutralize? • We make a salt and water, that does not necessarily work out to a pH of 7 (more later) • We could bring the solution to a pH of 7.

  29. Neutralization Reactions and Titration [H+] = [OH-] 1 x 10-7 M = 1 x 10-7 M Ex. If you react a strong acid and a strong base, the ions in solution will cancel each other out, producing a neutral solution. The products are always, regardless of the product pH, a salt and water. HCl(aq) + NaOH(aq)NaCl(aq) + H2O(l)

  30. Neutralization Reactions and Titration • The point where the number of moles of hydronium ion equals the number of moles of hydroxide ion, is called the equivalence point. (not necessarily pH =7) • Salt • The compound formed by the cation of the base bonding with the anion of the acid.

  31. Neutralization Reactions and Titration • Titration • The process of adding a known amount of solution of known concentration to determine the concentration of another solution • When the color changes, this is called the end point, which is the point of neutralization • There is only salt and water at this point, if both were strong acids and bases.

  32. Neutralization Reactions and Titration • Animation of Titration • Lab Primer • The idea is to calculate the concentration or number of moles for an unknown solution, using the equivalence point to determine neutralization

  33. Neutralization Reactions and Titration • Titration Curves

  34. Salts • Salts are the combination of the cation from the base and the anion from the acid and are the products of neutralization reactions • Salts can be acidic, basic or neutral, depending on the strength of the acid and base that formed it. • Buffers are an equilibrium condition which consists of the weak acid and conjugate base (salt) in solution, or a weak base and its conjugate acid (salt) keeping pH stable

  35. Salts • Generally, if a strong acid reacts with a strong base, the resulting salt will be neutral (pH=7) (ie. The equivalence point is 7) • For salts formed from weak acids with a strong base, or weak bases with a strong acid, the salt will not be neutral • This is called by salt hydrolysis, as the cations or anions from a dissociated salt remove or add hydrogen ions to water, creating either H+, or OH- in solution

  36. Salts • Another way to determine the acidity or basicity of a salt is to look at the net ionic equation and remembering that strong bases and acids dissociate 100%, while weak acids and bases do not. • In general • Acidic salts produce positive ions that release protons into water • Basic solutions produce negative ions that attract protons from water

  37. Salts Strong Acid + Strong Base Neutral solution Strong Acid + Weak BaseAcidic Solution Weak Acid + Strong BaseBasic Solution

  38. Buffers • Buffer • A substance in which the pH remains constant, when small amounts of acid or base is added. • It contains the weak acid and one of its salts (anion – conjugate base) or the weak base and one of its salts (cation – conjugate acid) • Pure water is not a buffer, as when you add acid (H+) or base (OH-), the concentrations increase, changing the pH

  39. Buffers • A buffer is like a sponge • When hydrogen ions are added, they are absorbed by the negative ion, forming the “whole” weak acid, that does not dissociate 100%, lowering the acidity, and raising the pH to neutral • When hydroxide is added, they react with the acid to form the negative ion and water, lowering the basicity, and lowering the pH to neutral

  40. Buffers • Animation of Buffering • Common buffers, keep the pH at a stable level • Ethanoic acid maintain pH 4.76 • Carbonic acid maintain pH 6.5 (blood) • Ammonia maintain pH 9

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