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Acids and bases p pt

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  • 1.  1. General Definitions:  Acid: any substance which when dissolved into the water produces hydrogen ions [H+].  Base: any substance which when dissolved into the water produces hydroxide ions [OH- ].  2. Water dissociation: H2O(l) → H+ (aq) + OH- (aq)  equilibrium constant, KW = [H+][OH-] / [H2O]  Value for Kw = [H+][OH-] = 1.0 x 10-14
  • 2.  Note: The reverse reaction, H+ (aq) + OH- (aq) → H2O(l) is not equal to 1 x 10-14  [H+] for pure water = 1 x 10-7 [OH-] for pure water = 1 x 10-7  3. Definitions of acidic, basic, and neutral solutions based on [H+]  acidic: if [H+] is greater than 1 x 10-7 M  basic: if [H+] is less than1 x 10-7 M  neutral: if [H+] if equal to 1 x 10-7 M
  • 3.  Example 1: What is the [H+] of a sample of lake water with [OH-] of 4.0 x 10-9 M? Is the lake acidic, basic, or neutral?  Solution: [H+] = 1 x 10-14 / 4 x 10-9 = 2.5 x 10-6 M  Therefore the lake is slightly acidic
  • 4.  pH is a measurement of the H+ concentration in a liquid.  relationship between [H+] and pH  pH = -log10[H+]
  • 5.  Acids  1. Strong Acids:  A substance is strong acid when dissolved into the water or other solvent completely dissociate into proton( H+) and an anion.  example: HN03 dissociates completely in water to form H+ and N03 1-.  The reaction is  HNO3(aq) → H+ (aq) + N03 1- (aq)
  • 6.  example: HN03 dissociates completely in water to form H+ and N03 1-.  The reaction is  HNO3(aq) → H+ (aq) + N03 1- (aq)  A 0.01 M solution of nitric acid contains 0.01 M of H+ and 0.01 M N03 - ions and almost no HN03 molecules. The pH of the solution would be 2.0.
  • 7.  1.HCl 2.H2SO4 3.HNO3 4.HClO4 5.HBr 6.HI
  • 8.  Note: when a strong acid dissociates only one H+ ion is removed. H2S04 dissociates giving H+ and HS04 - ions( first ionization).  H2SO4 → H+ + HSO4 1-  A 0.01 M solution of sulfuric acid would contain 0.01 M H+ and 0.01 M HSO4 1- (bisulfate or hydrogen sulfate ion).  Because HS04 - is still having proton, it is also an acid and can dissociate into H+ and SO4 2-( Second ionization)
  • 9.  2. Weak acids:  a weak acid only partially dissociates in water or other solvents to give H+ and the anion  for example, HF dissociates in water to give H+ and F-. It is a weak acid. with a dissociation equation that is :
  • 10.  There are only 6 strong acids, the remainder of the acids therefore are considered weak acids.
  • 11. Strong Bases  Dissociate 100% into the cation and OH- (hydroxide ion).  example: NaOH(aq) → Na+ (aq) + OH- (aq)  a. 0.010 M NaOH solution will contain 0.010 M OH- ions (as well as 0.010 M Na+ ions) and have a pH of 12.  The strong bases are the hydroxides of Groups I and II.
  • 12.  Note: the hydroxides of Group II metals produce 2 mol of OH- ions for every mole of base that dissociates. These hydroxides are not very soluble, but what amount that does dissolve completely dissociates into ions.  exampIe: Ba(OH)2(aq) → Ba2+ (aq) + 2OH- (aq)  a. 0.000100 M Ba(OH)2 solution will be 0.000200 M in OH- ions (as well as 0.00100 M in Ba2+ ions) and will have a pH of 10.3. 
  • 13. Weak Bases  What compounds are considered to be weak bases?  Most weak bases are anions of weak acids.  Weak bases do not furnish OH- ions by dissociation. They react with water to furnish the OH- ions.  Note that like weak acids, this reaction is shown to be at equilibrium, unlike the dissociation of a strong base which is shown to go to completion.
  • 14.  When a weak base reacts with water the OH- comes from the water and the remaining H+ attaches itsef to the weak base, giving a weak acid as one of the products. You may think of it as a two-step reaction similar to the hydrolysis of water by cations to give acid solutions.
  • 15.  examples:  NH3(aq) + H2O(aq) → NH4 + (aq) + OH-(aq)  methylamine: CH3NH2(aq) + H20(l) → CH3NH3 + (aq) + OH- (aq)  acetate ion: C2H3O2 - (aq) + H2O(aq) → HC2H302(aq) + OH- (aq)  General reaction: weak base(aq) + H2O(aq) → weak acid(aq) + OH- (aq)
  • 16.  Since the reaction does not go to completion relatively few OH- ions are formed.
  • 17. Acid-Base Properties of Salt Solutions  A salt : an ionic compound made of a cation and an anion, other than hydroxide or the product besides water of a neutralization reaction.  determining acidity or basicity of a salt solution:
  • 18. 1.split the salt into cation and anion 2. add OH- to the cation a. if you obtain a strong base. the cation is neutral b. if you get a weak base, the cation is acidic 3. Add H+ to the anion a.if you obtain a strong acid, the anion is neutral b. if you obtain a weak acid. the anion is basic  Salt solutions are neutral if both ions are neutral  ( BPH WEEKEND)
  • 19.  Salt solutions are acidic if one ion is neutral and the other is acidic  Salt solutions are basic if one of the ions is basic and the other is neutral.  The acidity or basicity of a salt made of one acidic ion and one basic ion cannot be determined without further information.
  • 20. Examples: determine if the following solutions are acidic, basic, or neutral  KC2H3O2  NaHPO4  Cu(NO3)2  LiHS  KClO4  NH4Cl
  • 21. Acid-Base Reactions  Strong acid + strong base: HCl + NaOH → NaCl + H2O  net ionic reaction: H+ + OH- → H2O  Strong acid + weak base: •example: write the net ionic equation for the reaction between hydrochloric acid, HCl, and aqueous ammonia, NH3. What is the pH of the resulting solution?
  • 22.  when solution gets neutralized?  During this process, indicators will be used.  Indicators are chemical compounds that turn different colors when they're in solutions with different pH's.  Litmus, for example, is red in acid solutions and blue in basic solutions.  Phenolphthalein is clear in acid solutions and pink in basic solutions.  .The basic equation for titration or neutralization is:  M1V1 = M2V2 
  • 23. •M1 stands for the molarity of the acid •V1 stands for the volume of the acid you use •M2 stands for the molarity of the base •V2 stands for the volume of the base you use Example: If it takes 55 mL of 0.1 M NaOH solution to neutralize 450 mL of a HCl solution of unknown concentration, what's the molarity of the acid?  M1, in our equation, stands for the molarity of the acid.
  • 24.  Since that's what we're trying to find, we'll call that X.  V1 stands for the volume of the acid we use. Since HCl is an acid, the volume of acid is 450 mL  M2 stands for the molarity of the base. Since NaOH is a base, the molarity was stated in the problem to be 0.1 M  V2 stands for the volume of the base. The problem says that we used 55 mL of base, so that's M2.
  • 25.  Now, all we need to do is plug it into the equation:  (X)(450 mL) = (0.1 M)(55 mL) X = 0.12 M
  • 26. Buffers solutions  solutions that don't change pH very much when you add acid or base solutions to it.  For example, if you were to add a little bit of HCl to a glass of water, the pH might change from 7 to 3.  If you had the same amount of buffer solution, the pH might change from 7 to 6.8.
  • 27.  Buffers are formed : • a weak acid + its conjugate . Example1: acetic acid+ sodium acetate. → acidic buffer Example 2: a weak base+ its conjugate acid. → basic buffer
  • 28. Weak acid-strong base titrations Example: Titration curve for the titration of vinegar with NaOH. pH at end point- approximately 8.5 ; species present- H2O and NaC2H3O2 and appropriate indicator-phenolphthalein
  • 29.  Note: no matter what type of titration you do, at the equivalence (end) point the number of moles of H+ is equivalent to the number of moles of OH-.  This applies whether you have weak or strong acids and/or bases.  Problems: l. Citric acid (C6H807) contains a mole of ionizable H+/mole of citric acid. A sample containing citric acid has a mass of 1.286 g.
  • 30.  The sample is dissolved in 100.0 mL of water. The solution is titrated with 0.0150 M of NaOH. If 14.93 mL of the base are required to neutralize the acid. then what is the mass percent of citric acid in the sample?
  • 31. Models of acids • Arrhenius Model  The basis for the model is the action in water  The Arrhenius definition:  acids are compounds that give off H+ ions in water  bases are compounds that give off OH- ions in water.
  • 32.  As a result, all acids have hydrogen atoms on them that are ready to go jumping off in water.  Most common acids have the letter H in the beginning of the formula, with the exception of acetic acid.  Arrhenius and Bronsted-Lowry definitions are for most purposes identical. When you see the formula of a base, it's got "OH" in it.  The one exception to this is ammonia, NH3. (NH3 combines with water to form NH4OH, which is really the thing that's basic in ammonia.
  • 33. Strong base + weak acid: •example: write the net ionic equation for the reaction between citric acid (H3C6H507) and sodium hydroxide. What is the pH of the resulting solution? Titrations  Titration : method used in order to determine the concentration of an acidic solution( or basic solution) by adding amount of base( or acid) that you know the concentration.
  • 34.  You have an acidic solution and you want to figure out the molarity. You can't do that directly, because you can't count acid molecules. You can, however, make a basic solution with a concentration that you already know. If you keep adding base to the acid, eventually all of the acid molecules will be neutralized and the solution will turn from an acid to a base.
  • 35.  If you know how many base molecules you added to the solution before the solution gets neutralized (and you will, because you'll add the solution drop-by-drop), you can figure out how much acid was in the solution in the first place.
  • 36.  Indicators: chemical compounds that turn different colors when they're in solutions with different pH's.  Litmus, for example, is red in acid solutions and blue in basic solutions.  Phenolphthalein is clear in acid solutions and pink in basic solutions.
  • 37.  The basic equation for titration or neutralization is:  M1V1 = M2V2  M1 stands for the molarity of the acid  V1 stands for the volume of the acid you use  M2 stands for the molarity of the base  V2 stands for the volume of the base you use
  • 38.  Example: If it takes 55 mL of 0.1 M NaOH solution to neutralize 450 mL of a HCl solution of unknown concentration, what's the molarity of the acid?  M1, in our equation, stands for the molarity of the acid. Since that's what we're trying to find, we'll call that X.  V1 stands for the volume of the acid we use. Since HCl is an acid, the volume of acid is 450 mL
  • 39.  M2 stands for the molarity of the base. Since NaOH is a base, the molarity was stated in the problem to be 0.1 M  V2 stands for the volume of the base. The problem says that we used 55 mL of base, so that's M2.  Now, all we need to do is plug it into the equation:  (X)(450 mL) = (0.1 M)(55 mL) X = 0.12 M
  • 40.  Buffers solutions  Buffers are solutions that don't change pH very much when you add acid or base solutions to it.  For example, if you were to add a little bit of HCl to a glass of water, the pH might change from 7 to 3.  If you had the same amount of buffer solution, the pH might change from 7 to 6.8. 
  • 41.  Buffers are formed when you have a weak acid and its conjugate base present in the same place.  If you wanted to make an acidic buffer, you'd place some acetic acid into a container with some sodium acetate.  If you want a basic buffer, just put a weak base into a container with it's conjugate acid. Our blood is a buffered solution.
  • 42.  If it wasn't, our pH would be go way down every time we had a soda and way up whenever we took some Tums.
  • 43. Weak acid-strong base titrations
  • 44.  l. Arrhenius Model  The basis for the model is the action in water  The Arrhenius definition of acids says that they're compounds that give off H+ ions in water and that bases are compounds that give off OH- ions in water.  These definitions are the same. Basically, if you've got something that can give off H+ in water, it's an acid. As a result, all acids have hydrogen atoms on them that are ready to go jumping off in water.
  • 45.  As a result, all acids have hydrogen atoms on them that are ready to go jumping off in water.  Most common acids have the letter H in the beginning of the formula, with the exception of acetic acid. Bases, on the other hand, are compounds that give off OH- in water.
  • 46.  (The two definitions of a base are for our purposes identical, as OH- combine with H+ to form water -- the Arrhenius and Bronsted- Lowry definitions are for most purposes identical).  When you see the formula of a base, it's got "OH" in it. The one exception to this is ammonia, NH3.
  • 47.  (NH3 combines with water to form NH4OH, which is really the thing that's basic in ammonia. So our definition is sort of true).  Here are a couple of charts which show the most common acids and bases. Some are strong and some are weak, as indicated.
  • 48. Formula Name Strong? HCl hydrochloric acid yes HBr hydrobromic acid yes HI hydroiodic acid yes HF hydrofluoric acid no HNO3 nitric acid yes H2SO4 sulfuric acid yes H3PO4 phosphoric acid no CH3COOH acetic acid no
  • 49.  2. Bronsted-Lowry Model  The basis for the model is proton transfer  According to Bronsted-Lowry; acids are compounds that give off H+ ions when you stick them in water. This definition also says that bases are compounds that can accept H+ ions when you stick them in water.
  • 50.  Simply, acids have H+ in them and bases have OH- in them.  The conjugate base of an acid is whatever is formed when the acid loses its H+ or the base becomes the conjugate acid after it accepts the proton because it can now donate it back.
  • 51.  The acid becomes the conjugate base after it donates the proton because it can now accept it back.  As a general rule of thumb, the conjugate bases of strong acids are weak. For example, Cl- is the conjugate base of hydrochloric acid
  • 52.  3. Lewis Model  The basis for model is the electron pair transfer  The hydrogen requirement of Arrhenius and Brønsted–Lowry was removed by the Lewis definition of acid–base reactions, devised by Gilbert N. Lewis in 1923, in the same year as Brønsted–Lowry, but it was not elaborated by him until 1938. Instead of defining acid–base reactions in terms of protons or other bonded
  • 53.  substances, the Lewis definition defines a base (referred to as a Lewis base) to be a compound that can donate an electron pair, and an acid (a Lewis acid) to be a compound that can receive this electron pair.  In this system, an acid does not exchange atoms with a base, but combines with it.
  • 54.  Lewis definition can be applied to reactions that do not fall under other definitions of acid–base reactions. For example, a silver cation behaves as an acid with respect to ammonia, which behaves as a base, in the following reaction:  Ag+ + 2 :NH3 → [H3N:Ag:NH3]+  The result of this reaction is the formation of an ammonia–silver adduct.
  • 55. Formula Name Strong? NaOH sodium hydroxide yes LiOH lithium hydroxide yes KOH potassium hydroxide yes Mg(OH)2 magnesium hydroxide no Ca(OH)2 calcium hydroxide no NH3 (NH4OH) ammonia (ammonium hydroxide) no

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