  Acids  Produce H + ions when dissolved in water  Ionize into H + ions and negative ion  (Ex. HCl, HBr)  Bases  Produce OH - ions when dissolved in water  Ionize into OH - ion and positive ion  (Ex. NaOH, Mg(OH) 2 )  Does not apply to ALL acids and bases! ! Arrhenius Definition

  HCl  NaOH Examples

  Focus on hydronium (H 3 O + ) ion  Acids  H + ion donator/proton donator  Produces hydronium ion  Donates hydrogen ions to molecule acting as base  Bases  H + ion acceptor/proton acceptor  Accepts hydrogen ions from donating compound  Contains compounds with NONBONDING valence electrons to accept proton Bronsted-Lowry Definition

  CH 3 CO 2 H + H 2 O  CH 3 CO H 3 O +  C 6 H 5 NH 2 + HNO 3  C 6 H 5 NH NO 3 - Examples

  Deals with electron pairs  Acids  Electron pair acceptor  Bases  Electron pair donor  Covalent bond forms between acid and base  ALL Bronsted-Lowry bases are Lewis bases, NOT all Bronsted-Lowry acids act like Lewis acids Lewis Definition

  CaO + SO 2  CaSO 3 Example

  Monoprotic acids  Acid contains only ONE hydrogen ion that can be donated  Polyprotic acids  More than one hydrogen ion can be donated from acid  Amphiprotic  Chemical compound acting as either acid or base Monoprotic vs. Polyprotic Acids

  Acids and bases are related to each other through the addition/loss of hydrogen ions  Conjugate acid-base pairs  Acids produce conjugate bases  Bases produce conjugate acids Conjugate Acids/Bases

  HClO 2 + KOH  H 2 O + KClO 2  HNO 3 + NH 3  NH NO 3 - Conjugate Examples

  Stronger the acid, the weaker the conjugate base  Stronger the base, the weaker the conjugate acid  Weak acids/bases have strong conjugate bases/acids  Acid/Base reactions favor direction from  Stronger weaker of each conjugate acid/base Conjugate Acid/Base Strength

  Based on the concentration of H + or OH - ions in a solution.  Strong Acids/Bases: completely dissociate into ions in a solution.  Weak Acids/Bases: do NOT completely dissociate into ions in a solution. Strength of Acids and Bases

 Strong Acid Example: HCl

 Weak Acid Example: CH 3 CO 2 H

  HClO 4  HI  HCl  HNO 3  HBr  H 2 SO 4 “BIG 6”---Strong Acids (Know them!!)

  Group I metal hydroxides (NaOH, KOH, etc.)  Soluble/Slightly soluble Group II metal hydroxides ( Ca(OH) 2, Sr(OH) 2, Ba(OH) 2 )  Soluble metal oxides ( Li 2 O, Na 2 O, K 2 O, CaO ) Strong Bases (Know them!!)

  Strong Acids/Bases COMPLETELY dissociate with water.  Weak Acids/Bases PARTIALLY dissociate with water.

 Example 1:  HBr + H 2 O  H 3 O + + Br –  NH H 2 O  H 3 O + + NH 3

 Structural Factors Influencing Acidity 1) Polarity 2) Size of Anion 3) Acid/Base Charge

 Structural Factors Influencing Acidity 1)Polarity—How polar is H-X bond?  acid strength,  polarity of chemical bond  If anion (X) is very electronegative, less electron density around hydrogen ion  Ex. HF > H 2 O > NH 3 > CH 4 (similar in size)

 Structural Factors Influencing Acidity 2) Size of Anion (H-X)  acid strength as go down a column  Sizes affects H-X bond  Larger anions hold onto hydrogen ions LESS and hydrogen can be easily donated  Strong acids have weak H-X bonds  Bond strength greater influence than polarity  Example: HI and HF

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 Structural Factors Influencing Acidity 3) Acid/Base Charge  negative charge on compound,  acidity (more basic)  Greater charge separation with opposite charges so greater attraction between positive H ion charge and anion—less likely to lose H + ion  Ex. H 3 PO 4 H 2 PO 4 - HPO 4 -2 PO 4 -3

 Oxyacids 1)Hydroxide group bonded to element with NO other oxygen atoms attached. H—O—Y 2) Hydroxide group bonded to element WITH other oxygens attached. H—O—Y—O

 Oxyacids—how strong are they?  electronegativity of central atom,  acid strength  # of oxygens surrounding central atom,  acid strength  Oxygen is very electronegative—pulls electron density to it  Electron density moves away form O-H bond, bond weakens, H + ions easily pulled off  Ex. H 2 SO 4 > H 2 SO 3, HNO 3 > HNO 2

 Carboxylic Acids  Weak organic acids  General Form—  electronegativity of “R” elements,  acid strength  Examples: CH 3 COOH (acetic acid) and CF 3 COOH

  Water acts as both acid and base  Dissociates into H 3 O + and OH - ions  2H 2 O (l)  H 3 O + (aq) + OH - (aq)  Rewritten,  H 2 O (l) + H 2 O (l)  H 3 O + (aq) + OH - (aq) Water Dissociation

  Rate of reaction reaches equilibrium and allows [H 3 O + ] and [OH - ] to be determined  2H 2 O (l)  H 3 O + (aq) + OH - (aq)  K w = [H 3 O + ] [OH - ]  K w = water dissociation constant = 1x for pure water  [H 3 O + ] = 1x10 -7 M in pure water  [OH - ] = 1x10 -7 M in pure water Water Dissociation (cont.)

  K w = [H 3 O + ] [OH - ] = 1 x regardless of where ions come  More acid, increase [H 3 O + ] /decrease [OH - ]  2 sources for [H 3 O + ], so less water dissociation  More base, decrease [H 3 O + ] /increase [OH - ] Acid/Base Addition

  Measure of the concentration of [H 3 O + ] ions in an acidic/basic solution  Uses logarithmic scale  pH = -log[H + ] or –log[H 3 O + ]  [H 3 O + ] = 10 -pH  pOH = -log[OH - ]  pH + pOH = 14 pH

 pH Scale  Logarithmic scale.  Measures the concentration of hydrogen ions [H + ] in a solution.  Range from  NEUTRAL, pH=7. (pure water)  BASE, pH > 7. (ocean water, milk of magnesia, baking sodea)  ACID, pH < 7. (stomach acid/HCl, vinegar, soft drinks)

  Increase [H 3 O + ], decrease pH value, decrease [OH - ]  Decrease [H 3 O + ], increase pH value, increase [OH - ] pH scale (cont.)

 pH Scale

 How do we measure the pH of a solution?  Acid-base indicators  Weak acids/bases  (ex. litmus paper)  pH meter

 Equations  pH = -log[H + ]  pOH = -log[OH - ]  pH + pOH = 14  [H + ][OH - ] = 1x M 2

  Strong acids completely dissociate in solution  SOOO solution’s [H 3 O + ] ~ [H 3 O + ] in acid  All [H 3 O + ] ions resulting from acid  Dissociation of water is not a factor due to its small value pH and strong acids

  What is the pH of Pepsi Cola if the [H 3 O + ] in the solution is M? Example 1

  Find the pH of a 2.5M HCl solution. Example 2

  Find the pH of a 0.05M H 2 SO 4 solution Example 3

 1)[H 3 O + ] = 1x10 -7 M in water  not a big number 2)LeChatlier’s Principle  As more H 3 O + ions dissociate from water, reaction shifts to LEFT to compensate (reforms water)  If more H 3 O + ions added to solution, water reforms to try and compensate SOOO decrease in [H 3 O + ] from water dissociation Why is the [H 3 O + ] from water dissociation not a factor?

  Acid/Base Introductory Packet—due Wednesday Homework