1. Acids and Bases

2. Ionization of Water  Only happens to a small amount of water molecules  H 2 O separates into H + and OH -  Not the whole story  H+ never occurs on its own  In reality, another H 2 O molecule picks it up and becomes H 3 O + (hydronium ion)

3. Acids and Bases AcidsBases Taste sourTaste bitter Feel wateryFeel slippery Conduct electricity Change litmus to redChange litmus to blue pH = 0-7pH = 7-14 Neutralize basesNeutralize acids

4. Classifying Acids and Bases  Arrhenius  Acid- substance that dissociates into H + and an anion  For Example: HCl and H 2 SO 4  Base- substance that dissociates into cation and OH -  For example: NaOH and Mg(OH) 2  Does not explain bases without an OH ion

5. Classifying Acids and Bases (cont)  Brønsted-Lowry  Acid - Proton (H + ) donor  For example: HCl and H 2 SO 4  Base - Proton (H + ) acceptor  For example: NH 3 and OH -

6. Conjugate Acid and Bases  Occur on the other side of acid base equations.  Lets look again at  NH 3 is a base. It will accept a proton (H + )  H 2 O is an acid. It will donate a proton (H + )  NH 4 + is NH 3 ’s conjugate acid. It can donate a proton (H + ) to become NH 3 again  OH - is water’s conjugate base. It can accept a proton (H + ) to become H 2 O again

7. Amphiprotic  Amphiprotic –  Substances that can act like an acid or a base  Water is an amphiprotic substance.  H 2 O can accept a proton to become H 3 O +  H 2 O can donate a proton to become OH -

8. Strength of Acids and Bases  Depends on how much they dissociate in water  Strong  Considered to dissociate completely in water  Weak  Only partially dissociate in water  Reaction is reversible (  )  Conjugate pairs  Strength is inversely proportional  For example: Strong acids have weak conjugate bases

10. Acids  Strong acids  HI  HBr  HCl  HNO 3  H 2 SO 4  HClO 4  HClO 3  All have 100% of the molecules break apart. There is no reverse reaction.  Weak acids  All others

11. Polyprotic Acids  Have multiple H’s  H 2 SO 4  H 2 SO 4 gives up 1 H + to form HSO 4 -  This happens to 100% of the molecules since H 2 SO 4 is strong  HSO 4 - gives up another H + to form SO 4 -2  This only happens to some HSO 4 - because it is weak  Solution will contain  A lot of water molecules  H 3 O + molecules (mostly from the first H+ but some from the second and from ionization of water)  HSO 4 -  a little bit of SO 4 -2  A little bit of OH - (from the ionization of water)

12. Acid Names  Binary acids (H with an element)  Prefix hydro-  Root of element name  Suffix –ic  Add acid  For example: HCl is hydrochloric acid  Acids with Oxygen (H with a polyatomic)  Root name of polyatomic (with polyatomic prefix if applicable)  Some polyatomic roots are modified slightly to be easier to say  Suffix  -ic with polyatomics ending in –ate  -ous with polyatomics ending in -ite  Add acid  For example: H 2 SO 4 is sulfuric acid

13. Bases  Strong bases  Group 1 metals with OH -  Ca, Sr, and Ba with OH -  These three are not very soluble in water, but the amount that does dissolve ionizes completely.  Weak bases  All others

14. Chemical Equilibrium  Reversible reactions  Indicated with a   Both reactions are happening at the same time  System reaches equilibrium when both are happening at same rate  At equilibrium  Could have lots of reactant and little product  Could have lots of product and little reactant  Could have equal amounts of both  Changes to the system can shift equilibrium  Temperature  Pressure  Adding reactants or products

15. Equilibrium Expressions  Mathematical way to represent equilibrium  For the equation, aA + bB  cC + dD  K = [C] c [D] d [A] a [B] b  K is the equilibrium constant for the equation  [ ] indicates the concentration of each substance in mol/L (M)  Solid and pure liquids are not entered into the expression

16. Ionization of Water  2H 2 O(l)  H 3 O + (aq) + OH - (aq)  This equilibrium “lies to the left”  In other words, there is far more water molecules than hydronium and hydroxide ions in a sample  K w = [H 3 O + ] [OH - ]  K w = 1.0 x 10 -14  In pure water and neutral solutions, [H 3 O + ] and [OH - ] are 1.0 x 10 -7 M  In acidic solutions, [H 3 O + ] is greater than [OH - ]  In basic solutions [OH - ] is greater than [H 3 O + ]

17. pH  pH  Stands for potential of Hydrogen (really hydronium)  Logarithmic scale  pH = -log [H + ] or [H + ] = 10 -pH  Values between 0-14 with each number representing a 10-fold increase from the previous number  pH  7 is acidic  pH = 7 is neutral  pH  7 is basic

19. pOH  pOH = -log [OH - ] or [OH - ] = 10 -pOH  Opposite scale  pOH  7 is basic  pOH = 7 is neutral  pOH  7 is acidic

20. Indicators  Compounds that change color in the presence of different levels of pH

21. Soil pH 6.0-6.5 Soil pH 5.0-5.5

22. Neutralization (Acid-Base Reaction)  Special type of double displacement reaction  Acid + Base  Water + Salt

23. Titration  Process of neutralizing an acid (or base) with an unknown concentration with a base (or acid) of a known concentration  Moles of H 3 O + must equal moles of OH - for neutralization to occur  Often indicators are used to determine the end of the reaction  V a M a = V b M b  V a = volume in L of acid  M a = molarity of acid  V b = volume in L of base  M b = molarity of base