Chemistry Ch. 6

Front 1

Arrhenius acid

Back 1

-anything that produces hydrogen ions in aqueous solution

Front 2

Arrhenius base

Back 2

-anything that produces hydroxide ions in aqueous solution

Front 3

Bronsted Lowry acid

Back 3

-anything that donates a hydrogen in aqueous solution

Front 4

Bronsted Lowry base

Back 4

-anything that accepts a hydrogen in aqueous solution

Front 5

Lewis base

Back 5

-anything that donates a pair of electrons

Front 6

Lewis acid

Back 6

-anything that accepts a pair of electrons
-include molecules that have an incomplete octet of electrons around the central atom, like AlCl3 and BF3.
-also include all simple cations except the alkaline earth metal cations

Front 7

Strength of cation acids

Back 7

-the smaller the cation and the higher the charge, the stronger the acid strength.

Front 8

pH

Back 8

pH = -log[H+]

- [H+] = moles/liter
-a measure of the hydrogen ion concentration in a solution.
-At 25C a pH of 7 is neutral.

Front 9

Rule about acid base reaction

Back 9

-If there is an acid in a reaction, there must be a base; you can't have a proton donated without something to accept it.
-Every reaction with an acid has a conjugate base; every reaction with a base has a conjugate acid.

Front 10

Conjugate acids and conjugate bases: relative strength

Back 10

The stronger the acid the weaker its conjugate base and the stronger the base the weaker the conjugate acid.

Front 11

Amphoteric

Back 11

Substances that can be either acidic or basic depending on their environment.

ex. water

Front 12

7 strong acids

Back 12

1. HI
2. HBr
3. HCl
4. HNO3
5. HClO4
6. HClO3
7. H2SO4

Front 13

7 Strong bases

Back 13

1. NaOH
2. KOH
3. NH2-
4. H-
5. Ca(OH)2
6. Na2O
7. CaO

Front 14

Polyprotic

Back 14

-Acids that can donate more than one proton.

Front 15

Diprotic

Back 15

-Acids that can donate 2 protons.
-The second proton is typically so weak that its effect on pH is negligible. (If the Ka value differs by more than 10^3 than the second proton can be ignored)

Front 16

Amount of acid dissociation

Back 16

-The percent dissociation of an acid decreases with increasing acidity.
-Acids dissociate less in more concentrated solutions. Does NOT mean that concentrated solutions are less acidic.

Front 17

Definition of strong acid and base

Back 17

-Strong acid is stronger than H3O+
-Strong base is stronger than OH-
-Fully dissociate in water.

Front 18

3 factors in molecular structure that determine whether or not a molecule containing a hydrogen will release its hydrogen into solution

Back 18

1. the strength of the bond holding the hydrogen molecule
2. the polarity of the bond
3. the stability of the conjugate base

Front 19

Acid strength of oxyacids

Back 19

-electronegative oxygens draw electrons to one side of the bond, increasing polarity
-the oxygens in the conjugate of an oxyacid can share the negative charge spreading it over a large area and stabalizing the conjugate base.
-acidity increases with

Front 20

Hydrides

Back 20

-binary compounds containing hydrogen
-can be basic, acidic, or neutral
-on the periodic table, basic hydrides are on the left, acidic hydrides are on the right.

Front 21

Autoionization of water

Back 21

-Pure water reacts with itself to form a hydronium and hydroxide ion.
-The equilibrium constant for this reaction is Kw.
-The equilibrium of this reaction lies far to the left (Kw = 10^-14)

Front 22

Acid dissociation constant (Ka)

Back 22

-An acid will have its own equilibrium constant in water
-Dissociation constant for acid Ka:

[H+][A-]/[HA]

Front 23

Value of Kw at 25 degrees C

Back 23

Kw = 10^-14

Front 24

Base dissociation constant

Back 24

-for every Ka there is a corresponding Kb
-Kb is the equilibrium constant for the reaction of a conjugate base with water
-Kb for conjugate base A-:

[OH-][HA]/[A-]

Front 25

Salts as acids and bases

Back 25

-ionic compounds that dissociate in water
-often create acidic or basic conditions.
-the pH of the solution can be predicted qualitatively by comparing the conjugates of the respective ions.

Front 26

Titration

Back 26

-A drop by drop mixing of an acid and base.
-Performed in order to find the concentration of some unknown by comparing it with the concentration of a titrant.

Front 27

Equivalent or Stochiometric point

Back 27

-for a monoprotic acid, the equivalence point in the titration is the point where there are equal equivalents of acid and base in solution.
-If the concentrations differ, the equivalence point will not be where the volumes are equal (need molar equivalents).
-For equally strong acid-base titrations, the equivalence point will be pH 7.

Front 28

Titrant

Back 28

During titrations, the concentration of some unknown is compared with the known concentration of a titrant.

Front 29

Half equivalence point

Back 29

The point in a titration where exactly one half of the acid has been neutralized by the base.
The concentration of the acid is equal to the concentration of its conjugate base.
The area on the graph that represents a horizontal line-buffered.
The pH is equal to the pKa of the acid at this point.

Front 30

Buffered solution

Back 30

The spot on a titration curve where you can add the largest amount of base or acid with the least amount of change in pH.
The solution is said to be buffered at this point

Front 31

Henderson-Hasselbalch equation

Back 31

pH = pKa + log([A-]/[HA])

-predicts that when a solution is buffered, pH = pKa b/c concentration of acid and base is equal and log(10) = 0.

Front 32

How to make a buffer solution

Back 32

-Start with an acid whose pKa is closest to the pH at which we want to buffer the solution
-Mix equal amounts of acids with conjugate base
-Want the concentration of buffer solution to greatly exceed the concentration of the outside acid or base affecting our solution, so made from equal and copious amounts of acid and conj. base.

Front 33

What happens when we add water to a buffered solution?

Back 33

-Under ideal circumstances, adding a little water to an ideally dilute, buffered solution will have no effect on pH.
-Adding copious amounts of water will change the pH b/c it can act as an acid or base, but MCAT will most likely consider only the first scenario.

Front 34

Chemical Indicator

Back 34

-Used to find the equivalence point
-Usually a weak acid whose conjugate base is a different color.
-Can use the Henderson Hasselbalch equation for indicator because it never reaches an equivalence point.

Front 35

Endpoint

Back 35

The point where an indicator changes color.
(NOT the same as an equivalence point but typically choose an indicator whose range will cover the equivalence point

Front 36

Polyprotic titrations

Back 36

-Have more than one equivalence point and more than one half equivalence point

Front 37

pKa

Back 37

pKa = -log(Ka)

-A strong acid has a pKa less than -2.
-Want a LOW number for a strong acid.