UARK CHEM 2 Exam 3

Adding reactant (which makes Q < K) causes

the reaction to shift to the right

Adding product (which makes Q > K) causes

the reaction to shift to the left

Removing reactant (which makes Q > K) causes

the reaction to shift to the left

Removing product (which makes Q < K) causes 

the reaction to shift to the right

Decreasing volume causes

a reaction to shift in the direction with fewer moles of gas

Increasing volume causes

a reaction to shift in the direction with more moles of gas

Adding an inert gas to a mixture at a fixed volume

does not affect equilibrium (no shift)

When a reaction has equal number of moles of gas on both sides of a chemical equation,

a change in volume does not affect equilibrium (no shift)

In an exothermic reaction,

heat is a product

increasing the temp causes the reaction to shift left

decreasing the temp causes the reaction to shift right

In an endothermic reaction

heat is a reactant

increasing the temp causes the reaction to shift right

decreasing the temp causes the reaction to shift left

Binary acids

composed of hydrogen and a nonmetal

Oxyacids

contain hydrogen and an oxyanion

Oxyanions ending with -ate

-ate goes to -ic

ex. sulfate to sulfuric acid

Oxyanions ending with -ite

-ite goes to -ous

ex. sulfite to sulfurous acid

Acid-base reaction

an acid reacts with a base and the two neutralize each other, producing water or in some cases a weak electrolyte

Titration

a substance of known concentration is reacted with another substance of unknown concentration

Equivalence point

just enough titrant has been added to completely react with the analyte

indicator

used to identify the equivalence point

Acid

a substance that produces H+ ions in aqueous solution

Base

a substance that produces OH- ions in aqueous solution

Single arrow indicates

complete ionization

Double arrow indicates

partial ionization

Bronsted-Lowry Definition of Acids

donates a proton (H+ ion) and becomes a conjugate base

Bronsted-Lowry Definition of Bases

accepts a proton (H+ ion) and becomes a conjugate acid

For Binary acids

acidity increases as electronegativity increases

For binary acids pt. 2

acidity decreases as bond strength increases

In a column,

bond strength has a greater effect than electronegativity on acidity

For oxyacids,

electronegativity increases, acidity increases

For oxyacids pt. 2

as the # of oxygen atoms attached increases, acidity increases

What are the 6 strong acids

1. Hydrochloric acid (HCl)

2. Hydrobromic acid (HBr)

3. Hydriodic acid (HI)

4. Nitric acid (HNO3)

5. Perchloric acid (HClO4)

6. Sulfuric acid (H2SO4, diprotic)

What are the 6 weak acids?

1. Hydrofluoric acid (HF)

2. Acetic acid (HC2H3O2)

3. Formic acid (HCHO2)

4. Sulfurous acid (H2SO3, diprotic)

5. Carbonic acid (H2CO3, diprotic)

6. Phosphoric acid (H3PO4, triprotic)

The strength of an acid can be measured by

using its ionization equation

ex.

The weaker the acid,

the smaller the Ka value and the less the acid ionizes in the solution

Water is

amphoteric- it can act as an acid or a base

Water can undergo

autoionization- it acts as both an acid and bas to itself

Ion product constant for water

Kw

At 25 C Kw is equal to 1.0×10^-14

In a neutral solution, like pure water,

In an acidic solution,

In a basic solution,

The strenght of an acid and its conjugate base are

inversely related

the stronger the acid, the weaker the conjugate base

Equation for pH

Equation for pOH

The pH scale

The pOH scale

A strong base

completely dissociates to produce OH-

Ka x Kb =

Kw = 1 × 10^-14. at 25 C

The strength of an acid and its conjugate base are

inversely related

The stronger the acid,

the weaker the conjugate base

Anions

Conjugate base of a strong acid generally produces a neutral solution

Conjugate base of a weak acid produce a basic solution

Cations

Cations that are counterions of strong bases

ex. NaOH, Ca(OH)2, KOH

produce a neutral solution

Cations pt.2

Cations that are the conjugate acids or weak bases produce an acidic solution

Cations pt.3

Cations that are small, highly charges metals produce an acidic solution

ex. Fe3+, Al3+

Salts in which neither cation nor anion acts as an acid or base

form a neutral solution

ex. NaCl

Salts in which the cation does not act as an acid, but the anion does act as a base

form a basic solution

ex. KF

Salts in which the cation does act as an acid, but the anion does not act as a base

form an acidic solution

ex, NH4Br, FeCl3

Salts in which the cation does act as an acid, and the anion does act as a base

Acidity or basicity depends on relative strengths

Polyprotic acids

ionize in successive steps, each with its own Ka

successive Ka, values get smaller

Lewis acid

electron pair acceptor

doesn’t have lone pairs, so it accepts lone pairs

ex. BF3

Lewis base

electron pair donor

Has lone pairs, so it donates them

ex. NH2

Examples of Lewis Acids and Base

Buffer

resists pH change by neutralizing added acid or base

A buffer contains significant amounts of either

• a weak acid and it conjugate base

• a weak base and its conjugate acid

The Henderson-Hasselbalch Equation

• the equation holds as long as the x is small approximation is valid

the initial concentrations of acids and conjugate bases should be at least 100 to 1,000 times greater than the equilibrium constant

What color should your titration be?

A very light pink

Dark pink means you overshot

For titrations,

the acids and bases need to be in moles for the RICE table

What is the pH at the equivalence point in a SA and SB titration?

pH of 7

Before the equivalence point in a SA and SB titration,

H3O+ is in excess

After the equivalence point in a SA and SB titration,

OH- is in excess

Solubility-product constant

K_sp

a measure of the solubility of a compound

Molar solubility

moles of a compound that dissolve in 1 L of liquid

Know that…

molar solubility doesn’t equal K_sp

When comparing solubility, you cannot compare K_sp values directly unless

the two compounds produce the same number of ions from their formula unit

In general, the solubility of an ionic compound is…

lower in a solution containing a common ion than in pure water

In general, the solubility of an ionic compound with a strongly basir or weakly basic anion increases with

increasing acidity (decreasing pH)

If Q < Ksp,

the solution is unsaturated and more ionic compound candissolve in the solution

If Q = Ksp,

the solution is saturated; the solution is holding theequilibrium amount of dissolved ions

If Q > Ksp,

the solution is supersaturated