Chem 102 Study Deck

Relate the strengths of binary acids to their molecular structures

Acids get stronger with:
more electronegative atoms
more resonance structures

longer bond lengths

How is Ka related to the strength of an acid?

The bigger the Ka value, the stronger the acid
The smaller the Ka value, the weaker the acid

How can you recognize conjugate acid-base pairs and their complementary base or acid

conjugate bases and acids differ from each other by one proton (H+ atom)

Interconvert [H3O+ ] and pH

-log[H3O+]=pH

10^-pH=[H3O+]

Interconvert [OH-] and pOH

-log[OH-]=pOH

10^-pOH=[OH-]

Relate the pH, percent ionization, and the principal solute particles present in solution

to the Ka and Kb values of weak species

pH indicates whether more acid/base is present
Major species=determine if the original species was strong/weak

Calculate the pH of an aqueous solution containing strong acid or base

Use rice table, use K expression, then convert to pH/pOH

Use a RICE table governed by Ka or K b and initial acid or base concentration to

calculate the equilibrium concentrations of H 3O+ and/or OH- and report as pH

Convert to needed K value, find final concentration of H3O+/OH-, then use -log[H3O+/OH-] to find pH/pOH

Explain the 5% rule and use it when appropriate

5% rule is when you can ignore the -x in the reactant term
5% rule is appropriate when the K value is very small and the change in the initial concentration of the reactant will not be significant

For polyprotics acids, write equilibrium reactions and Ka expressions using stepwise

dissociation

Write dissociation equation and equilibrium expression for first proton, then write the dissociation equation and equilibrium expression for the molecule with the second proton

Relate Ka to Kb for a conjugate pair using Kw

Kw=Ka*Kb

Identify the major species in aqueous solution that will affect pH

The major species that will affect pH are the leftover conjugate base/acid (could be one)

Explain the common ion effect

The concentration of a conjugate base MUST take into account the concentrations of introduced salts
ie: .1M NaCl combined with .1M HCl = .2M Cl

Write the neutralization reaction that occurs when a strong acid is added to the buffer

A- + H3O+ —> HA + H2O

Write the neutralization reaction that occurs when a strong base is added to the buffer

HA- +OH —> H2O + A-

Write the equilibrium reaction that describes the chemistry of a buffer solution

A- + H2O —> HA + OH (can go in reverse too)

Use the Henderson Hasselbalch Equation to calculate the pH of a buffer solution

containing a weak acid or base and its conjugate base or acid

pH=pKa + log ([base]/[acid])

Evaluate a buffer’s capacity to resist changes in its pH with the addition of acid or base

and calculate pH after that addition

Use BCA table to determine how many moles are left after reaction, if there are equal amounts of weak base/acid and its conjugate left, then the solution is a buffer.
If the amounts of weak base/acid and its conjugate are not equal, then the buffer has reached its buffer capacity

Describe how to prepare a buffer with a desired pH

Find weak acid with pka within ±1 of the desired pH, then use hendersson hassel bach equation to find the molarity ratio

Identify the half-equivalence point (midpoint) and explain its significance

Half equivalence point occurs halfway to the equivalence point (where [titrant]=[analyte])
Half equivalence point a point in the buffer region where pH=pka (because [HA/A-]=[conjugate pair])

Calculate the pH during the titration of a monoprotic strong/weak acid (or

strong/weak base), from initial conditions to the equivalence point

Use a BCA table for the neutralization reaction, then use the final volume (of both the titrant and analyte) to calculate the molarity (moles/litres)

Relate the solubility of an ionic compound to its solubility product constant, Ksp

the bigger the ksp value, the more soluble

the smaller the ksp value, the less soluble

Determine how solubility changes with changes in common ion concentration and solution pH

solubility decreases with the common ion effect because of le chatelier’s principle
solubility increases with a lower solution pH because of le chatelier's principle

Relate the strength of bases to their molecular structure

Bases get stronger with:

greater electron density

less electronegative elements

Relate the strength of oxoacids with their molecular structure

Bond length does not matter
Oxoacids become stronger with:

more oxygens

weaker conjugate base(if the conjugate base is less stable)

more electronegative heteroatom (ie: HClO4 stronger than HBrO4 bc Cl is more electronegative)—inductive effect

Relate the strengths of polyprotic acids with their molecular structure

Less Acidic as more protons (H+) are given away (want to keep neutral charge)