All my love for Chem Pt.2

Arrhenius Concept

Acid - produce hydrogen ion in an aqueous solution
Base - produce hydroxide ion in an aqueous solution

Bronstead Lowry Concept

Acid - proton donor
Base - proton acceptor

Conjugate Acid & Base

Conjugate Acid - species formed when proton transferred to the base
Conjugate Base - what remains of acid molecule after proton is lost

Lewis Acid & Lewis Base

Acid - any substance that accepts a lone pair of electrons
Base - Any substance that donates a lone pair of electrons

Lewis Acid/Base Examples

Acid is usually a cation or something that has an incomplete octet or central atoms can hold more than 8 valence electrons
Base is usually anion (Polyatomic Ions)
Or unpaired electrons

Strong Acid Relative Strengths

Strong acid wants to give away proton; the more readily it wants to give it away, the strong the acid (higher dissociation)
Creates a weak C. Base
Strong Base creates weak C. Acid because it does not want to give it away

Common Strong Acids

HCl, HBr, HI, HNO3, HCLO4, H2SO4

Common Strong Bases

LiOH, KOH, NaOH, RbOH, CsOH, Sr(OH)2, Ba(OH)2

Determining Strength Using Equil Constant

Larger Ka value, stronger the acid & the weaker the base
Smaller Ka value, weaker the acid & the stronger the base

Binary Acid

acid containing hydrogen as well as nonmetallic element

Oxyacid

acid containing oxygen, hydrogen, and another element
General Formula: HnYOm
Y is a polyatomic anion that contains metal or nonmetallic atom

Binary Acid Strength

For elements in same group of periodic table: shorter bonds are stronger → directly related to atomic radius
For elements in same period of periodic table: more electronegative, stronger the acid

Oxyacid Strength

For oxyacids with same # of OH groups & same # of O atoms, acid strength increases as electronegativity of Y increases
Oxyacids with same Y atom but different # of oxygen atoms, acid strength increases as # of oxygen atoms increase

Significant difference between acid strength values

Know its significant if different by 3 order of magnitude or more for Ka values

Monoprotic

one dissociable proton

Diprotic

two dissociable protons

Calculating Polyprotic Acids

Ka gets significantly weaker (only consider 1st when calculating pH)

Relationship between Ka and Kb

Ka * Kb = Kw = 1×10^-14

Acid Base Properties of Salts

Acid + Base → Salt + H2O
Can go backwards to determine what acid/base created the salt

Salt of strong base + strong acid

No hydrolysable ions and gives neutral solution (neutral)

Salt of strong base & weak acid

Anion of salt is conjugate of weak acid so it will hydrolyze and give a basic solution (basic salt)

Salt of weak base and strong acid

Cation is conjugate of base, so it will hydrolyze and give an acidic solution (acidic salt)

Salt of weak base & weak acid

Both ions hydrolyze, so pH depends on relative acid-base strength of two ions (depends on Ka/Kb values)

Determining Salt Acidity from Weak Base and Weak Acid

Which ever value is larger, that is the ion that will dominate the equilibrium
EX: CN- > NH4+ (BASIC SOLUTION)
Kb > Ka

Acidic Buffer

Consists of weak acid and its salt (conjugate base)

Basic Buffer

Consists of a weak base and its salt (conjugate acid)

A-/HA Ratio & Buffering Capacity

The closer the A-/HA ratio is to 1, the buffering capacity increases; all solutions that have the same A/HA ratio will have the same pH

Choosing & Preparing Buffer Solutions

The best buffer solution is the one where the pKa of the weak acid is as close as possible to the desired pH

Strong Acid-Strong Base Titration regions

1. Initial pH

2. Between initial & equivalence point

3. Equivalence Point

4. After EP

Equivalence Point Equation

d²pH/dV² = 0

When pH = pKa

Vb = 1/2Vc

Weak Acid - Strong Base Titration Regions

1. Initial pH

2. Buffer region (HH equation)

3. Equivalence Point

4. After EP

Two distinct inflection points

Titration Curve Dependencies

Strength of acid/base

- As strength decreases, rapid rise decreases
- pKa > 8
Concentration of original solution
- As concentration increases, rapid rise increases
- HA EP becomes blurred when HA too diluted

Titrations in Multi-protic Systems

Starts out with low pH value
Go through titration curve multiple times

When titrating acid, how much In- must be present to detect color change

In-/HIn - 1/10

When titrating base, how much In- must be present to detect color change

In-/HIn = 10/1

Choosing Indicator

Ideally, want endpoint of indicator to be as close to EP as possible

Getting up Solubility Equations

Whole solid on reactants side, then dissociated ions on the products

Common Ion Effect

The shift in an equilibrium position caused by the addition or presence of an ion in the equilibrium reaction

All Group I OH- salts are

soluble

Formation Constant

Kf - used for complex ions

Solubility Constant

Ksp

Effect of Formation of Complex Ions on Solubility

The solubility of a sparingly soluble salt will INCREASE if a ligand that forms a stable complex ion is added to the solution

Q < Ksp

Solution of NOT saturated (more salt will dissolve)

Q > Ksp

Solution of OVER saturated (precipitation will occur)

Q = Ksp

System is saturated

Oxidation Reduction Reaction

a reaction in which one or more electrons are transferred

Oxidation

loss of electrons (increase in oxidation state)

Reduction

gain of electrons (decrease in oxidation state)

Oxidizing Agent

a reactant that accepts electrons from another reactant

Reducing Agent

a reactant that donates electrons to another substance to reduce the oxidation state of one of its atoms

Steps for Balancing Redox Reaction in Acid Solution

1. Write separate equations for oxidation and reduction ½ reaction
2. For each halt reaction
- Balance all elements except H and O
- Balance O by adding H2O
- Balance hydrogen by adding H+
- Balance charge with electrons
3. If number of electrons are unequal between two ½ reactions, multiply ½ reactions by an integer to make it equal
4. Add ½ reactions and cancel identical species
5. Check that elements and charges are balanced

Steps for Balancing Redox Reactions in Basic Solution

1. Same first two steps as acidic solution
2. To both sides of ½ reaction equations, add a number of OH- ions equal to the number of H+ ions
- This creates water and eliminates H+
3. Finish as usual

Galvanic Cell

Utilizes electron transfer for electron energy; spontaneous; always from anode → cathode

Salt Bridge

has salt not made up of chemical in reactions; allows for ion transfer to maintain electron neutrality

Anode

Oxidation ½ reaction

Cathode

Reduction ½ Reaction

Line Notation for Galvanic Cells

Anode components on the left; cathode on the right; anode & cathode are separated by a double line (indicates salt bridge); phase boundary indicated by single line

Battery

a group of galvanic cells connected in a series

Lead Storage Battery

a battery (in cars) where the anode of lead, the cathode is lead dioxide, and the electrolyte is a sulfuric acid solution

Dry Cell Battery

common battery used in calculator, watches, etc

Fuel Cell

galvanic cell where the reactants are continuously supplied

Alkaline Battery

Lasts longer than acidic batteries because zine anode corrodes less rapidly

Cell Potential

Driving force in a galvanic cell that pulls electrons from the reducing agent (anode) to the oxidizing agent (cathode) (Ecell)
Galvanic Cells always runs spontaneously in the direction that produces a (+) cell potential

Volt

one Joule of work/one Coulomb of charge transferred

Standard Reduction Potential

The higher the SRP, the stronger the oxidizing agent (cathode)

SHE Cell

Pt conductor in contact with 1M H+ ions and bathed in H(g) at one temperature

Electronegativity and Cathode Strength

The more electronegative a molecule is, the more it wants to “grab” electrons and be the cathode (higher V value)
F is the most electronegative, will force everything beneath it to oxidize

Electromotive Force (V) =

-work(J)/charge(C)

Change in Free Energy

The change in free energy equals the max number of useful work obtainable from that process
wmax = deltaG

Faraday (F) Constant

constant representing one mole of electrons
96485 C/mole e-

deltaG with Cell Potential

deltaG (J) = -nFE

deltaG with K

deltaG = -RTln(K)
R = 8.314

Negative or Positive E Value Indications

-E value = nonspontaneous
+E value = spontaneous

Electrolysis

A process that involves forcing a current through a cell to cause a non-spontaneous reaction to occur
Input electron current greater than Ecell

Ampere

C/sec

Stoichiometry of Electrolytic Process

current and time → quantity of charge in coulombs (F) → moles of electrons → moles of element → grams of element

Nerst Equation

Ecell = E0cell - (RT/nF)ln(Q)
Q = products/reactants
R = 8.314
Usually used when have concentrations of molecules

Ecell Shift when Adding reactants

Ecell will increase because eq shifts to the right

Ecell Shift when Adding products

Ecell will decrease because eq shifts to the left

Concentration Cells

A Galvanic cell in which both compartments contain the same components but different concentrations
Eventually concentrations will equal out

Reaction Rates

The increase in the concentration of a product per unit time OR the decrease in the concentration

(Average) Rate =

= M change/time change

Reaction Order

the value of the exponents of concentration terms in the rate law

Zeroth Order

Rate = K
M/s or Ms-1

First Order

Rate = K[A]
1/s or s-1

Second Order

Rate = K[A][B]
M^-1*s^-1

Third Order

Rate = K[A][B]²
M^-2*s^-1

Determining Order for Reactant

Have to experimentally change each reactant and see how that effects the rate law

Reaction Mechanisms

The sequence of reaction steps that describes the pathway from reactant to products

Intermediate

A molecule that is formed in one elementary reaction and used up in the next one; not found in overall reaction

Elementary Reaction

A reaction whose rate law can be written from its molecularity

Molecularity

number of species that must collide to produce the reaction indicated by that step

Unimolecular

A → Productsk
Rate = K[A]

Bimolecular

A+A → Products or A+B → Products
Rate = K[A]² or K[A][B]

Termolecular

A+A+B → Products or A+B+C → Products
Rate = K[A]²[B] or K[A][B][C]

Rate Determining Step

The slowest step in a reaction mechanism; determines the overall rate
All the ER have to wait for the slow one
Use the reactants to form the rate law?

Multistep Reactions with Initial Fast Step

No good way to figure out concentration of some molecules because they are all used up; need to go back to previous reaction (Step 1)

Catalyst

1st shows up as a reactant then it is completely reformed as a product
Speeds up overall reaction process

Temperature Dependence of Rate Constant

Chemical reactions speed up when temperature rises