chapter 8

Question 1

Leo the lion says Ger

Answer 1

• Loss of electrons = oxidatin
• gain of electrons = reduction

Question 2

a redox reaction is one in which?

Answer 2

• the oxidation states of compounds change

Question 3

what is oxidation state?

Answer 3

• a method for keeping track of how electrons are shared within a molecule
  
  • a model of electron distribution where we simplify things by assigning the electron to the mroe electronegative atom in a bond

Question 4

rules for calculating oxidation state:

Answer 4

• pure elements have an oxidation state of 0 and applies even if they are found in diatomic molecules such as O₂
• the oxidation state of monoatomic ions is equal to their charge so Fe²⁺ ion has an oxidation state of +2 and Cl^- ion has an oxidation state of -1
• the sum of the oxidation states of the components of molecules and polyatomic ions is equal to their charge (zero for neutral molecules and some integar value for polyatomic ions)

priority rules for individual atoms within a compound:

• the oxidation state of F is -1 because it is the most EN
  
  • other halogens will usually have an oxidation state of -1 unless they are bonded to a more EN halogen, N or O
  • the oxidation state of H is +1 except when it is bonded to a more electropositive element (NaH, NaBH₄, and LiAlH₄), in which case it will be -1.
  • the oxidation state of O is usually -2 with some important exceptions, such as peroxides, in which it is -1
  • the oxidation state of alkali metals (the first column in the PT is always +1 and that of alkaline earth metals (2nd column on the PT) is always +2

Question 5

what are examples of non-redox reactions?

Answer 5

• the oxidation states do not change because the overall distribution of electrons remains the same
  
  • acid base chemistry
  • precipitation reaction
  • many substitution reactions
  • many double displacement reactions

Question 6

what are examples of redox reactions?

Answer 6

• ‘classic’ redox (single displacement reactions)
• combustion
• co,bination reactins
• many metabolic reactions

Question 7

how to balance redox reactions:

Answer 7

1. split into half reactions
2. balance the non-O and non-H atoms
3. balance the O
4. balance the H
5. add electrons to balance the charge
6. multiply so that both half reactions have the same number of electrons
7. add and cancel like terms
   
   • this is for acidic conditions, but if basic conditions applied, wen use OH^- to balance out the oxygen molecules and H₂O to balance out the H⁺ molecules

Question 8

what is an oxidizing and reducing agent respectively?

Answer 8

• an oxidizing agent is a compound that you can add to a reaction mixture to cause another substance to be oxidized
• a reducing agent is a compound that you can add to a reaction mixture to cause another substance to be reduced
  
  • this means that oxidizing agents are themselves reduced and reducing agents are themselves oxidized.

Question 9

what are common oxidizing agents?

what are common reducing agents?

Answer 9

common oxidizing agents:

• CrO₃
• Na₂Cr₂O₇
• pyridinium chlorochromate (PCC)

common reducing agents:

• NaBH₄
• LiAlH₄

Question 10

how do we measure reduction potentials (E°)?

Answer 10

• measured in volts and are defined relative to the standard hydrogen electrode which is defined as being 0V
  
  • greater (more positive) reduction potentials indicate that a substance wants to be reduced more
  • smaller (more negative) reduction potentials indicate that a substance wants to be reduced less and would prefer oxidation if possible
    
    • reudction potentials start being useful when we compare them!

Question 11

reduction potentials are not affected by?

Answer 11

• stoichiometric constants

Question 12

how can oxidation potentials be found?

Answer 12

• flipping the sign on the reduction potentiral of the reduction half-reaction

Question 13

the electrode where oxidation occurs is known as the?

the electrode where reduction happens is known as the?

mnemonic?

Answer 13

• anode
• cathode
• An Ox Red Cat
  
  • ​therefore a surplus of electrons is generated at the anode (because electrons are lost during oxidation) and they travel to the cathode

Question 14

in a galvanic cell/ voltaic cell, a spontanous redox reaction is used to generate?

Answer 14

• a positive potential difference that can drive current

Question 15

the total standard potential generated by a cell, E_cell, can be calculted by:

Answer 15

• E_cell = E°_cathode - E°_anode

Question 16

how to solve when you are given a table of standard half reduction potentials and then asked to calculate the potential of the galvanic cell that would be set up using 2 specific half-reactions:

Answer 16

• identify which half reaction has the higher reduction potential (this will be the cathode because it wants to be reduced more) and the other half-reaction will be the anode
• plug those values into the equation E_cell = E°_cathode - E°_anode
• understand that the actual half-reaction that happens at the anode will be the reverse of the one shown in the table of reduction potentials because oxidation happens at the anode

Question 17

When there is a difference between the standard values of a parameter (which assumes Standard conditions) and the values of a paramenter in non-standard (real world) conditions, we use the Nernst equation which helps us account for how electrical potential of a cell is affected by conditions including temperature and the concentration of reactants:

Answer 17

• E’_cell = E°_cell - (RT/zF) ln Q
  
  • E’_cell refers to the actual cell potential under a given set of conditions
  • E°_cell is the standard cell potential
  • R is the ideal gas constant
  • T is temp in K
  • z is the number of moles of electrons transferred
  • F is the Faraday constant (1 x10⁵)
  • Q is the reaction quotient
    
    • simplified version is E’_cell = E°_cell - (0.05916)/z) log₁₀Q

Question 18

what is a Daniell cell?

Answer 18

• prototype of a galvanic cell
  
  • the half-reactions are carried out in 2 physically separated half-cells
  • the electrodes are connected by a conductive wire and the half cells are additionally connected by a salt bridge

Question 19

what is a common shorthand for denoting Daniell cells?

Answer 19

anode| anode solution || cathode solution| cathode

Question 20

In Daniell cells, the cathode is on the ? and the anode is on the ?

Answer 20

• the cathode is on the right (where electrons are taken out of solution via a reduction)
• the anode is on the left (source of electrons)

Question 21

in the Daniell cell, the 2 half-reactions are physically separated but it is possible to create a galvanic cell in which the 2 half-reactions take place in the same chamber. such cells are known as?

Answer 21

• concentration cells and they have to satisfy 2 conditions:
  
  • there needs to be a concentration difference between 2 regions of the cell
  • the electrodes need to be made out of the same material
    
    • applications in pH and on each side of the cell membrane

Question 22

what happens in an electrolytic cell?

Answer 22

• the idea is to apply energy to the system rather than to obtain energy from the system
  
  • this means that tehr edox reactions will be carried out in the nonspontaneous direction
    
    • the name reflects the fact that these cells are often used to break down compounds into their constituent parts through the application of electrical energy

Question 23

in electroyltic cells, the cathode (where reduction is carried out) is marked with a ?

while the anode (where oxidation is carried out is marked with a ?

Answer 23

• negative charge
• positive charge
  
  • the opposite of what occurs in galvanic cells
  • the convention used for an electrolytic cell is the same that is used for SDS-PAGE and isoelectric focusing so the Cathode is neg and the anode is pos (PANIC)

Question 24

galvanic and elecrtolytic. ells are combined in?

Answer 24

• batteries
  
  • discharge spontaneously (acting as a galvanic cell that produces current)
  • recharged (acting as an electrolytic cell)

Question 25

2 batteries to know: * lead-acid * nickel-cadmium batteries

Answer 25

Lead-acid: * in its charged state, a lead storage battery has a lead oxide electrode and a lead electrode seperated by an area that contains concentrated H₂SO₄ * _​2 redox half-reactions take place * the oxidation half-reaction takes place at the Pb electrode (Pb_(s) + HSO_4-(aq) → PbSO_4(s) + H⁺_(aq) + 2e^- * the reduction half-reaction is PbO_2(s) + HSO₄^- + 3H⁺ (aq) + 2e^- → PbSO₄(s) + 2H₂O (l) * the product of both half reaction is solid ( PbSO₄(s)) because it is very poorly soluble in water so we are left with some dilute leftover * once the lead-acid battery is discharged, current can be applied to recharge it,r eversing the reaction Nickel-cadmium * one cathode contianing nickel oxide-hydroxide and one anode containing cadmium * when discharging, the oxidation half-reaction is Cd + 2OH^- → Cd(OH)₂ + 2e^- * the reduction half reaction is 2 NiO(OH) + 2 H₂O + 2e^- → 2Ni(OH)₂ + 2OH^- * during recharging, the reactions are reversed

Question 26

the key to colorimetric redox titrations is to?

Answer 26

* find an indicator that changes colour between its oxidized and reduced form * ex. Mn. The polyatomic MnO₄^- ion, in which Mn has an oxidation state of +7 is purple in aqueous solution, whereas Mn²⁺ is colourless

Question 27

what equations to use in redox titrations?

Answer 27

M₁V₁ = M₂V₂ or N₁V₁ = N₂V₂ * in redox chemistry, normality related to the number of electrons that are either given up or absorbed by one mole of a species in the corresponding redox reaction so for Mn it is 5 because (+7 to +2) * normality can be calculated by multiplying each of these values by the molarity of the associated solution

Question 28

redox titrations have many applications. one example is Benedicts reagent:

Answer 28

* tests for reducing sugars, which after heating reduce Cu²⁺ to Cu⁺, causing a colour change

Question 29

the delta G° of a redox reaction is directly connectef with the E°_cell value as given in this equation:

Answer 29

deltaG° = -nFE°_cell * deltaG° is the standard Gibbs free energy change of a reaction (J) J = C x V * E°_cell is the standard cell reduction potential * n and F reger to the moles of electrons exchanged in the reaction and Faradays constant

Question 30

deltaG° = -RTlnK_eq is also:

Answer 30

nFE°_cell = RTlnK_eq

Question 31

how to know if it is spontanous or non spontaneous?

Answer 31

* delta G: * spontaneous: negative \<0 * nonspontaneous: positive \>0 * E°_cell * ​spontaneous: postive \>0 * nonspontaneous: negative \<0 * k_eq: * spontaneous: greater than 1 \>1 * nonspontaneous: less than 1 (0eq\<1)