Unit 6 Electrochemistry Flashcards Preview

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Flashcards in Unit 6 Electrochemistry Deck (29):
1

Oxidation number

(also known as oxidation state): the apparent net electric charge assigned to an atom or an element

-->helps keep track of electrons
In ionic compounds, the assigned charge on each ATOM reflects the loss or gain of electrons in making the ionic bond

In covalent compounds, this is the charge an atom would have if the electron pairs in a covalent bond belonged to the more electronegative atom

2

Rule 1: the oxidation number of any pure element or diatomic molecule is __
ex: O2, Fe

zero

3

Rule 2: the oxidation number of a monoatomic ion is ___
ex: Cl-, Mg2+

equal to its charge

4

Rule 3: the total sum of the oxidation numbers in a compound is ____ or ______ if an ion
ex: NaCl, NO3-

zero if neutral

equal to the charge if an ion

5

Rule 4: Fluorine always has an oxidation number of _

-1 in ALL compounds and ions
-->most electronegative

unless F2 = 0

6

Rule 5: alkali metals in a compound or ion always have an oxidation number of _
ex: NaCl

+1

7

Rule 6: alkaline earth metals in a compound or ion always have an oxidation number of __
ex: MgCl2

+2

8

Rule 7: In most compounds, Aluminum has an oxidation number of _, silver __, and zinc __

Al +3
Ag +1
Zn +2

9

Rule 8: In most compounds, hydrogen has an oxidation number of _

*Go to

+1

Exception: unless a hydride (use naming) with alkali or alkaline earth metals -1
ex: MgH2

10

Rule 9: In most compounds, oxygen has an oxidation number of __

*Go to

-2

Exception: unless a peroxide -1
Ex: H202 or Na2O2

11

*Rule 10: IN all compounds, a halogen as a halide (-ide) has an oxidation number of __

-1
ex: chloride, bromide, iodide

Exception: when combine with an element more electronegative than itself. Use F>O>Cl>others where the more electronegative element is assigned the negative oxidation number (determines which element gets the - and +)

ex: OCl2, OF2-

12

Rule 11: Sulfur as a sulfide has an oxidation number of __
Nitrogen as a nitride has an oxidation number of __

S -2
N -3

*In sulfates, S is usually +6
In nitrates, N is usually +5

13

Rule 12: when a compound contains two polyatomic ions,

break the compound into its individual ions (each with a charge or neutral)

Ex: (NH4)2SO4

(ignore coefficient/no effect)

14

Rule 13: All ionic compounds __

Can be split into their positive and negative ions if necessary to determine the oxidation state or charge

Metals follow the naming and formula writing rules

Ex: Fe(NO3)3
Iron (III) nitrate
Therefore, Fe is assigned +3

15

Rule 14: if any other rule does not work,

assume the oxidation number is the same as another member of the same family

16

How to balance redox reactions

1) Oxidation number method
2) half-reaction method (for acidic and basic solutions)

17

Redox reaction

Oxidation and reduction chemical reaction

Oxidation: increasing oxidation number or process whereby there is a lost of electrons
Reduction: decreasing oxidation number or process whereby there is a gain of electrons

^must occur together

18

Using the Standard Reduction Potentials Table to determine spontaneity

Use the Standard Reduction Potentials table (shows the relative strengths of oxidizing and reducing agents)
• Oxidizing agents on left side (strongest top); reducing agents on right side (strongest bottom)
-MER and MEO

Determine if the following reactions will occur spontaneously:
1) locate the REACTANTS (one will be the oxidizing agent and one will be the reducing agent) -
2) if the diagonal is from left to right (downward), then will react spontaneously
3) predict products (balance the electrons, then cancel electrons, reverse the oxidation reaction)

19

Electrochemical cell

A device that uses a SPONTANEOUS chemical reaction to produce an electric current (chemical to electrical energy) ex: a battery
-called voltaic or Galvanic cells
-Produces electrical energy via spontaneous chemical reaction

Electrode: a solid electrical conductor (i.e. metal zinc and copper)

Half-cell: an electrode and an electrolyte that form a half-cell

Cell: a system in which two connected electrodes are in contact with an electrolyte

Galvanic cell: an arrangement of two connected half-cells that spontaneously produces an electric current due to a redox reaction.
• Reaction occurs at the interface between the electrodes and the solutions where the electron transfer occurs
• Electrodes are connected by a wire. Solutions are connected by a salt bridge

-Anode is the site of oxidation (negative pole for electrochemical cells)
-Cathode is the site of reduction (positive pole)
AN OX RED CAT
• Use the half reactions to help (shows produces what)

Salt bridge: a U-shaped tube that contains an electrolyte solution and connects two half-cells in a galvanic cell
• The ends of the tube are covered with cotton balls so that the solution does not pour out. Cotton balls are porous and allows the ions to go through
• Soluble -->choose KNO3
• Ions do not react with the electrodes or electrolytes in the solution
• Positive ions in salt bridge flow into the side losing positive ions
• Negative ions in salt bridge flow into the side producing positive ions
• Result: the solutions in each half-cell remain electricially neutral (otherwise the build up of charges would stop the redox reaction)

Directions:
• In the conducting wire: Electrons flow from the anode to the cathode (always)
• In the salt bridge: + ions go to cathode, - ions go to anodes (reverse the ions Ex: NaCl becomes written as Cl- with anode on left and Na+ with cathode on right)

Observations:
• Oxidation electrode decreases in mass (producing ions)
• Reduction electrode increases in mass (producing solid)

Line notation: used to describe cells
Anode(s)/electrolyte//electrolyte/cathode(s)
Ex: Zn/Zn2+//Cu2+/Cu →(teacher does it another way)
• SOLIDs and its ion are in their own half cell
-> if given this notation, then electrochemical cell

*The potential difference between the electrodes (voltage) causes electrons to flow from the reductant to the oxidant through the external circuit, generating an electric current

20

Net Ionic equations


Net ionic equations:
• No spectator ions (i.e. NO3-) and electrons (electrons lost = electrons gained)
• Combination of the half reactions (metal and its ion are on OPPOSITE sides)
o Reduction reaction produces metal while the oxidation reaction produces ions
• Usually reverse the oxidation reaction (because on the table not written as a reactant), but once both reactions are combined, the REACTANTS combine to form products (metal and its ion are on opposite sides)

The net ionic equations still have to balance (charges and atoms)

21

Activity Series of Metals and Spontaneity

Arrange the metal ions in order of...
-Decreasing tendency to attract electrons (most affinity to least affinity; left top to bottom)
-Strongest oxidizing agent to weakest oxidizing agent (left top to bottom)
-Decreasing strength of oxidizing agent (strongest to weakest)
-Increasing strength of oxidizing agent (weakest to strongest)

1) List metal ions on the left and metals on the right
2) if reaction occurs, then metal ion is above metal (downward diagonal left to right). If reaction does not occur, then metal ion is below the metal (upward diagonal)
3) write all the ions of metals on the left
4) arrange according to the question (higher means stronger oxidizing agent)

22

Using the standard reduction potentials of half-cells Table

MER: most easily reduced -->top left-hand corner
-strongest oxidizing agent
-high affinity for electrons

MEO: most easily oxidized -->bottom right-hand corner
-strongest reducing agents
-easily gives up electrons

23

What do you do when there are two or more possible oxidizing more reducing reactions?

Ex: which Cu2+ will reduce?

Reduction reaction: choose MER (higher on left side)
Oxidation reaction: choose MEO (lower on right side)

24

Net cell potential (E voltages)

Each half cell reaction (oxidation and reduction) has a voltage (given on the table). Reverse the sign for the oxidation reaction (because writing reactants first)

Net cell potential is the sum of the E values.
If the net cell potential...
+ voltage = spontaneous reaction
- voltage = nonspontaneous reaction

*need sign of either + or -

^use to check answers

Net cell potential does not change with the multiplication of a factor that makes the electrons the same on both sides

25

Electrolytic Cell

The reverse reaction of the electrochemical cell is non-spontaneous and requires electrical energy to occur.

Electrolytic cell: a cell that uses electrical energy to produce a chemical change that would not occur spontaneously
• Anode and cathode are reversed (anode – becomes cathode -, cathode + becomes anode +)
• The direction of electron and ion flow are reversed
• MINIMUM voltage required
- use of inert electrodes to conduct electrons and complete the circuit

Electrolytic cells use Electrolysis: the application of current from an external source through a cell to produce a chemical change
Ex: electrolysis of water
• Twice as much hydrogen as oxygen is produced
• H2O → 2H2 + O2

Usually one beaker

Reduction still takes place at the cathode and oxidation occurs at the anode.

The sign (not the magnitude) of the cell potential has been reversed (minimum voltage required to initiate the reaction)

Molten and Aqueous salts
-ions in the beaker
-MUST CONSIDER ALL POSSIBLE SUBSTANCES (watch out for H20, Fe, Cr, Pb, Sn, Cu -->all can compete for oxidation and reduction). Choose based on MER and MEO
-positive ions are attracted to the negative cathode (reduced). Negative ions are attracted to the positive anode (oxidized).
-Pay attention to the products produced

26

Standard conditions for cells

1 M
Gases have a pressure of 1 atm
temperature 25 degrees celsius
pH 4 in the anode

27

Quantitative Electrolysis

There is always a simple relationship between the amount of substance produced or consumed at an electrode during electrolysis and the quantity of electrical charge Q which passes through the cell.

m = (mm)(I)(t)/(F)(# e-)
Faraday's Constant: 96500 C/mol
1 Ampere = 1 C/second

**mm of the metal! Time in seconds!

-find moles using C = n/V

28

Application of Electrolysis: Electroplating

• Cathode (- pole) →object to be plated
• Anode (+ pole) → plating metal

*the anode and cathode reactions are the same but reversed

Extra:
Electro refining metals (purification): anode is the impure metal; cathode is pure metal →metal is deposited on the cathode; impure metal dissolves

29

Application of Redox Reactions: Corrosion

Corrosion: the deterioration of a metal by a redox reaction (spontaneous reaction)
• Has negative effects on the environment
• Metals oxidize easily (are below oxygen on the SRP chart)
o Metals oxidize (lose electrons)
o Oxygen reduces (gains electrons)
• Cannot tell how fast the reaction occurs
• Very active metals are used as structural materials because of the oxide that forms on the surface. However the oxide that forms on iron and steel often peel off, exposing the surface to corrosion.
• Gold does not corrode because it has a standard reduction potential greater than oxygen
• Non-uniform surfaces causes metals to be easily oxidized at these stress points (anode regions)
• Moisture is important →water acts like the salt bridge and is involved in the reduction half reaction
• Salt accelerates the process by acting like a salt bridge and providing ions

• What happens during rusting: Electrochemical
o Electrons from the oxidation reaction go to the cathode where oxygen is reduced. At the cathode region, the ions of the metals react with oxygen to form rust →hydrated iron oxide

Prevention of corrosion:
• Coating:
o Protects the metal from oxygen and moisture
o I.e. plating steel with other metals, such as chromium and tin, with plastic, paint, grease

Alloys:
o Combines two metals
o These metals form oxide coatings that help steel resist corrosion
- ex: stainless steel is a metal alloy with chromium and iron

Galvanizing/Active Metal Contact: the process in which steel is coated with a thin layer of zinc to protect the steel from corrosion.
o More active metal oxidizes
• Sacrificial anode – a form of cathodic protection in which the oxidation of a more active metal is attached to steel and prevents the iron in the steel from being oxidized
• Must replace the sacrificial anode because it gets oxidized
**Choose more active metal