Unit 4 Notes Electrochemistry.pdf

  Unit 4 Electrochemistry

• _- Electrochemistry the interconnection between electricity and chemical reactions

  Oxidation – Reduction Reactions

  Oxidation – a process in which chemical entities lose electrons _{– 2+} ex. Zn (s) (aq) + 2e  Zn

  Reduction – a process in which chemical entities gain electrons _{– 2+} ex. Cu (aq) + 2e (s)  Cu

  Memory Aids

  LEO the lion say GER Losing Electrons: Oxidation Gaining Electrons: Reduction

  Oxidation Is Loss Reduction Is Gain

  Redox Reaction – a reaction in which one reactant is oxidized and the other reactant is reduced _{2+ 2+} Zn (s) + Cu (aq) (aq) + Cu (s)

   Zn In this reaction Zn loses electrons oxidized

  ∴ Cu gains electrons reduced ∴ To identify the oxidized and reduced reactants in a redox reaction:

  1. write the total ionic equation 2. write the net ionic equation 3. assign “0” charge to uncombined elements 4. compare charges to identify loss/gain of electrons or 1. if reaction is between two elements (no ions), assign charge of “0” to elements 2. determine ion charges in compound

  3. compare charges to identify loss/gain of electrons

  

ex. Identify the oxidized and reduced reactants in the following reaction

  Pb (s) + Cu(NO ^{3 ) 2 (aq) (s) + Pb(NO 3 ) 2 (aq)}  Cu Recall that the aqueous reactants and products are ions in solution.

  Step 1 – total ionic equation _{– 2+ 2+ –}

  Pb (s) + Cu (aq) + 2 NO ^{3 (aq) (s) + Pb (aq) + 2 NO 3 (aq)}  Cu

  Step 2 – net ionic equation _{– – 2+ 2+}

  Pb (s) + Cu (aq) + 2 NO ^{3 (aq) (s) + Pb (aq) + 2 NO 3 (aq)} _{(s) (aq) (s) (aq) 2+ 2+}  Cu Pb + Cu + Pb

   Cu

  Step 3 – assign charges _{2+ 2+}

  Pb (s) + Cu (aq) (s) + Pb (aq)  Cu

  Step 4 – compare charges _{– 2+}

  Pb becomes Pb – Pb losses 2e oxidized _{2+ –} ∴ Cu becomes Cu – Cu gains 2e reduced

  ∴

ex. Identify the oxidized and reduced reactants in the following reaction

  2 Mg (s) + O ^{2 (g) (s)}  2MgO

  No ions formed so use second method

  Step 1 – assign charges to elements

  2 Mg (s) + O ^{2 (g) (s)}  2MgO

  Step 2 – assign charges in compound _{(s) 2 (g) (s) 2+ 2–}

  2 Mg + O O  2Mg

  Step 3 – compare charges _{– 2+}

  Mg becomes Mg – Mg losses 2e oxidized

• _{– 2–}

  ∴

  O becomes O – O gains 2e reduced

  ∴ Homework Practice Oxidation/Reduction Use p.374-376 to help p.377 #1,2,3a,c,e,g,5

  Redox Reactions of Nonmetals

  The redox reactions studied up to this point all involve a metal with a non metal, also known as ionic compounds. Many redox reactions do involve non-metals that share electrons ex. sulfur dioxide can produce acid rain if released into the atmosphere S (s) + O 2 (g) 2 (g)

   SO

  Oxidation Numbers

  Non-metals will share electrons but often it is an unequal sharing because one element is more electronegative (stronger pull on electrons) than the other. Consider carbon dioxide – CO ² When we examine the electronegativity difference between carbon and oxygen in carbon dioxide, we notice that oxygen is more electronegative that carbon. Because of this chemists find it useful to assign an apparent charge on each atom. This apparent charge is called an oxidation

  number.

  2

  ex. #1 Determine the oxidation number of carbon and oxygen in CO

• _-

  according to rule #4, oxygen is always –2 _- there is 1 carbon and 2 oxygens the sum of the oxidation numbers must equal 0 for a compound C + 2 O = 0 C + 2(–2) = 0 C – 4 = 0 C = +4

  the oxidation number for carbon is +4 and oxygen is –2 ∴

  ex. #2 Determine the oxidation number of the nitrogen atom in lithium nitrate, LiNO ^{3 .}

  Li + N + 3 O = 0 1 + N + 3(–2) = 0 1 + N – 6 = 0 N – 5 = 0 N = +5

  the oxidation number for nitrogen is +5 ∴ Homework p.381 #1,2 Identifying Redox Reactions Using Oxidation Numbers

  Perhaps the one of the most useful applications of oxidation numbers is to determine if a redox reaction has occurred or not. In order for a redox reaction to have occurred, there must have been…

• an oxidation number decrease (reduction)
• an oxidation number increase (oxidation) ex. #1 – Use oxidation numbers to showthat the reaction of zinc metal with sulfur is a redox reaction. The chemical equation is…

  Zn (s) + S (s) (s)  ZnS

  Write all known oxidation numbers

  0 0 2+ 2– zinc oxidation number changes from 0 to 2+ sulfur oxidation number changes from 0 to 2–

  there are changes in oxidation numbers so this is a redox reaction ∴

• – ex. #2 Is the following reaction of sulfur trioxide with water a redox

  reaction? SO 3 (g) + H ^{2 O (l) 2 SO 4 (aq)}  H

  Write all known oxidation numbers

  S O ^{3 + H 2 O 2 S O 4}  H

  S –2(3) +1(2) –2(1) +1(2) S –2(4) S –6 +2 –2 +2 S –8 Solve for unknown S S – 6 = 0 2 + S – 8 = 0 S = +6 S – 6 = 0

  S = +6

  

there is no changes in oxidation numbers for any of the element so this

∴ is not a redox reaction Try

Which of the following equations represent redox reactions? Which do not

represent redox reactions? Prove your answer with oxidation numbers.

  b. CaCO + CO _{3(s) (s) 2(g)}  CaO c. 2H O + O _{2 (l)  2H 2(g) 2(g)}

  d. 2Li + 2H O + H _{(s) 2 (l) (aq) 2(g)}  2LiOH e. Fe O + 3CO + 3CO _{2 3(s) (g)  2Fe (s) 2(g)}

• – Homework p.383 #3 a j

  The Activity Series

• a list of metals arranged in order of their tendency to react (become oxidized), based on observations gathered from single displacement reactions

  Recall that a single displacement reactions involves one metal displacing another metal.

  A + BC AC + B The most reactive metals (the most easily oxidized) are at the top of the list, the least reactive metals are at the bottom ex. the metals at the top are so reactive that they even react with relatively unreactive substances like water, where metals like copper, silver and gold are unreactive even in the strongest acids _- Predicting Chemical Reactions if a metal is found higher in the activity series it will displace an

• _- metal that is lower the lower metal is always left as a pure metal ex. Use the activity series to predict whether a chemical reaction will occur and write the balanced chemical equation

  a) aluminum foil is placed in a solution of silver nitrate (AgNO ^{3 )} Al (s) + AgNO ^{3 (aq) 3 ) 3 (aq) + Ag (s)}

   Al(NO Al (s) + 3 AgNO ^{3 (aq) 3 ) 3 (aq) + 3 Ag (s)}

   Al(NO

  b) copper is placed in a solution of iron (II) sulfate (FeSO ^{4 )} Cu (s) + FeSO 4 (aq)

   no reaction

  c) zinc metal is placed in a solution of hydrochloric acid (HCl) _{(s) (aq) 2 (aq) 2 (g)} Zn + HCl + H

   ZnCl Homework p.388 a–e p.389 #1–4

  Galvanic Cells

• _- How do batteries work? next couple lessons will cover this

  Converting Chemical Energy to Electrical Energy

  Consider the redox reaction _{2+ 2+} Zn (s) + Cu (aq) (s) + Zn (aq)

   Cu _{– 2+} Zn becomes Zn , so zinc loses 2e (oxidation) _{2+ – -} Cu becomes Cu, so copper gains 2e (reduction) by separating the zinc and copper ions by placing a metal wire between them, electrons that lost by the zinc are forced to travel _- through the metal wire to the copper solution _- moving electrons have energy energy can be used to power a device such as a radio or a clock

  Galvanic Cell

• _- a galvanic cell converts chemical energy from redox reactions into _-

  electrical energy _- batteries contain galvanic cells these reactions are spontaneous, meaning they require no outside

• _- assistance or energy input the oxidation reaction of one metal and the reduction reaction of _-

  another metal are in separate compartments called half–cells _- each half–cell contains a solid conductor called an electrode _- the electrode where the oxidation occurs is called the anode the electrode where the reduction occurs is called the cathode

• _-

  the zinc half–cell consists of a strip of zinc in a zinc solution the copper half–cell consists of a strip of copper in a copper

• _- solution _-

  the half–cells are connected using a wire and a salt bridge a salt bridge is a tube that contains a concentrated solution of an electrolyte that replaces the ions in solution (salt bridges discussed

  Cell Reactions

• _- Consider the zinc/copper galvanic cell zinc is higher on the activity series so it is oxidized _{– 2+}

  Anode half–reaction: Zn (s) (aq) + 2e (oxidation)  Zn _{– 2+ -}

  Cathode half–reaction: Cu (aq) + 2e (s) (reduction)  Cu atoms from the zinc electrode lose electrons and dissolve into _{2+ - 2+} solution as Zn

  Cu ions in solution gain electrons and become neutral copper atoms

  as the cell operates, the mass of zinc electrode decreases, and the ∴ mass of copper electrode increases

The overall reaction for the zinc/copper galvanic cell is the addition of the

two half reactions

• _{– 2+}

  Anode half–reaction: Zn (s) (aq) + 2e _{(aq) (s)}  Zn _{– 2+} Cathode half–reaction: Cu + 2e . _{2+ 2+}  Cu Overall cell reaction Cu (aq) + Zn (s) (aq) + Cu (s)

   Zn ex. Write the anode, cathode, and overall cell reactions that occur when each pair of half–reactions is combined to form a galvanic cell #1 – A copper strip is a solution of copper (I) nitrate, Cu(NO ^{3 ) 2 (aq) , and a} _- tin strip in a solution of tin (II) chloride, SnCl 2 (aq) tin is higher on the activity series so it is oxidized

• _{– 2+} Anode half–reaction: Sn (s) (aq) + 2e

   Sn _{– 2+} Cathode half–reaction: Cu (aq) + 2e (s) . _{2+ 2+}  Cu Overall cell reaction Cu (aq) + Sn (s) (aq) + Cu (s)

   Sn #2 – An aluminum strip is a solution of aluminum nitrate, Al(NO ^{3 ) 3 (aq) , and} a silver strip in a solution of silver nitrate, AgNO 3 (aq)

• aluminum is higher on the activity series so it is oxidized _{– 3+}

  Anode half–reaction: Al (s) (aq) + 3e  Al _{– +} Cathode half–reaction: Ag (aq) + 1e (s) .

   Ag

• * Balance number of electrons *
_{(s) (aq) – 3+}

  Anode half–reaction: Al + 3e _{+ –}  Al Cathode half–reaction: 3Ag (aq) + 3e (s) . _{+ 3+} 

  3Ag Overall cell reaction 3Ag (aq) + Al (s) (aq) + 3Ag (s)

   Al

• _- Purpose of the Salt Bridge

  the salt bridge provides ions to prevent charge buildup from _- occurring around the electrodes basically it will replace ions being removed from the solution to

• _- maintain a neutral solution if the salt bridge is removed ions will build up at the electrodes and
• _- prevents electrons from being transferred in the galvanic cell

  removing the salt bridge will stop the flow of electrons and has ∴ the same effect as disconnecting the wires _-

  Cell Potential (Voltage)

• _- a measure of the potential difference across the electrodes it is measured in volts ex. 9V battery has 9V of difference between _-

  the electrodes

  as the galvanic cell operates, the potential difference gradually decreases until it becomes 0, meaning the electrons are no longer moving and the cell is ‘dead’

  Homework p.397 a,b, p.400 #1-7

  Batteries

  Many products today require mobile, long-lasting, and lightweight power

• _- Battery

  a battery is a group of two or more galvanic cells connected in _- series to each other _- many different types and sizes available now _- becoming more specialized never touch a leaking battery, they are very corrosive to your skin

  How they are made:

• _- Outside outside casing protects us from harmful chemicals and provides information about the type of battery it is and the voltage it can _-

  produce both the positive and negative terminals of the battery are

• _- exposed the size of the battery determines the amount of chemicals it can

  hold and how long it will last ex. D batteries will last longer than AA batteries

• _- Inside 3 basic parts, two electrodes (an anode and a cathode), and an _-

  electrolyte paste in the center is a brass pin surrounded by an anode of powdered

• _- zinc the cathode chemicals of manganese dioxide MnO ^{2 and carbon} _-

  surround the anode a thin fabric separates the two and the electrolyte paste allows ions to move back and forth between the two

  Cell Reactions

  Two types of batteries

• – –

  Cathode half–reaction: 2 MnO ^{2 + H 2 O 2 + 2e}  Mn ^{2 O 3 + 2 OH} _{– –} Anode half–reaction: Zn + 2 OH ^{2 O + 2e}

   ZnO + H

  Nickel-Cadmium Cell

• _{– –}

  Cathode half–reaction: 2 NiOOH + 2 H ^{2 O + 2e 2 + 2 OH} _{– –}  2 Ni(OH) ₂ Anode half–reaction: Cd + 2 OH + 2e  Cd(OH) _{- -} electrons only moving when switch is on _- batteries can last up to 5 years and still work electrons from oxidation of zinc travel through pin, out negative terminal of battery, through electrical device and back into positive _- terminal of battery if larger voltages are required the batteries can be connected in series in increase their overall voltage ex. 1.5V + 1.5V = 3V

  Primary and Secondary Cells

  Primary cells – cannot be recharged, generally short life span ex. all alkaline dry cell batteries, lithium cells in cell phones Secondary cells – can be recharged, difficult to dispose of because metals are very toxic ex. nickel-cadmium cells

• a list of other primary/secondary cells is included on p.406-407 of the text * Homework p.408 #1–3,5–7

  Corrosion

• _- Corrosion
• without protection from corrosion, most metals are oxidized by _-

  their environment some metals protect themselves by forming a protective layer of corroded metal on the surface but the metal below remains intact ex. the copper roof on the parliament forms a layer of corroded copper that is green in colour, as opposed to the normal reddish brown colour of copper. This layer of oxidized copper protects the other layers of copper below and can last well over 50 years zinc also forms a protective layer, which explains why galvanized (meaning zinc covered) substances, like fences last much longer than iron fences aluminum is another metal that forms a protective layer, Al ^{2 O 3 , one} of the hardest compounds known which gives aluminum products incredibly long lifespans

  Rusting

• _- the reddish-brown flaky material that is produced when iron- _-

  containing metals, like steel corrode _- rust is a mixture of oxides and hydroxides rust does not stick to the layer of metal below it and flakes off

• _- exposing the layer below it rust is a redox reaction where oxidation and reduction are happening on the same surface of metal
• the anode is located where the metal has been dented or _-

  scratched the cathode can be almost anywhere

• the presence of moisture (water) and oxygen help iron releases 2 _-

  electrons rusting is just like a galvanic cell

  Factors that Affect Rusting

• _- Moisture corrosion cannot occur without water _- a relative humidity of at least 40% is required for corrosion to take place _-

  Electrolytes _- salts do not cause corrosion but they help they provide a conductive solution to move the electrons and act like a salt bridge to prevent the buildup of ions at the cathode and anode

• _- Contact with Less Reactive Metals corrosion can occur when 2 different metals come in contact with _-

  each other more reactive metals (higher on activity series) lose electrons to

• _- less reactive metals which starts the corrosion process this is why it is important to make sure the proper nails, rivets, or screws are used _-

  Mechanic Stresses bending, shaping, or cutting metals can introduce stresses in the metal that could become corrosion sites

  Homework p. 416 #1-6

• _- Preventing Corrosion
• other times it can be disastrous ex. oil pipeline, bridge, or artificial bone replacement

  Methods to Prevent Corrosion

• _- Protective Coatings _-

  cover the metal with a rust-inhibiting paint

• _- simplest, and one of the cheapest ways to prevent corrosion some problems are that it only works for aboveground structures and the whole surface has to be covered completely or a chip or scratch in the coating could be a location for corrosion to begin _-

  and could extend under the paint without even knowing

• _- galvanize the metal by coating it with zinc

  this can be done by dipping the steel in molten zinc or by _- electroplating (will look at tomorrow) the coating is very resistant to corrosion

  Corrosion-Resistant Metals

• _- removing undesirable impurities and adding corrosion-resistant elements that usually makes the metal stronger _- these mixtures of metals are called alloys and have specific _-

  desirable properties like strength, and unreactivity by making alloys with aluminum and magnesium, they can form protective oxide layers and be corrosion-resistant

• _- Cathodic Protection

  forces the metal object you wish to protect into the cathode of a _- corrosion cell 2 ways

• _- Sacrificial Anode involves covering the surface of a metal with another one that is more likely to corrode, like zinc, so even if the metal becomes _-

  scratched, it is only the zinc that corrodes not the iron underneath electrons are sent to the iron and it remains a cathode and does not corrode

• gas stations do this by connecting their gas tanks to magnesium blocks _- Impressed Current involves pumping electrons into the metal object to turn it into a _-

  cathode, this is done by connecting it to a power source most of the Canadian oil and natural gas pipeline network is protected this way p.425 #1-4,6,7