17/08/13 Week 6 Sunday 2013 Home

The Beginning of the EEI
The EEI report will have the following layout:

Abstract
Topic Report
Introduction
Research Question
Discussion
Conclusion
Practical
Title
Aim
Hypothesis
Equipment
Method
Results
Analysis of Results
Conculsion
References
Appendix

The topic report will consist of all background information on tin can corrosion and galvanic cells with reference to the variables pH and surface area. The Research Questions and Hypothesis have already been decided on and created in an earlier post, but I will include them again for reading conveniences. Below you will find the Aim, Method and Equipment with a diagram of the setup of the Fe and Sn galvanic cell.

Research Question

What effect will the pH of the solution within a galvanic cell and surface area of metals have on the rate of corrosion?

Hypothesis
It was hypothesized that as the pH decreased and the surface area increased within a Fe and Sn galvanic cell then the corrosion rate of the Fe would increase when the amount of solution, type of container, multimeter, temperature and type of joint are kept constant.

Aim
To investigate the effect that pH of the electrolyte solution and surface area of electrode have on the corrosion rate of Fe in a Fe and Sn galvanic cell when type of container, amp meter, temperature and type of joint are kept constant. 

Equipment

• Tin Metal
• Iron Nails
• 6x 250mL Beaker
• 1M HCL
• Tin Chloride in solution (SnCl)
• Iron Chloride in solution (FeCl)
• Multimeter
• Sodium Chloride (NaCl)
• Paper Towel
• Safety Glasses
• Gloves
• Lab Coat
• Blu-Tack

Method
The method will appear at a later date as the electrolyte solutions will need to be mixed up.

Below is a diagram of the setup of the Galvanic Cell I have made on Word.

       

17/08/13 Week 6 Saturday 2013 Home

Background Research - pH dependence of corrosion
Normal galvanic corrosion occurs more rapidly in acidic solutions than in alkaline ones(taken from Chemistry in Use Book 2 p375).

Link
Standard electrode potential: http://en.wikipedia.org/wiki/Standard_electrode_potential_(data_page)

Example:

In galvanic corrosion the half reaction of O_2(g) + 2H₂O_(l) + 4e­^-  → 4OH^-_(aq) has a standard electrode potential of 0.40V where the value is [OH^-] = 1.00mol/L. If this half reaction occurs in a neutral solution of pH of 7 ([OH^-] = 1.0 x 10^-7 mol/L) the non standard electrode potential changes to 0.81V and at pH = 4.00 a further 0.99V. This information suggests that electrode potential increases as acidity increases. If  this is the situation then this half reaction has a greater tendency to occur if the pH decreases (becomes more acidic).

8/08/13 Week 5 Thursday 2013 Period 2

Background Research - Surface area's affect on galvanic corrosion

When dealing with surface area as a variable in galvanic corrosion the area ratio of the anode:cathode is an important variable that affects the corrosion rate pertaining to the anode. This area ratio is also important when cosidering the amount of current available that arises from the cathodic reation.

 

The image below taken from the Corrosion-Club.com demonstrates the affect surface area has on galvanic corrosion:

The diagram suggests that the lower the ratio of anode:cathode the higher the corrosion rate at the anode is likely to occur. When the same situation's ratio is flipped the higher the ratio of anode:cathode results in little influence to help corrode.  If the cathode is larger than the anode, this allows for more oxygen reduction and other cathodi reaction which as whole results in a greater galvanic current. t would beinteresting to find imformation on what results would occur if the ratio was 1:1.

 

When putting this information into the context of tin can corrosion the Fe is considered the anode as it is the metal oxidising. Therefore is the ratio of Sn is higher than Fe then the corrosion rate should increase. This information presented here assists in proving the validity of the hypothesis made in the first post.

 

Link to valuable resource:

http://corrosion-doctors.org/Corrosion-Forms/area-effects.htm

 

 

4/08/13 Week 4 Sunday 2013 Home 2

Background Research - Corrosion
Corrosion is the degradation of metal. In this EEI the corrosion of tin cans will be looked at, in more detail the galvanic corrosion of Fe in a Fe and Sn galvanic cell.

The most common from of rust that we encounter is Iron rust. This is a result of the Fe metal oxidising as seen in these half reactions both for Iron 1 and Iron 2.

Fe_(s) → Fe²⁺_(aq) + 2e^-

Fe²⁺→ Fe³⁺ + e^-

The second half reaction containing Iron 2 is because the  Fe³⁺ ion is the ion in Fe₂O₃ and Fe₂O₃ is the hydrated iron oxide with H₂O.

Therefore if the more easily the metal is oxidised the higher its corrosion rate will be.

Facts about rusting
(taken from the Chemistry in Use Book 2  p372)

1. Both oxygen and water are necessary for rust to form.

2. Salt water accelerates rusting.

3. Impure iron rusts more rapidly that pure iron.
4. Iron rusts more rapidly when attached to a less reactive metal such as Cu or Sn when on its own (the EEI being conducted involves Fe in Contact with Sn)

5. Rust occurs most readily where iron is under mechanical stress-at bends in sheets, points of nails, sharp edges of knives and razor blades, and around bolts and rivets under tension.

4/08/13 Week 4 Sunday 2013 Home

Background Research - Galvanic Cells continued.

The salt bridge is important component in galvanic cells. As I talked about in the last diagram with Cu and Ag electrodes in solution, each container containing an electrode is called a half cell. The salt bridge is not only there to link the half cells, however balance out the change of electrons.

Cu(s) → Cu²⁺_(aq) + 2e^-_­

Ag⁺_(aq) + e^-_­ → Ag_(s)

If these two reactions were the only ones to occur in this galvanic cell, then the container containing Cu and NO₃ would end up with an excessive amount on positive ions due to the Cu ions being produced by the electrode reaction (metal to ion). Also the container containing Ag and NO₃ would end end up with an excessive amount of negative ions due to the electrode reaction (ion to metal). As a whole this migration and balance of electrons preserves neutrality in both half cells.

In summary the electrode process that releases electrons in called the anode and the electrode process absorbs electrons is called the cathode. The migration of negative electrons transfers from the cathode to the anode and the migration of positive ions transfers from the anode to the cathode.

A galvanic cell is an electron pump. It pumps electrons from the negative terminal into the external circuit and the back into the positive terminal due to the redox reaction occurring within the cell.

Anode: The electrode at which oxidation occurs

Cathode: The electrode at which reduction occurs.

1/08/13 Week 4 Thursday 2013 Home

Background Research - Galvanic Cells

A galvanic cell is a device in which a chemical reaction occurs in such a way that it generates energy. The conductors of the cell, metals in the activity series, are called the electrodes; they are connected to the external circuit. For a galvanic cell to function in must be in an electrolyte solution. The electrolyte is the substance which in solution conducts electricity and is connect by a salt bridge.

In the diagram below the copper nitrate, silver nitrate and potassium nitrate solutions are electrolyte solutions. The salt bridge in this diagram is a U tube filled with KNO3. The salt bridge in a galvanic cell is used to make electrical contact between the two solutions and therefore must contain a conducting substance; in this case KNO3. The electrodes in the diagram below are Cu (Cu electrode) and Ag (Ag electrode) and give off two half reactions which result in the electron transfer that gives the electrical current; this electricity again being produced by a chemical reaction from the electrodes and electrolytes. The electrons flow out of the Cu electrode and into the voltmeter (external circuit) then into the Ag electrode. The half reactions are as followed:

Cu(s) → Cu²⁺_(aq) + 2e^-_­

Ag⁺_(aq) + e^-_­ → Ag_(s)

In these half reaction the Cu is losing electrons, therefore oxidising and the Ag is gaining electrons, therefore reducing. When combining these two half reaction, you get the full redox reaction of:

Cu_(s) + 2Ag⁺_(aq)→ Cu²⁺_(aq) + 2Ag_(s)

In the diagram above the reason the Cu is oxidising to Cu ions and Ag ions are reducing to Ag metal is due to their position on the activity series for metals.

The activity series of metals is a listing of metals in order from left to right of decreasing reactivity. This means that metals to the left are more likely to oxidise or decrease and give up electrons to the metals on the right. Below is the activity series of metals.

K    Na    Li    Ba    Ca    Mg    Al    Zn    Fe    Sn    Pb    (H)    Cu    Ag    Pt   Au