Monday, 27 February 2017

GCSE OCR Gateway C4:2a Tests for gases: identifying the products of chemical reactions

GCSE OCR Gateway C4:2a Tests for gases: identifying the products of chemical reactions
Learning Objective:

C4.2a To be able to describe tests to identify selected gases
oxygen, hydrogen, carbon dioxide and chlorine

Test for Oxygen (O2)

This is one of I think the easier GCSE chemistry topics.  I would hope that you have been aware of tests for simple common gases for a couple of years before you get to GCSE stage.

I’d be disappointed of you had not already carried out these tests before you come to GCSE work.

So I’m expecting that you know pretty well the test for oxygen gas.

Oxygen gas relights a glowing spill because it accelerates combustion reactions.

What you might not be so certain of is a reaction at school chemistry level that produces oxygen gas because there are not that many.

One is the catalytic decomposition of hydrogen peroxide (H2O2(aq))solution.

2H2O2          2H2O    +    O2

The catalysts here are many and varied such as blood, potato, celery and MnO2 (and other transition metal oxides) to name four.

And another reaction, one incidentally you ought not to try seeing as it can have explosive consequences, is the effect of heat on potassium manganate(VII) containing a trace of manganese (IV) oxide catalyst.

2KMnO4              K2O    +   2MnO2    +  1½O2

Finally there is the classic demonstration called “Elephnt’s Toothpaste”!!

Just add finely ground potassium iodide crystals to a mix of 10ml 100vol hydrogen peroxide and fairy liquid in a very large (say 500 -1000ml) measuring cylinder and stand back!!

Plunge a glowing spill into the resultant foam and see that it is an oxygen foam.

Here is the reaction:

2KI   +   2H2O2             I2    +   ½O2   +  2KOH  +  H2O

The iodine stains the foam brown.

Tests for Hydrogen (H2)

Hydrogen is another of those common gases you ought to have met long before you start a GCSE course.

Its explosive properties are exploited in the test since it will ignite with a lighted spill and produce what we all call a squeaky pop!!

So in this picture, we have what looks like magnesium ribbon in hydrochloric acid releasing hydrogen to be tested.

Mg   +   2HCl        MgCl2    +    H2(g)

Follwed by the test for the hydrogen

2H2    +    O2          2H2O   (+  pop!!  )

When you do this test look for the product, water vapour, on the sides of the test tube and look too for bands in the vapour.  I think the bands in the water vapour are due to the standing wave from the noise of the squeaky pop.

Test for Carbon dioxide (CO2)

Here is yet another gas whose test you ought to be familiar with because you should have been testing for this gas in Biology classes and Chemistry classes from your earliest days in your school.

Carbon dioxide (CO2) turns limewater (calcium hydroxide solution: Ca(OH)2) cloudy or chalky.

Ca(OH)2    +    CO2         CaCO3     +   H2O

Best not to say it turns limewater milky since calcium carbonate (CaCO3) is not milk even if it might look like it in the test tube.

The chalkiness is a calcium carbonate precipitate.

Carbon dioxide is formed in many reactions at GCSE level such heating some carbonates or reacting them with an acid.

ZnCO3         ZnO    +   CO2

ZnCO3      +    2HCl        ZnCl2     +   H2O +    CO2

Test for Chlorine (Cl2)

Now here is a gas that you may not have come across before.  Nor might you have come across its test.

Do you know how to test for chlorine?

Chlorine is poisonous and coloured green but those properties won’t do as tests.

But chlorine will bleach damp indicator paper.  Bleach means chlorine will turn the coloured paper white.

There is no equation for this test.

And chlorine is not usually made in chemical reactions in the Lab.

In my next post, I’m going to discuss the tests for cations and anions.

Saturday, 25 February 2017

Redox (II): Standard Electrode Potential Eo(6) The electrochemical series

Edexcel A level Chemistry (2017)
Topic 14: Redox (II): Standard Electrode Potential Eo(6)

Here is a further learning objective:

14/13 To know that standard electrode potentials can be listed as an electrochemical series

Electrode potentials and the electrochemical series

So what do you think this list is really telling you?

You will see lists like this is data books and examination questions but unless you are aware of the significance of the values then much of the meaning of electrode potential will pass you by.

So I thought I would just take this opportunity to re–emphasise the meaning of the values and use this learning objective to lay more stress on you understanding the significance of these values.

In fact, I can say that I know from my own personal experience of trying to get my head round the significance of electrode potentials and lists of them that it is not easy, at least it was not for me.

Here is a typical list:

Best oxidant

E° V
Best reducing agent
Li+   +  e    Li
K+   +  e    K
Ba2+   +  2e    Ba
Ca2+   +  2e    Ca
Na+   +  e    Na
Mg2+   +  2e    Mg
Al3+   +  3e    Al
Mn2+   +  2e    Mn
2H2O +  2e   2OH   +   H2(g)
Zn2+   +  2e    Zn
Cr2+   +  2e    Cr
Fe2+   +  2e    Fe
Cr3+   +  3e    Cr
Cd2+   +  2e    Cd
Co2+   +  2e    Co
Ni2+   +  2e    Ni
Sn2+   +  2e    Sn
Pb2+   +  2e    Pb
Fe3+   +  3e    Fe
2H+   +  2e    H2(g)
S + 2H+  +2e     H2S (g)
Sn4+   +  2e    Sn2+
Cu2+   +  e    Cu+
SO42—  +   4H+  +  2e     SO2 (g)  +   2H2O
Cu2+   +   2e    Cu
2H2O +  O2  + 4e     4OH
Cu+   +  e    Cu
I2  +2e     2I
O2 + 2H+ + 2e   H2O2
Fe3+   +  e    Fe2+
NO3  +   2H+  +  e     NO2   +   H2O
Hg2+   +  2e    Hg(l)
Ag+   +  e    Ag
NO3  +   4H+  +  3e     NO   +   2H2O
Br2  +2e     2Br
O2 + 4H+ + 4e   2H2O
MnO2  +  4H+ +  2e    Mn2+  +  2H2O
Cr2O72— + 14H+ + 6e   2Cr3+ + 7H2O
Cl2  +2e     2Cl
Au3+   +  3e    Au
MnO4  +  8H+ +  5e    Mn2+  +  4H2O
Co3+   +  e    Co2+
F2  +2e     2F

We call the list an electrochemical series.

But what does it all really mean?

First, notice how the redox equilibria are written.

Each one is written as a gain of electrons, a reduction

So for example the equation:

Li+ (aq) +  e      Li (s)

represents the half–cell:

Li+(aq) | Li(s)

And the electrode potential was measured when the half–cell was connected to the standard hydrogen electrode; the reference electrode. Here is the cell diagram:

Pt | [H2 (g) ,2H+(aq)] || Li+(aq) | Li(s)

So this is how things are set up to measure these standard electrode potential values

but what does —3.04v mean?

It means this: of all the species on the right hand side of the list, lithium (Li) metal at the top is the best reducing agent or reductant. Or you can say it is most easily oxidized.

Similarly, fluoride ions (F) near the bottom of the list with an electrode potential of +2.87v are the poorest reductants.  Those ions are least easily oxidized.  Fluoride ions do not give up the electron that completes the fluorine outer shell at all easily, as we know.

Now you can look at the other left hand species on the list and make similar statements about them too.

So for example at Eo = +1.52v manganate (VII) ions (MnO4) are very good oxidizing agents, they easily gain electrons and are reduced to manganese (II) ions (Mn2+) in the presence of dilute acid.

So looking at that list again, the best oxidizing agents (oxidants) are at the bottom right.  And you can see there there is Fluorine, Manganate (VII) ions and Chlorine—all good oxidizing agents. 

You can see I hope the implications of this layout of half–cell reactions.

If a reaction is going to be feasible then only top right species (e.g. Lithium (Li), Sodium (Na) etc.) will reduce bottom left species such as manganate (VII) ions (MnO4) or cobalt (III) ions (Co3+).

This conclusion is vital to understand if you are going to progress using electrode potential values as I have posted earlier here.

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