Redox Reactions Test

Introduction

For today's practical session, the aim was to find out if the different substances were oxidising or reducing agents. For this test, we used:

• Aqueous potassium manganate (VII) --> Used to test for reducing agent
• Aqueous potassium iodide --> Used to test for oxidising agent
• Dilute sulfuric acid
• Solids P and Q, and solutions R and S

Procedure

1. Prepare solutions of substances P and Q by transferring half a spatula of each substance in separate test tubes and dissolving half a test tube of distilled water. Divide the solutions of P and Q into two portions to be used for Test 1 and 2 respectively.
   
2. For Test 1, add 1cm^3 of dilute sulfuric acid to each solution. Then add KMnO4 solution.
• Test 1 is testing for the presence of reducing agents. A positive test will involve a change in colour of solution from purple to colourless.
3. For Test 2, add 1cm^3 of dilute sulfuric acid to each solution. Then add KI solution.
• Test 3 is testing for the presence of oxidising agents. A positive test will involve a change in colour of solution from colourless to brown.
4. Record your observations in a table as shown below.

Observations

Solution S, Test 2: Brown solution (left) from original colourless solution (right)

From left to right: Solution P test 1 and 2, Solution R test 1 and 2

Close up of dark brown solution when KI is added to solution P.

Close up of purple solution when KMnO4 was added to solution P.

It is important to first add your sulfuric acid, otherwise your results would look something like that.

Conclusions

P is an oxidising agent based on the results from test 2.

Q is a reducing agent based on the results from test 1.

R is unreactive based on the results from both tests.

S is disproportionated based on the results from both tests.

Test for Anions

The tests for anions are relatively more complicated as compared to those for the cations. But nonetheless, let's attempt to tackle them all in an organised manner.

If We're Classifying by Reactants Added...

• Point to note: To differentiate between carbonates and sulfates, it is important that you add barium nitrate or lead (II) nitrate first. Why? 
• The problem with carbonates is that they form precipitate with any acid added. But carbonates don't react with nitrates. That's why you should add barium or lead (II) nitrate first, because these would produce precipitates BaSO4 or PBSO4 which are insoluble.

If We're Classifying by Reactions...

So yeah, the most memorable traits of each anion:

• Carbonate — Insoluble with any acid added, CO2 gas produced
• Sulfate — White ppt with barium or lead (II) nitrate; soluble in nitric acid
• Chloride —  White ppt with Pb or Ag; forms white crystals when heated then cooled
• Iodide — Yellow ppt ftw
• Nitrate — Devarda's alloy, heating, pungent NH3 gas produced

Tests for Cations

To understand tests for cations, we must first know: What are cations? Cations are basically positively charged ions, if you still don't know what I'm talking about, check out my post on ions here.

To test for cations, you usually use aqueous sodium hydroxide or aqueous ammonia. To identify the cation, you look at three things:

• The colour of the precipitate formed  
• The solubility in excess alkali (base)
• Reactivity to a flame test (only if required for further identification)

COLOUR OF PRECIPITATE

When NaOH is combined with cations, they form metal hydroxides. Soluble metal hydroxides are all those in Group I and ammonia. Meanwhile, any other metal hydroxide is insoluble. Insoluble means? THEY'LL FORM A PRECIPITATE. <-- Great.

Adding NaOH to any salt solutions containing transition metals usually form coloured solutions, while salt solutions with other metal ions(except Group I and ammonia) form white precipitate.

For ammonia, however, since it has a lower concentration of hydroxide ions, slightly soluble hydroxides such as calcium and barium hydroxides are not precipitated. Can't remember your solubility table? Here it is:

SOLUBILITY IN EXCESS ALKALI/BASE

So there is this trio of metal hydroxides that are fortunately or unfortunately complicated. Aluminium (II) Oxide, Zinc (II) Oxide and Lead (II) Oxide are amphoteric in nature, meaning they react with both acids and bases. When this happens, they form complex salts that are soluble to give colourless solutions.

So what happens exactly when you add excess alkali (base)?

Acid-base neutralisation takes place to form salt and water. That's why when you add excess base, the white precipitate you saw would dissolve and form a colourless solution. 

For ammonia, only two metal ions form precipitates which then dissolve to form soluble complexes when excess ammonia is added: Zinc and Copper. Copper would form a deep blue solution whilst zinc would form a colourless solution.

REACTIVITY TO FLAME TEST

After all those troublesome things you did, you still have more?? Yeah.

Fortunately or unfortunately, this isn't covered in the syllabus, so this'll be for another post another day...To sum up, here is everything in a pretty table:

Atomic Structure: Ions

All atoms are neutral, but when they have a higher or lower number of electrons compared to protons, they become CHARGED IONS. These are divided into two groups:

Cations — Positively charged ions
Anions — Negatively charged ions

How to remember which is which? Just remember that ANIONS are A-Negative-ION. Wallah, so genius, I know.

TRENDS...

• Elements in group 1 to 4, all their atoms tend to lose an electron to complete their shells. Meanwhile, Elements in group 5 to 7, all their atoms tend to gain an electron.
• All metals are cations. All non-metals are Anions.

To help remember their charge, we can always go back to that beautiful valency table we had in the last post!

Atomic Structure: Valency

So remember in the last post when we talked about ELECTRON SHELLS? Remember that they could have many many many shells, and yet a maximum number of electrons in each shell? If you don't it's okay, we'll recap.

So basically, the way electrons are arranged in an atom is also called electronic configuration. If you want to represent it visually, it'd look like this:

If you want it in orbital notation it'd look like this:

I would continue with spdf and all but that's for another day. ANYWAYS...

The circles you see with the x's on them? That's called a shell. And those x's are electrons. So the first shell can hold up to 2 electrons and the second and subsequent ones can hold up to 8 electrons. Now, now, here's the complicated part.

A VALENCE ELECTRON is an electron that lies on the outermost shell of an atom. The outermost shell is also called the "valence shell". So what is this complicated "valency"?

Valency --> No. of electrons needed to be gained or lost to complete the shell

"Complete the shell?" Simi lanjiao. 

"Completing the shell" means making the shell is full. So for example, iodine, has seven valence electrons yes? To complete the shell, you need to gain another electron, to ensure you have the maximum number of electrons in the shell (8).

Here's the complicated part. Sometimes an ion will also lose electrons to complete the shell. In the case of lithium for example, a lithium ion has only one valence electron. So what does it do to complete the shell? It loses that electron. 

Why does it do that? Because losing one electron is easier than gaining seven more. (informal answers you must not use in exam) BUT YES, VALENCY IS COMPLICATED. So here's a little valency table for you:

Things to note:

Down the group, elements all have the same number of valence electrons. 

However, number of valence shells increase down the group.

Across the period, number of valence electrons increase

Why does all this happen? Moving on to electron affinity...

Atomic Structure: Parts of an Atom

So before we delve into the deep and confusing processes, we need to know the parts of an atom, also known as SUB-ATOMIC PARTICLES.

Okay so parts of the atom...

THE NUCLEUS
This is made up of protons and neutrons.

THE ELECTRON SHELL
The electron is just a rough gauge for where the electrons move. The inner most shell can carry up to 2 electrons, while the ones further away can carry up to 8 electrons. And you can have many many many electron shells, but we'll get to that later.

ALL THE LITTLE 'ON'S
Protons — Determine the identity of the atom (like on the periodic table)
Electrons — Determine the charge of the atom (the less you have, the more positive the charge. Vice versa)
Neutrons — Determine the mass of the atom

---

Okay, so maybe I'm not very good at describing and you might still be wondering:

So let's put a little link to the periodic table. So on the periodic table they usually give you a little bit of information on each element. For example, carbon here, on this lovely image I found on Google.

So like the diagram says, the number above the letter, that is the atomic number, AKA proton number. This is, as you might've guessed, the number of protons in that atom. And if you haven't noticed, this is how elements in the periodic table are arranged: By proton number.

The number below that is the nucleon number or atomic mass. You might wonder, why is this the mass?? Remember that ATOMS ARE ELECTRICALLY NEUTRAL. Sooooo... in every atom, there will be an equal number of protons and electrons, cancelling each other out. So what's the only thing that actually has mass? That's right, your good old buddy neutrons.

If an atom does not have an equal amount of protons and electrons, it will be an ion, which we'll be covering in the next post!

Acids and Bases: Concentration

What is concentration?

Concentration is the amount of substance dissolved in a fixed volume of solution. In other words it's: 

Amount of substance / Volume of solution

Main Points:

• The strength of the acid is relative to the % dissociated

That's why even if the acid has higher strength, if the concentration is higher, it can have a lower pH. This is the formula for pH by the way:

Anything with square brackets mean it's "concentration of" apparently.

• The concentration of H+ ions determines the pH

There are always H+ ions and OH- ions in every acid. The weaker the acid, the higher the amount of OH- ions and the lower the amount of H+ ions. Therefore, your H+ ions are relative to your acidity and the OH- ions to your basicity. 

If you still don't get it it's fine. Just remember that the higher the concentration of H+ ions, the lower the pH. 

Other points:

• pH is only dependant on the concentration of H+ ions.
• pH is non-linear