Ionic Bonding (2) Dot and cross diagrams/Lewis structures

Ionic Bonding (2) Representing ionic bonds with dot-cross diagrams (Lewis structures)

So here we are again trying to give a few words of wisdom about ionic bonds.

Let’s start by being controversial and say that the ionic bond does not exist in its raw form. 

Unlike say hydrogen gas where we can see a discrete covalent bond of two electrons perfectly shared between two atoms you don’t get that sort of thing with ionic bonding.

Ionic bonds exist in massive clusters called structures.

An ionic bond never occurs in isolation.

Rather its better to talk about moles of ionic bonds.

Which reminds me to mention in the blogs I’m going to write for the next few weeks that Mole Day will be upon us soon.

Nevertheless you are going to find the ionic bond spoken about pretty frequently. 

You’ll have been taught that its formed by the donation of electrons from a metal atom to a non metal atom. 

Here’s a typical diagram illustrating this process which you can find everywhere on the Internet:

Here’s your first diagram that shows why you shouldn’t believe the Internet ever.

All is going fine with this diagram until we get to the second line “to form a molecule of sodium chloride” 

No, 1000 times no!!!! someone has to speak out here at this …p.

I suggest you try that line (about the molecule!) on your chemistry tutor and see what their reaction is!!

I would annihilate you actually if it were me. 

And here is another dodgy diagram:

I just don’t like the red splodge labelled ionic bond.

It smacks too much of a covalent bond picture like an σ bonding molecular orbital.

 Let’s leave this stuff off our Lewis Structures and dot and cross diagrams.

Although this next diagram doesn’t give enough information at least it gives us the accompanying structure of the ionic compound formed. 

But let’s have the sodium labelled as an “Ion” and same with the chloride it’s also an “ion”.

The colours are immaterial.

There’s no colour at the atomic level. 

But the structure is correct for sodium chloride as its has 6:6 coordination. 

What that means is that around each sodium ion are 6 chloride ions and vice versa. 

So you don’t just get one positive ion attached to one negative ion.

The formula of an ionic compound just gives you the overall ratio of the numbers of positive to negative ions. 

The stoichiometric ratio…..

That’s 1:1 in the case of sodium chloride hence NaCl.

And as the ratio is 1:1, there are no subscript stoichiometric numbers required. 

What you find with these diagrams is that they don’t show ionic bonds but the process of how they are formed. 

So you have the transfer of outer shell electrons from the Group 1, 2 or 3 element (usually a metal) to the outer shell of the Group 5, 6 or 7 element in order to fill the outer shell with sufficient electrons. 

Here is a classic diagram you have probably seen more times than sliced bread. 

Now the reason I like this diagram is not for the pink and pale blue blobs that are supposed to represent electrons. 

Nor the distinguishing dots and crosses.  Very sad!!

Nor the  correct numbers of electrons in each shell of the atoms and the ions, good one that!!

But I like the brackets (or parentheses for some of you) that encase the ion arrangements with the appropriate charge outside the bracket. 

I also like the electron in the chloride in outer shell keeping its original colour and cross because that’s what examiners like to see. 

I also like the statement about 1 electron transferred.

There’s none of this nonsense about atoms “wanting” electrons to fill their outer shells.

Atoms are inanimate objects without personality or personal traits so they can’t want anything.  Its crazy!!!

You really need to watch your use of language in this topic. 

But best of all is the comment in the middle of the diagram between the ions.

Here the attraction between ions is labelled correctly with a word not well known or remembered by students of chemistry at school or college level.

The attraction between ions is termed electrostatic attraction.  Yes!!!

Now, we like this!!

We can also see that each ion has a Noble gas electron configuration once it has formed. 

Sodium that of Neon and the chloride ion that of Argon.

So finally, let me challenge you to draw up a few of these dot and cross diagrams or Lewis structures to illustrate the ionic bonding in some simple binary ionic compounds like magnesium chloride, potassium oxide and lithium fluoride.

Some humorous diagrams from the Internet with the odd sexist joke thrown in.  

More on Ionic bonding in the next blog especially on the Born Haber Cycle.

Ionic Bonding (1) Evidence for the existence of Ions

Ionic Bonding (1) Evidence for the existence of Ions

In this blog, I want to show you the evidence that exists to suggest that ions are real particles.

You might have carried out an experiment like this in your laboratory:—

•Take a microscope slide and cover with wet filter paper.

•At the mid point place a crystal of potassium manganate(VII).

•Connect each of the short sides of the slide to a 20v dc power source and wait.

You can see a similar experiment done here

You can find instructions and further discussion of the manganate(VII) experiment here. Or here and a good example of the experiment from the Royal Society of Chemistry here.

Here is a picture (its fairly stylised) of the kind of result you might expect:

If the purple colour moves towards only one of the electrodes (the positive electrode) what does that tell us about the particles that make up the purple colour?

Yes, that's right, they could only be negatively charged.

And we are pretty sure potassium particles are colourless given how many colourless salts there are of potassium.   

Here is another experiment you can do fairly easily though it takes time to develop and see the results:—

•Place a paste of copper(II)chromate in the bottom of a U tube.

•Top up each arm of the U tube with conducting liquid and insert charcoal electrodes

•Set up a 20v dc potential difference across the electrodes and watch.

After 20mins or so you might see movement of colour blue to the negative terminal and yellow to the positive terminal.

These diagrams illustrate the results of this demonstration:

Here’s further evidence, as if we needed it, to show that copper ions (blue) are positively charged species and that chromate(VI) ions (yellow coloured) are negatively charged species.

One final piece of evidence to suggest that charges atoms called ions do really exist is the phenomenon of the electrolysis of molten salts like lead bromide, PbBr₂

You can find details of the experiment and how to carry it out here

And you can see the experiment itself on video here

Here is a diagram of what we think is happening in the electrolysis of lead bromide:

What I like about this picture is that it is not too fussy and crowded with detail.

Lead is produced at the negative cathode suggesting that lead ions carry a positive charge.

Bromine gas (highly toxic) is evolved at the anode suggesting bromide ions carry a negative charge,

The magnified diagram on the left shows that we think bromide ions lose electrons to the anode and the resulting atoms pair up to form bromine molecules:

2Br^—     =      Br₂(g) +     2e^–

On the right, the magnified diagram shows positive lead ions each accepting two electrons.

Pb²⁺      +    2e^–      =   Pb(l)

The evidence for the existence of ions is pretty convincing.

The word ION (he derived it from the Greek for “going”) was first conceived by Michael Faraday in 1834 in his early experiments with electrolysis in the 19^th century. 

Cations (positive ions) Faraday so named because they go to the negative cathode and anions (negatively charged) because they go to the positive anode.

This is a brief introduction into the existence of ions but there is much more to say about how they enable the formation of compounds and all that will be found in forthcoming blogs. 

Best now go and do an ion experiment and see them move for yourself. 

Happy experimenting!!

Ionization Energy (5) Orbitals and the Pauli Exclusion Principle

Ionization energy (5) Orbitals and the Pauli Exclusion Principle.

Before I finish posting on ionisation energy and related areas I thought I ought to discuss atomic orbitals. 

What have atomic orbitals to do with ionisation energy?

Let’s go back to the way we can represent the electron configuration of an atom using the electrons in boxes approach.

Here is how sodium’s electron configuration is represented:

The question we want to ask is what do the boxes represent and why are the three boxes in the 2p subshell labelled p_x, p_y and p_z?

The box represents an atomic orbital.

Atomic orbitals contain a maximum of two electrons. 

These electrons spin in opposite directions.

The Pauli exclusion principle dictates that each electron in an atom has to have an exclusive quantum number i.e. a unique identifier.

The attribution of opposite spins facilitates that unique identity.

For example each 2s electron is unique because they have opposite spins.

So how can each of the 2p electrons be unique?

The answer lies in the distribution of the electrons around the atom.

The shape of the atomic orbitals shows how the electrons are distributed around the atom.

The best place on the net to investigate atomic orbitals is here at the Orbitron.

For s orbitals that pattern of the distribution of electrons is spherical.

But p orbitals are not spherical but essentially dumbbell shaped.

As you can see the distribution is symmetrical about the 3D axis. 

Another way of thinking about the orbital is to realise it is a space in the atom where there is a 95% probability of finding the electron. 

The space is the result of calculations using the electron’s wave function ψ in the Schrodinger wave equation.

Plotting the square of the wave function ψ² against r the distance from the nucleus generates the different orbital shapes. 

If we extend our description of the shapes of atomic orbitals to the d subshell then there are 5 and each has a unique position in space around the nucleus. 

What these models of electron subshells do is to de bunk the idea that atoms are spherical in shape. 

Studies of atomic orbitals are important when we discuss the bonding in organic molecules (see here) and transition metal complexes. 

You can see more about Schodinger’s Cat here