Exploring Electrostatics: Part 1

In the series ‘exploring electrostatics’ I want to delve into the interesting workings behind the electrostatic relationships found in different sorts of bonding. I will start at the very beginning and then explore more complicated chemistry (and physics)

The world around us consists of small particles that form materials we have the wonderful pleasure of investigating. Some are metals, others have covalent bonds and make up molecules. Others form ionic lattices between metals and non metals… There are all kinds of interesting forces working between particles that hold our physical world in place — but what forces particles to ‘stick’ to each other?

Chemistry is the science that investigates the relationships between valence electrons, that is, an atom’s electrons on the outermost shell.

These valence electrons are the ‘currency’ of chemical reactions: atoms exchange valence electrons or share them in covalent bonds.

This means that we can differentiate between three main types of bonding, which are dependent on how valence electrons of two species interact:

covalent bonds what is this?

ionic bonds

metallic bonds what is this?

Of which the latter is the one we will explore in this post.

All of these three types of bonding rely on the (theoretically) simple principle of electrostatic attraction/repulsion. What is that you may ask?

In an ionic bond, a metallic element binds to a non metallic element. Metallic elements (e.g.: Na, Mg, Fe) often have few valence electrons. These electrons aren’t attracted strongly by the nucleus.

This means that they prefer giving off an electron that gaining an additional one during a chemical reaction. We say that they have a low electronegativity. Non metallic elements (e.g.: O, H, S, C) have almost entirely filled shells, meaning they are desperate to take in an additional electron in order to fill their shell. They have a high electronegativity. Which brings me to my next point.

Travelling Electrons

Electronegativity is essentially just a term which describes how eager an atom is to gain electrons. If we glance at the periodic table, we can observe that nonmetals have a high electrongeativity whereas metals typically have a lower electronegativity. This comes from the number of valence electrons they have.

In an ionic relationship, a stark contrast between high and low electronegativity, in other words, a contrast between metals and nonmetals, is leveraged to build an ionic lattice.

One species has few valence electrons and prefers to lose them whilst the other is desperate to gain electrons. When combined, one species loses all its electrons whilst the other fills its shell.

Now, two ions are left. Let’s take the example in Fig. 2:

If we mix sodium (Na) and chlorine (Cl) in a beaker. The sodium loses it’s electron when surrounded by chlorine (Cl) atoms. This process is described by:

Na → Na⁺ + e⁻

This forms a positive Na ion (a Na cation) which is expressed as Na⁺. Now, because the chlorine and sodium are reacting in the same beaker, chlorine gains the electron that is lost by Na. This is described by:

Cl + e⁻ → Cl^–

This forms a negative Cl ion (a Cl anion).

We now have a beaker with two ions that have opposite charges. Just as electrons are attracted by protons, negative charges are attracted to positive ones (and vice versa). This means that the two ions will attract each other.

Ionic lattices

Up until now we’ve only looked at what happens to two ions of opposite charges. However, in real life we’re not just adding two ions together and observing them: what happens if several millions of ions of opposite charges meet?

The answer is simple: the ions arrange in the only possible way they can to be surrounded by ions of opposite charges and as far away as possible from ions of the same charge: they form an ionic lattice.

An ionic lattice is a gigantic arrangement of oppositely charged species. It looks like this:

Energy

When talking about ionic bonding and lattices, a fundamental part to understand is the energetic aspects involved in these processes. Imagine energy as the electron ‘currency’ I described before. It is needed to make many chemical reactions happen including the ionization of atoms.

Ionization energy refers to the energy needed to remove an electron from the outermost shell of its atom. For an element X this means the energy required for the reaction to the ion X⁺:

X → X⁺ + e^–

Let’s take a look at the element magnesium (Mg). Magnesium has two electrons on its outermost shell:

We can define E₁ as the energy needed to remove one electron from its shell.

Mg → Mg⁺ + e^–

However, now it still has one electron left on its outer shell:

If we want to go ahead and remove that last electron, we need a second ionization energy E₂:

Mg⁺ → Mg²⁺ + e^–

This goes on for however many times you’d like. The only important principle to respect is the following:

E₁ < E₂ < E₃ <…. < E_n

Imagine an atom with two shells. Once you have removed the first electron from the outer shell, the second electron you want to remove is attracted even more by the positive nucleus. As you work your way around the atom and reach the last electron on any given shell, the ‘energy jump’ to go from the outer shell to the inner one is so immense that it will take you an extreme amount of energy to separate the last electron from its shell. (However atomic electron transitions are a topic entirely by themselves)

FAQ

With this understanding of ionic bonding you are ready to hear all about the specifics behind these bonds in my series ‘exploring electrostatics’. If this peaked your interest, you might be curious to understand how scientists solved a fundamental biochemical problem and won a Nobel prize for it.

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