Lecture 6 – Chemical Bonding and Electron Movements

So, in the last lecture, we learned about electron configurations and periodic table trends. We discussed a bit of reactivity. Now, let’s talk about atoms form compounds.

Elements can form compounds by interacting with other elements. They bond with each other to form ionic compounds or molecules. We will discuss how this is achieved in this lecture. It will be another heavy lecture, so be prepared.

Also, remember, we have a handy periodic table popup on the bottom right of the screen! An additional polyatomic ion table is included in the popup as well.

Types of Chemical Bonding

Why do we want to learn about chemical bonding? Chemical bonding is the building block of molecules and ionic compounds. The bonding energy and bond length affect chemical reactivity. And, in terms of electrons, molecules have molecular geometry, which positions valence electrons in a specific way. How these electrons are positioned can affect reactivity as well.

There are 3 types of bonding.

• Ionic bonding
• Covalent bonding
• Metallic bonding

We will discuss covalent bonding the most because it has most relevance in organic chemistry and biochemistry.

Ionic bonding and covalent bonding can be explained by electrons being donated and accepted or shared, with electron donation and acceptance being a more ionic characteristic while electron sharing being more covalent.

One of the factors influencing whether the bond is ionic or covalent is electronegativity differences .

• High leads to ionic bond formation.
• Low leads to covalent bond formation.

This is because, with a high , one element is so much better at retaining the electrons it has that it can just grab electrons from the other element, forming ions. The resulting ions then attract through electrostatic forces (opposite charges attract), forming ionic bonds.

With a low , the elements aren’t good at taking each other’s electrons, so they are content with simply sharing those electrons. The shared electrons then become a covalent bond.

So, if we have a moderate , we have this interesting situation that, despite having shared electrons, those electrons can be a bit biased towards the atom with a higher value. This is called a polar covalent bond.

The exact value of electronegativity difference that determines ionic or covalent bonds differs throughout sources. Some say is enough to warrant a designation of ionic bonds, while others ask for a higher value such as . The exact number is not important, but we will still provide it as a reference.

• Electronegativity differences is one of the many factors that influence bond formation. Other factors include chemical and physical properties. However, for most cases, just looking at is sufficient to predict bond formation.
  • If , it generally forms an ionic bond.
  • If , it generally forms a pure covalent bond.
  • If is in between, it generally forms a polar covalent bond.
• Again, there are exceptions to the boundaries, so don’t just plug-and-chug.

Metallic bonds, on the other hand, occurs mostly in metals, like sodium, iron, etc, though metallic bonds aren’t the only bond type metals exhibit. We’ll talk about it when we get there.

Ionic Bonding

Let’s discuss ionic bonding. Ionic bonds are formed between ions. More specifically, they are formed between cations (positively charged) and anions (negatively charged).

Ionic bonding is a type of chemical bonding that involves the electrostatic attraction between oppositely charged ions, or between two atoms with sharply different electronegativities.

An ionic bond is likely to form if the difference in electronegativity is larger than 1.8 between the elements.

Cation and anion formation is explained in the previous lecture.

• Alkali metal ions mostly have a +1 charge.
• Alkaline earth metal ions mostly have a +2 charge.
• Chalcogen ions mostly have a -2 charge.
• Halides have mostly a -1 charge.
• Transition metals vary.

Ionic Bond Formation

Let’s look at some examples. Blue denotes potential cations, while green denotes potential anions.

A typical feature of ionic bonds is that the species form into ionic crystals, in which no ion is specifically paired with any single other ion in a specific directional bond. Rather, each species of ion is surrounded by ions of the opposite charge, and the spacing between it and each of the oppositely charged ions near it is the same for all surrounding atoms of the same type. The chemical formula for ionic compounds exists for stoichiometric reasons rather than structural ones. Molecules formed from covalent bonds (including polar covalent), on the other hand, have specific structures that can be reflected from chemical formulae.

Other ionic compounds can be formed from metal cations and polyatomic anions (bottom right popup). Same thing can be said for their ionic bonding structures.

Ionic Compounds in Medicine

There are a lot of medicine that are ionic compounds. Let’s go through some examples. While we’re going through those examples, we should get familiar with the etymology of medical conditions.

How can calcium, a highly reactive metal, be replaced by lead, a relatively unreactive metal compared to calcium? Reactivity does not solely depend on electron configuration.

Firstly, both elements are ions, so we need to take that into account. Secondly, reactivity depends on a myriad of factors. In biochemistry and other related fields, reactivity also depends on:

• The reaction environment (concentrations of solvents, presence of interesting molecules, pH values, etc.)
• The state of products (are they more favored over the reactants by being structurally stable?)
• The presence of inhibitors (are there other molecules that can stop or slow the reaction down?)
• The presence of catalysts (are there other molecules that can speed the reaction up?)
• etc. etc.

We will discuss those topics in future units and courses. For now, it is important to be aware that:

Reactivity depends on a lot of factors, and electron configuration is just one of many.

“Exceptions” of Ionic Bonding

There are some exceptions to the electronegativity difference values we provided, because electronegativity alone does not completely (albeit accurately, for the most time) determine bond type. Here are some common exceptions.

• has but is covalently bonded.
• has but is ionically bonded.

In summary, multiple factors can determine bond type, not exclusively electronegativity differences.

With ionic bonds explained, let’s turn our attention to covalent bonds.

Covalent Bonding

In this section, we will discuss covalent bonds and polar covalent bonds. so we can officially understand why there are lines between elements.

As we discussed in the introduction, covalent bond utilizes a more friendly electron-sharing model rather than the “donation and acceptance” model. So, what does sharing electron mean in a chemistry sense?

A covalent bond is a chemical bond that involves the sharing of electrons from both elements to form electron pairs between atoms. These electron pairs are known as shared pairs or bonding pairs. When electrons are shared, we can count the shared electrons as belonging to both elements.

• An important point about electron sharing is that: both atoms need to share electrons for covalent bonds to work.
• When 2 electrons are shared (1 from each side), it forms 1 electron pair.

Covalent bonds can be predicted by electronegativity differences , though other factors also affect bond type.

, the bond is likely purely covalent.
, the bond is likely polar covalent.

Also, covalent is basically “co-valent” – pertaining a characteristic of sharing valence electrons.

Sharing Electrons between Similar Elements

An electron can be shared if it occupies in a half-filled orbital.

• If an orbital of an element is not filled, that orbital cannot participate in electron sharing because there is nothing to share with.
• If an orbital of an element is fully filled (paired electrons in an orbital), those electrons cannot participate in electron sharing because they are already paired, and oversharing, similar to real life, can be overbearing, which is not stable.

Let’s look at the following examples.

Now that we understand which electrons can be shared and which cannot, let’s talk about how they actually share their electrons. Introducing the sharing diagram (formally and more comprehensively the molecular orbital diagram, or MO diagram). The sharing diagram essentially combines atomic orbital diagrams and more or less systematically describes how electrons are shared.

Observe the following in the sharing diagram:

• We basically take the sharable valence electrons from their respective atomic orbital diagram and fill them into sharing (bonding) orbitals.
• If somehow unsharable valence electrons are involved, we fill them into the oversharing (antibonding) orbitals.
• For each pair of valence electron orbitals, there are 2 orbitals stemming from it.
  • Bonding orbital – essentially electron sharing (stabilizing effect).
  • Antibonding orbital – essentially oversharing (destabilizing effect).
• The same 3 rules that apply to atomic orbitals also apply to sharing orbitals.
  • The energy rule (Aufbau)
  • The exclusion rule (Pauli)
  • The single occupancy rule (Hund)

Using this diagram, we can calculate bond order.

Covalent Bond Formation

So, in truth, our version of molecular orbitals is not the most accurate or comprehensive. Our version is highly simplified, but it is more than enough to explain covalent bonds.

So, now that we know how electrons are shared, we can understand how covalent bonds are formed. Here are the steps.

• Step 1: Draw out the atomic orbital diagram of the atoms. We only need to draw the orbitals for valence electrons, but it’s never a bad thing to draw the entire thing.

• Step 2: For each pair of valence electron orbitals, create the 2 orbitals:
  • The lower-energy bonding (sharing) orbital.
  • The higher-energy antibonding (oversharing) orbital.

• Step 3: Fill the orbitals from the low energy levels to high energy levels (energy rule). If an electron pair forms, they should have opposite spin (exclusion rule). If the orbitals have the same energy level, the single occupancy rule still stands.

Let’s go through a lot of examples.

We know that noble gases are very unreactive. As a result, we know that shouldn’t exist. From an electron configuration standpoint, already has a fully filled shell – there is no point sharing electrons. What about from a sharing orbital standpoint? To summarize, overshares, so bond formation is not feasible.

However, the helium hydride ion , consisting of ion and , is feasible because there is no oversharing.

Here’s a fun thing. According to our sharing orbital diagram, diatomic halogen molecules () aren’t the only molecules we can theoretically produce. It is theoretically possible to create molecules with 2 different halogen atoms because our sharing orbital diagram determines that covalent bonding with bond order 1 can be formed between 2 halogen atoms – it didn’t specify the halogen atoms need to be the same.

In real life, interhalogens do exist. An interhalogen compound is a molecule which contains two or more different halogen atoms. For our example, we can have interhalogens of the following:

It isn’t much considering how many molecules there are (and the examples we’ve given). So, let’s expand our repertoire.

Sharing Electrons between Different Elements

Let’s mix it up even more. Previously, we only mixed and matched elements of similar valence electron configuration (as evident by looking at the number of sharable electrons). Now, let’s mix and match elements of different valence electron configurations.

The easiest examples involve hydrogen and halogens.

Lecture Content (Part 2)

With our minds refreshed, let’s look at the atomic orbitals again.

Visualizing Orbital Sharing

Before we start talking about atomic orbitals, here’s a general rule.

If a double or triple bond is formed, the electrons that contribute to the 2nd and 3rd bond cannot be in the s-orbital.

It’s a weird rule, but we will get there when we start visualizing. We can see it in examples such as . If we look at their sharing orbital diagrams, the bonds always come from the p-orbital electrons. Single bond molecules such as , on the other hand, can have electrons from s- and p-orbitals that contribute in sharing.

Sigma Bonding

Sigma bonding is a covalent bond that features the strongest bonding. In a way, it’s like the electrons are hugging each other. For example, all single-bonded molecules are formed from sigma bonds, like . Let’s look at the visualization.

As we can see, sigma bonds can form between and orbitals, and orbitals, and and orbitals. Other orbital combinations like and , and , and and orbitals can also yield a sigma bond.

Pi Bonding

Suppose we have a double bond and triple bond. The first bond is almost always a sigma bond. The second (and, if applicable, third) bond will be pi bonds. It’s called “pi” because we need at least 2 p-subshells (a.k.a, ) to obtain pi bonds.

Pi bonds can be found in molecules like . It is weaker than sigma bonds. Rather than electron orbitals “hugging” each other to merge into one molecular orbital, pi bonds are like electron orbitals “shaking hands” to merge into one molecular orbital.

As we can see, pi bonds can be formed between and orbitals. Other orbital combinations like and , and and orbitals can yield pi bonds.

Note: the visualization above is not accurate. However, it provides a good visualization of pi bonds. are actually hybridized, as we will discuss soon.

Delta Bonding

Delta bonding happens when we encounter quadruple bonds, which, in this course, will never happen. So, we’ll just summarize it in a paragraph or two.

Delta bonding is called “delta” because we need at least 2 d-subshells (a.k.a ) to obtain delta bonds. That’s the extent we will talk about delta bonds.

So, quick summary:

If we look at (a triple bond), the 1st bond is sigma, and the 2nd and 3rd bonds are pi.

Using this newfound knowledge, we can update our sharing orbital diagram.

Hybridization

Now that we have a clear picture of what orbital sharing looks like, let’s talk about hybridization.

Hybridization is the merging of electron orbitals of different energy levels in the same shell to create a new kind of hybridized orbitals, moving electrons in the process. The creation of those hybridized orbitals allow the electron configurations to change, making the element have more sharable electrons.

Let’s break this down into steps.

• Step 1: An electron in a paired orbital (basically an unsharable electron) moves to another empty orbital closest to it but with a higher energy (because the lower energy orbitals are already occupied). Because the shell has higher energy, this movement requires energy.
  • Think about 2s to 2p, 3s to 3p, etc.

• Step 2: Once the electron is moved, it causes a mixing of orbitals. There are different types of mixing depending on how we want to bonds to form, but after the mixing, we yield hybridized orbitals. This stabilizes the atom so much that it offsets the energy required to move the atom, making this process feasible.
  • Think about s orbitals mixing with p orbitals to form some sp hybridized orbitals.
  • Thinking about s, p, and d orbitals mixing to form some spd hybridized orbitals.

• Step 3: The shapes of the orbitals change.

In a way, orbital hybridization is a means to make previously unsharable electrons sharable. More sharable electrons means more possible bonding.

That’s a lot of words, so let’s look at the best example there is – carbon. Carbon can hybridize its orbitals in many different ways.

Hybridizing s and p Orbitals

Carbon, according its atomic orbital, can only share 2 electrons. But carbon is a very versatile element, forming molecules like , etc. This versatility can be attributed to its ability to hybridize.

Since there is only s and p orbitals at , carbon undergoes hybridization among s and p orbitals.

• What if we hybridize 1 s orbital and 3 p orbitals?
• What if we hybridize 1 s orbital and 2 p orbitals?
• What if we hybridize 1 s orbital and 1 p orbital?

As we can observe from the animations above, hybridization allows more electrons to be sharable.

• Carbon used to have only 2 sharable electrons. Now, it has 4.
• Depending on the hybridization, we are presented with different bonding possibilities.
• Geometry is determined by electron repulsion – given the orbitals, they want to be as far away as possible because electrons repel each other.

This opens up a lot more possibilities, because hybridization also happens with nitrogen and oxygen. Considering that carbon, nitrogen, and oxygen are key elements in biology, their versatility in bonding is expected.

Let’s look at some examples. Here’s a quick reference of some atoms and how many electrons they can share:

• 1 sharable electron:
• 2 sharable electrons: (rarely but possible)
• 3 sharable electrons:
• 4 sharable electrons: due to hybridization

As we look at the correlation between the structure and sharable electrons, we can conclude that:

• Atoms with 1 sharable electron (like ) can only participate in 1 covalent bond.
• Atoms with 2 sharable electrons (like ) can participate in 2 covalent bonds.
• Atoms with 3 sharable electrons (like ) can participate in 3 covalent bonds.
• Atoms with 4 sharable electrons (like ) can participate in 4 covalent bonds.
• We will encounter exceptions, but this serves as a general guideline on how we can draw chemical structures.

The conclusions above can be abbreviated into “HONC 1234”, a mnemonic device for remembering covalent bonding for those elements.

Hybridizing s, p, and d Orbitals

For elements like in period 3, we notice that they can potentially hybridize to the d orbitals because at . So, for the sake of understanding molecules like , we’ll generously assume that can hybridize to the 3d orbitals.

Note: this assumption is technically incorrect when it comes to molecular orbitals, but it still provides us with enough insight that we will still use it to explain molecular geometry. We will not ask questions regarding this assumption except for its molecular geometry.

To reflect this inaccuracy, we will add quotation marks around inaccurate descriptions (yet helpful) of these atoms.

So, p-block elements in period 3 can “hybridize” among s, p, and d orbitals. Here are some examples.

Metallic Bonding

All this covalent bonding theories flying everywhere is a bit confusing. Let’s reorient ourselves. We’re talking about chemical bonding and how electrons behave under different types of bonding.

• For ionic bonds, we have valence electron donating and accepting.
• For covalent bonds, we have valence electron sharing, which opens up a huge can of worms.
• For metallic bonds, valence electrons are essentially free from the confines of the nucleus.

We won’t be talking about metallic bonding as much because metals in medicine usually involve ionic bonding (as demonstrated in the examples) and covalent bonding with other nonmetals. Some pure metals can be used as equipment (like prosthetics) because of their lack of reactivity with the body.

This phenomenon, where electrons can freely move around within a substance, is called electron delocalization. For example, metallic bonds express delocalization because electrons are free to move around. Ionic bonds, on the other hand, do not express delocalization because it follows a donating-accepting model. Covalent bonds become more complicated, and we will talk about it next lecture.

Electron delocalization means that electrons in a molecule, ion, or solid metal are not associated with a single atom or a covalent bond.

• In a pure metallic substance (like a block of sodium), all of the valence electrons are not associated with a single atom because they are freely roaming, thus expressing delocalization.
• In a molecule (covalently bonded), some electrons can be shared among multiple atoms (as we will see in the next lecture), thus expressing delocalization. In fact, some polyatomic ions do express some electron delocalization.