unexkae peverel waidah

[Orestes Hardy]

In this chapter learners will explore the concept of a covalent bond in greater detail. In grade ten learners learnt about the three types of chemical bond (ionic, covalent and metallic). A great video to introduce this topic is: Veritasium chemical bonding song. In this chapter the focus is on the covalent bond. A short breakdown of the topics in this chapter follows.

As revision you can ask learners to draw Lewis diagrams for the first \(\text20\) elements and give the electronic structure (this was covered in grade \(\text10\)). This then leads into thinking how the elements can share electrons in a bond. Learners should recognise that there are unpaired electrons in atoms that can be shared to form the bonds.

It is important to note that when drawing Lewis diagrams, we first place single electrons around the central atom and only once four electrons have been placed, do we pair electrons up. This will avoid the need to explain hybridisation. It is also important for learners to realise that the placement of electrons is arbitrary and the electrons can be placed anywhere around the atom.

This part of the chapter is interleaved with electron structure and Lewis diagrams as these two concepts play a key role in understanding why hydrogen is diatomic and helium is monatomic. In this part learners are introduced to the idea that when two atoms come close together there is a change in the potential energy. This forms a strong foundation for explaining the energy changes that occur in chemical reactions and will be seen again in chapter \(\text12\) (energy changes in chemical reactions).

Four cases are looked at to try to understand why bonds form. This is all about the covalent bond, so all the examples you use must be of covalent molecules (and you must also only pick examples of covalent molecular structures as covalent network structures are more like ionic networks and do not form simple molecular units). It is also important to help learners realise that a lone pair of electrons is very much dependant on the molecule that they are looking at. Lone pairs of electrons can be used under special circumstances to form dative (or coordinate) covalent bonds.

You can build the different molecular shapes before starting to teach VSEPR from large polystyrene balls and kebab sticks or you can give your learners jellytots or marshmallows and toothpicks and get them to build the molecular shapes. Remember that the shapes with lone pairs need more space for the lone pairs and so it is not as simple as just removing the toothpick for the lone pair.

This topic covers the shapes that molecules have. This is only the shapes of covalent molecular compounds, covalent network structures, ionic compounds and metals have very different three dimensional forms. This topic is important to help learners determine polarity of molecules. Two approaches are used to determine the shape of a molecule. The first one looks at molecules matching up to a general formula while the second one considers how many electron pairs are around a central atom. These two approaches can be used together to help learners fully understand this topic.

It is important to note that CAPs does not give a definitive source for electronegativity values. You should use the ones found on the periodic table in the matric exams (these are the same values as the ones on the periodic table at the front of this book). Learners should be aware that they may see different values on other periodic tables. Learners must not think of the different types of bonding as being exactly defined. Also, the values for where the types of bonding transition are not exact and different sources quote different cut-off points.

The simplest examples of polarity are the ideal shapes with all the end atoms the same and so you should stick to this in your explanation. You can explain this for trigonal planar molecules by using your learners. Get three girls or three boys to link hands (they all put their right hand into the centre and hold the other two learners right hands). Then they try to pull away (all learners pull equally). This is the even sharing of electrons. Now replace a girl with a boy (or vice versa) and tell the new learner to pull a bit less. This shows the uneven sharing of electrons.

In this final part of the chapter we return to our energy diagram and add two pieces of information: bond energy and bond length. The bond length is the distance between the two atoms when they are at their minimum energy, while the bond energy is this minimum energy. The bond energy comes up again in chapter \(\text12\) (energy and chemical change) when the topic of exothermic and endothermic reactions is covered.

Coloured text has been used as a tool to highlight different parts of Lewis diagrams. Ensure that learners understand that the coloured text does not mean there is anything special about that part of the diagram, this is simply a teaching tool to help them identify the important aspects of the diagram, in particular the unpaired electrons.

We live in a world that is made up of many complex compounds. All around us we see evidence of chemical bonding from the chair you are sitting on, to the book you are holding, to the air you are breathing. Imagine if all the elements on the periodic table did not form bonds but rather remained on their own. Our world would be pretty boring with only \(\text100\) or so elements to use.

Imagine you were painting a picture and wanted to show the colours around you. The only paints you have are red, green, yellow, blue, white and black. Yet you are able to make pink, purple, orange and many other colours by mixing these paints. In the same way, the elements can be thought of as natures paint box. The elements can be joined together in many different ways to make new compounds and so create the world around you.

In Grade \(\text10\) we started exploring chemical bonding. In this chapter we will go on to explain more about chemical bonding and why chemical bonding occurs. We looked at the three types of bonding: covalent, ionic and metallic. In this chapter we will focus mainly on covalent bonding and on the molecules that form as a result of covalent bonding.

As we begin this section, it's important to remember that what we will go on to discuss is a model of bonding, that is based on a particular model of the atom. You will remember from the discussion on atoms (in Grade \(\text10\)) that a model is a representation of what is happening in reality. In the model of the atom that you are learnt in Grade \(\text10\), the atom is made up of a central nucleus, surrounded by electrons that are arranged in fixed energy levels (sometimes called shells). Within each energy level, electrons move in orbitals of different shapes. The electrons in the outermost energy level of an atom are called the valence electrons. This model of the atom is useful in trying to understand how different types of bonding take place between atoms.

There are two cases that we need to consider when two atoms come close together. The first case is where the two atoms come close together and form a bond. The second case is where the two atoms come close together but do not form a bond. We will use hydrogen as an example of the first case and helium as an example of the second case.

Let's start by imagining that there are two hydrogen atoms approaching one another. As they move closer together, there are three forces that act on the atoms at the same time. These forces are described below:

Let us imagine that we have fixed the one atom and we will move the other atom closer to the first atom. As we move the second hydrogen atom closer to the first (from point A to point X) the energy of the system decreases. Attractive forces dominate this part of the interaction. As the second atom approaches the first one and gets closer to point X, more energy is needed to pull the atoms apart. This gives a negative potential energy.

For hydrogen the energy at point X is low enough that the two atoms stay together and do not break apart again. This is why when we draw the Lewis diagram for a hydrogen molecule we draw two hydrogen atoms next to each other with an electron pair between them.

Now if we look at helium we see that each helium atom has a filled outer energy level. Looking at Figure 3.6 we find that the energy minimum for two helium atoms is very close to zero. This means that the two atoms can come together and move apart very easily and never actually stick together.

For helium the energy minimum at point X is not low enough that the two atoms stay together and so they move apart again. This is why when we draw the Lewis diagram for helium we draw one helium atom on its own. There is no bond.

Now that we understand a bit more about bonding we need to refresh the concept of Lewis diagrams that you learnt about in Grade \(\text10\). With the knowledge of why atoms bond and the knowledge of how to draw Lewis diagrams we will have all the tools that we need to try to predict which atoms will bond and what shape the molecule will be.

In grade \(\text10\) we learnt how to write the electronic structure for any element. For drawing Lewis diagrams the one that you should be familiar with is the spectroscopic notation. For example the electron configuration of chlorine in spectroscopic notation is: \(1\texts^22\texts^22\textp^5\). Or if we use the condensed form: \([\textHe]2\texts^22\textp^5\). The condensed spectroscopic notation quickly shows you the valence electrons for the element.

Using the number of valence electrons we can easily draw Lewis diagrams for any element. In Grade \(\text10\) you learnt how to draw Lewis diagrams. We will refresh the concepts here as they will aid us in our discussion of bonding.

A Lewis diagram uses dots or crosses to represent the valence electrons on different atoms. The chemical symbol of the element is used to represent the nucleus and the core electrons of the atom.

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