draw the lewis structure for the molecule ch2chch3. how many sigma and pi bonds does it contain?

Affiliate three: Atomic combinations

In this affiliate learners will explore the concept of a covalent bond in greater detail. In form ten learners learnt well-nigh the iii types of chemic bond (ionic, covalent and metallic). A bully video to introduce this topic is: Veritasium chemical bonding song. In this chapter the focus is on the covalent bail. A brusque breakdown of the topics in this chapter follows.

• Electron structure and Lewis diagrams (from grade \(\text{x}\))

  As revision yous can ask learners to depict Lewis diagrams for the commencement \(\text{twenty}\) elements and give the electronic structure (this was covered in grade \(\text{10}\)). This so leads into thinking how the elements can share electrons in a bond. Learners should recognise that at that place are unpaired electrons in atoms that tin can be shared to form the bonds.

  It is important to note that when cartoon Lewis diagrams, we outset place single electrons effectually the cardinal atom and only once four electrons take been placed, practice nosotros pair electrons up. This will avoid the demand 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 effectually the atom.

• Why hydrogen is a diatomic molecule only helium is a monatomic molecule

  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 office learners are introduced to the thought that when ii atoms come close together at that place is a change in the potential energy. This forms a strong foundation for explaining the energy changes that occur in chemical reactions and will exist seen once more in chapter \(\text{12}\) (energy changes in chemic reactions).

• Deducing simple rules about bail formation (and cartoon Lewis diagrams for these molecules)

  Four cases are looked at to effort to sympathise 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 but pick examples of covalent molecular structures as covalent network structures are more like ionic networks and do not form simple molecular units). It is besides important to aid learners realise that a solitary pair of electrons is very much dependant on the molecule that they are looking at. Lone pairs of electrons can be used nether special circumstances to class dative (or coordinate) covalent bonds.

• The basic principles of VSEPR and predicting molecular shape

  You tin 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. Call back that the shapes with lone pairs need more space for the lone pairs and so it is not every bit simple as just removing the toothpick for the lone pair.

  This topic covers the shapes that molecules have. This is simply the shapes of covalent molecular compounds, covalent network structures, ionic compounds and metals take very different 3 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 ane looks at molecules matching up to a general formula while the second ane considers how many electron pairs are around a central atom. These two approaches tin exist used together to assistance learners fully understand this topic.

• Electronegativity and polarity of bonds

  Information technology is important to note that CAPs does not requite a definitive source for electronegativity values. You should use the ones establish on the periodic table in the matric exams (these are the same values as the ones on the periodic tabular array at the front of this book). Learners should be aware that they may see different values on other periodic tables. Learners must not call up of the different types of bonding every bit beingness exactly defined. Also, the values for where the types of bonding transition are not verbal and unlike sources quote different cutting-off points.

  The simplest examples of polarity are the ideal shapes with all the stop 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 3 boys to link hands (they all put their right manus into the heart and hold the other two learners right hands). And then they endeavour to pull abroad (all learners pull every bit). This is the fifty-fifty sharing of electrons. Now replace a daughter with a boy (or vice versa) and tell the new learner to pull a chip less. This shows the uneven sharing of electrons.

• Bond length and bond energy

  In this final part of the chapter we render to our energy diagram and add two pieces of information: bond energy and bond length. The bond length is the altitude between the two atoms when they are at their minimum free energy, while the bond energy is this minimum energy. The bond energy comes upwards again in chapter \(\text{12}\) (energy and chemical modify) when the topic of exothermic and endothermic reactions is covered.

Coloured text has been used equally a tool to highlight unlike parts of Lewis diagrams. Ensure that learners understand that the coloured text does not mean there is annihilation special well-nigh that part of the diagram, this is just a educational activity tool to assistance them place the of import aspects of the diagram, in particular the unpaired electrons.

Nosotros alive in a globe that is made upwards of many complex compounds. All effectually u.s. nosotros see evidence of chemic bonding from the chair you lot are sitting on, to the book yous are holding, to the air you lot are breathing. Imagine if all the elements on the periodic table did not grade bonds but rather remained on their own. Our world would exist pretty irksome with merely \(\text{100}\) or so elements to utilise.

Imagine y'all were painting a picture and wanted to evidence the colours effectually you. The simply paints you have are crimson, green, yellow, blue, white and black. Yet you are able to brand pinkish, purple, orange and many other colours by mixing these paints. In the same way, the elements can be thought of equally natures paint box. The elements tin exist joined together in many different ways to make new compounds and then create the world effectually you.

In Course \(\text{10}\) nosotros started exploring chemical bonding. In this chapter we will become on to explicate more than about chemical bonding and why chemical bonding occurs. Nosotros looked at the three types of bonding: covalent, ionic and metallic. In this chapter we volition focus mainly on covalent bonding and on the molecules that form as a effect of covalent bonding.

In this chapter we will employ the term molecule to hateful a covalent molecular structure. This is a covalent chemical compound that interacts and exists as a single entity.

3.1 Chemic bonds (ESBM4)

Why do atoms bond? (ESBM5)

As we brainstorm this section, information technology'due south important to recollect that what we will go on to discuss is a model of bonding, that is based on a particular model of the atom. You lot will remember from the discussion on atoms (in Grade \(\text{10}\)) that a model is a representation of what is happening in reality. In the model of the atom that you are learnt in Grade \(\text{x}\), the atom is made up of a central nucleus, surrounded by electrons that are arranged in fixed free 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 cantlet are called the valence electrons. This model of the atom is useful in trying to empathize how different types of bonding accept place between atoms.

A model takes what we see in the world effectually us and uses that to make certain predictions about what we cannot see.

Figure 3.ane: Electron organization of a fluorine cantlet. The black electrons (small circles on the inner ring) are the cadre electrons and the white electrons (small circles on the outer ring) are the valence electrons.

The post-obit points were made in these earlier discussions on electrons and free energy levels:

• Electrons ever try to occupy the lowest possible energy level.
• The noble gases take a full valence electron orbital. For example neon has the post-obit electronic configuration: \(i\text{s}^{2}2\text{s}^{2}2\text{p}^{6}\). The second energy level is the outermost (valence) beat out and is full.
• Atoms form bonds to try to achieve the same electron configuration every bit the noble gases.
• Atoms with a full valence electron orbital are less reactive.

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Energy and bonding (ESBM6)

There are two cases that we demand to consider when two atoms come close together. The first case is where the 2 atoms come close together and course a bond. The 2d instance is where the two atoms come shut together but do non form a bail. We will use hydrogen every bit an case of the beginning case and helium as an instance of the second case.

Case 1: A bond forms

Let's first past imagining that there are two hydrogen atoms approaching 1 another. As they move closer together, there are three forces that act on the atoms at the same fourth dimension. These forces are described below:

1. repulsive force betwixt the electrons of the atoms, since like charges repel

   

   Figure three.ii: Repulsion betwixt electrons

2. attractive force between the nucleus of 1 atom and the electrons of another

   

   Figure 3.three: Attraction between electrons and protons.

3. repulsive strength betwixt the two positively-charged nuclei

   

   Figure 3.4: Repulsion betwixt protons

These three forces work together when two atoms come close together. Every bit the total force experienced by the atoms changes, the amount of free energy in the system besides changes.

Now look at Figure 3.five to understand the energy changes that take place when the 2 atoms move towards each other.

Effigy 3.five: Graph showing the alter in energy that takes identify as 2 hydrogen atoms move closer together.

Allow us imagine that nosotros have fixed the one atom and nosotros will movement the other atom closer to the first cantlet. Every bit we move the second hydrogen atom closer to the outset (from point A to signal X) the energy of the system decreases. Bonny forces boss this part of the interaction. As the 2nd atom approaches the first ane and gets closer to point X, more free energy is needed to pull the atoms apart. This gives a negative potential energy.

At betoken X, the attractive and repulsive forces acting on the 2 hydrogen atoms are balanced. The energy of the organisation is at a minimum.

Further to the left of point Ten, the repulsive forces are stronger than the bonny forces and the free energy of the system increases.

For hydrogen the energy at point 10 is low plenty 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 side by side to each other with an electron pair between them.

We also annotation that this arrangement gives both hydrogen atoms a full outermost free energy level (through the sharing of electrons or covalent bonding).

Example two: A bond does not form

Now if we await at helium nosotros run into that each helium atom has a filled outer energy level. Looking at Figure 3.six 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 movement autonomously very hands and never actually stick together.

Figure three.6: Graph showing the modify in energy that takes place every bit 2 helium atoms move closer together.

For helium the energy minimum at point X is non low plenty that the two atoms stay together and then 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.

We also see that helium already has a full outermost energy level then no chemical compound forms.

Valence electrons and Lewis diagrams (ESBM7)

Now that we understand a bit more about bonding nosotros demand to refresh the concept of Lewis diagrams that you learnt about in Grade \(\text{10}\). With the cognition of why atoms bail and the knowledge of how to describe Lewis diagrams nosotros will accept all the tools that we need to effort to predict which atoms will bond and what shape the molecule volition exist.

In grade \(\text{x}\) we learnt how to write the electronic construction for whatsoever chemical element. For cartoon Lewis diagrams the i that you should exist familiar with is the spectroscopic annotation. For example the electron configuration of chlorine in spectroscopic notation is: \(1\text{s}^{two}2\text{southward}^{ii}2\text{p}^{5}\). Or if we apply the condensed form: \([\text{He}]2\text{due south}^{2}2\text{p}^{v}\). The condensed spectroscopic notation rapidly shows you the valence electrons for the element.

Using the number of valence electrons we can hands draw Lewis diagrams for whatsoever chemical element. In Grade \(\text{10}\) yous learnt how to draw Lewis diagrams. We will refresh the concepts here as they will aid u.s. 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 stand for the nucleus and the core electrons of the atom.

Lewis diagrams for the elements in period \(\text{2}\) are shown beneath:

Element	Grouping number	Valence electrons	Spectroscopic notation	Lewis diagram
Lithium	\(\text{ane}\)	\(\text{1}\)	\([\text{He}]2\text{due south}^{one}\)
Beryllium	\(\text{two}\)	\(\text{2}\)	\([\text{He}]2\text{due south}^{2}\)
Boron	\(\text{thirteen}\)	\(\text{iii}\)	\([\text{He}]2\text{southward}^{2}2\text{p}^{1}\)
Carbon	\(\text{14}\)	\(\text{4}\)	\([\text{He}]2\text{southward}^{2}2\text{p}^{2}\)
Nitrogen	\(\text{xv}\)	\(\text{5}\)	\([\text{He}]ii\text{due south}^{2}2\text{p}^{3}\)
Oxygen	\(\text{sixteen}\)	\(\text{6}\)	\([\text{He}]ii\text{south}^{ii}2\text{p}^{4}\)
Fluorine	\(\text{17}\)	\(\text{seven}\)	\([\text{He}]2\text{southward}^{2}2\text{p}^{five}\)
Neon	\(\text{18}\)	\(\text{8}\)	\([\text{He}]2\text{due south}^{2}2\text{p}^{6}\)

Y'all can place the unpaired electrons anywhere (tiptop, bottom, left or right). The exact ordering in a Lewis diagram does not matter.

Lewis diagrams

Textbook Exercise 3.ane

magnesium

\([\text{Ne}]3\text{south}^{two}\)

sodium

\([\text{Ne}]three\text{s}^{1}\)

chlorine

\([\text{Ne}]3\text{s}^{ii}3\text{p}^{5}\)

aluminium

\([\text{Ne}]three\text{s}^{ii}3\text{p}^{1}\)

argon

\([\text{Ne}]iii\text{s}^{two}3\text{p}^{half dozen}\)

Covalent bonds and bond germination (ESBM8)

Covalent bonding involves the sharing of electrons to form a chemical bond. The outermost orbitals of the atoms overlap then that unpaired electrons in each of the bonding atoms tin be shared. By overlapping orbitals, the outer free energy shells of all the bonding atoms are filled. The shared electrons move in the orbitals effectually both atoms. Every bit they movement, at that place is an attraction between these negatively charged electrons and the positively charged nuclei. This attractive force holds the atoms together in a covalent bond.

Covalent bond

A form of chemical bond where pairs of electrons are shared between atoms.

Covalent bonds are examples of interatomic forces.

We will look at a few unproblematic cases to deduce some rules well-nigh covalent bonds.

Remember that it is only the valence electrons that are involved in bonding, and so when diagrams are drawn to show what is happening during bonding, it is only these electrons that are shown. Dots or crosses represent electrons in different atoms.

Case 1: 2 atoms that each accept an unpaired electron

For this instance we will look at hydrogen chloride and methyl hydride.

Worked example 1: Lewis diagrams: Simple molecules

Stand for hydrogen chloride (\(\text{HCl}\)) using a Lewis diagram.

For each atom, determine the number of valence electrons in the atom, and represent these using dots and crosses.

The electron configuration of hydrogen is \(1\text{s}^{1}\) and the electron configuration for chlorine is \([\text{He}]2\text{s}^{2}2\text{p}^{5}\). The hydrogen atom has \(\text{1}\) valence electron and the chlorine atom has \(\text{7}\) valence electrons.

The Lewis diagrams for hydrogen and chlorine are:

Discover the single unpaired electron (highlighted in blue) on each atom. This does not mean this electron is different, we use highlighting hither to assist you run into the unpaired electron.

Arrange the electrons so that the outermost free energy level of each atom is total.

Hydrogen chloride is represented below.

Find how the two unpaired electrons (one from each atom) form the covalent bail.

The dot and cross in between the two atoms, represent the pair of electrons that are shared in the covalent bond. We tin can also show this bond using a single line:

Notation how we still testify the other electron pairs around chlorine.

From this we tin can conclude that whatever electron on its own will effort to pair up with another electron. So in practise atoms that have at least one unpaired electron can form bonds with any other atom that also has an unpaired electron. This is not restricted to only two atoms.

Worked instance ii: Lewis diagrams: Unproblematic molecules

Stand for marsh gas (\(\text{CH}_{4}\)) using a Lewis diagram

For each cantlet, determine the number of valence electrons in the atom, and stand for these using dots and crosses.

The electron configuration of hydrogen is \(1\text{s}^{1}\) and the electron configuration for carbon is \([\text{He}]2\text{s}^{2}ii\text{p}^{2}\). Each hydrogen atom has \(\text{1}\) valence electron and the carbon atom has \(\text{iv}\) valence electrons.

Remember that we said we tin can identify unpaired electrons at whatsoever position (superlative, lesser, left, right) effectually the elements symbol.

Adjust the electrons so that the outermost energy level of each atom is full.

The methyl hydride molecule is represented below.

Or:

Textbook Exercise 3.two

chlorine (\(\text{Cl}_{two}\))

boron trifluoride (\(\text{BF}_{3}\))

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Case 2: Atoms with lonely pairs

We will use h2o every bit an example. Water is fabricated up of one oxygen and 2 hydrogen atoms. Hydrogen has i unpaired electron. Oxygen has two unpaired electrons and two electron pairs. From what nosotros learnt in the first examples we run across that the unpaired electrons can pair up. But what happens to the two pairs? Tin can these form bonds?

Worked case 3: Lewis diagrams: Uncomplicated molecules

Represent water (\(\text{H}_{2}\text{O}\)) using a Lewis diagram

For each atom, determine the number of valence electrons in the cantlet, and correspond these using dots and crosses.

The electron configuration of hydrogen is \(i\text{due south}^{i}\) and the electron configuration for oxygen is \([\text{He}]2\text{southward}^{ii}2\text{p}^{4}\). Each hydrogen atom has \(\text{1}\) valence electron and the oxygen atom has \(\text{6}\) valence electrons.

Adapt the electrons and then that the outermost energy level of each atom is full.

The water molecule is represented below.

or

Notice how in this example we wrote a \(\text{2}\) in front of the hydrogen? Instead of writing the Lewis diagram for hydrogen twice, nosotros only write it in one case and use the \(\text{2}\) in front of it to indicate that two hydrogens are needed for each oxygen.

And now we can answer the questions that we asked earlier the worked example. We see that oxygen forms two bonds, one with each hydrogen cantlet. Oxygen however keeps its electron pairs and does not share them. Nosotros tin can generalise this to any cantlet. If an atom has an electron pair it will normally not share that electron pair.

A lone pair is an unshared electron pair. A lone pair stays on the atom that it belongs to.

A lone pair can be used to course a dative covalent bond.

In the example in a higher place the lonely pairs on oxygen are highlighted in cerise. When we draw the bonding pairs using lines it is much easier to meet the lonely pairs on oxygen.

Textbook Exercise iii.3

ammonia (\(\text{NH}_{3}\))

oxygen difluoride (\(\text{OF}_{ii}\))

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Case 3: Atoms with multiple bonds

We will utilise oxygen and hydrogen cyanide as examples.

Worked example 4: Lewis diagrams: Molecules with multiple bonds

Represent oxygen (\(\text{O}_{ii}\)) using a Lewis diagram

For each atom, determine the number of valence electrons that the atom has from its electron configuration.

The electron configuration of oxygen is \([\text{He}]2\text{s}^{2}2\text{p}^{4}\). Oxygen has \(\text{vi}\) valence electrons.

Arrange the electrons in the \(\text{O}_{2}\) molecule so that the outermost energy level in each atom is full.

The \(\text{O}_{2}\) molecule is represented below. Notice the two electron pairs between the ii oxygen atoms (highlighted in blue). Because these ii covalent bonds are between the same two atoms, this is a double bond.

or

Each oxygen cantlet uses its two unpaired electrons to form two bonds. This forms a double covalent bond (which is shown by a double line betwixt the two oxygen atoms).

Worked example 5: Lewis diagrams: Molecules with multiple bonds

Correspond hydrogen cyanide (\(\text{HCN}\)) using a Lewis diagram

For each atom, decide the number of valence electrons that the atom has from its electron configuration.

The electron configuration of hydrogen is \(1\text{s}^{i}\), the electron configuration of nitrogen is \([\text{He}]2\text{s}^{two}ii\text{p}^{iii}\) and for carbon is \([\text{He}]ii\text{due south}^{2}2\text{p}^{ii}\). Hydrogen has \(\text{1}\) valence electron, carbon has \(\text{iv}\) valence electrons and nitrogen has \(\text{5}\) valence electrons.

Arrange the electrons in the \(\text{HCN}\) molecule so that the outermost free energy level in each atom is full.

The \(\text{HCN}\) molecule is represented beneath. Discover the iii electron pairs (highlighted in cherry-red) betwixt the nitrogen and carbon atom. Considering these iii covalent bonds are betwixt the same 2 atoms, this is a triple bail.

or

As we have merely seen carbon shares i electron with hydrogen and three with nitrogen. Nitrogen keeps its electron pair and shares its three unpaired electrons with carbon.

Textbook Exercise 3.four

acetylene (\(\text{C}_{two}\text{H}_{ii}\))

formaldehyde (\(\text{CH}_{ii}\text{O}\))

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Example 4: Co-ordinate or dative covalent bonds

Dative covalent bond

This type of bail is a description of covalent bonding that occurs between two atoms in which both electrons shared in the bail come from the same atom.

A dative covalent bond is also known equally a coordinate covalent bond. Earlier we said that atoms with a pair of electrons volition normally not share that pair to course a bond. Only now nosotros volition see how an electron pair tin be used by atoms to form a covalent bond.

One example of a molecule that contains a dative covalent bond is the ammonium ion (\(\text{NH}_{4}^{+}\)) shown in the figure below. The hydrogen ion \(\text{H}^{+}\) does not contain whatever electrons, and therefore the electrons that are in the bond that forms between this ion and the nitrogen atom, come simply from the nitrogen.

Notice that the hydrogen ion is charged and that this charge is shown on the ammonium ion using square brackets and a plus sign outside the square brackets.

We can also show this as:

Notation that we do not use a line for the dative covalent bond.

Another example of this is the hydronium ion (\(\text{H}_{3}\text{O}^{+}\)).

To summarise what we take learnt:

• Whatsoever electron on its own will try to pair upward with some other electron. And then in theory atoms that take at least one unpaired electron can form bonds with whatever other atom that as well has an unpaired electron. This is non restricted to only ii atoms.

• If an cantlet has an electron pair it will ordinarily not share that pair to form a bond. This electron pair is known as a lone pair.

• If an atom has more one unpaired electron it can form multiple bonds to another atom. In this way double and triple bonds are formed.

• A dative covalent bail can be formed between an atom with no electrons and an cantlet with a alone pair.

Atomic bonding and Lewis diagrams

Textbook Exercise 3.v

calcium

lithium

phosphorous

potassium

silicon

sulfur

bromine (\(\text{Br}_{2}\))

carbon dioxide (\(\text{CO}_{ii}\))

nitrogen (\(\text{Northward}_{2}\))

hydronium ion (\(\text{H}_{3}\text{O}^{+}\))

sulfur dioxide (\(\text{SO}_{2}\))

the number of valence electrons for each of the atoms involved in the reaction

Nitrogen: \(\text{v}\), hydrogen: \(\text{i}\) Carbon: \(\text{4}\), hydrogen: \(\text{i}\)

the Lewis diagram of the product that is formed

\(\text{NH}_{three}\)

\(\text{CH}_{4}\)

the proper name of the product

\(\text{NH}_{iii}\): ammonia

\(\text{CH}_{4}\): methane

How many valence electrons does element Y take?

\(\text{six}\). There are \(\text{vi}\) dots around element Y and from our knowledge of Lewis diagrams we know that these stand for the valence electrons.

How many valence electrons does element X have?

\(\text{1}\). X contributes 1 electron (represented by a cross) to the bond and Ten has no other electrons.

How many covalent bonds are in the molecule?

\(\text{2}\) unmarried bonds. From our knowledge of Lewis diagrams we look at how many cross and dot pairs at that place are in the molecule and that gives united states the number of covalent bonds.

These are single bonds since there is only 1 dot and cross pair between next atoms.

Suggest a name for the elements X and Y.

The most likely atoms are: Y: oxygen and X: hydrogen.

Annotation that Y could besides be sulfur and Ten hydrogen and the molecule would and so exist hydrogen sulfide (sulfur dihydride).

Complete the following tabular array:

Compound	\(\text{CO}_{2}\)	\(\text{CF}_{4}\)	\(\text{How-do-you-do}\)	\(\text{C}_{2}\text{H}_{ii}\)
Lewis diagram
Total number of bonding pairs
Total number of non-bonding pairs
Unmarried, double or triple bonds

Compound	\(\text{CO}_{2}\)	\(\text{CF}_{4}\)	\(\text{Hullo}\)	\(\text{C}_{2}\text{H}_{ii}\)
Lewis diagram
Total number of bonding pairs	\(\text{4}\)	\(\text{iv}\)	\(\text{1}\)	\(\text{five}\)
Full number of not-bonding pairs	\(\text{4}\)	\(\text{12}\)	\(\text{3}\)	\(\text{0}\)
Single, double or triple bonds	Ii double bonds	Four unmarried bonds	1 single bond	One triple bond and two single bonds

Source: https://www.siyavula.com/read/science/grade-11/atomic-combinations/03-atomic-combinations-01

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