Chemistry

1. Drawing and Bonding

In this section, we will go over ionic, covalent, metallic bonds, and also how to draw atoms, ions, etc.

Let's start by drawing atoms. There's a few things you need to include in your drawing:

• The nucleus (and protons and neutrons).
• The electrons (how many electrons and how many electron shells).

Note:

• Where you place the electron inside the shell doesn't matter.
• When drawing ions, you must enclose the drawing with square brackets and put the charge of the ion on top.

Let's use a couple of examples to demonstrate how to draw.

Important tips and notes to remember when drawing ▼

1. Sometimes, the exam question will ask you to draw only the valence electrons or the valence shell. In that case, only draw the outermost shell with its electrons. Never draw the inside shells or electrons.
2. Sometimes, the question will tell you the number of neutrons that you are supposed to have. In that case, don't use the nucleon number in the periodic table.
3. In some periodic tables, the nucleon number is not a whole number (this is because the periodic table typically shows the average nucleon number of that element, since many elements have isotopes with differing nucleon numbers). If you have to use the periodic table to decide how many neutrons you should have, round up/down the nucleon number to have a whole number.
4. Again, when drawing ions, remember to enclose your ion in square brackets with the ion charge on top, same way I have drawn it in the chloride example.
5. More details regarding drawing will be mentioned in future sections, when appropriate.

Chemical Bonds

I recommend you revise the section regarding the goal of chemical bonds before reading forward.

Ionic bonding

Ionic bonding is the first major concept we will cover. As I said previously, ionic bonds occur when a transfer of electrons happens.

Let us apply these rules in the following examples.

Example 1: Sodium Chloride.

• Identify the electron donor and acceptor.

Sodium is the electron donor because it is the metal, chloride (or chlorine) is the electron acceptor because it is the non-metal.

• How many electrons are given away by 1 atom of the electron donor, and how many electrons are accepted by 1 atom of the electron acceptor?

Since sodium is in group 1, then it gives away 1 electron. Chlorine is in group 7, and from the rules above, electrons accepted = 8-group number. Hence, 8-7=1, every chlorine atom will accept 1 electron.

• What is the chemical formula of sodium chloride

I will use the shortcut I showed you above.

• Draw the resulting ionic compound.

Please check with your syllabus or teachers how exactly they want the diagram to be drawn. Sometimes, they don't want the inner shells, only the outermost shell. I will draw the complete compound.

The rules of drawing are the same as they were before (just put the symbol of the element inside the nucleus instead of protons/neutrons). Remember to keep track of how many electrons each element has in their shells before you start drawing. You must always have a full outer shell.

This is known as a dot-and-cross diagram. It is given this name because the electrons are dots and crosses. Moreover, notice that sodium now has 2 shells instead of 3. This is because it lost the electron that was in the third shell.

Example 2: Aluminum oxide.

• Identify the electron donor and acceptor.

Aluminum is the metal, hence it will be the electron donor. Oxide (the ion of oxygen) is the non-metal, hence it will be the electron acceptor.

• How many electrons are given away by 1 atom of the electron donor, and how many electrons are accepted by 1 atom of the electron acceptor?

Aluminum is in group 3, which means that each aluminum atom will donate 3 electrons. Oxygen is in group 6, and electrons accepted=8-group number = 8-6=2. Therefore, each oxygen atom will accept 2 electrons.

• What is the chemical formula of aluminum oxide?

I will use the same shortcut I used previously.

Let's assume that I'm not sure whether Al₂O₃ is the correct formula or not. Here's how to check. ▼

1. One aluminum atom will give 3 electrons, so 2 aluminum atoms will give 6 electrons. Electrons donated = 6.
2. One oxygen atom will accept 2 electrons, so 3 oxygen atoms will accept 2×3=6 electrons in total. Electrons accepted = 6.
3. Electrons donated = electrons accepted.
4. Moreover, the numbers are in their simplest ratio (2 and 3, you can't simplify it more than that).
5. Hence, it is correct!

• Draw the resulting ionic compound.

New drawing ruling to keep in mind (in addition to the dot-and-cross): It will be tedious to draw 2 complete aluminum ions and 3 complete oxygen ions. Instead, we will draw just one ion of each type and indicate how many of each ion are present by writing the number to the left of the brackets, as shown below:

Each aluminum atom lost 3 electrons, while each oxygen atom gained 2 electrons, giving them both full outer shells.

Example 3: Calcium sulfide

• Identify the electron donor and acceptor.

Calcium is the metal, so it will be the electron donor. Sulfide (the ion of sulfur) is the non-metal, so it will be the electron acceptor.

• How many electrons are given away by 1 atom of the electron donor, and how many electrons are accepted by 1 atom of the electron acceptor?

Calcium is in group 2, so it will give away 2 electrons. Sulfide (or sulfur) is in group 6, so electrons accepted = 8-group number =8-6=2.

• What is the chemical formula of calcium sulfide?

We will use the same trick we used in the previous examples.

As you can see, the shortcut we used initially told us: for every 2 calcium ions, you will need 2 sulfide ions. But that is the same as saying: for every 1 calcium ion, you will need 1 sulfide ion. And since the latter is the simplest ratio, we should pick that. Moreover, as we stated previously, if we have a 1, we can simply make it disappear, giving us a final answer of CaS.

If you are unsure, check: does electron donated=electron accepted? You will notice that it does.

• Draw the resulting ionic compound.

So are ionic compounds just 2 or 3 ions floating around in space?

Characteristics of ionic compounds

• High melting/boiling points
  This is because of the multiple strong ionic bonds in the lattice, which require a lot of energy to break.
• Good conductors of electricity in molten or aqueous states, but not solid state
  
  

  In the solid state → the ions are stuck to each other and don't move → hence preventing any electric current from flowing.

  When they are in the molten/aqueous state → the ionic bonds are broken → allowing the ions to roam freely and transfer the electric current.

• Brittle
  They break fairly easily. This is because of their ionic lattice structure. When a force is applied to an ionic compound, some of the rows shift their position, causing the following arrangement. As you can see, now positive ions are now directly next to each other. Same thing with the negative ions. This causes extreme repulsion to occur. Recall the rule: same-charge particles repel each other.

Covalent bonding

This is the second type of bonding that you should be familiar with. A compound made up of covalent bonds is called a molecular compound.

Covalent bonding usually occurs between non-metals. It is based on atoms sharing electrons. Covalent bonding happens so that each atom has a full outer shell, which makes them very stable.

Why don't non-metals completely transfer their electrons to each other similar to metals in ionic bonding (recall that during ionic bonding metals completely lose their electrons and give them away to non-metals)? It's because it is more beneficial for non-metals to gain electrons rather than lose them to reach stability, and covalent bonding allows both of the participating atoms to "gain by sharing."

To understand the general idea of covalent bonding, I will first give a simple example for you to imagine.

In this imaginary example, the apples are the electrons and the 2 boys are the atoms. There are a few things I want you to notice in this example:

• The most common case of covalent bonding (which is what I'm covering) is that each atom contributes some electrons to share with the other atom in the covalent bond.

  This means that half of the shared electrons will be from Atom 1 and the other half will be from Atom 2.

  This also means that the number of shared electrons must be an even number.

• The shared electrons are owned by both atoms in the covalent bond. Keep in mind: a covalent bond is 1 shared pair of electrons.

Single, double, and triple bonds

Between 2 atoms, covalent bonds come in 3 levels:

1. Single bonds: each atom contributes a single electron to the sharing (so you have 1 pair of shared electrons). In structural formulae, it is made up of 1 dash (-).
2. Double bonds: each atom contributes 2 electrons to the sharing (so you have 2 pairs of shared electrons). This means that a double bond is made up of 2 covalent bonds, so in structural formulae, it is made up of 2 dashes on top of each other (=).
3. Triple bonds: each atom contributes 3 electrons to the sharing (so you have 3 pairs of shared electrons). This means that a triple bond is made up of 3 covalent bonds, so in structural formulae, it is made up of 3 dashes on top of each other (≡).
4. Quadruple bonds exist, but you aren't expected to know them for beginner-level chemistry courses.

Drawing molecules - Example 1: Fluorine

Now, we will look at group 7 and draw covalent bonds using dot-and-cross diagrams and also their structural formula.

Group 7 is also known as the halogens group (each element can be called a "halogen.") For your general information, "halogen" is a Greek term that means "salt-former."

The halogens (and some other elements we will discuss) are known as a diatomic elementDiatomic elementsElements that are found in nature as 2 atoms bonded together. (The prefix di- means 2, atomic refers to atoms. So diatomic means 2 atoms). This basically means that they are found in nature as molecules made from 2 atoms of that element, bonded together by a covalent bond.

Covalent bonding with more than 2 atoms:

It is extremely common for covalent structures to have more than 2 atoms. Here is an example, CH₄ or methane (1 carbon atom bonding with 4 hydrogen atoms).

Naming convention of typical covalent compounds

This is very important. Most of the time, you will only be given the name of the chemical compound, and you would be expected to know the chemical formula. Add this to our list of naming conventions:

1. First word: Write the name of the first element in the chemical formula.
2. Add a prefix to that name (do not write mono. For the first element, the prefixes start from di-)

3. Next word: Write the monoatomic-ion name of the second element.
4. Add a prefix. Keep in mind that of the "vowel drop" rule: If your prefix ends with a vowel and the monoatomic-ion name starts with a vowel, then you will drop the ending vowel of the prefix. The only exception to this rule is if your prefix ends with "i," in that case don't drop anything.

   List of vowels in English: A, E, I, O, U, and (sometimes) Y.

   Examples -
   Mono (prefix) + oxide (monoatomic-ion name of element) = monoxide (not monooxide)

   Penta + oxide = pentoxide (not pentaoxide)

   Exception example: di + oxide = dioxide

A mix between ionic and covalent bonding - Polyatomic ions

As you hopefully know by now, monoatomic ions are ions made from a single atom. However, we can also have ions made from multiple atoms, called polyatomic ions (poly means many).

These are ions formed from multiple atoms covalently bonded together that act as if they are a singular ion.

For example, in NaCl, Na⁺ is the positive ion and Cl^- is the negative ion. But how about NaNO₃? The positive ion is still Na⁺, but now the negative ion is NO₃^- as a whole. And if you were to look at NO₃^-, you would notice it is a covalent compound.

Your goal in this section is to memorize the common polyatomic ions by name, charge, and formula. I organized the ions into groups for ease of memorization.

Tips on memorizing in the end.

NOTE: The polyatomic ions section will be continued in the next update to these notes.

Simple molecules and Giant covalent structures:

There are 2 types of covalent chemical structures: simple molecular compounds and giant covalent structures.

Simple molecular compounds are the most common type you will see. These molecules are made up of a relatively small number of covalent bonds, which makes the molecule "simple."

1. Low melting/boiling points: Between each molecule, there are specific forces responsible for holding the molecules together, called "intermolecular forces." These forces are relatively weak and don't require much energy to break, giving simple molecular molecules low melting and boiling points.

   Important note: the covalent bonds are not broken when you melt or boil something. These covalent bonds remain and they are relatively quite strong. Instead, what is broken is the intermolecular forces between the molecules.

2. Poor electrical conductivity: This is because they do not have any delocalized electrons or ions.
3. Examples: The majority of covalent molecules you see are all simple covalent structures. For example, all the covalent molecules I drew above when explaining covalent bonding are simple covalent structures.

Giant covalent structures are a network of many atoms connected together by covalent bonds. There are 3 giant covalent structures you should know about: graphite, diamond, and silicon dioxide. We will go over all 3 of them (you need to be able to recognize these structures if you were shown a picture in an exam).

Graphite

First, let's look at its structure:

Blue - carbon atom
Red - covalent bond
White - intermolecular forces

As you can see, graphite is made up of carbon atoms covalently bonded together in a hexagonal arrangement to form a layer, and each layer is weakly connected to layers above and below them via the weak intermolecular forces (obviously my diagram is limited to a few layers, in reality there's a bunch). You can deduce a couple of things from this structure:

1. Lubricant - Because the layers are weakly connected together via intermolecular forces, they can slide over each other quite easily. This makes graphite a good lubricant. In fact, this is why graphite is used in pencils - because the layers slide off the pencil and onto the paper.
2. Good electrical conductor, allowing it to be used as electrodes - A carbon atom has 4 valence electrons. In graphite, notice how every carbon atom is forming 3 covalent bonds, which means every carbon atom is sharing 3 of its 4 electrons. This means that for every carbon atom, there is 1 electron that is not involved in bonding, instead it is delocalized and freely roaming around the structure. This allows graphite to be a good conductor of electricity, and hence used as an electrode (an electrode is an electrical conductor).

Diamond

Silicon Dioxide (silica)

NOTE: The silicon dioxide section will be continued in the next update to these notes.

First, what is the chemical formula of Silicon dioxide? Let's recall our naming conventions.

The structure of silicon dioxide is very similar to the structure of diamond:

Blue - carbon
Yellow - oxygen
Red - covalent bond