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Read and discuss the article about Chemical Bonding

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Chemical Bond Introduction

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Chemical bonds | Chemistry of life | Biology (article) | Khan Academy

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The difference between ionic and covalent bonds

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Ionic and covalent bonds are the two main types of chemical bonding. A chemical bond is a link formed between two or more atoms or ions. The main difference between ionic and covalent bonds is how equally the electrons are shared between atoms in the bond.

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Polar & Non-Polar Molecules: Crash Course Chemistry #23

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Watch the video to practice Ionic, Non polar Convelent, and Polar Convelent bonds

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Watch the video to practice Ionic, Non polar Convelent, and Polar Convelent bonds

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Ionic Bonds, Polar Covalent Bonds, and Nonpolar Covalent Bonds

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Play a game with students to correctly categorize different chemical bonds as polar, non polar, or ionic.

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Practice determining different types of Chemical Bonds

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Practice determining different types of Chemical Bonds

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Read the article on the Lewis Structure for writing chemical bonds

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Lewis structures, also known as Lewis-dot diagrams, show the bonding relationship between atoms of a molecule and the lone pairs of electrons in the molecule. Lewis structures can also be useful in predicting molecular geometry in conjuntion with hybrid orbitals. A compound may have multiple resonance forms that are also all correct Lewis structures. This section will discuss the rules for writing out Lewis structures correctly.

Introduction

Writing out Lewis structures can be at times, tricky and somewhat difficult. A compound can have multiple Lewis Structures that contribute to the shape of the overall compound, so one Lewis structure of a compound may not necessarily be exactly what the compound looks like. But before we begin, there are a few things to know. An electron is represnted as a dot. A bond, which is made up of 2 shared electrons, is represented by two dots between the bonded atoms or a line ( Kekulé structures lines).Double bonds and triple bonds are represented as two and three lines/(pairs of electrons), respectively. Lone pairs on the outer rims of an atom are represented as two dots. The electrons represented in a lewis stucture are the outer-shell electrons, which are called valence electrons. This is because they are the ones involved in chemical reactions. The examples listed below may or may not be consistent for each step for the sake of showing multiple examples, since all steps apply to every compound. Now we are ready to begin.

Step 1: Draw out the molecular backbone of the compound. This is probably the most difficult step in drawing Lewis structures because we must pick one or multiple atoms to be the central connector to terminal (outside) atoms. Usually, the central atom is already known as in the case with many organic compounds containing Carbon as the central atom. But at the basics, the atom with the lowest electronegativity is the central atom. Other factors come into play such as the octet rule, which will be discussed in this section. One can find the electronegativity values of atoms by looking at a periodic table of elements that contains the electronegativity values for each element. Electronegativity typically increases from left to right and decreases from top to bottom in a periodic table.

Correct Correct Wrong

Step 2: Add up the valence electrons for each atom in the molecule.

For example,

H ₂ O 2 H: 2 x 1 electron = 2 electrons

1 O: 1 x 6 electrons = 6 electrons

--------------------------------------------------

Total: 8 electrons

Step 3: (Octet Rule) All covalent bonds are shown by two shared electrons. Place a pair of electrons between two elements that are connected to each other with a single bond. Next, give as many atoms as possible 8 electrons, except for Hydrogen, which has a duet (2-shared electrons). Do this by placing lone electron pairs on the outer atoms until each outer atom fulfills the octet. Place any remaining lone electron pairs on the central atom. Usually, electronegative atoms will have lone pairs on them. The number of electrons placed on the molecule should correspond to the total number of valence electrons of the atoms in the molecule.

Hydrochloric Acid

Many times, satisfying the octet rule and total number of valence electrons in the molecule with only single bonds is not the case, meaning that there could be less valence electrons than needed to satisfy all octets with only single bond connections. In these situations, double or triple bonds may be needed to obtain octets. Rectify these situations by moving a lone pair of electrons from one atom to the bond between those two atoms. By doing this, one atom that is deficient of an electron pair of its octet can be filled, while having no effect on the other atom's octet. Each compound is a case-by-case scenario and thus must be thought out carefully to solve the octet/total # of valence electron problems.

Ethene Ethyne

****A simple way to find the need for double or triple bonds is by counting the number of electrons to fill all octets, then finding the difference with the total number of valence electrons available in that compound. The difference will show exactly how much electrons that need to be shared to meet the octet (duet for hydrogen) for all atoms.*****

-Note: Ethene has 2 C + 4 H amounting to 12 available electrons. If every bond was single electron paired, there would be 10 electrons. The hydrogen would have its duet met but both Carbons will not have their octet met. Placing the extra pair of electrons in between the two carbons, making a double bond amends the problem. Ethyne has 2 C + 2 H amounting to 10 available electrons. Placing the two extra pair of electrons between both carbons, making a triple bond fixes the problem with the octet rule and allows for all available electrons to be placed on the compound.

Step 4: Assign Formal Charges (FC) to each atom in the molecule.The formal charge of an atom is pretty much the difference between the number of valence electrons that a neutral atom would have and the number of electrons that is contained on the atom of the Lewis structure. If a charge on an atom is zero, you don't need to put any symbols on that atom. The total of the formal charges on a neutral molecule should be equal to zero. An atom is charged if the total charge of the atom is different from the valence electrons of the free, non-bonded atom.

Formal Charge

FC = # of valence electrons - # of lone pair electrons - ½(number of bonding electrons)

Charge is distributed over the surroundings and thus called, a formal charge.

Hydroxide Anion Methyl Cation

Exceptions

There are some exceptions to the Lewis structure rules that one always needs to keep in mind. Regardless, these exceptions are still correct Lewis structures.

Exception 1: Some compounds have odd # of electrons.

For example,

Ethyl Radical

Exception 2: Some compounds don't have enough valence electrons for an octet. The octet rule only works for elements in the second row given enough valence electrons and usually, some of these compounds are very reactive.

For example,

Lithium Hydride

-Note: Lithium and hydrogen each have 1 electron to share. Hydrogen meets its duet, but Lithium is still 6 electrons deficient from meeting octet.

Exception 3: After the second row of table of elements, it is not uncommon to see an "expanded valence shell," where atoms have more than 8 valence electrons. The usual explanation is that the bonding involves d-orbital electrons, and the resulting Lewis structures look like that shown below.

For example,

Sulfuric Acid

Notice that sulfur atom has 12 electrons around it.

However, expanded octets are not required to draw Lewis structures for these types of molecules. For example, another Lewis structure of sulfuric acid is shown here, and does not include any expanded octets.

By the rules of drawing Lewis structures, both of these structures have undesirable features. The top structure has expanded octets, whereas the bottom structure contains separated formal charge. However, quantum mechanical electronic structure calculations find little d-orbital participation in the bonding, as would be expected for expanded octet bonding, suggesting that the formal charge separation is preferable to expanding the octets. Therefore, there is generally no reason to draw structures with expanded octets, because they usually don't contribute significantly to the overall structure of the molecule. Even molecules such as XeF ₆ can be drawn without expanded octets.

Consider Ozone (O ₃)

---------------------------------------------

Total: 18 electrons

For example,

Problems

Draw out a correct Lewis Structure for the following compounds.

1. HCN

2. LiF

3. C₃H₆ (two possibilities)

4. CO₃ ²⁻

5. CH₃NO₂

Answers:

1.

2.

3. or

4.

5.

Another simple and general procedure to draw Lewis structures has been proposed by A.B.P. Lever (see reference 5). Before beginning this procedure it is necessary to know the basic geometry of the molecule, i.e. whether it is cyclic or noncyclic, and which atoms are connected to which. Several worked examples for the determination of the Lewis structures of simple and more complicated species using the above described method can be found in the original paper and in Chemistry Net.

References

1. Schore, Neil E. and Vollhardt, K. Peter C. Organic Chemistry: Structure and Function. 5th Edition. New York: Bleyer, Brennan. 2007. (14-19)

2. Klein, David R. Organic Chemistry as a Second Language. New York: Wiley, 2007.

3. Miburo, Barnabe B. "Simplified Lewis Structure Drawing for Nonscience Majors." Journal of Chemical Education 75 (1998): 317-19.

4. Reed, J.L. The Lewis Structure: An Expanded Perspective. Journal of Chemical Education 71, (1994): 98-100.

5. Lever, A.B.P. "Lewis Structures and the Octet Rule" Journal of Chemical Education 49, (1972): 819 - 821

Contributors

Daniel S. Kim, Paul Wenthold

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Lewis structures, also known as Lewis-dot diagrams, show the bonding relationship between atoms of a molecule and the lone pairs of electrons in the molecule. Lewis structures can also be useful in predicting molecular geometry in conjuntion with hybrid orbitals. A compound may have multiple resonance forms that are also all correct Lewis structures.

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Lewis Diagrams Made Easy: How to Draw Lewis Dot Structures

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Practice writing Chemical Formulas

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Complete Chemical Bonding lab

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Some people work best on their own. They get on with the task with minimal input from others. But other people work best in a group. They achieve their best outcomes when they combine forces; sharing ideas, knowledge, and tasks. Neither way is better than the other - it simply depends on which method suits you the best.

Chemical bonding is very similar to this. Some atoms are much happier by themselves, while some prefer to join up with others. They do this by forming chemical bonds.

Chemical bonding is the attraction between different atoms that enables the formation of molecules or compounds. It occurs thanks to the sharing, transfer, or delocalization of electrons.

This article is an introduction to the types of bonding in chemistry.

We will look at why atoms bond.

We'll explore the three types of chemical bonds.

We'll then look at factors affecting the strength of bonding.

Why Do Atoms Bond?

At the start of this article, we introduced you to a chemical bond: the attraction between different atoms that enables the formation of molecules or compounds. But why do atoms bond to each other in this way?

Simply put, atoms form bonds in order to become more stable. For the majority of atoms, this means obtaining a full outer shell of electrons. An atom's outer shell of electrons is known as its valence shell; these valence shells typically require eight electrons to fill them up completely. This gives them the electron configuration of the noble gas closest to them in the periodic table. Achieving a full valence shell puts the atom in a lower, more stable energy state, which is known as the octet rule.

The octet rule states that the majority of atoms tend to gain, lose, or share electrons until they have eight electrons in their valence shell. This gives them the configuration of a noble gas.

But to get to this more stable energy state, atoms might need to move some of their electrons around. Some atoms have too many electrons. They find it easiest to get a full valence shell by getting rid of surplus electrons, either by donating them to another species, or by delocalizing them. Other atoms don't have quite enough electrons. They find it easiest to gain extra electrons, either by sharing them or accepting them from another species.

When we say 'easiest', we really mean 'most energetically favorable'. Atoms don't have preferences - they're simply subject to the laws of energy that govern the whole universe.

You should also note that there are some exceptions to the octet rule. For example, the noble gas helium has just two electrons in its outer shell and is perfectly stable. Helium is the noble gas nearest to a handful of elements such as hydrogen and lithium. This means that these elements are also more stable when they have just two outer shell electrons, not the eight that the octet rule predicts. Check out The Octet Rule for more information.

Moving electrons around creates differences in charges, and differences in charges cause attraction or r epulsion between atoms. For example, if one atom loses an electron, it forms a positively charged ion. If another atom gains this electron, it forms a negatively charged ion. The two oppositely charged ions will be attracted to each other, forming a bond. But this is just one of the ways of forming a chemical bond. In fact, there are a few different types of bonds that you need to know about.

Types of Chemical Bonds

There are three different types of chemical bonds in chemistry.

These are all formed between different species and have different characteristics. We'll start by exploring the covalent bond.

Covalent Bonds

For some atoms, the simplest way to achieve a filled-up outer shell is by gaining extra electrons. This is typically the case with non-metals, which contain a large number of electrons in their outer shell. But where can they get extra electrons from? Electrons don't just appear out of nowhere! Non-metals get around this in an innovative way: they share their valence electrons with another atom. This is a covalent bond.

A covalent bond is a shared pair of valence electrons.

A more accurate description of covalent bonding involves atomic orbitals. Covalent bonds form when valence electron orbitals overlap, forming a shared pair of electrons. The atoms are held together by electrostatic attraction between the negative electron pair and the atoms' positive nuclei, and the shared pair of electrons counts towards the valence shell of both bonded atoms. This enables them both to effectively gain an extra electron, bringing them closer to a full outer shell.

In the example above, each fluorine atom starts with seven outer shell electrons - they're one short of the eight needed to have a full outer shell. But both fluorine atoms can use one of their electrons to form a shared pair. In this way, both atoms seemingly end up with eight electrons in their outer shell.

There are three forces involved in covalent bonding.

The repulsion between the two positively charged nuclei.

The repulsion between the negatively charged electrons.

The attraction between the positively charged nuclei and the negatively charged electrons.

If the total strength of the attraction is stronger than the total strength of the repulsion, the two atoms will bond.

Multiple Covalent Bonds

For some atoms, such as fluorine, just one covalent bond is enough to give them that magic number of eight valence electrons. But some atoms might have to form multiple covalent bonds, sharing further pairs of electrons. They can either bond with multiple different atoms, or form a double or triple bond with the same atom.

For example, nitrogen needs to form three covalent bonds in order to achieve a full outer shell. It can either form three single covalent bonds, one single and one double covalent bond, or one triple covalent bond.

Covalent Structures

Some covalent species form discrete molecules, known as simple covalent molecules, made up of just a few atoms joined with covalent bonds. These molecules tend to have low melting and boiling points. But some covalent species form giant macromolecules, made up of an infinite number of atoms. These structures have high melting and boiling points. We saw above how a fluorine molecule is made up of just two fluorine atoms covalently bonded together. Diamond, on the other hand, contains many hundreds of atoms covalently bonded together - carbon atoms, to be precise. Each carbon atom forms four covalent bonds, creating a giant lattice structure that stretches in all directions.

Check out Covalent Bonding for a more detailed explanation of covalent bonds. If you want to know more about covalent structures and the properties of covalent bonds, head over to Bonding and Elemental Properties.

Ionic Bonds

Above, we learned how non-metals effectively 'gain' extra electrons by sharing an electron pair with another atom. But bring metaland anon-metaltogether, and they can do one better - they actually transfer an electron from one species to the other. The metal donates its extra valence electrons, bringing it down to eight in its outer shell. This forms a positive cation. The non-metal gains these donated electrons, bringing the number of electrons up to eight in its outer shell, forming a negative ion, called an anion. In this way, both elements are satisfied. The oppositely charged ions are then attracted to each other by strong electrostatic attraction, forming an ionic bond.

An ionic bond is an electrostatic attraction between oppositely charged ions.

Here, sodium has one electron in its outer shell, whilst chlorine has seven. In order to achieve a complete valence shell, sodium needs to lose one electron whilst chlorine needs to gain one. Sodium, therefore, donates its outer shell electron to chlorine, transforming into a cation and an anion respectively. The oppositely charged ions are then attracted to each other by electrostatic attraction, holding them together.

When the loss of an electron leaves an atom with no electrons in its outer shell, we consider the shell below as the valence shell. For example, the sodium cation has no electrons in its outer shell, so we look to the one below - which has eight. Sodium, therefore, satisfies the octet rule. This is why group VIII is often called group 0; for our purposes, they mean the same thing.

Ionic Structures

Ionic structures form giant ionic lattices made up of many oppositely charged ions. They don't form discrete molecules. Each negatively charged ion is ionically bonded to all of the positively charged ions around it, and vice versa. The sheer number of ionic bonds gives ionic lattices high strength, and high melting and boiling points.

Covalent bonding and ionic bonding are actually closely related. They exist on a scale, with completely covalent bonds at one end and completely ionic bonds at the other. Most covalent bonds exist somewhere in the middle. We say that bonds that behave a little like ionic bonds have an ionic'character'.

Metallic Bonds

Now we know how non-metals and metals bond with each other, and how non-metals bond with themselves or with other non-metals. But how do metals bond? They have the opposite problem to non-metals - they have too many electrons, and the easiest way for them to achieve a full outer shell is by losing their extra electrons. They do this in a special way: by delocalizing their valence shell electrons.

What happens to these electrons? They form something called a sea of delocalization. The sea surrounds the remaining metals centers, which arrange themselves into an array of positive metal ions. The ions are held in place by electrostatic attraction between themselves and the negative electrons. This is known as a metallic bond.

Metallic bonding is a type of chemical bonding found within metals. It consists of the electrostatic attraction between an array of positive metal ions and a sea of delocalized electrons.

It is important to note that the electrons aren't associated with any one metal ion in particular. Instead, they move freely between all of the ions, acting both as a glue and a cushion. This leads to good conductivity in metals.

We learned earlier that sodium has one electron in its outer shell. When sodium atoms form metallic bonds, each sodium atom loses this outer shell electron to form a positive sodium ion with a charge of +1. The electrons form a sea of delocalization surrounding the sodium ions. The electrostatic attraction between the ions and the electrons is known as a metallic bond.

Metallic Structures

Like ionic structures, metals form giant lattices that contain an infinite number of atoms and stretch in all directions. But unlike ionic structures, they are malleable and ductile, and they usually have slightly lower melting and boiling points.

Bonding and Elemental Properties contains all you need to know about how bonding affects the properties of different structures.

Summarising Types of Bonds

We've made you a handy table to help you compare the three different types of bonding. It summarises all you need to know about covalent, ionic, and metallic bonding.

Covalent

Ionic

Metallic

Description

Shared pair of electrons

Transfer of electrons

Delocalization of electrons

Electrostatic forces

Between the shared pair of electrons and the atoms' positive nuclei

Between oppositely charged ions

Between positive metal ions and the sea of delocalized electrons

Structures formed

Simple covalent moleculesGiant covalent macromolecules

Giant ionic lattices

Giant metallic lattices

Diagram

The Strength of Chemical Bonds

If you had to guess, which type of bonding would you label as the strongest? It is actually ionic > covalent > metallic bonding. But within each type of bonding, there are certain factors that influence the bond's strength. We'll start by looking at the strength of covalent bonds.

Strength of Covalent Bonds

You'll remember that a covalent bond is a shared pair of valence electrons, thanks to the overlap of electron orbitals. There are a few factors that affect the strength of a covalent bond, and they all have to do with the size of this area of orbital overlap. These include the type of bond and the size of the atom.

As you move from a single covalent bond to a double or triple covalent bond, the number of overlapping orbitals increases. This increases the strength of the covalent bonding.

As the size of the atoms increases, the proportional size of the area of orbital overlap decreases. This decreases the strength of the covalent bonding.

As polarity increases, the strength of the covalent bonding increases. This is because the bond becomes more ionic in character.

Strength of Ionic Bonds

We now know that an ionic bond is an electrostatic attraction between oppositely charged ions. Any factors that affect this electrostatic attraction affect the strength of the ionic bond. These include the charge of the ions and the size of the ions.

Ions with a higher charge experience stronger electrostatic attraction. This increases the strength of the ionic bonding.

Ions with a smaller size experience stronger electrostatic attraction. This increases the strength of the ionic bonding.

Visit Ionic for a deeper exploration of this topic.

Strength of Metallic Bonds

We know that a metallic bond is an electrostatic attraction between an array of positive metal ions and a sea of delocalized electrons. Once again, any factors that affect this electrostatic attraction affect the strength of the metallic bond.

Metals with more delocalized electrons experience stronger electrostatic attraction, and strongermetallic bonding.

Metals ions with a higher charge experience stronger electrostatic attraction, and stronger metallic bonding.

Metal ions with a smaller size experience stronger electrostatic attraction, and stronger metallic bonding.

You can find out more at Metallic .

Bonding and Intermolecular Forces

It is important to note that bonding is completely different from intermolecular forces. Chemical bonding occurs within a compound or molecule and is very strong. Intermolecular forces occur between molecules and are much weaker. The strongest type of intermolecular force is a hydrogen bond.

Despite its name, it is not a type of chemical bond. In fact, it is ten times weaker than a covalent bond!

Types of Chemical Bonds - Key takeaways

Chemical bonding is the attraction between different atoms that enables the formation of molecules or compounds. Atoms bond to become more stable according to the octet rule.

A covalent bond is a shared pair of valence electrons. It typically forms between non-metals.

An ionic bond is an electrostatic attraction between oppositely charged ions. It typically occurs between metals and non-metals.

A metallic bond is an electrostatic attraction between an array of positive metal ions and a sea of delocalized electrons. It forms within metals.

Ionic bonds are the strongest type of chemical bond, followed by covalent bonds and then metallic bonds. Factors affecting the strength of bonding include the size of atoms or ions, and the number of electrons involved in the interaction.

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Some people work best on their own. They get on with the task with minimal input from others. But other people work best in a group. They achieve their best outcomes when they combine forces; sharing ideas, knowledge, and tasks. Neither way is better than the other - it simply depends on which method suits you the best.

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Bonding Models and Lewis Structures: Crash Course Chemistry #24

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