1.3 - Bonding

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Bonding intro

The physical and chemical properties of compounds depend on the way they are held together, by chemical bonds or intermolecular forces. Scientists use theories of bonding and structure to engineer new materials with desirable properties. These new materials can offer many new applications in a range of modern technologies and advancements.

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Ionic Bonding

Metals have either 1, 2 or 3 electrons in their outer shells, therefore, easiest way for a metal to gain a full outer shell, is to lose these electrons. An ionic bond is formed when one atoms loses electrons, and one atom gains electrons.

Ionic bonds are between metals and non-metals

Electrons are transferred from the metal to the non-metal

Ions are formed, both positive and negative

Consider the reaction between an atom of sodium and an atom of chlorine. Looking at the periodic table, you can determine how many electrons are in the outer shell of each atom. Sodium has 1, and chlorine has 7.

Therefore, the easiest way for both atoms to become stable, is for the sodium atom to 'give' one of its electrons to the atom of chlorine. This results in two ions:

Na+ (Because it has lost a negatively charged electron, leaving it positively charged)

Cl- (Because it has gained a negatively charged electron, leaving it negatively charged)

These ions are of opposite charge and therefore form a very strong electrostatic bond. We can therefore say that an ionic bond is the result of electrostatic attraction between oppositely charged ions within a lattice.

Properties of an Ionic Bond

Solid at room temperature

Giant structures

High melting points

Conduct electricity when in solution or molten, but not when in solid state

Ionic compounds are brittle solids

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Covalent Bonding

Covalent bonds are similar to ionic bonding as both are based on electrostatic attraction. In covalent bonds however, the electrostatic attraction is between the positively charged nucleus of each atom, and a one of more pairs of shared, negatively charged, electrons.

The sharing of a single pair of electrons is known as a single bond, but you can also have double and triple bonds where 2 or 3 pairs of electrons are shared between atoms.

Consider the simplest example: Hydrogen. Hydrogen has one electron in its outer shell, but needs two in order to complete its outermost shell. To do this, a hydrogen atom will simply find another, and will share an electron pair with the other atom, forming H2.

Now consider Methane. It has the chemical formula of CH4. Carbon needs 4 electrons in order to fill its outer shell, and therefore will form a bond with 4 atoms of Hydrogen.

A Nitrogen atom has an electron configuration of 2,5, and needs 8 to fill its outer shell. This means it needs 3 extra electrons, and that it will bond with another Nitrogen atom.

We mentioned earlier that covalent bonds could involve more than a single pair of electrons, and can be double or even triple bonds. In the case of Nitrogen, a triple bond is formed.

Properties of a Covalent Bond

Unless the substance is a Giant Covalent Structure, they are either gases, liquids or solids with low melting and boiling points. This is because there is only weak forces of attraction between molecules

They are poor conductors of electricity

Solutions do not conduct electricity

Coordinate Bonding

A covalent bond consists of the sharing of a pair (or more than one pair) of electrons, with a single electron being provided by each atom forming the bond. In some cases however, both electrons in the pair may be provided by the same atom. When this happens it's called a coordinate bond, or sometimes a dative covalent bond.

In the diagram below, the nitrogen atom is providing both atoms to the bond between itself and one of the oxygen atoms. This is the coordinate bond.

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Metallic Bonding

Most metals are solid at room temperature, and are hard. They also tend to have quite high melting points, which are a result of the strong bonds within the structure, holding the atoms together.

Metallic bonding involves attraction between delocalised

electrons and positive ions arranged in a lattice.

A metal is very similar in structure to graphite. Every atom (or ion) in the metal has a spare electron in its outer shell. In this solid state, this outer electron is free to move around the structure. As it does this, it is described as being 'delocalised'. A metallic crystal, therefore, is a structure of atoms surrounded by a 'sea' of delocalised electrons,

This 'sea' of delocalised electrons is free and is able to travel throughout the structure, which means that they can carry and conduct and electrical charge.

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Ionic Crystals

All ionic compounds exist as giant, three-dimensional crystalline structures of tightly packed, and well ordered, positive and negative ions.

These three-dimensional lattices are not molecules. Molecules always have a fixed number of every atom, whereas the size of an ionic lattice varies and depends on the amount of material present. (i.e, 5g or 5kg of the same ionic compound will produce different sized lattices)

This lattice is held together by the strong electrostatic attraction of the oppositely charged ions. This attraction is very strong and leads to very hard structures with high melting and boiling points.

Below is the lattice that would form when Sodium and Chlorine atoms bond together, forming Sodium Chloride.

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Metallic Crystals

Metals exist as a lattice of positive ions in a sea of delocalised electrons. The attraction of positive to negative exists throughout the lattice. Metals generally have a high melting point, due to a result of these strong metallic bonds. They can also conduct electricity due to the freely moving sea of electrons.

The structure of a metallic crystal is very regular.

Metals tend to be malleable. If enough force is applied to a metal it will become misshapen, but then when the force is removed, will return to its original shape. Therefore, metals can be said to have elastic properties. This can be useful. Consider a tall building with steel beams acting as reinforcement: The building may bend slightly in strong winds, but when the winds die down, the building will return to its original shape.

If a larger force is applied however, the shape of the metal can change permanently. Metals can be worked into many different shapes, from wires to the chassis in a car.

This change in shape is brought about because the force that is applied is strong enough that it will cause some of the particles in the structure to permanently slide past their neighbours.

The regularity of the structure makes this easier. This change in structure would now be possible in diamond, for example, as the atoms of carbon are far too interconnected and would be unable to slide past one another in the same way.

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Molecular Crystals

Molecular crystals consist of molecules held together in a regular structure by one, or more, of the three types of intermolecular forces.

In molecular crystals, covalent bonds within the molecules hold the atoms together, but they do not act between molecules.

Given that intermolecular forces are much weaker than covalent, ionic or metallic bonds, molecular crystals have low melting points and low enthalpies of fusion (the energy required to change the state of a substance from solid into liquid).

Molecular crystals generally have the following properties:

Crystals that are soft and easy to break apart

Low melting points

Often sublime to form gases (Sublimation: Changing from a solid into a gas, without passing through the liquid phase)

They do not conduct electricity

A good example of a molecular crystal would be Iodine.

Macromolecular Crystals

These are formed in substances with large molecules where the covalent bonds extend throughout the compounds, creating a giant structure held together by strong bonds.

Macromolecular crystals generally have high melting points.

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Structure of Diamond

Some molecules are very simple structures, and some are huge and extremely complex.

Molecules always contain a fixed number of atoms bonded together, either by ionic or covalent bonds. Many of these are small but some can be very large indeed, like DNA and some plastics. No matter how big the molecule, however, it will always contain a specific number of atoms.

Despite this, some giant structures can exist whereby the number of atoms (or ions) is not fixed. One example would be diamond. Diamond is pure Carbon and only contains Carbon atoms. Each Carbon atom is bonded with 4 other atoms of Carbon, forming a tetrahedral structure.

While bonded, each individual Carbon atom will continue to seek a stable outer shell and therefore will continue bonding with other Carbon atoms. As a result of this, the structure can quickly become very large. It is not a molecule however because the number of atoms is not fixed or constant.

Carbon-Carbon bonds are generally very strong, which contributes to the main properties of diamond, being high melting points and extreme strength. The fact that every bond within the structure is the same Carbon-Carbon bond is what makes diamond remarkable.

Properties of Diamond

Extremely hard

Very high melting point

Does not conduct electricity (electrons are held tightly together)

Cannot be dissolved by any solvent yet discovered

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Structure of Graphite

Graphite is similar to diamond in that it only contains atoms of Carbon, but its structure, and therefore its properties, are very different. Instead of being arranged in a three-dimensional lattice, as in diamond, graphite is a series of Carbon layers.

It is similar in structure to a giant stack of paper, with many sheets all placed on top of one another. Each individual layer is very strong, because of the Carbon-Carbon bonds, but the layers can be easily separated. This is because the layers are held together by Van der Waals bonds, which are extremely weak.

Because of its structure, graphite is a very soft material. So soft in fact that it has a very greasy feel, and is often mixed with something else before it is of any use. (E.g being mixed with clay to form the 'graphite' we use in pencils)

Graphite retains the high boiling point and insolubility of diamond, despite its softness. This is because while the layers of graphite come apart easily, the bonds between the Carbon atoms are still very strong.

Unlike diamond, graphite can conduct electricity. This is because in graphite, each carbon atom is only bonded with three others, whereas in diamond it would be 4. This means that not all atoms have complete outer shells. This results in a number of electrons with free movement throughout each layer. This is how graphite is able to conduct electricity.

Interestingly, the delocalised electrons help contribute to stronger forces of attraction between the Carbon atoms, meaning graphite is actually more stable than diamond.

Properties of Graphite

Very soft

Greasy feel

Very high melting point

Insoluble

Conducts electricity

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Structure of Other Materials

Sodium Chloride:

Sodium Chloride forms an ionic crystal structure, with strong electrostatic attractions between the Na(+) and Cl(-) ions.

Magnesium:

Magnesium forms a metallic crystalline structure, as it is a lattice of positively charged ions that exist within a sea of negatively charged, delocalised electrons

Iodine:

Iodine forms a molecular crystal structure, as the iodine atoms are covalently bonded, and weak intermolecular forces hold the molecules together in a crystalline structure.

LEARN STRUCTURE OF ICE

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Shapes of Simple Molecules and Ions (Bonding angles)

We can predict the shape of simple molecules composed of covalent bonds by understanding how they arrange themselves.

In a molecule, each pair of electrons will repel other pairs of electrons because they have the same charge. This means that electrons pairs will seek to be as far apart from one another as possible. This is known as the Electron Pair Repulsion Theory.

Electrons pairs can exist either as lone pairs, or as shared pairs, which are involved in bonds.

Ammonia for example, contains 3 shared pairs and 1 lone pair of electrons. (See below).

Bond Angles in Substances with Shared Pairs

No. of electron pairs: 2

Structure name: Linear

Bond angles: 180°

No. of electron pairs: 3

Structure name: Trigonal Planar

Bond angles: 120°

No. of electron pairs: 4

Structure name: Tetrahedral

Bond angles: 109.5°

No. of electron pairs: 5

Structure name: Trigonal Bipyramidal

Bond angles: 90° & 120°

No. of electron pairs: 6

Structure name: Octahedral

Bond angles: 90°

Bond Angles in Substances with Lone Pairs

The main difference between lone pairs and shared pairs, is that lone pairs of electrons exert a greater 'push' on the shared pairs, pushing them further away from themselves. This is because the lone pairs have a greater attraction to the nucleus of the central atom than the shared pairs have.

Ammonia contains 1 lone pair, and 3 shared pairs of electrons. As a result, it forms what's known as a pyramidal structure, with a bond angle of 107.5°.

Water contains 2 lone pairs, and 2 shared pairs of electrons. Water forms a bent line structure, with a bond angle of 104.5°

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Bond Polarity

The force that holds molecules together is essentially an electrical attraction. It is the difference in charge between positive and negative ions.

With an ionic bond, this is easy to see. An electron is transferred from one atom to another, created two ions that are then attracted to one another.

Electronegativity occurs within a covalent bond. This is because, despite the electrons being shared, they will not be shared evenly. Bigger atoms will exert a stronger pull on electrons, due to the larger positive charge from having a greater number of protons in their nucleus. This means that electrons will be slightly displaced towards the bigger atom, creating electronegativity in the bond.

Consider the example of water. The oxygen atom is much larger than the hydrogen atoms, and will therefore attract the electrons of the covalent bonds towards itself, and away from the hydrogen atoms. This means that oxygen has a greater electronegativity.

When considering electrons in terms of bonding, we should think of charge clouds. The term electron density is used to describe how negative charge is distributed in the molecule.

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The Pauling Scale

The Pauling scale is used to measure electronegativity and runs from 0 to 4.

The noble gases do not have a Pauling Scale number, as they, generally speaking, do not form covalent bonds.

Electronegativity depends on a number of factors:

The charge of the nucleus

The distance between the nucleus and outer electrons

The shielding effect of the nuclear charge by the inner electrons

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Polarity of Covalent Bonds

Polarity will occur in a covalent bond when there are atoms of two different elements, with different Pauling numbers, bonded together.

If we have a bond between 2 Chlorine atoms, for example, they both have the same electronegativity so the electrons in the covalent bond will be shared equally.

When the covalent bond is between two different atoms with different Pauling numbers, the bond will appear closest to the atom with the greatest electronegativity.

This can be seen in water. The Oxygen atom has a greater electronegativity than the Hydrogen atoms, and will therefore attract electrons towards it. This creates a polar covalent bond.

The important thing about this bond is that it will create a small negative charge on the Oxygen atom (because the electrons are closer to it) and a small positive charge on the Hydrogen atoms (because the electrons are further from it).

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Forces Between Molecules

Atoms within molecules are held together by strong covalent, ionic or metallic bonds.

Molecules, and other individual atoms, are attracted together by much weaker forces. These are known as intermolecular forces.

If these forces are strong enough, then the substance will exist as either a solid or a liquid.

3 main types of intermolecular forces:

Van der Waals forces (Weakest)

Dipole-dipole forces

Hydrogen bonding (Strongest)

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Van der Waals Forces

There are 3 Van der Waals forces. These forces dictate how atoms are attracted to one another, and can be used to explain the state of many elements.

Even if the overall charge of a substance is neutral, all atoms and molecules consist of positive and negative charges. These positive and negative charges will create attractions between molecules.

In order to grasp Van der Waals forces, you must be comfortable with the concept of electron orbitals. At any given moment the electrons can be located anywhere around their orbit of an atom, and because they are in constant motion, their position is always changing.

Wherever they are in their orbit, they will produce temporary dipoles, meaning there will be more of a negative charge where the electrons are located, and a slight positive charge away from them.

At any given instant, one atom or molecule will be attracted to another because of this temporary dipole, but the next instant that dipole will change position and the attraction will cease. After this however, another attraction will form between that molecule and another. This process is constantly repeating, as the electrons are in constant motion.

Van der Waals forces:

They act between all atoms and all molecules, all of the time

They act in addition to other intermolecular forces

The dipole effect is caused by the movement of the electron cloud, so the more electrons present, the greater the forces involved

Larger atoms and molecules have higher Van der Waals forces, meaning they will have higher boiling points. Van der Waals forces can also be known as induced dipole bonds.

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Dipole-Dipole Forces

These forces only exist between certain molecules, and are stronger than Van der Waals forces.

These forces exist between molecules which have permanent dipoles.

Hydrogen Chloride has a permanent dipole because Chlorine is more electronegative than Hydrogen. This difference in electronegativity is what causes the permanent dipole.

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London Dispersion Forces

The London dispersion force is the weakest of all intermolecular forces. It is a temporary attractive force that results when the electrons in two adjacent atoms occupy positions that make the atoms form temporary dipoles.

Unequal sharing of electrons causes rapid polarisation and counter-polarisation of the electron cloud, forming short lived dipoles. These dipoles interact with the electron clouds of neighbouring molecules, forming more dipoles.

The attractive interaction of these dipoles are called dispersion or London dispersion forces. These forces are weaker than other intermolecular forces. They do not extent over long distances. The strength of these interactions within a given molecule depends directly on how easily the electrons in the molecule can move. Large molecules in which the electrons are far from the nucleus are relatively easy to polarise, and therefore possess greater dispersion.

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Hydrogen Bonds

Hydrogen bonds have some characteristics of a Dipole-Dipole bond, and some of a covalent bond. They are the strongest of intermolecular forces.

Hydrogen bonds form when a hydrogen atom is located between two strongly electronegative atoms.

We also need a strongly electronegative atoms with a lone pair of electrons covalently bonded to a Hydrogen atom. Water is a perfect example. Oxygen is far more electronegative than Hydrogen.

There are two reasons why this bond is much stronger than a normal dipole-dipole bond:

The oxygen atoms have a lone pair of electrons

The hydrogen atoms are very deficient in electrons because hydrogen is much less electronegative than oxygen

Hydrogen bonds form between water molecules because the lone pair of electrons on one water molecule, will be very attracted to the heavily electron deficient hydrogen atoms of another water molecule.

The importance of Hydrogen bonds

Hydrogen bonds are only around 10% the strength of a covalent bond, but despite this, they have a very significant effect. Their relative weakness means that they can form and break under conditions where a covalent bond would be unable to do so.

When we consider water, the hydrogen bonds that form in its liquid state are able to form and break easily as the molecules are in motion.

When water freezes to form ice, the molecules are no longer in motion, and it is the up to hydrogen bonds to hold the molecules in place. When this happens, a three dimensional structure similar to diamond forms.

When this structure forms, the molecules need slightly more room than when they're in water, and this is why water expands when it freezes. This also means that ice becomes less dense than water as the molecules are more spaced out.

Some reasons as to why hydrogen bonds are vital to life on earth:

They attribute to the relatively high boiling point of water, keeping it as a liquid at normal temperatures

They attribute to the specific heat capacity of water

They attribute to the formation of ice, which can insulate rivers and lakes and can help underwater life survive.

Gives water its surface tension

Bonds hold DNA together