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Structures, Bonding and Periodicity @ AS Level

AQA: AS Level Chemistry Unit 1

Kate Woodhead

on 4 November 2014

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Transcript of Structures, Bonding and Periodicity @ AS Level

Structures, Bonding and Periodicity @ AS Level
Intermolecular Forces
There are weak interactions acting between molecules and atoms - these are called "intermolecular forces". There are three main types, and they vary in strength:
van der Waals forces - these act between all atoms and molecules
Dipole-dipole forces - these only act between certain types of molecules (ones that contain permanent dipoles)
Hydrogen bonds - these are particularly strong dipole-dipole interactions that act between certain types of molecules.
Types of Structures
Types of Bonding
Chemical bonds between atoms always involve their outer electrons.

Atoms share or transfer electrons in order to achieve a more stable electron arrangement, often similar to those of the noble gases.

There are three types of strong chemical bonds: ionic, covalent & metallic.
In a covalent bond, the electrons may not be shared evenly between the atoms. One atom may be better at attracting the electrons to itself than the other.

Electronegativity is defined as the power of an atom to attract electron density in a covalent bond towards itself. (Electron density is the term used to describe the way that negatively charged electron clouds are distributed in a molecule).
Trends in Electronegativity
Electronegativity is often given a number on a scale of 0 to 4 to gauge how electronegative an atom is - this is known as the Pauling scale, and a higher number indicates a greater electronegativity.
States of Matter
One of the most basic ideas in Science is that all matter is made up of tiny particles. These particles move, and therefore have kinetic energy. There are 3 states of matter that Chemists are interested in: solid, liquid and gas. We need to explain the energy changes between these physical states in order to understand them.
Energy Changes on Heating
Heating a substance supplies energy to that substance. This affects the arrangement, movement and spacing of the particles.

Heating solids
- particles vibrate more about a fixed position, the average distance between particles increases, so the solid expands.
Melting (also known as "enthalpy change of fusion")
- the energy supplied weakens the forces acting between the particles so they are no longer held in fixed positions, but are able to move over each other. The temperature remains constant during this process.
Heating liquids
- particles move more quickly. The distance between particles is also greater on average.
Boiling (also known as "enthalpy change of vaporisation")
- as for melting, the energy supplied greatly weakens the forces between the particles so that they are far apart and moving independently. Again, there is no change in temperature during this process.
Heating gases
- particles gain more kinetic energy, move faster and hence are further apart. Gases expand a great deal on heating.
Shapes of Molecules
Molecules and complex ions (ions comprised of more than one element) are 3D structures, and may adopt a number of different shapes.

At this level, only molecules consisting of a central atom surrounded by a number of other atoms will be considered. Remember that the atoms are joined by covalent bonds i.e. a shared pair of electrons, and that a central atom may have lone pairs of electrons that are not involved in covalent bonds.
Electron Pair Repulsion Theory
The electron pairs that surround a central atom will
repel other electron pairs
near them. Therefore, bonding pairs and lone pairs of electrons will arrange themselves
as far away from each other as possible
around a central atom, in order to reduce the repulsion between them. The shape of the molecule depends on the
number of electron pairs
surrounding the central atom. A
"dot-cross" diagram
must be used to work out the number of electron pairs.
Dot-Cross Diagrams
Periodicity explores trends and similarities in the properties of the elements of the Periodic Table. The structure of the Periodic Table can be shown in terms of the electronic structure of the elements: "s-block" elements all have their highest energy electrons in s-orbitals, "p-block" in p-orbitals, and so on.
When metal atoms form chemical bonds with each other, their outer electron levels "merge" so that the outermost electrons are no longer associated with one particular atom. The electrons form a "sea" of delocalised electrons, in which the positive metal ions sit, forming a regular lattice.
Metallic Bonding
The number of delocalised electrons depends on how many have been lost by each metal atom.
Ionic bonds occur between metals and non-metals, and are formed when electrons are transferred from the metal to the non-metal. This results in the formation of positive and negative ions.
Ionic Bonding
Positive and negative ions are attracted to each other in a compound through electrostatic attractions.
Non-metal atoms share some of their electrons so that both atoms have stable noble gas electron configurations.
Covalent Bonding
The Nature of a Covalent Bond
A covalent bond is a shared pair of electrons.
Dative Covalent Bonding
In a dative covalent bond (also called a co-ordinate bond), one atom provides both the electrons that are used in the covalent bond.
The atom accepting the electrons is electron-deficient; the atom donates its lone pair of electrons to form the bond.
Bond Polarity
"Polarity" is the property of a covalent bond - it describes the unequal sharing of electrons within the bond.

If two atoms are the same element, then electron density is shared equally between them - the bond is non-polar.
Dipole-dipole forces act between molecules that have a dipole moment. A dipole moment is the net effect of all the polar bonds within a molecule.

In a molecule such as HCl, there is a permanent dipole as Cl is more electronegative than H.

Two molecules both with dipoles will attract one another. They will "flip" to ensure they attract each other, no matter what their original orientation.
Dipole-dipole forces
All atoms and molecules contain positive and negative charges, even though they are neutral overall. The distribution of these charges is not even throughout the molecule - it is constantly fluctuating.

When the charge is unevenly distributed, a temporary dipole will occur within the molecule. However, it is instantaneous, and will change again.
van der Waals forces
Hydrogen bonds are the strongest type of intermolecular force. They form when two conditions are fulfilled:
a hydrogen atom is covalently bonded to a highly electronegative atom (N, O, F)
a neighbouring molecule has a highly electronegative atom with a lone pair of electrons
Hydrogen Bonding
Electronegativity depends on:
the nuclear charge
the distance between the nucleus and the outer shell of electrons
the shielding provided by the inner electron shells
Two atoms that have different electronegativities will have unequal shares of the electrons in the covalent bond between them. These molecules are called "polar", and the bigger the difference in electronegativities, the more polar the bond. These covalent bonds have some ionic character.
Polar Molecules
At any one time, an instantaneous dipole may influence the electron distribution within a nearby molecule, and induce a dipole in that molecule. The dipoles will then attract each other.

van der Waals forces are often called "instantaneous dipole-induced dipole forces". They occur between all atoms and molecules at all times, regardless of any other intermolecular forces that may also be present.

As the van der Waals forces depend on the fluctuation in the electron cloud of a molecule, the larger the molecule the stronger the forces, as there are more electrons.
Attracting Other Molecules
The boiling points of the noble gases show a gradual increase in boiling points as the van der Waals forces increase. However, the hydrides show a different pattern: ammonia, HF and water all have H-bonds acting between the molecules, and so have higher boiling points than the hydrides of other elements in their respective groups.
Boiling Points of the Hydrides
When water freezes, the water molecules are held in fixed positions by H-bonds. The 3D structure shares some similarities with diamond.

In order to fit the molecules into the structure, the molecules are slightly further apart than in liquid water. For this reason, ice is less dense than cold water. This has allowed life to evolve.
Structure and Density of Ice
Electrostatic forces hold the oppositely charged ions together in a giant lattice. The attractive and repulsive forces balance, and extend throughout the 3D structure.
Ionic Structures
Ionic compounds have high melting and boiling points as a lot of energy is required to break up the giant lattice of ions. They are brittle (shatter easily) when hit as like charged ions align due to the blow. They conduct electricity when molten or in solution as the ions are free to move.
Metallic Properties
Metals are good conductors of heat and electricity as the delocalised electrons are free to move throughout the structure.

Metals have high melting points, and are also strong - this depends on the charge on the metal ion, and the size of the ion.

Metals are ductile and malleable - the metal ions are still surrounded by the delocalised electrons, even after they have been hit.
These crystals consist of molecules that are held in place by one of three types of intermolecular force (van der Waals, dipole-dipole, H-bonding). These forces are relatively weak, and so molecular crystals have low melting points.
Molecular Crystals
Iodine is an example of a molecular crystal. It is soft and breaks easily, sublimes easily to form gaseous iodine molecules, and does not conduct electricity as there are no free particles (ions or electrons to carry a charge).
Diamond consists of carbon atoms only, connected via covalent bonds only, and has a giant structure. Each carbon atom uses its 4 outer shell electrons to form 4 single covalent bonds - this gives each carbon atom a tetrahedral structure (see Shapes...).
It is very hard, has a very high melting point and does not conduct electricity as there are no free charged particles (ions or electrons).
Graphite also consists only of carbon atoms, however each carbon atom is only connected to 3 other carbon atoms via covalent bonds, forming a 2D sheet of interlinked hexagons. Each carbon atom has a trigonal planar shape (see Shapes...). Between the layers, there are weak van der Waals forces, and delocalised electrons which add to the strength of the interactions.
Graphite is soft, has high melting and boiling points, and conducts electricity as the delocalised electrons are free to move between the layers.
Lone Pairs
Trends in the Properties of Period 3 Elements
Across Period 3, the properties of the elements vary.
Atomic radii
- across a period, the atomic radii of the elements
. This is because more protons are added to the nucleus, so the
nuclear charge increases
across the period. The increased charge
attracts the outer electrons more
, and as there are no additional shells to provide more shielding, the size of the atoms decrease across a period.
Ionisation Energies Across Period 3
This trend is also explored in more detail in the Atomic Structure topic.
First Ionisation Energy
- the first ionisation energies
generally increase
across a period. This is because the number of
in the nucleus
, but the
are added to the
same principle level
. The increased charge on the nucleus attracts the outer electrons more strongly, meaning more energy is required to remove one.
Melting Points of Period 3 Elements
The trend in melting points across a period can be explained by the structures adopted by each of the elements.
Metallic Elements
- the melting points increase from Na to Al as the strength of the metallic bonds increase. This is because the charge on the ion increases, so there is a greater attraction for the sea of delocalised electrons.
- this has a giant covalent structure similar to diamond, and hence has a very high melting point.
Non-Metal Elements
- these elements have simple covalent molecular structures, and their melting points are governed by the strength of the van der Waals between the molecules, which in turn depends on the size of the molecules: S8>P4>Cl2
- this has an atomic structure, and hence has the lowest melting point due to weak van der Waals forces between individual atoms,
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