Bonding and Periodicity

  • Created by: Pope1912
  • Created on: 11-03-15 11:54

Ionic Bonding

Ions are formed when atoms are transferred from one atom to another.

Electrostatic attraction holds positive and negative ions together. when atom are held like this it is ionic bonding.

the formula for sodium chloride is NaCl. this tells us that it is made from Na+ and Cl- ions as the ratio is one to one.

The positive charge in the compound is balanced exactly by the negative charge. This means the overall charge is zero.

Ionic crystals form giant lattices. A lattice is a regular structure.

The sturcture is made of the same basic unit repeated over and over again.

Different ionic compunds have different shaped structures but are all giant lattices.

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Ionic Bonding

Ionic compounds conduct electricity when they're molton or dissolved - but not when they are solid. The ions in a liquid are free to move and carry charge where as in a solid they are fixed by the strong ionic bonds.

Ionic compunds have high melting points. The giant lattices are held together by electrostatic forces. It takes large amounts of energy to overcome these forces resulting in very high melting points.

Ionic compounds tend to dissolve in water. Water molecules are polar - part of the molecule has a small negative charge and the other bits have a small positive charge. The water molcules pull the ions away from the lattice and causes it to dissolve.

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Covalent Bonding

Covalent bonds form between two non-metals.

Two atoms share electrons so they both have full outer shells of electrons. Both the positive nuclei are attracted electrostaically to the shared electrons.

Single, double and triple covalent bonds can form.

Giant covalent structures have a huge network or covalently bonded atoms.

Carbon atoms can form this type of structure because they can form four strong covalent bonds.

Dative covalent bonding is where both electrons come from the same atom. For example, Ammonium Ion NH4+

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Covalent Bonding

Graphite - sheets of hexagons with delocalised electrons.

Sheets of hexagons are onded together by weak Van der Waals. These are easily broken and the sheets can slide over each other.

Each Carbon only has three bonds.

The delocalised electrons in graphite aren't attatched to any particular carbon atoms and are free to move. This means electric current can flow.

The layers are quite far apart relative to the lengh of the covalent bonds. Graphite has low density and is used to make strong lightweight sports equipment.

Strong covalents bonds cause very high melting points.

Insoluble in any solvent. The covalent bonds in the sheets are to difficult to break.

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Covalent Bonding

Diamond - carbon atoms arranged in a tetrahedral shape.

Diamond has a very high melting point.

Diamond is extremely hard.

Good thermal conductor.

Cannot conduct electricity as all the outer electrons are held in localised bonds.

Won't dissolve in any solvent.

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Shapes of Molecules

Molecular shape depends on electron pairs around the central atom.

Bonding pairs and lone pairs of electrons exist as charge clouds. A charge cloud is an area where you hae a really big chance of finding an electron pair.

Electrons are all nagatively charged so the clouds repel each other as much as they can.

The shape of the charge cloud affects how much it repels other charge clouds. Lone-pair charge clouds repel more than bonding-pair charge clouds.

The greatest angles are between lone pairs of electrons, and bond angles between bonding pairs are often reduced because they are pushed together by lone pair repulsion.

Lone-pair / Lone-Pair angles are the biggest.

Lone-pair / Bonding-pair angles are the middle.

Bonding-pair / Bonding-pair angles are the smallest.

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Shapes of Molecules

Linear                                                                    Trigonal pyramidal (1lone-pair)

Trigonal Planar                                                      Triagonal bipyramidal

Tetrahedral                                                             Octrahedral

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Polarisation and Intermolecular Forces

Some atoms attract bonding electrons more than other atoms.

Electronegativity - The ability to attract the bonding electrons in a covalent bond.

Fluorine is the most electronegative element. Oxygen, Nitrogen and Chlorine are also very electronegative.

In a covalent bond between two atoms of different electronegativities, the bonding electrons are pulled towards the more electronegative atom. This makes the bond polar.

The covalent bonds in diatomic gases are non-polar because the atoms have equal electronegativities. eg / H-H Cl-Cl

Some elemants like carbon and hydrogen have pretty similar electronegativities so the bonds are essentially non-polar.

In a polar bond, the difference in electronegativity between two atoms casues a dipole. A dipole is a difference in charge caused by a shift in electron density in the bond. Greater difference in electronegativity - more polar!

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Polarisation and Intermolecular Forces

Polar molecules have permanent dipole-dipole forces.

The d+ and d- charges on the polar molecules cause weak electrostatic forces of attraction betwen molecules.

Intermolecular are forces between molecules. They are very weak.

Van der Waals or Induced dipole-dipole:

They cause all atoms and molecules to be attracted to each other.

Electrons in charge clouds are always moving really quickly. At any moment, the electrons in an atom are likely to be more to one side than the other.

The dipole can cause another temporary dipole in the opposite direction on a neighbouring atom. the two dipoles are then attracted to each other. The second dipoles are then attracted to each other.

Because the electrons are constantly moving, the dipoles are being created and destroyed all the time. The overall effect is for the atoms to be attracted to each other.

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Polarisation and Intermolecular Forces

Van der Waals forces can hold molecules in a lattice.

They are responsible for holding iodine molecules together in a lattice.

Iodine atoms are held together in pairs by strong covalent bonds to for I2 molecules.

The molecules are held together in a molecular lattice arrangement by weak van der waals.

Stonger van der waals mean higher boiling points.

Larger molecules have larger electron clouds which means stronger van der waals.

Molecules with a larger surface area have stronger van der waals as they have a more exposed electron cloud.

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Polarisation and Intermolecular Forces

Hydrogen bonding is the strongest intermolecular force.

Hydrogen bonding only takes place when hydrogen is covalently bonded to fluorine, nitrogen and oxygen. These are very electronegative and draw the bonding electrons away from the hydrogen atoms.

The bond is so polarised and hydrogen has such a high charge density because its so small, that the hydrogen atoms form weak bonds with lone pairs of electrons on the F, N or O atoms of the other molecules.

Molecules with hydrogen bonds are usually organic containing -OH or -NH groups.

Substances with hydrogen bonds have higher melting and boiling points because of the extra energy needed to break the hydrogen bonds.

Ice has more hydrogen bonds than liquid water and hydrogen bonds are relatively long. The H2O molecules in ice are further apart on average, making ice less dense than water.

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Metallic Bonding and Properties of Structures

Metal elements exist as giant metallic lattice structures.

The outermost shell of electrons of metal atom is delocalised leaving a positive metal ion.

The positive metal ions are attracted to the delocalised negative electrons. They form a lattice of closely packed positve ions in a sea of delocalised electrons.

The number of delocalised electrons per atom is the number of bonds. Therefore, the higher the number the higher the melting and boiling point as more energy is required to break the bonds.

As there are no bonds holding specific ions together the ions can slide over each other when the structure is pulled out making it malleable and ductile.

Delocalised electrons pass kinetic enegy and electrical energy making them good thermal and electrical conductors.

They are insoluble except in liquid metals.

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Metallic Bonding and Properties of Structures

Melting and boiling points depend o attraction between particles.

Closer the particles, greater the density.

Charged particles that can move freely, allow it to conduct electricity.

If a solid has a regular sructure it is called a crystal. The structure is a crystal lattice.

Covalent bonds do not break during melting or boiling.

Except in giant covalent structure like diamond and graphite.

To melt or boil a simple covalent compound you only have to overcome the intermolecular forces that hold the molecules together. This is why simple covalent compounds have relatively low melting and boiling points.

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Metallic Bonding and Properties of Structures

Bonding models match observations.

eg. the physical properties of ionic compounds provide evidence that supports the theory

They have high melting points. This tells us that the atoms are held together by a strong attraction. Positive and negative ions are strongly attracted.

They are often soluble in water but not in non-polar substances. This tells us that the particles are charged. Ions are pulled apart by polar molecules but not non-polar molecules.

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The periodic table arranged elemants by their proton number. The table is arranged in periods (rows) and groups (columns).

All the elements within a period have the same number of electron shells.

All the elements within a group have the same number of electrons in the outer shell.

The s-block elements have an outer shell electron configuration of s1 or s2.

The p-block elements have an outer shell electron configuration of s2p1 to s2p6.

The d-block elements have an outer shell electron configuration where the d shells are being filled.

Atomic radius decreases across a period. As the number of protons increases the positive charge of the nucleus increases. This means electrons are pulled closer to the nucleus making the atomic radius smaller.

The extra electrons that the elements gain across a period are added to the outer energy level so they don't really provide any extra sheilding effect.

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Melting and boiling points are linked to bond strength and structure.

Na, Mg and Al are all metals. Their melting and boiling points increase across the period because the metal-metal bonds are stronger. The bonds get stronger because the metal ions have an increasing number of delocalised electrons and a decreasing radius. This leads to a higher charge density which attracts the ions together more strongly.

Si is macromolecular, with a tetrahedral structure. The atoms are linked toogther by strong covalent bonds. A lot of energy is required to break these bonds resulting in a high melting and boiling point.

Phosphorus, sulphur and chlorine are all molecular substances. The melting and boiling points depends on the strength of the Van der Waals. These are easily overcome and therefore the melting and boiling points are low.

More atoms in a molecule mean stronger van der waals.

The ionisation energy generally increases across a period.

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