Chemistry Topic 2 – Chemical bonding and structure
GIANT STRUCTURE
2.1.1 Metallic bonding
Electrostatic forces of attraction between nuclei of cations and delocalised e –.
Electrons are said to be delocalised as they are free to move throughout structure.
Melting temperature
To melt metals, forces of attraction between cations and electrons need to be broken
Metals have giant lattice structure where many of these forces must be overcome.
Energy required to do this is usually high, so mtp are typically high.
Number of delocalised e– per cation affects mtp; group 1 – low; group 2 – high.
Metals in d-block have high mtp as they have more delocalised e – per ion.
Size of cation affects mtp; smaller cation delocalised e– closer to nucleus.
Results in increase in forces of attraction between nuclei and delocalised e –.
Electrical conductivity
When PD is applied across ends of metal, delocalised e – will be attracted.
They will move towards +ve terminal; movement of charge constitutes current.
Electrical conductivity generally increases as no. of outer-shell e – increases.
Thermal conductivity
Two factors contribute to ability of metals to transfer heat energy:
free-moving delocalised e– pass KE along metal.
cations are closely packed and pass KE from one cation to another.
Malleability and ductility
Metals can be hammered/pressed into different shapes and drawn into wire.
Properties depend on ability of delocalised e– and cations to move.
When stress is applied to metal, layers of cations may slide over one another.
Delocalised e– move with cations, prevent repulsion forces forming between layers.
2.1.1 Ionic bonding
Nature of ionic bonding
Ionic bonding – strong electrostatic attraction between oppositely charged ions.
Attractions between oppositely charged ions > repulsions between same charge ions
Strength of ionic bonding
Determine strength by calculating energy required in one mole of solid to separate
ions to infinity – at infinite distance from one another, ions can no longer interact.
Generally, smaller the ions, larger the charge on ions, stronger the ionic bond.
Trends in ionic radii
Ionic radii are difficult to measure with certainty, vary according to environment.
Several different ways of measuring ionic radii, all produce slightly different values.
To make reliable comparisons using ionic radii, they must come from same source.
, As you go down each group, ions have more electron shells; ions get larger.
Isoelectric – same number of electrons and so same electronic configuration.
Ionic radius decreases as number of protons increases.
As positive charge of nucleus increases, electrons are attracted more strongly.
High melting temperatures
Ionic solids consist of network of giant lattice network of oppositely charged ions.
Combined electrostatic forces of attraction among ions is large.
Large amount of energy required to overcome forces for ions to break free and slide.
Brittleness
If stress is applied to crystal of ionic solid, layers of ions may slide over one another.
Ions of same charge are now side by side and repel one another, crystals break.
Electrical conductivity
Solid ionic compounds don’t conduct electricity – ions aren’t free.
Molten will conduct; ions can migrate to opposite sign electrode when PD is applied.
If DC is used, compound will undergo electrolysis, ions are discharged at electrodes.
Aq will conduct and undergo electrolysis; lattice breaks down into separate ions.
Solubility
Many ionic compounds are soluble in water.
Energy required to break lattice and separate ions can be supplied by hydration.
+ve and –ve ions are attracted to H2O due to polarity that H2O possess.
Evidence for existence of ions
Most convincing evidence for existence of ions is ability of ionic compound to
conduct electricity and undergo electrolysis when molten/in aq solution.
+ve ions migrate towards –ve electrode where they gain electrons.
-ve ions migrate towards +ve electrode where they lose electrons.
2.1.3 Covalent bonding
Covalent bond – strong electrostatic attraction between nuclei of two atoms and
bonding pair of electrons, overlap of two atomic orbitals each containing single e –.
Ways in which orbitals may overlap:
end on overlap of two s-orbitals – sigma bond
end on overlap of two p-orbitals – sigma bond
sideways overlap of two p-orbitals – pi bond
end on overlap of s- and p-orbital – must be two different atoms polar covalent.
End on overlap leads to σ bond single covalent bond between two atoms.
Sideways overlap of two p-orbitals leads to π bond, can’t form until σ bond forms.
π bonds only exist between atoms that are joined by double/triple bonds.
Bond length and bond strength
Bond length – distance between nuclei of two covalently bonded atoms.
Bond strength – energy required to break one mole of bond in gaseous state.
Shorter bond, greater strength – for bonds of similar nature.
GIANT STRUCTURE
2.1.1 Metallic bonding
Electrostatic forces of attraction between nuclei of cations and delocalised e –.
Electrons are said to be delocalised as they are free to move throughout structure.
Melting temperature
To melt metals, forces of attraction between cations and electrons need to be broken
Metals have giant lattice structure where many of these forces must be overcome.
Energy required to do this is usually high, so mtp are typically high.
Number of delocalised e– per cation affects mtp; group 1 – low; group 2 – high.
Metals in d-block have high mtp as they have more delocalised e – per ion.
Size of cation affects mtp; smaller cation delocalised e– closer to nucleus.
Results in increase in forces of attraction between nuclei and delocalised e –.
Electrical conductivity
When PD is applied across ends of metal, delocalised e – will be attracted.
They will move towards +ve terminal; movement of charge constitutes current.
Electrical conductivity generally increases as no. of outer-shell e – increases.
Thermal conductivity
Two factors contribute to ability of metals to transfer heat energy:
free-moving delocalised e– pass KE along metal.
cations are closely packed and pass KE from one cation to another.
Malleability and ductility
Metals can be hammered/pressed into different shapes and drawn into wire.
Properties depend on ability of delocalised e– and cations to move.
When stress is applied to metal, layers of cations may slide over one another.
Delocalised e– move with cations, prevent repulsion forces forming between layers.
2.1.1 Ionic bonding
Nature of ionic bonding
Ionic bonding – strong electrostatic attraction between oppositely charged ions.
Attractions between oppositely charged ions > repulsions between same charge ions
Strength of ionic bonding
Determine strength by calculating energy required in one mole of solid to separate
ions to infinity – at infinite distance from one another, ions can no longer interact.
Generally, smaller the ions, larger the charge on ions, stronger the ionic bond.
Trends in ionic radii
Ionic radii are difficult to measure with certainty, vary according to environment.
Several different ways of measuring ionic radii, all produce slightly different values.
To make reliable comparisons using ionic radii, they must come from same source.
, As you go down each group, ions have more electron shells; ions get larger.
Isoelectric – same number of electrons and so same electronic configuration.
Ionic radius decreases as number of protons increases.
As positive charge of nucleus increases, electrons are attracted more strongly.
High melting temperatures
Ionic solids consist of network of giant lattice network of oppositely charged ions.
Combined electrostatic forces of attraction among ions is large.
Large amount of energy required to overcome forces for ions to break free and slide.
Brittleness
If stress is applied to crystal of ionic solid, layers of ions may slide over one another.
Ions of same charge are now side by side and repel one another, crystals break.
Electrical conductivity
Solid ionic compounds don’t conduct electricity – ions aren’t free.
Molten will conduct; ions can migrate to opposite sign electrode when PD is applied.
If DC is used, compound will undergo electrolysis, ions are discharged at electrodes.
Aq will conduct and undergo electrolysis; lattice breaks down into separate ions.
Solubility
Many ionic compounds are soluble in water.
Energy required to break lattice and separate ions can be supplied by hydration.
+ve and –ve ions are attracted to H2O due to polarity that H2O possess.
Evidence for existence of ions
Most convincing evidence for existence of ions is ability of ionic compound to
conduct electricity and undergo electrolysis when molten/in aq solution.
+ve ions migrate towards –ve electrode where they gain electrons.
-ve ions migrate towards +ve electrode where they lose electrons.
2.1.3 Covalent bonding
Covalent bond – strong electrostatic attraction between nuclei of two atoms and
bonding pair of electrons, overlap of two atomic orbitals each containing single e –.
Ways in which orbitals may overlap:
end on overlap of two s-orbitals – sigma bond
end on overlap of two p-orbitals – sigma bond
sideways overlap of two p-orbitals – pi bond
end on overlap of s- and p-orbital – must be two different atoms polar covalent.
End on overlap leads to σ bond single covalent bond between two atoms.
Sideways overlap of two p-orbitals leads to π bond, can’t form until σ bond forms.
π bonds only exist between atoms that are joined by double/triple bonds.
Bond length and bond strength
Bond length – distance between nuclei of two covalently bonded atoms.
Bond strength – energy required to break one mole of bond in gaseous state.
Shorter bond, greater strength – for bonds of similar nature.
