Ionic and covalent bonding
Definition: An Ionic bond is the electrostatic force of attraction between oppositely charged ions formed by electron transfer. 1.3 Bonding Metal atoms lose electrons to form +ve ions. Non-metal atoms gain electrons to form -ve ions. Mg goes from 1s2 2s2 2p63s2 to Mg2+ 1s2 2s2 2p6 O goes from 1s2 2s2 2p4 to O2- 1s2 2s2 2p6 Definition: covalent bond A covalent bond is a shared pair of electrons
Ionic bonding is stronger and the melting points higher when the ions are smaller and/ or have higher charges. E.g. MgO has a higher melting point than NaCl as the ions involved (Mg2+ & O2- are smaller and have higher charges than those in NaCl , Na+ & Cl-
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3.1.3.1 Ionic bonding
Ionic bonding involves electrostatic attraction between oppositely charged ions in a lattice. The formulas of compound ions eg sulfate, hydroxide, nitrate, carbonate and ammonium.
Students should be able to:
• predict the charge on a simple ion using the position of the element in the Periodic Table
• construct formulas for ionic compounds.
3.1.3.2 Nature of covalent and dative covalent bonds
A single covalent bond contains a shared pair of electrons. Multiple bonds contain multiple pairs of electrons.
Students should be able to represent:
• a covalent bond using a line
Dative covalent bonding
A Dative covalent bond forms when the shared pair of electrons in the covalent bond come from only one of the bonding atoms. A dative covalent bond is also called co-ordinate bonding. Common examples you should be able to draw that contain dative covalent bond (e.g. NH4+ , H3O+ , NH3BF3 ) The direction of the arrow goes from the atom that is providing the lone pair to the atom that is deficient Dative Covalent bonding The dative covalent bond acts like an ordinary covalent bond when thinking about shape so in NH4+ the shape is tetrahedral
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3.1.3.2 Nature of covalent and dative covalent bonds
A co-ordinate (dative covalent) bond contains a shared pair of electrons with both electrons supplied by one atom.
• a co-ordinate bond using an arrow.
Metallic bonding
Definition: A metallic bond is the electrostatic force of attraction between the positive metal ions and the delocalised electrons The three main factors that affect the strength of a metallic bond are: 1. Number of protons/ Strength of nuclear attraction. The more protons the stronger the bond 2. Number of delocalised electrons per atom (the outer shell electrons are delocalised) The more delocalised electrons the stronger the bond 3. Size of ion. The smaller the ion, the stronger the bond.Mg has stronger metallic bonding than Na and hence a higher melting point. The Metallic bonding gets stronger because in Mg there are more electrons in the outer shell that are released to the sea of electrons. The Mg ion is also smaller and has one more proton. There is therefore a stronger electrostatic attraction between the positive metal ions and the delocalised electrons and higher energy is needed to break
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3.1.3.3 Metallic bonding
Metallic bonding involves attraction between delocalised electrons and positive ions arranged in a lattice.
Bonding and structure
Ionic : electrostatic force of attraction between oppositely charged ions Sodium chloride Magnesium oxide Covalent : shared pair of electrons Simple molecular: With intermolecular forces (van der Waals, permanent dipoles, hydrogen bonds) between molecules Iodine Ice Carbon dioxide Water Methane Covalent : shared pair of electrons Macromolecular: giant molecular structures. Diamond Graphite Silicon dioxide Silicon Metallic: electrostatic force of attraction between the metal positive ions and the delocalised electrons Magnesium, Sodium (all metals) Only use the words molecules and intermolecular forces when talking about simple molecular substances Giant metallic lattice Property Ionic Molecular (simple) Macromolecular Metallic boiling and melting points high- because of giant lattice of ions with strong electrostatic forces between oppositely charged ions. low- because of weak intermolecular forces between molecules (specify type e.g van der waals/hydrogen bond) high- because of many strong covalent bonds in macromolecular structure. Take a lot of energy to break the many strong bonds high- strong electrostatic forces between positive ions and sea of delocalised electrons Solubility in water Generally good generally poor insoluble insoluble conductivity when solid poor: ions can’t move/ fixed in lattice poor: no ions to conduct and electrons are localised (fixed in place) diamond and sand: poor, because electrons can’t move (localised) graphite: good as free delocalised electrons between layers good: delocalised electrons can move through structure conductivity when molten good: ions can move poor: no ions poor (good) general description crystalline solids mostly gases and liquids solids shiny metal Malleable as the positive ions in the lattice are all identical. So the planes of ions can slide easily over one another -attractive forces in the lattice are the same whichever ions are adjacent. Four types of crystal structure: ionic, metallic, molecular and giant covalent (macromolecular). You should be able to draw the following diagrams or describe the structure in words to show the four different types of crystal. You should also be able to explain the properties of these solids. The tables earlier in the revision guide explain these properties. Ionic: sodium chloride Giant Ionic lattice showing alternate Na+ and Clions Metallic: magnesium or sodium Giant metallic lattice showing close packing magnesium ions Molecular: Iodine Regular arrangement of I2 molecules held together by weak van der Waals forces Molecular: Ice The molecules are held further apart than in liquid water and this explains the lower density of ice Macromolecular: diamond Tetrahedral arrangement of carbon atoms. 4 covalent bonds per atom Macromolecular: Graphite Planar arrangement of carbon atoms in layers. 3 covalent bonds per atom in each layer. 4th outer electron per atom is delocalised. Delocalised electrons between layers. Both these macromolecular structures have very high melting points because of strong covalent forces in the giant structure. It takes a lot of energy to break the many strong covalent bonds
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3.1.3.4 Bonding and physical properties
The four types of crystal structure:
• ionic
• metallic
• macromolecular (giant covalent)
• molecular.
The structures of the following crystals as examples of these four types of crystal structure:
• diamond
• graphite
• ice
• iodine
• magnesium
• sodium chloride.
Students should be able to:
• relate the melting point and conductivity of materials to the type of structure and the bonding present
• explain the energy changes associated with changes of state
• draw diagrams to represent these structures involving specified numbers of particles.
Shape of molecules
How to explain shape 1. State number of bonding pairs and lone pairs of electrons. 2. State that electron pairs repel and try to get as far apart as possible (or to a position of minimum repulsion.) 3. If there are no lone pairs state that the electron pairs repel equally 4. If there are lone pairs of electrons, then state that lone pairs repel more than bonding pairs. 5. State actual shape and bond angle. Remember lone pairs repel more than bonding pairs and so reduce bond angles (by about 2.5o per lone pair in above examples) Occasionally more complex shapes are seen that are variations of octahedral and trigonal bipyramidal where some of the bonds are replaced with lone pairs. You do not need to learn the names of these but ought to be able to work out these shapes using the method below e.g XeF4 e.g. BrF5 e.g I3 – e .g.ClF3 e.g. SF4 & IF4+ :X X X: : : X: : Xe has 8 electrons in its outer shell. 4 F’s add 4 more electrons. This makes a total of 12 electrons made up of 4 bond pairs and 2 lone pairs. The means it is a variation of the 6 bond pair shape (octahedral) Cl has 7 electrons in its outer shell. 3 F’s add 3 more electrons. This makes a total of 10 electrons made up of 3 bond pairs and 2 lone pairs. The means it is a variation of the 5 bond pair shape (trigonal bipyramidal) I has 7 electrons in its outer shell. 4 F’s add 4 more electrons. Remove one electron as positively charged. This makes a total of 10 electrons made up of 4 bond pairs and 1 lone pair. The means it is a variation of the 5 bond pair shape (trigonal bipyramidal)
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3.1.3.5 Shapes of simple molecules and ions
Bonding pairs and lone (non-bonding) pairs of electrons as charge clouds that repel each other.
Pairs of electrons in the outer shell of atoms arrange themselves as far apart as possible to minimise repulsion.
Lone pair–lone pair repulsion is greater than lone pair–bond pair repulsion, which is greater than bond pair–bond pair repulsion.
The effect of electron pair repulsion on bond angles.
Students should be able to explain the shapes of, and bond angles in, simple molecules and ions with up to six electron pairs (including lone pairs of electrons) surrounding the central atom.
Electronegativity and intermediate bonding
Electronegativity and intermediate bonding Definition Electronegativity is the relative tendency of an atom in a covalent bond in a molecule to attract electrons in a covalent bond to itself. F, O, N and Cl are the most electronegative atoms Factors affecting electronegativity Electronegativity increases across a period as the number of protons increases and the atomic radius decreases because the electrons in the same shell are pulled in more. It decreases down a group because the distance between the nucleus and the outer electrons increases and the shielding of inner shell electrons increases A compound containing elements of similar electronegativity and hence a small electronegativity difference will be purely covalent Formation of a permanent dipole – (polar covalent) bond A polar covalent bond forms when the elements in the bond have different electronegativities . (Of around 0.3 to 1.7) When a bond is a polar covalent bond it has an unequal distribution of electrons in the bond and produces a charge separation, (dipole) δ+ δ- ends. The element with the larger electronegativity in a polar compound will be the δ- end H – Cl + – A compound containing elements of very different electronegativity and hence a very large electronegativity difference (> 1.7) will be ionic e.g. CCl4 will be non-polar whereas CH3Cl will be polar A symmetric molecule (all bonds identical and no lone pairs) will not be polar even if individual bonds within the molecular ARE polar. Symmetric molecules The individual dipoles on the bonds ‘cancel out’ due to the symmetrical shape of the molecule. There is no NET dipole moment: the molecule is NON POLAR C H H H Cl δ+ δ- Electronegativity is measured on the Pauling scale (ranges from 0 to 4) The most electronegative element is fluorine and it is given a value of 4.0 Factors affecting electronegativity Electronegativity increases across a period as the number of protons increases and the atomic radius decreases because the electrons in the same shell are pulled in more. It decreases down a group because the distance between the nucleus and the outer electrons increases and the shielding of inner shell electrons increases
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3.1.3.6 Bond polarity
Electronegativity as the power of an atom to attract the pair of electrons in a covalent bond.
The electron distribution in a covalent bond between elements with different electronegativities will be unsymmetrical. This produces a polar covalent bond, and may cause a molecule to have a permanent dipole.
Students should be able to:
• use partial charges to show that a bond is polar
• explain why some molecules with polar bonds do not have a permanent dipole.
Van der waals’ forces
The increasing boiling points of the halogens down the group 7 series can be explained by the increasing number of electrons in the bigger molecules causing an increase in the size of the van der Waals between the molecules. This is why I2 is a solid whereas Cl2 is a gas. The shape of the molecule can also have an effect on the size of the van der Waals forces. Long chain alkanes have a larger surface area of contact between molecules for van der waals to form than compared to spherical shaped branched alkanes and so have stronger van der waals.
The increasing boiling points of the alkane homologous series can be explained by the increasing number of electrons in the bigger molecules causing an increase in the size of the van der Waals between molecules.
These are also called transient, induced dipole-dipole interactions. They occur between all simple covalent molecules and the separate atoms in noble gases. In any molecule the electrons are moving constantly and randomly. As this happens the electron density can fluctuate and parts of the molecule become more or less negative i.e. small temporary or transient dipoles form. These instantaneous dipoles can cause dipoles to form in neighbouring molecules. These are called induced dipoles. The induced dipole is always the opposite sign to the original one. Main factor affecting size of Van der waals The more electrons there are in the molecule the higher the chance that temporary dipoles will form. This makes the van der Waals stronger between the molecules and so boiling points will be greater
Van der Waals forces occur between all molecular substances and noble gases. They do not occur in ionic substances.
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3.1.3.7 Forces between molecules
Forces between molecules:
• induced dipole–dipole (van der Waals, dispersion, London) forces
The melting and boiling points of molecular substances are influenced by the strength of these intermolecular forces.
Students should be able to:
• explain the existence of these forces between familiar and unfamiliar molecules
• explain how melting and boiling points are influenced by these intermolecular forces.
Permanent dipole-dipole forces
•Permanent dipole-dipole forces occurs between polar molecules •It is stronger than van der waals and so the compounds have higher boiling points •Polar molecules have a permanent dipole. (commonly compounds with C-Cl, C-F, C-Br H-Cl, C=O bonds) •Polar molecules are asymmetrical and have a bond where there is a significant difference in electronegativity between the atoms.
Permanent dipole forces occurs in addition to van der waals forces
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3.1.3.7 Forces between molecules
Forces between molecules:
• permanent dipole–dipole forces
The melting and boiling points of molecular substances are influenced by the strength of these intermolecular forces.
Students should be able to:
• explain the existence of these forces between familiar and unfamiliar molecules
• explain how melting and boiling points are influenced by these intermolecular forces.
Hydrogen bonding
It occurs in compounds that have a hydrogen atom attached to one of the three most electronegative atoms of nitrogen, oxygen and fluorine, which must have an available lone pair of electrons. e.g. a –O-H -N-H F- H bond. There is a large electronegativity difference between the H and the O,N,F Hydrogen bonding Always show the lone pair of electrons on the O,F,N and the dipoles and all the δ – δ + charges
Hydrogen bonding occurs in addition to van der waals forces
Hydrogen bonding is stronger than the other two types of intermolecular bonding. The anomalously high boiling points of H2O, NH3 and HF are caused by the hydrogen bonding between the molecules The general increase in boiling point from H2S to H2Te is caused by increasing van der Waals forces between molecules due to an increasing number of electrons. Alcohols, carboxylic acids, proteins, amides all can form hydrogen bonds
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3.1.3.7 Forces between molecules
Forces between molecules:
• hydrogen bonding.
The melting and boiling points of molecular substances are influenced by the strength of these intermolecular forces.
The importance of hydrogen bonding in the low density of ice and the anomalous boiling points of compounds.
Students should be able to:
• explain the existence of these forces between familiar and unfamiliar molecules
• explain how melting and boiling points are influenced by these intermolecular forces.
Credits: Neil Goalby
