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Atoms bond in order to gain the electronic structure of the chemically stable noble gases. The noble gases, except for helium, have eight electrons in the valence shell of their atoms, which is known as a stable octet. This electronic configuration and that of helium (1s2) render the noble gases unreactive. Ionic bonding occurs between metals and non-metals. Metals tend to have low ionization energies, meaning that it does not require a great deal of energy to remove the electrons in their valence shell, of which they generally have no more than three. Non-metals usually have more than three electrons in their valence shell. When ionic bonding takes place, the valence shell electrons of a metal are attracted away from the atom by a non-metal. The metal attains a noble gas structure by losing its valence shell electrons, and the non-metal attains a noble gas structure by adding to its outer shell the electrons from the metal ion.
Sodium and chlorine
Magnesium and oxygen
What is meant by covalent bonding? Using a dot and cross diagram show how a covalent is formed using the following examples:- H2 hydrogen molecule, Cl2 chlorine molecule, HCl hydrogen chloride, CH4 methane, O2 oxygen molecule and CO2 carbon dioxide molecule.
Covalent bonding occurs when the energy required to remove the valence shell electrons of one of the atoms involved is too high. This happens in bonding between two non-metals. In order to achieve the electronic structure of a noble gas, the atoms share electrons.
HCl hydrogen chloride
CO2 carbon dioxide
Explain what is meant by a metallic bond. Why does a silver wire/element conduct electricity and a sulphur element does not.
The structure of metals is crystalline, meaning that the substance is rigid and the particles are ordered in a repeating pattern. A metal crystal consists of positive ions that are arranged in a lattice and surrounded by a sea of freely moving electrons. The shape of the metal is held together by the attraction between the negatively charged electron cloud and the positively charges ions. The electrons in the cloud come from the valence electrons of the metal atoms, which give them up and become positive ions. Electrons that form part of the electron cloud move freely and as such do not belong to the atom that has released them.
It is the freely moving electrons in the structure of a metal such as silver that allows it to conduct electricity. Ordinarily the movement of the electrons appears to be random. However, the negatively charged cloud of electrons, on application of a potential difference to the metal, is attracted towards the positive potential. In order to conduct electricity elements must have charged particles that are free to move. Sulphur atoms are covalently bonded and consequently do not have any freely moving charged particles that can carry an electric current.
What is the electro-negativity of an atom? What is a polar-bond? Use the electronegativity tables to work out if the following compounds are polar or non-polar CH4, H2O, HF, NH3 and CClF3.
The electronegativity of an atom is its ability to attract a pair of shared electrons in a covalent bond. It is possible to determine whether the bonding between two elements is covalent or ionic by looking at the electronegativity values of the two elements. If the difference between the electronegativity values of the elements is greater than 1.7 then the compound formed will usually be ionic. If the difference is less than 1.7 then the compound will be covalent. The relative polarity of covalent bonds can also be determined by comparing the electronegativity values of the elements involved. If the difference between the electronegativity values of the elements is greater than 0.3 and less than 1.7 then the compound is said to be polar. If it is less than 0.3 then it is said to be non-polar.
A poplar bond is one in which the unequal sharing of electrons, due to greater electronegativity of one of the atoms, causes the molecule to have a slightly positive charge at one end of the bond and a slightly negative charge at the other. For example, the fluorine in a molecule of hydrogen fluoride is much more electronegative than the hydrogen. As a consequence the shared electrons are less attracted to the hydrogen atom, and more attracted to the fluorine. Overall the charge of the molecule is neutral, but there is a slight positive charge at the hydrogen end of the bond and a slight negative charge at the other.
2.5-2.1 =0.4 polar
What is a hydrogen bond? How does hydrogen bonding explain the high boiling points of water, ammonia and hydrogen fluoride?
Hydrogen bonding is a type of intermolecular bonding. The bonding that occurs between covalent molecules holds them together when they are in solid and liquid states. Molecules with overall dipoles have a charge distribution that is uneven due to their polar bonds and shapes. As a result, the positive and negative ends of the molecules are attracted towards each other. This type of dipole-dipole interaction takes place in hydrogen bonding; two electronegative atoms such as fluorine, oxygen or nitrogen are bridged by a hydrogen atom. The bond can be represented as, where A and B are fluorine hydrogen or oxygen. The electrostatic forces involved in hydrogen bonding tend to be stronger than those involved in other dipole-dipole interactions but not as strong as those in ionic and covalent bonds.
Water, ammonia and hydrogen fluoride have much higher boiling points than similar compounds with comparable formulae. H2S, H2Se and H2Te, for example are hydrides whose general formulas are the same as that of water. However water is a liquid at room temperature whereas they are gases. This is explained by the fact that more energy is required to separate the hydrogen bonds between the H2O molecules, than is required to separate the weaker intermolecular forces between the other molecules. The same is true of ammonia (NH3) when compared phosphine (PH3) and arsine (AsH3), or hydrogen fluoride (HF) when compared to hydrogen bromide (HBr) or hydrogen iodide (HI).
The VSEPER - electron pair repulsion theory is used to determine the shapes of molecules.
Lone pair - lone pair > lone pair - bond pair >bond pair - bond pair repulsion.
Use the above theory to work out the shapes and bond angles of the following compounds:- H2O, CO2, BF3, NH3, SF6 and XeF4.
Bent tetrahedral - 104.5-
Linear - 180-
Trigonal planar - 120-
Trigonal pyramid - 107-
Octahedral - 90-
Square planar - 90-
Explain with the aid of diagrams where possible the structure of:-
An ionic compound (NaCl or MgO)
A simple molecular substance (I2)
A giant covalent substance (diamond, graphite)
A metallic substance (Cu or Ag)
A hydrogen bonded substance (ice)
For each of the above state the physical properties based on the intermolecular forces of attraction.
An ionic compound - NaCl
The structure of ions in an ionic compound, such as Na+ and Cl- ions in a sodium chloride crystal, is a giant lattice arrangement. The pattern and arrangement of the ions in the structure depends on the ionic compound in question. Sodium chloride has relatively simple arrangement of ions. Six chloride ions surround each sodium ion, and six sodium ions surround each chloride ion. The ions are held in place by the electrostatic forces of the ions.
Ionic compounds are solid at room temperature and have high melting and boiling points because of the high amount of energy that is required to break down the lattice structure and separate the particles. The large amount of energy required to break down the bonds is due to the fact that the ions are attracted together by strong electrostatic forces.
It is easy to shatter ionic crystals with an application of force. The crystal structure is made up of ions in layers that can slip, causing ions with like charges to align and repel each other. The repulsion of like charges causes the crystal to shatter.
They are usually soluble in water due to the fact that the ions are attracted to the positive and negative ends of the water molecule which have a dipole charge separation. The anions are attracted toward the water molecule's positive end to which they become attached, and the cations are attracted toward the water molecule's negative end. The ions are pulled away from the lattice structure by the water molecules causing the compound to dissolve.
Solid ionic compounds do not conduct electricity because they do not have freely moving charged particles. The ions are held in place by their strong electrostatic forces and cannot carry an electric current. However, in aqueous solution and when molten, ionic substances can conduct electricity because their charged particles are released from the structure and can move and carry a current.
A simple molecular substance - I2
The structure of iodine, which is solid at room temperature, is described as being face centred cubic. A unit cell of iodine comprises of a lattice structure with an iodine molecule in the centre of each of the faces of the cube and one in each corner.
Iodine is a non-polar molecule with no overall dipole moment so it is given its structure by Van der Waals' intermolecular forces. The varying density in the electron cloud of the particles causes a temporary dipole, during which the particle has a momentary, slight positive charge at one end and an equal negative charge at the other. Another temporary dipole is induced in a neighbouring particle by attracting some of its electron density at one end and giving it a slight positive charge at the other. The process of inducing temporary dipoles in neighbouring particles continues throughout the material.
The strength of the attraction due to Van der Waal's forces depends on the size of the electron density cloud, so stronger attractions are found in larger molecules. Iodine is a volatile substance because the forces holding it together are weak. It does not take much energy to break down the intermolecular forces and break up the crystal structure so solid iodine has a low melting and boiling point. However, when solid iodine is sublimated it is not the covalent bonds between the iodine atoms that are broken but the intermolecular attractive forces between the molecules.
In order for a substance to conduct electricity it must have charged particles that are free to move. Iodine does not conduct electricity in solid or molten states because its molecules have not got overall charges, nor do they possess delocalised electrons.
Iodine is almost insoluble in water due to the fact that it does not contain ions that can be pulled apart by the dipole charge separation found in the molecules of polar solvents. The attractive forces between the water molecules are stronger than the forces between the iodine molecules. The water structure is therefore impenetrable to the iodine molecules meaning that they do not dissolve readily. However, iodine is soluble in non-polar organic solvents such as tetrachloromethane and benzene. They contain covalent molecules which have weak intermolecular forces that are comparable to those found in iodine. Therefore it is easier for the non-polar solvents to solvate the iodine by penetrating its crystal structure more readily.
A giant covalent substance - diamond, graphite
Diamond is a giant covalent substance in which each carbon atom forms a single covalent bond with four other carbon atoms using each of its four valence shell electrons to make a tetrahedron. This pattern repeats itself with every carbon atom contained in the crystal. The arrangement of covalently bonded atoms in a crystal is known as a network solid.
The fact that the crystal structure is held together by strong covalent bonds makes the substance very hard. The necessity to break down the very strong carbon-carbon bonds in order to melt the substance means that it also has a very high melting point of approximately 4000- C.
Diamond cannot conduct electricity because all of the carbon valence shell electrons of each atom are involved in bonding to other carbon atoms. Without having charged particles that are free to move, substances cannot carry an electrical charge.
Diamond is also insoluble in either water or organic solvents because the electrostatic forces that hold the carbon atoms together are much stronger than the forces that exist between solvent molecules. The solvent molecules are unable to penetrate the diamond structure and solvate the carbon atoms
Graphite is another structural arrangement of carbon atoms. Like diamond, graphite is a network solid in which the carbon atoms are covalently bonded to each other. However, the carbon atoms in graphite are arranged in layers that are held together by Van der Waals' dispersion forces. Each carbon atom in a layer bonds to three other neighbouring carbon atoms using three of its valence shell electrons. The remaining un-bonded fourth valence shell electron of each carbon atom is delocalised in a similar way to the outer shell electrons in metallic bonding. The delocalised electrons form a cloud that can move throughout a layer, though electrons from different layers do not come into contact. It is the presence of the electron cloud in graphite that induces the van der Waal's dispersion forces that hold the layers of the substance together. The bonds holding the carbon atoms together in graphite are stronger than those in diamond because they have the additional strength provided by the electron cloud as well as their strong covalent bonds.
Graphite, like diamond, has a high melting point due to the high amount of energy that is required to break down the strong covalent bonds between the carbon atoms throughout a layer. Like diamond, graphite is also insoluble because the forces holding the carbon atoms together are far stronger than the forces that exist between solvent molecules.
The fact that the layers of carbon are held together by weak van der Waals' forces means that they can slide over each other, giving graphite a feel that is described as slippery. It is the slipperiness of the substance that allows for graphite to be used in pencils where the layers are rubbed off one another and onto the paper.
Graphite is less dense than diamond because the spacing between its layers (two and a half times the distance between the atoms in a layer) makes the atoms less tightly spaced than those in diamond.
The presence of a cloud of freely moving delocalised electrons in graphite allow for it to conduct electricity. The free electrons also give graphite its shiny appearance, by absorbing and re-emitting light.
A metallic substance - Cu
Copper, like silver and gold, has a face centered cubic structure in which each copper atom in the crystal is surrounded by 12 more copper atoms .
A hydrogen bonded substance - ice
(a) Draw the shapes of an s,p and d orbital for the quantum numbers (energy levels) 1 to 4. What are the maximum number of electrons that can be found in the s, p, d and f levels.
(b) Write the electronic configuration for the following using the s, p, d, f notation:- copper, copper (II), zinc, iron, iron (III), sodium, sodium ion, chlorine, chloride ion, fluoride ion, oxide ion, oxygen.
Copper - 1s2 2s2 2p6 3s2 3p6 3d10 4s1
Copper (II) - 1s2 2s2 2p6 3s2 3p6 3d9
Zinc - 1s2 2s2 2p6 3s2 3p6 3d10 4s2
Iron - 1s2 2s2 2p6 3s2 3p6 3d6 4s2
Iron (III) 1s2 2s2 2p6 3s2 3p6 3d5
Sodium - 1s2 2s2 2p6 3s1
Sodium ion - 1s2 2s2 2p6
Chlorine - 1s2 2s2 2p6 3s2 3p5
Chloride ion - 1s2 2s2 2p6 3s2 3p6
Fluoride ion - 1s2 2s2 2p6
Oxide ion - 1s2 2s2 2p6
Oxygen - 1s2 2s2 2p4
Define the term 'ionisation energy' and explain the factors that influence ionisation energy of an element. Why does the ionisation energy decrease as you go down the group, but it increases as you go across a period?
Ionisation energy is a measure of the quantity of energy that is needed to remove an electron from the outer shell of an atom in the gaseous state. Specifically it is the energy needed to remove a mole of electrons from a mole of gaseous atoms forming a mole of positive ions. The positively charged nucleus of the atom attracts the negatively charged electrons. The amount of energy required to remove an electron from the outer shell for an atom depends on the strength of the attraction between the electron and the nucleus. The greater the attraction between the electron and the nucleus, the more energy will be required to remove the electron.