In my essay on the structure and bonding in different substances, I am going to focus on electronic arrangement and their effects on the properties of the substance. METALIC BONDING Metallic bonding occurs between atoms with low electronegativity (. i. e. 1, 2 or 3 valence electrons); therefore there are many vacancies in valence shell. The crystal lattice of metals consists of ions NOT atoms.
When electron clouds overlap, electrons can move into electron cloud of adjoining atoms. The outer electrons (-) from the original metal atoms are free to move around between the positive metal ions formed (+). The metal is held together by the strong forces of attraction between the positive nuclei and the! yendelocalised! | electrons. This is sometimes described as "an array of positive ions in a sea of electrons." There is a strong electrostatic force of attraction that two neighbouring nuclei have for the shared electrons between them. - This is the metallic bond.
Each atom becomes surrounded by a number of others in a 3-D lattice, where valence electrons move freely from one valence shell to another. Each positive centre in the diagram represents all the rest of the atom apart from the outer electron, but that electron hasn't been lost - it may no longer have an attachment to a particular atom, but it's still there in the structure. Sodium metal is therefore written as Na - not Na+. The physical properties of metals This strong bonding and the fact that the particles present in metals are tightly packed in the lattice results in dense, strong materials with high melting and boiling points. The strength, elasticity and ability to bend without breaking of many metals make them useful in construction. Metals are good conductors of electricity because these 'free' or delocalised electrons carry the charge of an electric current when a potential difference (voltage) is applied across a piece of metal.
Copper, for example, is used to make the wire which goes inside electrical cables. Copper is chosen because it can be drawn into long thin wires very easily (it is ductile) and because it is a good conductor of electricity. Metals are also good conductors of heat. This is also due to the free moving electrons. Non-metallic solids conduct heat energy by! yen hotter! | more strongly vibrating atoms, knocking against! yen cooler! | less strongly vibrating atoms to pass the particle kinetic energy on. In metals, as well as this effect, the 'hot' high kinetic energy electrons move around freely to transfer the particle kinetic energy more efficiently to 'cooler' atoms.
Metals have high melting and boiling points. Strong forces of attraction exist between particles. A large amount of thermal energy is required to overcome the strong electrical forces between the positive ions and the delocalised electrons. These forces operate throughout the lattice. Metals are lustrous or have a! yen silvery surface! | but this may be easily tarnished by corrosive oxidation in air and water. The presence of free electrons causes most metals to reflect light.
Metals are malleable and ductile. The distortion does not disrupt the metallic bonding. Aluminium is very strong and can be beaten into very thin sheets. It is therefore used in aircraft manufacture.
A metal alloy is a mixture of two or more metals. Adding another element disturbs the pattern in the lattice so that the layers will not slide past each other so easily. Alloys are usually less malleable and ductile than pure metals and they tend to have lower melting points. They do, however, have other properties which make them more useful than pure metals. Some alloys that we use everyday include steel (used for a wide range of things, including knives and forks, building materials and in cars and ships. The coins that we use are all made of copper alloys.
Heat treatment: By means of heating and cooling process the strength and hardness of the metal can be changed. IONIC (ELECTROVALENT) BONDING Metal atoms are large and hold outer electrons weakly, but non-metal atoms are small and hold outer electrons strongly. This means that when metal and non-metal atoms meet, electrons are lost by the metal atom and gained by the non-metal atom. The non-metal atom gains new outer electrons making a negative ion. Electrons are gained from the metal atom until the outer shell is full. The ions formed have full outer shells.
They have the same electronic structure as a noble gas. The oppositely charged ions attract each other and an ionic bond is formed. Ionic bond: the electrostatic force of attraction between two oppositely charged ions formed as the result of electron transfer. Ionic bonds are sometimes called electrovalent bonds. The physical properties of ionic compounds All ionic substances have high melting and boiling points. A large amount of thermal energy is required to separate the ions which are bound by strong electrical forces.
Brittleness is again typical of ionic substances. Imagine what happens to the crystal if a stress is applied which shifts the ion layers slightly. Ions of the same charge are brought side-by-side and so the crystal repels itself to pieces. Many ionic solids are soluble in water - although not all. It depends on whether there are big enough attractions between the water molecules and the ions to overcome the attractions between the ions themselves.
Positive ions are attracted to the lone pairs on water molecules and co-ordinate (dative covalent) bonds may form. Water molecules form hydrogen bonds with negative ions. Ionic solids are insoluble in organic solvents. The attractions between the solvent molecules and the ions aren't big enough to overcome the attractions holding the crystal together. Ionic compounds do not conduct electricity, because there are no electrons which are free to move. Molten or aqueous form undergoes electrolysis, which involves conduction of electricity because of the movement of the ions.
Once the substance has melted or dissolved, the ions can move and carry charge. Positive ions are attracted towards the cathode: they are called cations. Negative ions are attracted towards the anode: they are called anions. This is a chemical change rather than a physical process.
The structure of a typical ionic solid - sodium chloride Compounds like this consist of a giant ionic lattice of ions. There could be billions of sodium ions and chloride ions packed together, or trillions, or whatever - it simply depends how big the crystal is. That is different from, say, a water molecule which always contains exactly 2 hydrogen atoms and one oxygen atom - never more and never less. A small representative bit of a sodium chloride lattice looks like this: the sodium ions and chloride ions alternate with each other in each of the three dimensions. Only those ions joined by lines are actually touching each other. The sodium ion in the centre is being touched by 6 chloride ions.
Sodium chloride is described as being 6: 6-co-ordinated. The pattern repeats in this way over countless ions. CO-VALENT Molecular BONDING Covalent bond: the electrostatic bonds between the shared pairs of electrons and the neighbouring nucleus. (+) "AE (-) "AE (+) Cl: Cl As well as achieving noble gas structures by transferring electrons from one atom to another as in ionic bonding, it is also possible for atoms to reach these stable structures by sharing electrons to give covalent bonds. For example, two chlorine atoms could both achieve stable structures by sharing their single unpaired electron as in the diagram. The two chlorine atoms are said to be joined by a covalent bond.
The reason that the two chlorine atoms stick together is that the shared pair of electrons is attracted to the nucleus of both chlorine atoms. Hydrogen atoms only need two electrons in their outer level to reach the noble gas structure of helium. Once again, the covalent bond holds the two atoms together because the pair of electrons is attracted to both nuclei. A double covalent bond is where two pairs of electrons are shared between the atoms rather than just one pair as above. Two oxygen atoms can both achieve stable structures by sharing two pairs of electrons as in the diagram. The double bond is shown conventionally by two lines joining the atoms.
Each line represents one pair of shared electrons. Ethane C 2 H 4 has a double bond between the two carbon atoms. Co-ordinate (dative covalent) bonding A covalent bond is formed by two atoms sharing a pair of electrons. The atoms are held together because the electron pair is attracted by both of the nuclei. In the formation of a simple covalent bond, each atom supplies one electron to the bond - but that doesn't have to be the case. A co-ordinate bond (also called a dative covalent bond) is a covalent bond in which both electrons come from the same atom.
In simple diagrams, a co-ordinate bond is shown by an arrow. The arrow points from the atom donating the lone pair to the atom accepting it. The weak bonding and the fact that the particles present in covalent molecular substance are loosely packed in the lattice results in low density, weak and soft materials. Weak forces of attraction that exist between particles require very little amount of thermal energy to be broken and so have low melting and boiling points Covalent molecular substances are bad conductors of electricity and heat because there no mobile charged particles. Molecules are not charged and electrons tightly bound to atoms or shared by atoms in covalent bonds. Electro negativity Electro negativity is a measure of the tendency of an atom to attract a bonding pair of electrons.
Consider a bond between two atoms, A and B. If the atoms are equally electronegative, both have the same tendency to attract the bonding pair of electrons, and so it will be found on average half way between the two atoms. To get a bond like this, A and B would usually have to be the same atom. (For example, H 2 or Cl 2 molecules. ) This sort of bond could be thought of as being a "pure" covalent bond - where the electrons are shared evenly between the two atoms. If B is slightly more electronegative than A, B will attract the electron pair rather more than A does.
That means that the B end of the bond has more than its fair share of electron density and so becomes slightly negative. At the same time, the A end (rather short of electrons) becomes slightly positive. This is described as a polar bond. A polar bond is a covalent bond in which there is a separation of charge between one end and the other - in other words in which one end is slightly positive and the other slightly negative. Examples include most covalent bonds. The hydrogen-chlorine bond in HCl or the hydrogen-oxygen bonds in water are typical.
They are soluble in polar solvents. H 2 or Cl 2 described as a non - polar bond is no separation of charge and the electrons are physically in the centre. They are soluble in non polar solvents. If B is a lot more electronegative than A, the electron pair is dragged right over to B's end of the bond. To all intents and purposes, A has lost control of its electron, and B has complete control over both electrons.
Ions have been formed (see ionic bonding). There are patterns of electronegativity in the Periodic Table. The most electronegative element is fluorine. If you remember that fact, everything becomes easy, because electronegativity must always increase towards fluorine in the Periodic Table. The shape of a molecule or ion is governed by the arrangement of the electron pairs around the central atom. All we need to do is to work out how many electron pairs there are at the bonding level, and then arrange them to produce the minimum amount of repulsion between them.
Two electron pairs around the central atom The only simple case of this is beryllium chloride, Bell 2. It is forming 2 bonds so there are no lone pairs. The two bonding pairs arrange themselves at 180 cX to each other, because that's as far apart as they can get. The molecule is described as being linear.
Three electron pairs around the central atom The simple cases of this would be BF 3 or BCl 3. Because it is forming 3 bonds there can be no lone pairs. The 3 pairs arrange themselves as far apart as possible. They all lie in one plane at 120 cX to each other. The arrangement is called trigonal planar. In the diagram, the other electrons on the fluorine have been left out because they are irrelevant.
Four electron pairs around the central atom There are lots of examples of this. The simplest is methane, CH 4. Four electron pairs arrange themselves in space in what is called a tetrahedral arrangement. A tetrahedron is a regular triangularly-based pyramid. The carbon atom would be at the centre and the hydrogen at the four corners.
All the bond angles are 109. 5 cX. GIANT COVALENT Network Substance are substances like diamond, graphite and silicon dioxide. The physical properties of giant covalent substances: They have a very high melting point.
Very strong carbon-carbon covalent bonds have to be broken throughout the structure before melting occurs. They very strong and hard. This is again due to the need to break very strong covalent bonds operating in 3-dimensions. Covalent network substances are brittle. If sufficient force is applied to a crystal, covalent bonds are broken as the lattice is distorted. Shattering occurs rather than deformation of a shape.
They do not conduct heat and electricity (except graphite! V see below). All the electrons are held tightly between the atoms, and aren't free to move. They are insoluble in water and organic solvents. There are no possible attractions which could occur between solvent molecules and carbon atoms which could outweigh the attractions between the covalently bound carbon atoms. The giant covalent structure of Diamond Carbon has an electronic arrangement of 2, 4. In diamond, each carbon shares electrons with four other carbon atoms - forming four single bonds.
This is a giant covalent structure - it continues on and on in three dimensions. It is not a molecule, because the number of atoms joined up in a real diamond is completely variable - depending on the size of the crystal. The structure of Graphite Graphite has a layer structure. The diagram below shows the arrangement of the atoms in each layer, and the way the layers are spaced. Carbon atom uses three of its electrons to form simple bonds to its three close neighbours. That leaves a fourth electron in the bonding level.
These "spare" electrons in each carbon atom become delocalised over the whole of the sheet of atoms in one layer. They are no longer associated directly with any particular atom or pair of atoms, but are free to wander throughout the whole sheet. The delocalised electrons are free to move anywhere within the sheet - each electron is no longer fixed to a particular carbon atom. This allows it to conduct electricity. The delocalised electrons are free to move throughout the sheets. If a piece of graphite is connected into a circuit, electrons can fall off one end of the sheet and be replaced with new ones at the other end.
Graphite has a soft, slippery feel, and is used in pencils and as a dry lubricant for things like locks. When you use a pencil, sheets are rubbed off and stick to the paper. It has a lower density than diamond. This is because of the relatively large amount of space that is "wasted" between the sheets. The structure of silicon dioxide, SiO 2 Silicon dioxide is also known as Silicon (IV) Oxide. Crystalline silicon has the same structure as diamond.
To turn it into silicon dioxide, we need to modify the silicon structure by including some oxygen atoms. Each Silicon atom is bridged to its neighbours by an oxygen atom. In conclusion, the electronic arrangement in atom and its interaction with other electrons of the compound (substance) can affect everything, ranging from its reactivity to its structure. Introduction to Atomic Structure All substances are made up of molecules made up of atoms. All atoms are made up of! yen sub atomic! | articles. The structure or all atoms follows a simple predictable structure.
All atoms have a very dense central part called the nucleus. Orbiting around this nucleus at a relatively large distance are electrons. NAME ELECTRIC CHARGE MASS LOCATION Protons +1 1 Nucleus Neutrons 0 1 Nucleus Electrons -1 0 (1/1800) (negligible) Orbiting the Nucleus.