C2.1 Structure and Bonding

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  • Created by: StephBea
  • Created on: 10-04-15 16:35

1.1 Chemical Bonding

  • Elements react to form compounds by gaining or losing electrons or by sharing electrons. 
  • The metal atoms of group 1 can combine the the non-metal atoms of group 7 by transferring electrons to form ions that have the electronic structures of noble gasses (very stable due to full outer shell).
  • All elements are trying to gain, lose or transfer electrons to complete a full outer shell and become stable.
  • When two non-metallic elements react together, they share electrons in their outer shell so that both have a full number of electrons in their outer shell making them stable. This is called covalent bonding.
  • When metallic elements react with non-metallic it is calld ionic bonding, the metal atom loses electrons to form a positive ions (because it now hos more protons of charge 1+ than electrons withcharge of 1-), the non-metal atom gains these electrons to form negative ions (becuase they contain more electrons than protons). The oppositly charged ions are then attract each other.
  • When two metallic elements react it is known simply as metallic bonding, they form neat layers of electrons with electrons from the outer shell of each atom forming a sea around the electrons creating strong electrostatic forces of attraction between them.
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1.2 Ionic Bonding

  • Ionic compounds are held together by strong forces between the oppositly charged ions. They form giant ionic structures (sometimes known as lattices), and the strong electrostatic forces of attraction act throughout the whole structure.
  • Giant ionic structures of ionic compounds are very regular because of the ions packet together so tightly like marbles in a box. This means that every ion in the structure is surrounded by oppositly charged ions, so is held firmly in place.
  • Sodium choloride, for example, has an equal number of sodium and chloride ions becasue each has lost or gained just 1 electron so have equal opposite charges of 1+ and 1-. This means the ions alternate making a cubic lattice.
  • The ratio of ions in the structure depends upon the charges on the ions, so in calcium chloride, where the calcium ions have a charge of Ca2+ compared to the chloride ions charges of 1-, there will need to be twice as many chloride ions for the overall charge to be neutral. This is why it has a formula of CaCl2.
  • Dot and cross diagrams can be used to represent the atoms and ions involved in forming an ionic compound. To simplify these, only the electrons in the outermost shell are shown.
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Ionic Bonding Diagrams

  • The first part of the diagrams show each atom with the structure of its electrons of its outer shell shown by either dots or crosses and its structure written below in brackets.
  • The second part of the diagrams, shown in square brackets show how each atom has become an ion by losing or gaining an electron/electrons, the charge is shown outside each bracket.
  • The electrons do not necessarily have to be shown as this can be understood from the resulting charge on each ion.

(http://mypchem.wikispaces.com/file/view/mgs.jpg/242217305/349x347/mgs.jpg)

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1.3 Formulae of Ionic Compounds

  • Charges on the ions in an ionic compound always cancel each other out, the formula of an ionic compound shows the ratio of ions within a compound. Sometimes brackets are needed to show this ratio. For example in magnesium hydroxide where the formula is; Mg(OH)2. 
  • The formula of an ionic compound is the simplest ratio of the ions in the compound and does not represent an actual molecule - ionic compounds are always giant ionic structures.
  • Ionic compounds are always neutrally charged, so if we know the charge on each ion in a compound then the formula can be worked out by balancing the charges so for every one 2+ ions in a compound there needs to be two 1- ions or for every 3+ ions there would need to be three 1- ions and so on.
  • The charge on simple elements in the main periodic table groups can be worked out from the number of the group, for transition metals (located between groups 2 and 3), the charge of the io is shown by a Roman numeral in the name of the compound, for example Fe(ii) would form an ion of charge 2+.
  • Sometimes ions are made up of more than one element, for example carbonate ions are CO3 2- and hydroxide ions are OH-. If we need to multiply these ions to form a ratio formular, they have to be first put into brackets, for example; Ca(OH)2.
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1.4 Covalent Bonding

  • A covalent bond is formed when twon atoms, one metallic and one non-metallic, share electrons. The number of covalent bonds depends on the number electrons it needs to achieve a stable electronic structure.
  • Many substances containing covalent bonds form simple molecules, but some form giant structures. (the term 'molecule' should only be used when referring to a covalently bonded structure.)
  • The atoms of non-metals need to gain electrons to achieve stable electronic structures. They can do this by sharing electrons with another atom. Each shared pair of electrons strongly attract the two atoms. 
  • Most of the atoms we work with are trying to gain a full outer shell of 8 electrons, therefore elements in group 7 can form 1 bond to make up 8 electrons, group 6 can form 2 bonds, group 5 can form 3 bonds and so on.
  • A covalent bond acts only between the two atoms it bonds together, therefore many covalently bonded substances form small molecules with just a few atoms.
  • Some atoms that can form many bonds, such as carbon, can join together to form giant covalent structures (occasionally referred to as macromolecules).
  • Covalent bonds can be represented by dot and cross diagrams showing shared electrons, or lines representing bonds.
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Covalent Bonding Diagrams

  • The first diagram shows a dot and cross diagram representing the shared electrons within the molecule with dots being the electrons of one atom and crosses representing the electrons of the other.
  • In the second diagram of a hydrocarbon, the lines represent bonds, with the double line showing a double bond.

(http://alevelnotes.com/content_images/i43_lewisCovalentBonding.gif)                          

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1.5 Metallic Bonding

  • The atoms in metals are closely packed together and arranged regularly in layers. The electrons in the highest energy level (the outside shell) are delocalised, the stron electrostatic forces between these electrons and the positively charged metal ions hold the metal together.
  • The atoms in a metallic element are all the same size, therefore they form giant metallic structures that are neatly arranged in a regular pattern.
  • When metal atoms pack together the electrons in the outside shell delocalise and can move freely between the atoms this produces a lattice of positively charged ions in a 'sea' of moving electrons. The delocalised electrons strongly attract the positive ions and hold the giant structure together.
  • It is because of these delocalised electrons that metals can conduct heat and electrocity as they are able to move and carry charge.
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Metallic Bonding Diagram

  • Diagram showing the regular, layered structure formed by positive metallic ions and their delocalised electrons in metallic bonding.

(http://www.bbc.co.uk/staticarchive/fba2965c626a450042effd6174b49257d3b3a69f.gif)

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