C2a: Bonding and Calculations.

C2a: Bonding and Calculations.

1. Atoms, Compounds and Isotopes.

   • Atoms
     • The Mass Number (top number): Total number of protons and neutrons.
     • The Atomic Number (bottom number): Number of protons. Atoms of the same element all have the same number of protons (even in ions and isotopes).
   • Compounds: 2 or more atoms that are chemically bonded.
     • Example: Carbon dioxide is a compound formed from a reaction between carbon and oxygen.
     • It is difficult to separate the two atoms back out again.
   • Isotopes (real definition): Different atomic forms of the same element, which have the same number of protons but a different number of neutrons.
     • They have the same atomic numbers but a different mass number.
     • Carbon-12: 6 protons, 6electrons and 6 neutrons.
     • Carbon-14: 6 protons, 6 electrons and 8 neutrons.

2. Ions: Bonding, Formulas and Electronic Structure.

   • Ionic Bonding: Atoms either loose or gain electrons to form charged particles (called ions) which make them strongly attracted to one another due to attraction of opposite charges.
     • A shell with just one or two electrons wants to get rid of it. These are the atoms on the left hand side of the periodic table. This turns the atom into a positively charged ion because it has more protons that electrons.
     • A nearly full shell is keen to get the extra electron. These are the elements on the right hand side of the periodic table. When they gain some electron(s) they become negatively charged ions because there are more electrons that protons.
     • Ionic compounds always have a regular giant ionic lattices. They are very strong electrostatic forces of attraction between oppositely charged atoms.
       • A single crystal of sodium chloride (salt) is one giant ionic lattice which is why they tend to be in the shape of a cuboid.
     • Ionic compounds all have similar properties: High melting points and high boiling points die to the strong attraction between the ions. They dissolve easily in water because the ions separate and are free to move in the solution. When ionic compounds melt or dissolve they carry electric current because they are free to move.
   • To work out the formula of an ionic compound we look at the charges.
     • Ionic compounds are made up of a positively charged part and a negatively charged part but the overall charge is zero.
     • Sodium chloride contains Na(+1) and Cl(-1) ions. (+1) + (-1) = 0. The charges are all balanced with one of each ion so the formula for sodium chloride is NaCl.
     • Magnesium chloride contains Mg(+2) and cl(-1) ions. Because a chloride ion only has a (-1) charge you need two of them to balance out the (+2) charge of a magnesium ion.
   • You can show the electronic structure of simple ions using diagrams.
     • NaCl (Sodium Chloride)
     • CaCl2 (Calcium Chloride)

3. Covalent Bonding and Covalent Substances

   • Covalent bonds share electrons in their outer shell. This way both atoms feel that they have a full outer shell.
     • Each covalent bond provides one extra shared electron for each atom. Each atom has to make enough covalent bonds to fill up its outer shell.
     • 7 Important Examples
       • Hydrogen atoms have just one electron. They only need one to complete the first shell.
       • Chlorine atoms only need one more electron to fill their outer shell. Covalent bonds can also be shown using lines: Cl - Cl.
       • Carbon has four outer electrons which is a half full shell so it forms four covalent bonds with four hydrogen atoms to make methane.
       • Hydrogen Chloride: Both atoms only need one more electron to complete their outer shell.
       • Nitrogen has five outer electrons so it needs to form three covalent bonds. In this instance nitrogen reacts with hydrogen making ammonia.
       • Oxygen atoms have 6 outer electrons and can sometimes form ionic compounds by taking two electrons to complete their outer shell. But to form water they share two electrons with two hydrogen atoms.
       • Oxygen shares two electrons with another oxygen atom to get a full outer shell. A double covalent bond is formed. This is drawn with two lines as shown on the bottom diagram.
   • There are two kinds of covalent substances.
     • Simple Molecular Substances: The atoms form very strong bonds to form small molecules of several atoms. The forces of attraction between these molecules is very weak.
       • The melting a boiling points are very low because the molecules are easily parted from each other. This is because its the intermolecular forces that get broken not the stronger covalent bonds.
       • Most of these substances are gases or liquids at room temperature.
       • They don't conduct electricity because there are no ions so no electrical charge.
     • Macromolecules: Giant covalent structures.
       • These are similar to giant ionic structures (lattices) except there are no charged ions.
       • All The atoms are Bonded to each other by strong covalent bonds, this means they have very high melting and boiling points.
       • They don't conduct electricity, not even when they are molten (except graphite).
       • Examples
         • Diamond: Each carbon atom forms four covalent bonds in a very rigid giant covalent structure. This makes it the hardest natural substance.
         • Silicon Dioxide (Silica): What sand is made of; every grain of sand is one giant structure of silicon and oxygen.
         • Graphite: Each carbon atom only forms three covalent bonds. This creates layers which are free to slide over each other, this makes graphite soft and slippery. The layers are held together very loosely by weak intermolecular forces. Graphite is also the only non-metal that is a good conductor of heat and electricity because each carbon atom has a free electron.

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