C2: BONDING, STRUCTURE, AND THE PROPERTIES OF MATTER

HEATING:

• 1) Heat a solid - Particles gain more energy.
• 2) Particles vibrate more - weakens the forces that hold the solid together.
• 3) At a certain temperature called the 'melting point', particles have enough energy to break free from their positions - 'melting' - turns solid into liquid.
• 4) Heat a liquid - Particles gain even more energy. 
• 5) Energy makes particles move faster - weakens and breaks bonds holding the liquid together.
• 6) At a certain temperature called the 'boiling point', the particles have enough energy to break their bonds - 'boiling' (or 'evaporating') - liquid becomes gas.

COOLING:

• 7) As gas cools, particles no longer have enough energy to overcome the forces of attraction between them - bonds form between particles.
• 8) At the boiling point, so many bonds have formed between the pas particles so gas turns into liquid - 'condensing'.
• 9) Liquid cools - particles have less energy, so move around less.
• 10) Not enough energy to overcome the particles' attraction between them, so more bonds form between them.
• 11)At the melting point, so many bonds have formed between particles so they're held in place. Liquid turns into solid - 'freezing'.

• If a temperature's below the melting point of a substance, then it'll be a solid.
• If a temperature's above the boiling point of a substance, then it'll be a gas.
• If a temperature's between the melting point and the boiling point, the substance will be a liquid.

• A chemical reaction can be shown using a word equation, or a symbol equation.
• Symbol equations can also include the state symbols next to each substance - they tell you what physical state the reactants and products are in.
• (s) = Solid
• (l) = Liquid
• (g) = Gas
• (aq) = Aqueous Solution (dissolved in water)
• Aqueous hydrochloric acid reacts with solid calcium carbonate, to form aqueous calcium chloride, liquid water and carbon dioxide gas:

• Transfer of electrons
• When a metal and a non-metal react together, the metal atom loses electrons to form a positively charged ion and the non-metal gains these electrons to form a negatively charged ion.
• These oppositely charged ions are strongly attracted to one another by electrostatic forces.
• This attraction as called an 'ionic bond'.

DOT AND CROSS DIAGRAMS:

• Show the arrangement of electrons in an atom or ion.
• Each electron is represented by a dot or a cross.
• Show which atom(s) the electron(s) in an ion originally came from.

SODIUM CHLORIDE (NaCl):

• The sodium atom gives up its outer electron, becoming an Na+ ion.
• The chlorine atom picks up the electrom, becoming a Cl- (chloride) ion.
• Now, both atoms are stable as they have full outer shells.

• When atoms lose or gain electrons, they form ions (charged particles).
• They can be single atoms (e.g. Cl-) or groups of atoms (e.g. NO  -)
• Atoms lose or gain electrons to get a full outer shell of 8 electrons, like a noble gas. Atoms with full outer shells are very stable.
• When metals from ions, they lose electrons from their outer shell, to form positive ions (cations). They lose electrons instead of gaining them, because they are in groups 1 or 2, so only have 1 or 2 electrons in their outer shell. This means that they find it in more feasible to lose/share electrons, than gain them. They have to lose a less amount, than if they were to gain electrons.
• When non-metals from ions, they gain electrons into their outer shell, to form negative ions (anions). They gain electrons instead of losing them, because they are in groups 6 or 7, so have 6 or 7 electrons in their outer shell. This means that they find it in more feasible to gain/share electrons, than lose them. They have to gain a less amount, than if they were to lose electrons.
• The number of electrons lost or gained is the same as the charge on the ion. Group 1 elements form 1+ ions because they lose 1 electron, so they become positively charged. Group 2 electrons form 2+ ions. Group 6 elements form 2- ions, because they gain 2 electrons so become negatively charged. Group 7 elements form 1- ions.

• Have a structure called a 'giant ionic lattice'.
• The ions form a closely packed regular lattice arrangement and there are very strong electrostatic forces of attraction between oppositely charged ions, in the lattice (through ionic bonding).

IONIC COMPOUND PROPERTIES:

• All have high melting points and high boiling points due to the many strong bonds between the ions. It takes a lot of energy to overcome this attraction.
• When they're solid, the ions are held in place, so the compounds can't conduct electricity.
• When ionic compounds melt, the ions are free to move so they'll carry electric current.
• Some ionic compounds dissolve easily in water. The ions separate so are all free to move in the solution - they'll carry electric current.

FINDING THE EMPIRICAL FORMULA OF AN IONIC COMPOUND:

• 1) If it's dot and cross diagram, count up the amount of atoms there are of each element.
• 2) For a 3D diagram of ionic lattice, work out the ions in the compound.
• 3) Balance, to make charge neutral.

• Non-metal atoms bond together, they share pairs of electrons - makes covalent bonds.
• The positively charged nuclei of the bonded atoms are attracted to the shared pair of electrons by electrostatic forces making covalent bonds very strong.
• Atoms only share electrons in their outer shells (highest energy levels) and each atom generally makes enough covalent bonds to fill up its outer shell, making them stable. 
• Covalent bonding happens in compounds of non-metals (e.g. H  O) and in non-metal elements (e.g. Cl  ).

PRESENTING COVALENT BONDS:

• Dot and Cross diagrams: ~ Electrons drawn in the overlap between the outer orbitals of 2 atoms are shared between those atoms. ~ Shows which atoms the electrons in a covalent bond come from. ~ Don't show relative sizes of the atoms, or how atoms arranged in space.
• Displayed formula: ~Show covalent bonds as single lines between atoms. ~ Good at showing how atoms are connected in large molecules. ~ Don't show 3D structure, or which atoms electrons have come from in that covalent bond.
• 3D Models: ~ Show the atoms, the covalent bonds, and their arrangement in space next to each other. ~ Confusing for large molecules that have lots of atoms. ~ Don't show where the electrons in the bonds have come from. 

• You can find the molecular formula of a simple compound from any of these diagrams by counting up how many atoms of each element there are.

HYDROGEN        :

• Each atom has 1 electron.
• Need 1 more electron to complete 1st outer shell.
• Form single covalent bonds.

OXYGEN            :

• Atoms each need 2 more electrons to complete outer shell.
• Share 4 electrons all together (2 electrons each).
• Creates double covalent bond.

METHANE          :

• Chlorine atom needs 4 electrons.
• Hydrogen atom needs 1 electron.
• 4 single covalent bonds.

• Substances containing covalent bonds usually have simple molecular strucutres.
• The atoms within the molecules are held togeher by very strong covalent bonds. 
• In contrast, the forces of attraction between these molecules (intermolecular forces) are very weak.
• To melt or boil a simple molecular compound, you need to break these feeble intermolecular fores and not the covalent bonds. So, the melting and boiling points are very low, because the molecules are easily parted from each other.
• Most molecular substances are gases or liquids at room temperature.
• As molecules get bigger (increase in molecular mass), the strength of the intermolecular forces increases, so more energy is needed to break them - melting and boiling points increase.

• Metals consist of a giant structure.
• The electrons in the outer shell of the metal atoms are delocalised (free to move around). These are strong forces of electrostatic attraction between the positive metal ions and the shared negative electrons.
• These forces of attraction hold the atoms together in a regular structure and are known as 'metalic bonding'.
• Metalic bonding is very strong.
• Sunstances held together by metalic bonding include metallic elements and alloys.
• Most metals are solid at room temperature: The electrostatic forces between the metal atoms and the delocalised sea of electrons are very strong, so need lots of energy to be broken - most compounds with metalic bonds have very high meltic and boiling points, so they're usually solids at room temperature.
• Metals are good conductors of electricity and heat: The delocalised electrons carry electrical current and thermal (heat) energy through the whole structure, so metals are good conductors of electricity and heat.
• Most metals are malleable: Layers of atoms in a metal slide over each other, making metals malleable / they can be bent, hammered, or rolled into flat sheets.

METALIC BONDING DIAGRAM:

ALLOYS:

• Harder than pure metals.
• Pure metals are often too soft when pure, so are mixed with other metals to make them harder.
• Most metals we use are 'alloys' (mixture of metals, or metal + element).
• Harder - more useful.
• Different elements have different sized atoms. Mixing them together distorts layers, so they can't slide - harder.

• In a polymer, lots of small units are linked together to form a long molecule with repeating sections.
• All atoms are joined by covalent bonds.
• To find molecular formula of a polymer, write down molecular formular of repeating unit in brackets, and put 'n' on the outside.
• Intermolecular forces are larger than between simple covalent molecules, so more energy is needed to break them - most polymers are solids at room temperature.
• The intermolecular forces are still weaker than ionic or covalent bonds, so they usually have lower boiling points than ionic or giant molecular compounds.
•                                  
•                                     This polymer                                           Bonds through the 
•                                     is called                                                brackets join up to the 
•                                  'Poly(ethene)'                                             next repeating unit.
•                                    (C  H   )    .
•                                    
•                                  Bit in brackets                                              'n' tells you that
•                                 is the repeating                                        the unit's repeated lots
•                                         unit.                                                of times - large number.

• Macromolecules that have all atoms bonded by strong, covalent bonds.
• Have very high boiling and melting points, as lots of energy is needed to break the covalent bonds between the atoms.
• Don't contain charged particles, so they don't conduct electricity, not even when molten (except graphite and a few others).

DIAMOND:

• Each carbon atom forms 4 covalent bonds in a rigid giant covalent structure.

GRAPHITE:

• Each carbon atom forms 3 covalent bonds to create layers of hexagons.
• Each carbon atom has 1 delocalised (free) electron.

SILICON DIOXIDE:

• Also known as 'Silica'.
• What sand is made of.
• Each grain of sand is 1 giant structure of silicon and oxygen.

• Allotropes are different structural forms of the same element in the same physical state. CARBON HAS MANY...

DIAMOND:

• A giant covalent structure made up of carbon atoms that each form
• 4 covalent bonds - Makes diamond really hard.
• The strong covalent bonds take a lot of energy to break - give diamond
• a very high melting point.
• Doesn't conduct electricity because it has no free electrons or ions.

GRAPHITE:

• Each carbon forms 3 covalent bonds creating sheets of carbon atoms arranged in hexagons.
• No covalent bonds between layers - held together weakly, so they can move over each other - makes graphite soft and slippery (ideal as lubricating material).
• High melting point - covalent bonds in layers need lots of energy to break.
• Each carbon has 1 delocalised electron (as only 3 out of the 4 outer electrons are used in bonds) conducts electricity and thermal energy.

GRAPHENE:

• A sheet of carbon atoms joined together in hexagons.
• The sheet is just 1 atom thick, making it a 2D compound.
• The network of covalent bonds makes it very strong - 1 layer of graphite.
• Very light, so it can be added to composite materials to
• improve their strength without adding much weight. 
• Contains delocalised electrons - can conduct electricity
• through the whole structure - has the potential to be used in electronics. 

FULLERENES:

• Molecules of carbon shaped like closed tubes or hollow balls.
• Have huge surface area - can be used as catalysts.
• Mainly made up of carbon atoms arranged in hexagons, but can also contain pentagons (rings of 5 carbon atoms) or heptagons (rings of 7 carbons atoms).
• Can 'cage' other molecules. The fullerene structure forms around another atom/molecule, trapping it - can deliver drugs into body.
• BUCKMINSTER FULLERENE was 1st fullerene to be discovered. Molecular formula of         and forms a hollow sphere.

NANOTUBES:

• Fullerenes can form nanotubes, which are tiny carbon cylinders. High length:diameter ratio.
• Conduct electricity and thermal energy (heat).
• Have high tensile strength (they don't break when stretched).
• Technology that uses very small particles such as nanotubes, is called 'nanotechnology'.
• Used in electronics or to strengthen materials without adding much weight (e.g. tennis racket frames).