BONDING

Ionic bonding is when ions are held together by electrostatic attraction - holds positive and negative ions together (very strong)

Ions are formed when one/+ electrons is transferred from one atom to another.

The simplest ions are single atoms which have either lots or gained an electrons so that they've got a full outer shell.

Elements in the same group have the same no of outer electrons. They have to loose or gain the same no. to get to the full outer shell. (NOBLE GAS ARRANGEMENT)

SO4(2-) OH(-) NO3(-) CO3(2-) NH4(+) - COMPOUND IONS

Ionic compounds are made up of positively charged part & a negatively charged part.

The overall charge is 0. The charges must balance.

YOU CAN USE THE CHARGES ON THE INDIVIDUAL IONS TO WORK OUT THE FORMULA OF AN IONIC COMPOUND.

NaCl has a giant ionic lattice structure. The lattice is cube shaped - some have different shapes.

Ionic crysystals are giant lattices of ions. A lattice is a regular structure. The structures giant because its made up of the same basic repeated unit.

The structure of ionic compounds determines their physical properties..

1) Ionic compounds conduct electricity when theyre moleten/dissolved but not solid.

The ions in a liquid are free to move (carry a charge). In a solid they are in a fixed position by strong ionic bonds.

2) Ionic compounds have high melting point.

Giant ionic lattices are held together by strong electrostatic forces - lots of energy needed to overcome these forces.

3) Ionic compounds tend to dissolve in water.

Water molecules are polar. These charged parts pull ions away from the lattice, causing it to dissolve.

Molecules are groups of atoms bonded together by strong covalent bonds.

A single covalent bond contains a shared pair of electrons.

Some molecules contain double or triple bonds - these multiple bonds contain multiple shared pairs of electrons.

Covalent bonds happen betwenn non-metals.

In covalent bonding, two atoms share electrons so theyve both got full outer shells. Both the positive nuclei are attracted electrostatically to the shared electrons.

Giant covalent structures have a huge network of covalently conded atoms (macromolecular).

Carbon atoms form this because they can each form four strong, covalent bonds. There are 2 types of giant covalent carbon structures - graphite & diamond.

Dative covalent bonding - where both electrons come from one atom

The Ammonium (NH4+) ion forms when the N atom (in NH3) donates a pair of E to a proton (H+)

The structure of graphite explains its properties;

1)The weak bonds between the layers in graphite are easily broken, the sheets can easily slide over each other.

2) The delocalised electrons in graphite arent attached to any particular carbon atoms & are free to move and carry a charge. Graphite is an electrial conducter.

3) The layers are quite far apart compared to the length of the covalent bonds, graphite has a low density and is used to make strong, lightweight, sports equipment.

4) Because of the strong covalent bonds in the hexagon sheets, graphite has a v. high melting point.

5) Graphite is insoluble in any solvent - the covalent bonds in the sheets are too strong to break.

Diamond is the hardest known substance. Each carbon atom is covalently bonded to four other carbon atoms - the atoms arrange themselves in a tetrahedral shape.

Due to the strong covalent bonds, Diamond has the following properties;

1) A very high melting point

2) Its extremely hard - like Dan. Its used in diamond-tipped drills.

3) Vibrations travel easily through the stiff lattice, its a good thermal conductor.

4) It cant conduct electricity - all outer electrons are held in localised bonds.

5) Like graphite, diamond is insoluble in any solvent.

6) Diamond can be cut to form gemstones. Its structure makes it refract light - why it sparkles.

Molecular shape depends on electron pairs around the central atom - the number of pairs of electrons in the outer shell of the central atom.

Electron pairs (BP & LP) exist as charge clouds - an area where you have a really high chance of finding an electron. The electrons arent stationary, theyre constantly moving.

Electron charge clouds repel each other - they're all negatively charged.

The shape of the charge cloud affects how much it repels other charge clouds. The greatest angles are between lone pairs, bond angles between bonding pairs are often reduced (pushed together by lone pair repulsion)

• LP & LP angles are the biggest
• LP & BP angles are the second biggest
• BP & BP angles are the smallest.

To predict the shape of a molecule, youll need to know how many bonding & lone electron pairs there are on the central atom of the molecule.

1) Work out the central atom.

2) Using the periodic table, deduce the no. of electons in the outer shell of the central atom.

3) Add one to this no. for every atom the central atom is bonded to.

4) Divide by 2 to find to no. of electon pairs on the central atom.

5) Compare the number of electron pairs to the number of bonds to find the number of lone pairs & bonding pairs on the central atom.

2 EP - 0 LP = LINEAR        3 EP - 0 LP = TRIGONAL PLANAR

4 EP - 0 LP = TETRAHEDRAL      4 EP - 1 LP = TRIGONAL PYRAMIDAL    4 EP - 2 LP = BENT

5 EP - 0 LP = TRIGONAL BYPYRAMIDAL   5 EP - 1 LP = SEESAW   5 EP - 2 LP = T-SHAPED

6 EP - 0 LP = OCTAHEDRAL   6 EP - 2 LP = SQUARE PLANAR.

Electronegativity - An atoms ability to attract the electon pair in a covalent bond. Flourine is the most EN, Oxygen, Chlorine & Nitrogen are also strongly EN.

Covalent bonds may be polarised by differences in electronegativity - the bonding electrons will be pulled towards the more EN atom - this makes the bond polar.

• A covalent bond between 2 atoms of the same element is non polar - the atoms have equal EN, the E are equally attracted to both nuclei.
• Some elements, C & H have pretty similar EN - bonds between them are essentially non - polar.
• In a polar bond, diff in EN causes a permanent dipole. The greater the diff in EN, the more polar the bond.

A dipole is a difference in charge between the two atoms caused by a shift in the electron density in the bond.

If you have a molecule that contains polar bonds, the result is an uneven distribution of charge across the whole molecule - the whole molecule is polar. If the polar bonds are arranged symmetrically they cancel out & there is no resulting dipole.

Polar molecules have permanent dipole - dipole forces, theres also weak electrostatic forces of attraction between the +ve and -ve charges on neighbouring molecules. If you put a charges rod next to a jet of a polar liquid, like water, the liquid will move towards the rod - becuase polar liquids contain molecules with permanent dipoles.

Van dew Waals forces are found between all atoms and molecules -  they cause all atoms & molecules to be attratced to each other.

• E in charge clouds are always moving - at any moment the E are more likely to be on one side - a temporary dipole has been induced.
• This causes a second dipole in the opp direction - the two are then attracted to each other
• The second dipole causes yet another in a third atom, then a fourth etc etc.
• Because E are constantly moving, the dipoles are always being created & destroyed.

Iodine (I2) is a solid at room temp, VDW forces keep the I2 molecules in a lattice. Iodine atoms are held together in pairs by strong, covalent bonds to form I2 molecules. But the molecules are then held together in a molecular lattice arrangement by weak VDW attractions.

Stronger VDW forces means higher boiling points.

Larger molecules have larger electron clouds, meaning stronger van der Waals forces.

The shape of molecules also affects the strength of Van der Waals forces.

• Long, straight molecules can lie closer together than branched ones - The closer together two molecules get, the stronger the forces between them are.

You need more energy to overcome stronger intermolecular forces, so liquids with stronger VDW forces will have higher boiling points.

HYDROGEN BONDING IS THE STRONGEST INTERMOLECULAR FORCE

H bonding only occurs when hydrogen is covalently bonded to F,N or O. These are all very EN so they draw bonding E away from the H atom. The bond is so polarised & H has such high charge density (because its so small) that the H atoms form weak bonds with lone pairs of electrons on the F, N or O atoms of other molecules.

Metal elements exist as giant metallic lattice structures.

• The outermost shell of E of a metal atom is delocalised - the electrons are free to move. This leaves a postive metal ion (cation).
• The positive metal ions are attracted to the delocalised electrons. They form a lattice of closely packed +ve ions in a sea of delocalised electrons - this is metallic bonding.

Metallic bonding explains the propeties of metals

• They have high melting points because of the strong electrostatic attraction between the positive metal ion & the delocalised electrons
• The no. of delocalised E per atom affects melting point - the more there are, the stronger the bonding will be & the higher the melting point.
• The delocalised E can pass kinetic energy to each other, making metals good thermal conductors
• Metals are good electrical conductors because the delocalised electrons can move & carry a current
• Metals are insoluble (except in liquid metals) - strength of metallic bonds.

The physical properties of solids, liquids & gases depend on particles

• A typical solid has particles very close together - gives a high density & makes it incompressible. The particles vibrate about a fixed point & cant move freely.
• A typical liquid has a similar density to a solid & is virtually incompressible. The particles move about freely & randomly within the liquid allowing it to flow.
• In gases, the particles have loads more energy & are much further apart - density is pretty low & they're very compressible. The particles move about freely, with not a lot of attraction between them - they'll quickly diffuse to fill a container.

IN ORDER TO CHANGE FROM A SOLID TO LIQUID OR LIQUID TO GAS YOU NEED TO BREAK THE FORCES HOLDING THE PARTICLES TOGETHER.

COVALENT BONDS DONT BREAK DURING MELTING AND BOILING. To melt or boil a simple covalent compound you only have to overcome the intermolecular forces that hold the molecules together. You dont need to break the much stronger covalent bonds that hold the atoms together in the molecules.Simple covalent compounds have relatively low melting & boiling points. By contrast, Diamond is a giant covalent substance so you do have to break the covalent bonds between atoms to turn it into a liquid or gas.

Enthalpy change (delta H) is the heat energy transferred in a reaction at constant pressure. kJmol(-1)

Exothermic reactions give out energy (delta H is negative). Endothermic reactions absorb energy (delta H is positive)

• You need energy to break bonds, its endothermic - stronger bonds take more enrgy to break.
• Energy is released when bonds are formed, bond making is exothermic - stronger bonds release more enrgy when they form
• The enthalpy change for a reaction is the overall effect of these two changes. If you need more energy to break bonds than is released when bonds are made (delta H is +ve). If its less delta H is -ve.

Mean bond ethalpies are not exact. They are averaged out across a range of compounds.

Enthalpy changes can be calculated using mean bond enthalpies. Enthalpy change of reaction = total energy absorbed - total energy realeased.

Standard enthalpy of formation is the enthalpy change when 1 mole of a compound is formed from its elements in their standard states under standard conditions

Standard enthalpy of combustion is the enthalpy change when 1 mole of a substance is completely burned in O2 under standard conditions.

You use calorimetry to find out how much heat it given out by a reaction by measuring a temperature change.

• To find the enthalpy of combustion of a flammable liquid you burn it.
• As the fuel burns it heats the water. You can work out heat energy that has been absobred by the water if you know the mass of the water, the temp change and the specific heat capacity of water.

q=mct - formula used to calculate an enthalpy change from a calorimetry experiment

q = heat loss of gained (in joules). This is the same as the enthalpy chnage if the pressure is constant.

m = mass of water (or other solution) in the calorimeter (in g)

c = specific heat capacity of water (4.18)

t = the change in temperature (in Kelvin) of the water or solution.

Calorimtery can be used to calculate the enthalpy change for a reaction that happens in solution (neutralisation, dissolution - dissolving or displacement)

• To find the enthalpy change for a neutralisation reaction, add a knwon volume of acid to an insulated container & measure the temp.
• The add a knwon volume of alkali & record the temp change - stir the sol, to make sure its evenly heated.
• You can work out the heat needed to raise the temp of the sol. using q=mct.

You can usually assume that all solutions have the same density as water. Since 1cm3 has a mass of 1g, if you have 50cm3 of a sol then it has a mass of 50g.If you're trying to find the energy change per mole of reactant, you might need mol = conc X vol to find the no. of mol. of a substance in solution.

EXPERIMENTAL PROBLEMS WITH ALL CALORIMETRY

• Some heat will be absorbed by the container, rather than the water.
• Some heat is always lost to the surroundings

PROBLEMS WITH FLAMMABLE LIQUID CALORIMETRY

• Some combustiuon may be incomplete - less energy will be given out
• Some of the flammable liquid may escape through evaporation.

Hess's law states ; The total enthaply change of a reaction is independent of the route taken.

Enthalpy changes of formation;

• You'll need to know the HF for all the reactants and products that are compounds
• The value for HF for all elements is 0.

HR = the sum of HF (PRODUCTS) - the sum of HF (REACTANTS)

Enthalpy changes can be worked out from enthalpies of combustion.

HF + the sum of HC (PRODUCTS) = the sum of HC (REACTANTS) 

If you ever need to go along an arrow backwards in a Hess's law diagram - substract the enthalpy change that goes with that arrow