Element- substance cannot be broken down
atomic number - number of protons
atomic mass- sum of protons and neutrons
atom - smallest uncharged particle of an element.
isotopes - same protons, different neutrons. same atomic number, different atomic mass.
relative atomic mass (R.A.M) - average mass of an elements isotopes relative to one-twelfth of Carbon-12 atom.
compound - 2 or more elements chemically joined.
molecule- smallest part of covalent compound.
ion - charged particle. atom that has lost or gained one or more electrons.
relative molecular mass (R.M.M) or relative formula mass - sum of all RAM's.
vertical = groups
horizontal = periods
similar physical and chemical properties in groups.
Physical: conduct electricity, malleable, ductile, solid at room temp.
Chemical: form +ve Cations, more reactive down a group,
Chemical: form -ve Anions. form covalent compounds with other non-metals. more reactive up a group.
Between metals and non-metals.
Chemical: do not form ions. semi-conduct electricity
Force and Energy
1) Strong Nuclear - important in nucleus
2) Weak Nuclear - important in nucleus
3)Gravitational - weak
Strength of force depends on:
- size of charge. bigger=stronger
- distance between centres of objects. smaller=stronger.
Energy is needed to break bonds.
More is needed to seperate small largely charged ions, less for big lower charged ions.
Energy is always converted. never lost or made.
to make bonds energy is released.
Endothermic / Exothermic
- chemical energy to heat
- temperature increases.
- enthalpy change is -ve.
- heat is absorbed.
- temperature decreases.
- enthalpy change is +ve.
Calculating amounts - The mole
- 1 mole is the amount of substance that contains as many particles as atoms in exactly 12g of Carbon-12 atoms.
- number of atoms in 12 g of Carbon-12 is 6.02 x 10 to the power of 23. known as Avogadro's Constant (L).
- mass in g of one mole of substance is known as the Molar Mass.
- Molar Mass = R.A.M. or R.M.M in grams
- To work out number of moles = mass of substance / molar mass
- amount of substance in a solution = volume of solution x concentration of solution
- Concentration also expressed in parts per million (ppm). amount by volume or mass.
- ppm= (number of parts of chemical / number of parts it is contained in) x 10 to power of 6
- this is the simplest whole number ratio of elements in a compound.
- different compounds may have same empirical formula
- e.g. Ethyne (C2H2) and Benzene (C6H6) both have emprical formula (CH).
1) write out element symbols present.
2) write underneath mass or percentage given in question
3) divide by atomic number of each element.
4) divide all by smallest answer.
5) find the smallest whole ratio *if numbers are not whole numbers you must times by 2 or 3 to give whole numbers. do not round.*
6) write out the formula.
- The Molecular formula of a covalent or ionic compound shows how many atoms/ions of each element combine to make that compound. exact multiple of empirical formula.
Molecular formula = empirical formula x (molecular mass / empirical formula mass)
- Balanced equations tell us what is reacting and the proportions in which they are reacting.
- To balance put a number in front of the compound you want to multiply by. (Change coefficients)
State symbols are very important and should be included for all:
- (s) for Solids
- (l) for Liquids
- (g) for Gases
- (aq) for Aqueous solutions.
Ionic Compounds are made of Cations and Anions (CATions are PUSSitive).
Aqueous solutions of metal compounds are ionic.
Polyatomic ions contain more than one element. (e.g. NH4 +)
How to work out the formula of an ionic compound:
1)If numerical charges on the ions is the same, the ratio is one cation to one anion.
2)If the numerical charges are not the same, the numbers of each ion have to be worked out so that the total positive charge= the total negative charge.
- write down formulae of ions including their charges.
- find lowest common multiple of two charges. equals the charge on the ion multiplied by numbers of that ion in the formula.
- work out how many of each ion is needed to give this total charge.
e.g. Formula of Aluminium Oxide.
ions are Al 3+ and O 2-
lcm of 3 and 2 is 6
for charge of +6, two Al 3+ ions needed; for charge of -6, 3 O 2- ions needed.
so formula is Al2O3.
Alternative method is "swapping over method".
If charges not same, number of cations is equal to value of charge on the anion and the number of anions is equal to charge on cation.
Hydrated ionic compounds
- Many solid ionic compounds contain water of crystallisation.
- Number of water molecules in formula of a hydrated compound is given by a full stop then the number of molecules of water. e.g. Na2Co3.10H2O.
Molar Masses of ionic compounds
- calculated in same way as molar masses of molecular substances.
- if the compound contains water of crystallisation the mass of water must also be taken into account.
Dissolving ionic substances
- Many ionic compounds are soluble in water because of strong forces of attraction between positive cation and slightly negative oxygen atom in a water molecule and between the negative anion and slightly positive hydrogen atom in a water molecule.
- equations representing dissolving of ionic compounds have the formula of the compound on the left and seperate ions on the right.
e.g. NaCl (s) + aq --> Na+ (aq) + Cl- (aq)
Solubility of Ionic Solids
- All group 1 metal compounds are soluble in water.
- All ammonium compounds are soluble in water.
- All nitrates are soluble in water.
- All chlorides are soluble in water -Except from silver chloride and lead(II) chloride
- All Sulphates are soluble in water- Except from strontium sulphate, barium sulphate, lead(II) sulphate and Calcium sulphate is only very slightly soluble.
Some ionic compounds are mostly insoluble in water:
- All carbonates are insoluble in water- Except from all group 1 carbonates and ammonium carbonate.
- All hydroxides are insoluble in water - Except from group 1 hydroxides, ammonium hydroxide and barium hydroxide. Calcium and strontium hydroxides are slightly soluble.
The rules are as follows:
- write the full equation and balance it. then write another equation using rules 2 3 and 4.
- for dissolved ionic substances, write the ions seperately.
- for all solids (ionic or not) liquids and gases write the full formula
- cross out all "spectator ions" (ions that appear on both sides of an equation in the same state)
e.g. reaction of copper(II)sulphate and sodium hydroxide.
1) CuSO4(aq) + NaOH(aq) --> Cu(OH)2 (s) + Na2SO4 (aq)
CuSO4(aq) + 2NaOH(aq) -->Cu(OH)2 (s) + Na2SO4 (aq)
2 and 3)
[Cu]2+(aq) + [SO4]2-(aq) + [2Na]+(aq) + [2OH]-(aq) --> Cu(OH)2 (s) + [2Na]+ (aq) + [SO4]2- (aq)
4) Cu2+ (aq) + 2OH- (aq) --> Cu(OH)2 (s)
balanced chemical equations are used to work out how much reactant is needed to make a certain amount of product. there are 3 steps:
- calculate moles of reactant
- use stochiometry (ratio) of equations to calculate moles of product.
- convert moles to mass
- e.g. How many tonnes of limestone (calcium carbonate) have to be heated to produce 100 tonnes of quicklime (calcium oxide)?
- Balanced equation: CaCO3 (s) --> CaO (s) + CO2 (g)
- amounts involved: 1 mole ---> 1 mole
- molar masses: 100g --> 56.0 g
- divide both by 56: 1.79 g --> 1.0 g
- multiply by 100: 179 t --> 100 t
- So... 179 tonnes of limestone must be heated.
Chemical Equations from Experimental Data
By measuring masses of reactants and products, it is possible to confirm the balanced equation for the reaction.
- e.g When 11.2 g of iron reacted with excess chloride, 32.0g of iron(II) chloride were formed. what is the equation of this reaction?
- Amount of iron: 11.2 / 56 = 0.200 mol
- amount of iron (III) chloride: 32.0/ 162.5 = 0.197 mol
- The limitation of technique and accuracy of measurment leads to the conclusion that 2 moles of iron in excess produces 2 moles of iron (III) chloride.
- so the equation must be:
- 2Fe + 3Cl2 --> 2FeCl3
Reactions With Gases
- The molar volume of a gas is the volume occupied by 1 mol of the gas under specified conditions of temperature and pressure.
- normally 24 dm ^3 mol ^-1 at standard temp of 298 K (25 degrees celcius) and pressure (1 atm)
- moles of gas: volume of the gas / molar volume
- no atoms are gained or lost in a reaction but it is unusual to gain the theoretical maximum amount of a product:
-the reaction may be reversible
- some reactants or products may be left behind in apparatus e.g. when filtering
- some reactants may react in ways different to the expected reaction (side reactions producing by products).
- amount of product obtained = yield.
- percentage yield = (actual mass of product / theoretical mass of product) x 100%
Yields in Salt Preparation
- Salts are ionic compounds.
- cation is metal ion or ammonium ion, anion is acid.
- double salt has more than one cation or anion. form when a solution of two simple salts crystalises to form a single substance.
- salts can be produced by :
- neutralising acids with an alkali, carbonate or metal oxide or hydroxide.
- reacting acids with reactive metals
- A soluble salt must be crystallized from a saturated solution - concentrate the solution by driving off some water, leaving to evaporate and then filtering.
- An insoluble salt forms a precipitate and can be filtered off, washed and dried.
- A measure of the amount of starting materials that end up as useful products. NOT the same as yield.
- calculated using the balanced equation assuming it produces 100% yield.
- Atom Economy = (mass of atoms in required product / total mass of atoms in reactants) x 100%
- A reaction with a high atom economy makes use of most of the atoms of the reactions, with few wasted as by-products
- this reduces the amount of waste products a company deals with thus reducing the cost of waste treatment.
- atom economy can be improved by finding uses for any by-products.