1. Lose electrons easily as you go down the group because the outer shell is further from the nucleus - less attracted to the nucleus - therefore more reactive
2. need to gain one electron in its outer shell, and the further from the nucleus it is the harder it is to attract another electron
3. Chemical Bonding
3.1. Ionic bonding
3.1.1. Ions are charged particles - atoms have to lose or gain electrons to form them. They are trying to get a full outer shell
3.1.2. Positive ions - Cations
3.1.2.1. Groups 1, 2, 3
3.1.2.1.1. Gr 1 will for 1+
3.1.2.1.2. Gr 2 will form 2+
3.1.2.1.3. Gr 3 will form 3+
3.1.3. Negative ions - Anions
3.1.3.1. Groups 5, 6, 7
3.1.3.1.1. Gr 5 will form 3-
3.1.3.1.2. Gr 6 will form 2-
3.1.4. Tricky ions
3.1.4.1. Ag +
3.1.4.2. Cu 2+
3.1.4.3. Fe 2+
3.1.4.4. Pb 2+
3.1.4.5. Zn 2+
3.1.4.6. H +
3.1.4.7. Hydroxide OH -
3.1.4.8. Ammonium NH4 +
3.1.4.9. Carbonate CO3 2-
3.1.4.10. Nitrate NO3 -
3.1.4.11. Sulfate SO4 2-
3.1.5. Transfer of electrons
3.1.5.1. metal loses electrons to become a positive ion
3.1.5.2. non metal gains electrons to become a negative ion
3.1.5.3. these ions are oppositely charged and therefore are strongly attracted to each other by electrostatic
3.1.5.3.1. This attraction is an IONIC BOND
3.1.5.3.2. melting and boiling points increase as the relative molecular mass (Mr) increases
3.1.6. Ionic compounds
3.1.6.1. Overall charge is 0, but made up of a negative and a positive part
3.1.6.1.1. eg. Ca 2+ , NO3 - when bonded become Ca(NO3)2 to become neutral
3.1.6.2. All form in a similar way
3.1.6.2.1. all end up with full outer shells
3.1.6.3. Properties
3.1.6.3.1. Compounds with ionic bonding always have giant lattice structure
3.1.6.3.2. Electrostatic attractions between ions is very strong- therefore high melting and boiling points
3.1.6.3.3. not electrical conductors while solid
3.1.6.3.4. when melted or dissolved in water they conduct electricity
3.2. Covalent Bonding
3.2.1. Covalent Substances
3.2.1.1. Simple molecular
3.2.1.1.1. Weak intermolecular forces between multiple covalent bonds
3.2.1.1.2. Melting and boiling points low
3.2.1.1.3. intermolecular forces are stronger between particles with larger Mr s
3.2.1.1.4. Don't conduct electricity
3.2.1.1.5. Usually liquid or gas or easily melted solid
3.2.1.2. Giant covalent
3.2.1.2.1. All atoms bonded together with strong covalent bonds - Lattice structure
3.2.1.2.2. Very high melting and boiling points
3.2.1.2.3. Don't conduct electricity
3.2.1.2.4. usually insoluable
3.2.1.2.5. eg. Diamond and Graphite
3.2.2. When atoms share electrons to fill up their outer shells (to become stable)
3.2.3. strong electrostatic attraction between the negatively charged shared electrons and the positively charges nuclei
3.2.4. Atoms can share more than one electron with more than one other atom
4. Periodic table
4.1. Arrangement
4.1.1. In order of atomic number
4.1.2. Periods - The horizontal groups. in order of amount of electrons in the outer shell within each period
4.1.3. Groups - vertical columns, all of each group has the same amount of outer shell electrons (eg group 2 has 2 outer shell electron) and therefore share chemical properties
4.2. Group 0
4.2.1. Noble gases
4.2.1.1. Inert colourless gases
4.2.1.1.1. Don't react much (stable - have full electron shells.)
4.2.1.1.2. Exist as single gases, unlike oxygen and hydrogen
4.2.1.2. He, Ne, Ar, Kr, Xe, Rn
4.2.1.2.1. Non metals
4.3. Group 1
4.3.1. Alkali metals
4.3.1.1. Very reactive
4.3.1.1.1. Reactivity increases as you go down the group
4.3.1.1.2. React with water
4.3.1.1.3. React with oxygen to form metal oxides
4.3.1.1.4. soft to cut
4.3.1.2. Li, Na, K, Rb, Cs, Fr
4.3.1.2.1. 1 electron in outer shell
4.4. Group 7
4.4.1. Halogens
4.4.1.1. F, Cl, Br, I, At
4.4.1.2. As you go down the group, the elements have a darker colour and higher boiling point
4.4.1.2.1. boiling point increases going down the group
4.4.1.3. Reactivity decreases going down the group
5. Reactivity series
5.1. The reactivity series lists metals in order of reactivity
5.1.1. Potassium
5.1.2. Sodium
5.1.3. Lithium
5.1.4. Calcium
5.1.5. Magnesium
5.1.6. Aluminium
5.1.7. Carbon
5.1.8. Zinc
5.1.9. Copper
5.1.10. Silver
5.1.11. Gold
5.2. Displacement reactions
5.2.1. These occur when a more reactive metal reacts with a less reactive metal's oxide, and displace it from the oxide
5.2.1.1. eg. iron oxide + aluminium -> aluminium oxide + iron
5.2.1.2. This is a redox reaction- the metal is oxidised and the displaced metal ion is reduced
5.2.2. this works the same with metals and less reactive metal's salts.
5.2.2.1. if the more reactive metal is already in the oxide/salt, no reaction occurs
5.3. Redox
5.3.1. Reduction
5.3.1.1. Loss of oxygen
5.3.1.1.1. eg. copper oxide is reduced
5.3.1.1.2. 2CuO + C -> 2Cu + CO2
5.3.2. Oxidisation
5.3.2.1. Gain of oxygen
5.3.2.1.1. eg. magnesium is oxidised
5.3.2.1.2. 2Mg + O2 -> 2MgO
5.3.3. In terms of electron loss and gain
5.3.3.1. Reduction
5.3.3.1.1. when an element/ion LOSES one or more electrons in a chemical reaction
5.3.3.2. Oxidisation
5.3.3.2.1. when an element/ion GAINS one or more electrons in a chemical reaction
5.4. Rusting
5.4.1. Iron and steel are very prone to rusting. Rust is hydrated IRON (III) OXIDE. This forms when metallic iron is OXIDISED.
5.4.1.1. Weakens steel and iron.
5.4.1.1.1. Worsens with air, moisture and warmth
5.4.1.2. Fe ---> Fe 2+ + 2e-
5.4.2. Preventation
5.4.2.1. Sacrificial preventaition
5.4.2.1.1. attaching a more reactive metal (Zinc) to iron structures because it will get oxidised instead of iron.
5.4.2.2. Barrier methods
5.4.2.2.1. Using grease, oil, plastic coating or paint to create a barrier between the metal and the water and oxygen
5.4.2.3. Galvanising
5.4.2.3.1. Iron coated in a layer of Zinc
5.4.2.3.2. Works even if damaged because the zinc is more reactive
6. Chemical calculations
6.1. Balancing equations
6.1.1. both sides of a symbol equation need to have equal amounts of each element
6.1.1.1. eg. H2SO4 + NaOH -> Na2SO4 + 2H2O
6.1.1.2. becomes: H2SO4 +2NaOH -> Na2SO4 +2H2o
6.2. Relative formula mass
6.2.1. All the relative atomic masses of a compound added together
6.2.1.1. eg. MgCl2 = (24)+(35.5*2)= 95
6.2.1.2. Just add up all the mass numbers from the periodic table
6.3. Moles
6.3.1. A mole is a unit for the amount of a substance.
6.3.1.1. A mole is 6.023x10^23 molecules. This is Avogadro's number.
6.3.1.1.1. One mole of atoms or molecules of any substance will have a mass equal to its Ar or Mr in grams
6.4. Calculating masses in reactions
6.4.1. Calculating the mass of product formed from a given mass of reactant (theoretical yeild)
6.4.1.1. 1. Write out the balanced equation
6.4.1.2. 2. Work out the Mr of the reactant and product you're interested in
6.4.1.3. 3. Find out how many moles there are in the substance you know the mass of
6.4.1.4. 4. Use the balanced equation to work out how many moles there'll be of the other substances
6.4.1.5. 5. use number of moles to calculate mass
6.4.1.5.1. this is your theoretical yield- how much product you would get if you got 100% of the yield
6.4.2. Percentage yeild
6.4.2.1. the actual amount of product you get compared to the theoretical yield
6.4.2.1.1. percentage yield = (actual yield / theoretical yield )* 100
6.4.2.1.2. this can be used to find how efficient (or profitable) a reaction is
6.5. Empirical and molecular formulae
6.5.1. The empirical formula gives you the smallest whole number ratio of atoms in a compound
6.5.1.1. 1. list all the elements in the compound
6.5.1.2. 2. write their experimental masses under them
6.5.1.3. 3. Find the moles of each element
6.5.1.4. 4. Divide the numbers by the smallest amount of moles
6.5.1.5. 5.Get the ratio into its simplest whole number form
6.5.2. The molecular formula gives you the real amount of each element in a compound. This can be the same as the empirical formula
6.5.2.1. Converting from the empirical formula to the molecular formula
6.5.2.1.1. 1.Find the mass of the empirical formula
6.5.2.1.2. 2. divide the given relative molecular mass of the molecule by the mass of the empirical formula
6.5.2.1.3. 3. multiply the numbers of each element in the molecule by your answer
6.6. Water of crystallisation
6.6.1. Salts can be anhydrous or hydrated
6.6.2. All salts are a lattice of positive and negative ions, and sometimes water molecules are incorporated into the lattice (hydrated)
6.6.3. the water in a lattice is called water of crystallisation
6.6.3.1. Calculating how much water of crystallisation a salt contains
6.6.3.1.1. Perform an experiment where the salt is heated to remove the water and weigh it before and afterwards
6.6.3.1.2. 1. find the difference in mass between the hydrated and anhydrous salt (therefore the mass of water lost)
6.6.3.1.3. 2. calculate the amount of moles of water lost
6.6.3.1.4. 3. calculate the amount of moles of anhydrous salt made
6.6.3.1.5. 4. Find the ratio between the anhydrous salt and water, and round the result
7. Dangerous because it combines with the haemoglobin in your red blood cells and stops them from carrying oxygen, which can lead to fainting, comas and death.
8. Because the particles of ammonia are smaller and lighter than the HCl molecules, there therefore diffuse faster
9. Diffusion
9.1. The gradual movement of particles from an area of high concentration to an area of low concentration
9.2. Ammonia and Hydrogen Chloride
9.2.1. NH3 (ammonia) and HCl in a tube experiment. When they react they form a white ring, closer to the HCl
10. Crude oil
10.1. Crude oil
10.1.1. Crude oil is a mixture of hydrocarbons
10.1.1.1. can be separated for use through fractional distillation, where the gases enter a fractioning column and individually condense at their boiling point and are collected for use
10.1.1.2. Longer hydrocarbons have higher boiling points. they condense early on, near the bottom of the column, where it is the hottest
10.1.1.3. Hydrocarbon
10.1.1.3.1. a compound made out of hydrogen and carbon only
10.1.2. The fractions are
10.1.2.1. Refinery Gases
10.1.2.1.1. used in domestic heating and cooking (shortest chain)
10.1.2.2. Gasoline (petrol)
10.1.2.2.1. used as a fuel in cars
10.1.2.3. Kerosene
10.1.2.3.1. used as a fuel in air crafts
10.1.2.4. Diesel
10.1.2.4.1. used as a fuel in larger cars , vehicles and trains
10.1.2.5. Fuel oil
10.1.2.5.1. used as fuel for large ships and power stations
10.1.2.6. Bitumen
10.1.2.6.1. used to surface roads and roofs (longest chain)
10.2. Cracking
10.2.1. Really long chain hydrocarbons aren't that useful- they can be made useful by cracking (there is higher demand for short term hydrocarbons than for long chain ones)
10.2.1.1. Cracking produces alkenes - useful to make polymers
10.2.2. Cracking is a form of thermal decomposition
10.2.2.1. Using heat + a powdered CATALYST
10.2.2.1.1. Catalysts - Silica (SiO2) or Alumina (Al3O2)
10.2.3. During this reaction the alkane (hydrocarbon) is vaporised and breaks down when it comes into contact with the catalyst, forming a shorter alkane molecule and an alkene
10.3. Burning Hydrocarbons
10.3.1. Fuels release energy in combustion reactions
10.3.2. When you burn hydrocarbons with plenty of oxygen, they give out lots of energy
10.3.2.1. Exothermic reaction
10.3.2.1.1. hydrocarbon + oxygen -> carbon dioxide + water
10.3.2.1.2. Complete combustion
10.3.3. When there is not enough energy for complete combustion, you get incomplete combustion
10.3.3.1. as well as CO2 and water, incomplete combustion produces carbon monoxide and and soot
10.3.3.1.1. Carbon monoxide
10.3.3.2. these mix with water vapour in the clouds and when this falls as rain it causes lake to become acidic, ruining habitats, and dissolves limestone cliffs and buildings.
10.3.4. Acid rain
10.3.4.1. When fractions from crude oil are burnt they sometimes produce sulfur dioxide and nitrogen oxides
10.4. Homologous series
10.4.1. a group of compounds that can be represented by the same general formula.
10.4.1.1. molecules of a homologous series contain the same functional group, meaning they react similarly.
10.5. Alkanes
10.5.1. Alkanes are saturated hydrocarbons
10.5.1.1. General formula C n H 2n+2
10.5.1.2. Chains of carbon atoms surrounded by hydrogen atoms.
10.5.1.3. Have formed bonds with as many other atoms as they can (no double bonds or spaces), therefore are saturated.
10.5.2. Alkanes burn in combustion reactions
10.5.2.1. make up the majority of hydrocarbons in crude oil and combust completely with good oxygen supply
10.5.2.2. Halogens react with alkanes to make haloalkanes
10.5.2.2.1. Chlorine and Bromine react with alkanes in the presence of UV light.
10.5.2.2.2. In these reactions a hydrogen atom from the alkane is substituted with chlorine or bromine
10.6. Eg bromomethane
10.7. Alkenes
10.7.1. Alkenes are another type of hydrocarbon, but they contain a C=C double bond
10.7.1.1. they have a double bond between two of the carbon atoms in their chain, therefore are unsaturated as there is space for more atoms in the molecule
10.7.1.2. General formula C n H 2n
10.7.1.3. Halogens react with alkenes , forming haloalkenes
10.7.1.3.1. ethene + bromine -> dibromethane
10.7.1.3.2. This is an addition reaction because the C=C bond is split and a halogen atom is added to each of the carbons
10.8. Isomers
10.8.1. Molecules with the same molecular formula but the structural formula is different
10.8.1.1. eg. the atoms are arranged differently
10.9. Names of Alkanes/enes
10.9.1. No. of carbons
10.9.1.1. 1
10.9.1.1.1. 2
10.9.2. Stem
10.9.2.1. meth-
10.9.2.1.1. eth-
11. States of matter and diffusion
11.1. Solid
11.1.1. Strong forces of attraction
11.1.2. Vibrate in their position
11.2. Fixed position
11.3. Liqiud
11.3.1. Weak Force of attraction
11.3.2. Randomly arranged, close together
11.3.3. Move around constantly and can flow
11.3.4. Expand when heated
11.4. Gas
11.4.1. Very weak forces of attraction
11.4.2. Free to move, far apart
11.4.3. Move constantly and randomly
11.4.4. Pressure increase when heated
11.5. Changes of state
11.5.1. Solid to Liquid = Melting
11.5.2. Liquid to solid = freezing
11.5.3. Liquid to gas = Boiling or evaporation
11.5.4. Gas to liquid = Condensation
11.5.5. Solid to gas = sublimation
11.5.6. Gas to solid = reverse sublimation (deposition)
11.6. Definitions
11.6.1. Solution = mixture of solute and solvent
11.6.2. Solvent = liquid that solute dissolves in
11.6.3. Solute = substance being dissolved
11.6.4. Solubility = how much solute will dissolve in a solvent
11.6.5. Saturated solution - a solution where the maximum amount of solute has been dissolved
12. Atomic structure
12.1. Isotopes
12.1.1. different atomic forms of the same element with different amounts of neutrons
12.2. Subatomic particles
12.2.1. Electrons
12.2.1.1. Negative charge
12.2.1.2. Relative mass of 0.0005
12.2.1.3. in the shells
12.2.2. Neutrons
12.2.2.1. Neutral charge
12.2.2.2. Relative mass of 1
12.2.2.3. in the nucleus
12.2.3. Protons
12.2.3.1. Positive charge
12.2.3.2. Relative mass of 1
12.2.3.3. in the nucleus
12.2.4. Atoms are neutral overall - same number of protons and electrons
12.2.4.1. Atomic number = how many protons there are
12.2.4.2. Mass number = amount of protons + neutrons
12.3. Relative atomic mass (Ar)
12.3.1. How heavy different atoms are compared with the mass of an average carbon-12 isotope (when carbon-12 = 12)
12.3.2. Ar = (Relative mass of an isotope x relative abundance) + (Relative mass of an isotope x relative abundance)
12.3.2.1. Divided by the sum of the relative abundances
12.4. Defenitions
12.4.1. Elements = consist of one type of atom only
12.4.2. Compound = two or more elements chemically bonded
12.4.3. Mixtures = no chemical bond between parts of a mixture, can be separated (not pure)
13. Acids Bases and Salts
13.1. Acids
13.1.1. an acid is a source of hydrogen ions (H+). They are proton donors
13.1.1.1. Reactions of acids
13.1.1.1.1. acid + base -> salt + water
13.1.1.1.2. acid + metal oxide -> salt + water
13.1.1.1.3. acid + ammonia -> Ammonium salt
13.1.1.1.4. acid + metal carbonate -> salt + water + carbon dioxide
13.1.1.2. Examples
13.1.1.2.1. HCl = Hydrochoric acid
13.1.1.2.2. H2SO4 = Sulphuric acid
13.1.1.2.3. HNO3 = nitric acid
13.2. Bases
13.2.1. A substance that can neutralise an acid. Alkalis are soluble bases.
13.2.1.1. an alkali is a source of hydroxide ions (OH-) and are proton acceptors
13.3. Salts
13.3.1. Making soluble salts
13.3.1.1. Use an acid and an insoluble base (metal oxide/hydroxide)
13.3.1.1.1. Heat the acid in a water bath in a fume cupboard
13.3.1.1.2. add the base until all the acid has been neutralised and you have excess base
13.3.1.1.3. Filter off the excess base to get a solution of the salt and water
13.3.1.1.4. Heat the solution gently to evaporate off some water and then leave the salt to crystallise
13.4. Indicators
13.4.1. a dye that changes colour depending on whether its above or below a certain pH
13.4.1.1. Universal indicator
13.4.1.1.1. add the indicator to an aqueous solution and match it to a colour on the pH scale
13.4.1.1.2. pH scale
13.4.1.2. Litmus paper
13.4.1.2.1. Red in acidic solutions
13.4.1.2.2. Purple in neutral
13.4.1.2.3. Blue in alkaline
13.4.1.3. Phenolphthalein
13.4.1.3.1. Colourless in acidic solutions
13.4.1.3.2. Bright pink in alkaline
13.4.1.4. Methyl orange
13.4.1.4.1. Red in acidic
13.4.1.4.2. Yellow in alkaline
13.5. Neutralisation
13.5.1. The reaction between an acid and a base (or alkali) is called a neutralisation reaction
13.5.1.1. H+(aq) + OH-(aq) -> H2O(l)
13.5.1.2. the acid donates protons which are accepted by the base
13.5.1.3. the products are neutral (pH 7)