Chemistry

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Chemistry by Mind Map: Chemistry

1. Diffusion

1.1. The gradual movement of particles from an area of high concentration to an area of low concentration

1.2. Ammonia and Hydrogen Chloride

1.2.1. NH3 (ammonia) and HCl in a tube experiment. When they react they form a white ring, closer to the HCl

1.2.1.1. Because the particles of ammonia are smaller and lighter than the HCl molecules, there therefore diffuse faster

2. States of matter and diffusion

2.1. Solid

2.1.1. Strong forces of attraction

2.1.2. Fixed position

2.1.3. Vibrate in their position

2.2. Liqiud

2.2.1. Weak Force of attraction

2.2.2. Randomly arranged, close together

2.2.3. Move around constantly and can flow

2.2.4. Expand when heated

2.3. Gas

2.3.1. Very weak forces of attraction

2.3.2. Free to move, far apart

2.3.3. Move constantly and randomly

2.3.4. Pressure increase when heated

2.4. Changes of state

2.4.1. Solid to Liquid = Melting

2.4.2. Liquid to solid = freezing

2.4.3. Liquid to gas = Boiling or evaporation

2.4.4. Gas to liquid = Condensation

2.4.5. Solid to gas = sublimation

2.4.6. Gas to solid = reverse sublimation (deposition)

2.5. Definitions

2.5.1. Solution = mixture of solute and solvent

2.5.2. Solvent = liquid that solute dissolves in

2.5.3. Solute = substance being dissolved

2.5.4. Solubility = how much solute will dissolve in a solvent

2.5.5. Saturated solution - a solution where the maximum amount of solute has been dissolved

3. Atomic structure

3.1. Isotopes

3.1.1. different atomic forms of the same element with different amounts of neutrons

3.2. Subatomic particles

3.2.1. Electrons

3.2.1.1. Negative charge

3.2.1.2. Relative mass of 0.0005

3.2.1.3. in the shells

3.2.2. Neutrons

3.2.2.1. Neutral charge

3.2.2.2. Relative mass of 1

3.2.2.3. in the nucleus

3.2.3. Protons

3.2.3.1. Positive charge

3.2.3.2. Relative mass of 1

3.2.3.3. in the nucleus

3.2.4. Atoms are neutral overall - same number of protons and electrons

3.2.4.1. Atomic number = how many protons there are

3.2.4.2. Mass number = amount of protons + neutrons

3.3. Relative atomic mass (Ar)

3.3.1. How heavy different atoms are compared with the mass of an average carbon-12 isotope (when carbon-12 = 12)

3.3.2. Ar = (Relative mass of an isotope x relative abundance) + (Relative mass of an isotope x relative abundance)

3.3.2.1. Divided by the sum of the relative abundances

3.4. Defenitions

3.4.1. Elements = consist of one type of atom only

3.4.2. Compound = two or more elements chemically bonded

3.4.3. Mixtures = no chemical bond between parts of a mixture, can be separated (not pure)

4. Chemical Bonding

4.1. Ionic bonding

4.1.1. Ions are charged particles - atoms have to lose or gain electrons to form them. They are trying to get a full outer shell

4.1.2. Positive ions - Cations

4.1.2.1. Groups 1, 2, 3

4.1.2.1.1. Gr 1 will for 1+

4.1.2.1.2. Gr 2 will form 2+

4.1.2.1.3. Gr 3 will form 3+

4.1.3. Negative ions - Anions

4.1.3.1. Groups 5, 6, 7

4.1.3.1.1. Gr 5 will form 3-

4.1.3.1.2. Gr 6 will form 2-

4.1.4. Tricky ions

4.1.4.1. Ag +

4.1.4.2. Cu 2+

4.1.4.3. Fe 2+

4.1.4.4. Pb 2+

4.1.4.5. Zn 2+

4.1.4.6. H +

4.1.4.7. Hydroxide OH -

4.1.4.8. Ammonium NH4 +

4.1.4.9. Carbonate CO3 2-

4.1.4.10. Nitrate NO3 -

4.1.4.11. Sulfate SO4 2-

4.1.5. Transfer of electrons

4.1.5.1. metal loses electrons to become a positive ion

4.1.5.2. non metal gains electrons to become a negative ion

4.1.5.3. these ions are oppositely charged and therefore are strongly attracted to each other by electrostatic

4.1.5.3.1. This attraction is an IONIC BOND

4.1.6. Ionic compounds

4.1.6.1. Overall charge is 0, but made up of a negative and a positive part

4.1.6.1.1. eg. Ca 2+ , NO3 - when bonded become Ca(NO3)2 to become neutral

4.1.6.2. All form in a similar way

4.1.6.2.1. all end up with full outer shells

4.1.6.3. Properties

4.1.6.3.1. Compounds with ionic bonding always have giant lattice structure

4.1.6.3.2. Electrostatic attractions between ions is very strong- therefore high melting and boiling points

4.1.6.3.3. not electrical conductors while solid

4.1.6.3.4. when melted or dissolved in water they conduct electricity

4.2. Covalent Bonding

4.2.1. Covalent Substances

4.2.1.1. Simple molecular

4.2.1.1.1. Weak intermolecular forces between multiple covalent bonds

4.2.1.1.2. Melting and boiling points low

4.2.1.1.3. intermolecular forces are stronger between particles with larger Mr s

4.2.1.1.4. melting and boiling points increase as the relative molecular mass (Mr) increases

4.2.1.1.5. Don't conduct electricity

4.2.1.1.6. Usually liquid or gas or easily melted solid

4.2.1.2. Giant covalent

4.2.1.2.1. All atoms bonded together with strong covalent bonds - Lattice structure

4.2.1.2.2. Very high melting and boiling points

4.2.1.2.3. Don't conduct electricity

4.2.1.2.4. usually insoluable

4.2.1.2.5. eg. Diamond and Graphite

4.2.2. When atoms share electrons to fill up their outer shells (to become stable)

4.2.3. strong electrostatic attraction between the negatively charged shared electrons and the positively charges nuclei

4.2.4. Atoms can share more than one electron with more than one other atom

5. Periodic table

5.1. Arrangement

5.1.1. In order of atomic number

5.1.2. Periods - The horizontal groups. in order of amount of electrons in the outer shell within each period

5.1.3. 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

5.2. Group 0

5.2.1. Noble gases

5.2.1.1. He, Ne, Ar, Kr, Xe, Rn

5.2.1.1.1. Non metals

5.2.1.2. Inert colourless gases

5.2.1.2.1. Don't react much (stable - have full electron shells.)

5.2.1.2.2. Exist as single gases, unlike oxygen and hydrogen

5.3. Group 1

5.3.1. Alkali metals

5.3.1.1. Very reactive

5.3.1.1.1. Reactivity increases as you go down the group

5.3.1.1.2. React with water

5.3.1.1.3. React with oxygen to form metal oxides

5.3.1.1.4. soft to cut

5.3.1.2. Li, Na, K, Rb, Cs, Fr

5.3.1.2.1. 1 electron in outer shell

5.4. Group 7

5.4.1. Halogens

5.4.1.1. F, Cl, Br, I, At

5.4.1.2. As you go down the group, the elements have a darker colour and higher boiling point

5.4.1.2.1. boiling point increases going down the group

5.4.1.3. Reactivity decreases going down the group

5.4.1.3.1. 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

6. Crude oil

6.1. Crude oil

6.1.1. Crude oil is a mixture of hydrocarbons

6.1.1.1. 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

6.1.1.2. Longer hydrocarbons have higher boiling points. they condense early on, near the bottom of the column, where it is the hottest

6.1.1.3. Hydrocarbon

6.1.1.3.1. a compound made out of hydrogen and carbon only

6.1.2. The fractions are

6.1.2.1. Refinery Gases

6.1.2.1.1. used in domestic heating and cooking (shortest chain)

6.1.2.2. Gasoline (petrol)

6.1.2.2.1. used as a fuel in cars

6.1.2.3. Kerosene

6.1.2.3.1. used as a fuel in air crafts

6.1.2.4. Diesel

6.1.2.4.1. used as a fuel in larger cars , vehicles and trains

6.1.2.5. Fuel oil

6.1.2.5.1. used as fuel for large ships and power stations

6.1.2.6. Bitumen

6.1.2.6.1. used to surface roads and roofs (longest chain)

6.2. Cracking

6.2.1. 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)

6.2.1.1. Cracking produces alkenes - useful to make polymers

6.2.2. Cracking is a form of thermal decomposition

6.2.2.1. Using heat + a powdered CATALYST

6.2.2.1.1. Catalysts - Silica (SiO2) or Alumina (Al3O2)

6.2.3. 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

6.3. Burning Hydrocarbons

6.3.1. Fuels release energy in combustion reactions

6.3.2. When you burn hydrocarbons with plenty of oxygen, they give out lots of energy

6.3.2.1. Exothermic reaction

6.3.2.1.1. hydrocarbon + oxygen -> carbon dioxide + water

6.3.2.1.2. Complete combustion

6.3.3. When there is not enough energy for complete combustion, you get incomplete combustion

6.3.3.1. as well as CO2 and water, incomplete combustion produces carbon monoxide and and soot

6.3.3.1.1. Carbon monoxide

6.3.4. Acid rain

6.3.4.1. When fractions from crude oil are burnt they sometimes produce sulfur dioxide and nitrogen oxides

6.3.4.1.1. 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.

6.4. Homologous series

6.4.1. a group of compounds that can be represented by the same general formula.

6.4.1.1. molecules of a homologous series contain the same functional group, meaning they react similarly.

6.5. Alkanes

6.5.1. Alkanes are saturated hydrocarbons

6.5.1.1. General formula C n H 2n+2

6.5.1.2. Chains of carbon atoms surrounded by hydrogen atoms.

6.5.1.3. Have formed bonds with as many other atoms as they can (no double bonds or spaces), therefore are saturated.

6.5.2. Alkanes burn in combustion reactions

6.5.2.1. make up the majority of hydrocarbons in crude oil and combust completely with good oxygen supply

6.5.2.2. Halogens react with alkanes to make haloalkanes

6.5.2.2.1. Chlorine and Bromine react with alkanes in the presence of UV light.

6.5.2.2.2. In these reactions a hydrogen atom from the alkane is substituted with chlorine or bromine

6.6. Alkenes

6.6.1. Alkenes are another type of hydrocarbon, but they contain a C=C double bond

6.6.1.1. 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

6.6.1.2. General formula C n H 2n

6.6.1.3. Halogens react with alkenes , forming haloalkenes

6.6.1.3.1. ethene + bromine -> dibromethane

6.6.1.3.2. This is an addition reaction because the C=C bond is split and a halogen atom is added to each of the carbons

6.7. Isomers

6.7.1. Molecules with the same molecular formula but the structural formula is different

6.7.1.1. eg. the atoms are arranged differently

6.8. Names of Alkanes/enes

6.8.1. No. of carbons

6.8.1.1. 1

6.8.1.1.1. 2

6.8.2. Stem

6.8.2.1. meth-

6.8.2.1.1. eth-

7. Acids Bases and Salts

7.1. Acids

7.1.1. an acid is a source of hydrogen ions (H+). They are proton donors

7.1.1.1. Reactions of acids

7.1.1.1.1. acid + base -> salt + water

7.1.1.1.2. acid + metal oxide -> salt + water

7.1.1.1.3. acid + ammonia -> Ammonium salt

7.1.1.1.4. acid + metal carbonate -> salt + water + carbon dioxide

7.1.1.2. Examples

7.1.1.2.1. HCl = Hydrochoric acid

7.1.1.2.2. H2SO4 = Sulphuric acid

7.1.1.2.3. HNO3 = nitric acid

7.2. Bases

7.2.1. A substance that can neutralise an acid. Alkalis are soluble bases.

7.2.1.1. an alkali is a source of hydroxide ions (OH-) and are proton acceptors

7.3. Salts

7.3.1. Making soluble salts

7.3.1.1. Use an acid and an insoluble base (metal oxide/hydroxide)

7.3.1.1.1. Heat the acid in a water bath in a fume cupboard

7.3.1.1.2. add the base until all the acid has been neutralised and you have excess base

7.3.1.1.3. Filter off the excess base to get a solution of the salt and water

7.3.1.1.4. Heat the solution gently to evaporate off some water and then leave the salt to crystallise

7.4. Indicators

7.4.1. a dye that changes colour depending on whether its above or below a certain pH

7.4.1.1. Universal indicator

7.4.1.1.1. add the indicator to an aqueous solution and match it to a colour on the pH scale

7.4.1.1.2. pH scale

7.4.1.2. Litmus paper

7.4.1.2.1. Red in acidic solutions

7.4.1.2.2. Purple in neutral

7.4.1.2.3. Blue in alkaline

7.4.1.3. Phenolphthalein

7.4.1.3.1. Colourless in acidic solutions

7.4.1.3.2. Bright pink in alkaline

7.4.1.4. Methyl orange

7.4.1.4.1. Red in acidic

7.4.1.4.2. Yellow in alkaline

7.5. Neutralisation

7.5.1. The reaction between an acid and a base (or alkali) is called a neutralisation reaction

7.5.1.1. H+(aq) + OH-(aq) -> H2O(l)

7.5.1.2. the acid donates protons which are accepted by the base

7.5.1.3. the products are neutral (pH 7)

8. Chemical calculations

8.1. Balancing equations

8.1.1. both sides of a symbol equation need to have equal amounts of each element

8.1.1.1. eg. H2SO4 + NaOH -> Na2SO4 + 2H2O

8.1.1.2. becomes: H2SO4 +2NaOH -> Na2SO4 +2H2o

8.2. Relative formula mass

8.2.1. All the relative atomic masses of a compound added together

8.2.1.1. eg. MgCl2 = (24)+(35.5*2)= 95

8.2.1.2. Just add up all the mass numbers from the periodic table

8.3. Moles

8.3.1. A mole is a unit for the amount of a substance.

8.3.1.1. A mole is 6.023x10^23 molecules. This is Avogadro's number.

8.3.1.1.1. One mole of atoms or molecules of any substance will have a mass equal to its Ar or Mr in grams

8.4. Calculating masses in reactions

8.4.1. Calculating the mass of product formed from a given mass of reactant (theoretical yeild)

8.4.1.1. 1. Write out the balanced equation

8.4.1.2. 2. Work out the Mr of the reactant and product you're interested in

8.4.1.3. 3. Find out how many moles there are in the substance you know the mass of

8.4.1.4. 4. Use the balanced equation to work out how many moles there'll be of the other substances

8.4.1.5. 5. use number of moles to calculate mass

8.4.1.5.1. this is your theoretical yield- how much product you would get if you got 100% of the yield

8.4.2. Percentage yeild

8.4.2.1. the actual amount of product you get compared to the theoretical yield

8.4.2.1.1. percentage yield = (actual yield / theoretical yield )* 100

8.4.2.1.2. this can be used to find how efficient (or profitable) a reaction is

8.5. Empirical and molecular formulae

8.5.1. The empirical formula gives you the smallest whole number ratio of atoms in a compound

8.5.1.1. 1. list all the elements in the compound

8.5.1.2. 2. write their experimental masses under them

8.5.1.3. 3. Find the moles of each element

8.5.1.4. 4. Divide the numbers by the smallest amount of moles

8.5.1.5. 5.Get the ratio into its simplest whole number form

8.5.2. The molecular formula gives you the real amount of each element in a compound. This can be the same as the empirical formula

8.5.2.1. Converting from the empirical formula to the molecular formula

8.5.2.1.1. 1.Find the mass of the empirical formula

8.5.2.1.2. 2. divide the given relative molecular mass of the molecule by the mass of the empirical formula

8.5.2.1.3. 3. multiply the numbers of each element in the molecule by your answer

8.6. Water of crystallisation

8.6.1. Salts can be anhydrous or hydrated

8.6.2. All salts are a lattice of positive and negative ions, and sometimes water molecules are incorporated into the lattice (hydrated)

8.6.3. the water in a lattice is called water of crystallisation

8.6.3.1. Calculating how much water of crystallisation a salt contains

8.6.3.1.1. Perform an experiment where the salt is heated to remove the water and weigh it before and afterwards

8.6.3.1.2. 1. find the difference in mass between the hydrated and anhydrous salt (therefore the mass of water lost)

8.6.3.1.3. 2. calculate the amount of moles of water lost

8.6.3.1.4. 3. calculate the amount of moles of anhydrous salt made

8.6.3.1.5. 4. Find the ratio between the anhydrous salt and water, and round the result

9. Reactivity series

9.1. The reactivity series lists metals in order of reactivity

9.1.1. Potassium

9.1.2. Sodium

9.1.3. Lithium

9.1.4. Calcium

9.1.5. Magnesium

9.1.6. Aluminium

9.1.7. Carbon

9.1.8. Zinc

9.1.9. Copper

9.1.10. Silver

9.1.11. Gold

9.2. Displacement reactions

9.2.1. These occur when a more reactive metal reacts with a less reactive metal's oxide, and displace it from the oxide

9.2.1.1. eg. iron oxide + aluminium -> aluminium oxide + iron

9.2.1.2. This is a redox reaction- the metal is oxidised and the displaced metal ion is reduced

9.2.2. this works the same with metals and less reactive metal's salts.

9.2.2.1. if the more reactive metal is already in the oxide/salt, no reaction occurs

9.3. Redox

9.3.1. Reduction

9.3.1.1. Loss of oxygen

9.3.1.1.1. eg. copper oxide is reduced

9.3.1.1.2. 2CuO + C -> 2Cu + CO2

9.3.2. Oxidisation

9.3.2.1. Gain of oxygen

9.3.2.1.1. eg. magnesium is oxidised

9.3.2.1.2. 2Mg + O2 -> 2MgO

9.3.3. In terms of electron loss and gain

9.3.3.1. Reduction

9.3.3.1.1. when an element/ion LOSES one or more electrons in a chemical reaction

9.3.3.2. Oxidisation

9.3.3.2.1. when an element/ion GAINS one or more electrons in a chemical reaction

9.4. Rusting

9.4.1. Iron and steel are very prone to rusting. Rust is hydrated IRON (III) OXIDE. This forms when metallic iron is OXIDISED.

9.4.1.1. Weakens steel and iron.

9.4.1.1.1. Worsens with air, moisture and warmth

9.4.1.2. Fe ---> Fe 2+ + 2e-

9.4.2. Preventation

9.4.2.1. Sacrificial preventaition

9.4.2.1.1. attaching a more reactive metal (Zinc) to iron structures because it will get oxidised instead of iron.

9.4.2.2. Barrier methods

9.4.2.2.1. Using grease, oil, plastic coating or paint to create a barrier between the metal and the water and oxygen

9.4.2.3. Galvanising

9.4.2.3.1. Iron coated in a layer of Zinc

9.4.2.3.2. Works even if damaged because the zinc is more reactive