CHEMISTRY

Chemical changes

Acids and alkalis

• For example, stomach acid is around pH 2, and bleach is around pH 12

• Indicators are chemical dyes that change to laine with the pH colour

  • Universal indicator is the most common, and changes to the same colours as above 🙂

  • A pH probe and meter give a number reading, and are more accurate and precise

• Acids are any substance that forms an aqueous solution with a pH less than 7

• A base is any substance with a pH greater than 7

  • Alkalis are any substances that dissolves in water to form a solution with a pH more than 7

• Neutralisation reactions are when you react an acid and a base to form a solution CLOSER to pH 7 (pure water)

  • They form a salt and water

  • HCl + NaOH $\rightarrow$ NaCl + H₂O

• Acids I need to know

  • Hydrochloric acid is HCl

  • Sulfuric acid is H₂SO₄

  • Nitric acid is HNO₃

• Alkalis I need to know

  • Sodium hydroxide is NaOH

  • Calcium carbonate is CaCO₃

Titration practical

• A titration is an experimental technique used to find an unknown concentration of an acid or alkali

• Neutralising an alkali with acid

  • Use a pipette to add 25cm³ of alkali to a conical flask

  • Add a few drops of an indicator (can’t be universal indicator) and place on a white tile

  • Fill a burette with acid and note the starting volume

  • Slowly add acid from the burette to the conical flask, swirling slowly, so that it is evenly distributed

  • Stop when the end point is reached - the acid has neutralised the alkali and the indicator has changed colour

  • Note the final volume reading and calculate the total acid added

  • Repeat until you get concordant results - within 0.10 cm³ of each other

  • Calculate a mean

• Indicators that can be used:

  • Litmus is RED in acids, and BLUE in alkaline solutions

  • Phenolphthalein is COLOURLESS in acids, and PINK in alkaline solutions

  • Methyl orange is RED in acid, and YELLOW in alkaline solutions

  • Universal indicator can’t be used as it gradually changes, and you can’t clearly see where it has reached the end point

Strong and weak acids

• Acids are ionised in aqueous solutions to release H⁺ ions

  • For example, HCl exists as H⁺ and Cl^- ions

• Strong acids ionise completely - all particles will dissociate

  • Like HCl or H₂SO₄

• Concentration is the amount of acid per amount of volume

  • The amount of H ions per unit of volume

• Weak acids don’t fully ionise - their reactions are reversible and stop reacting when they reach an equilibrium (net reactions)

  • Like CH₃COOH - ethanoic acid

• pH is a measure of concentration of H⁺ ions in a solution

  • The higher the concentration, the lower the pH

• HCl is a strong acid and could easily be pH0 at most concentrations

  • If it was a weak acid, to be pH0, it would have to be very very concentrated

Neutralisation reactions

• Metal oxide + acid $\rightarrow$ salt + water

  • Metal oxides end in O

• Metal hydroxides + acid $\rightarrow$ salt + water

  • Metal hydroxides end in OH

• Metal carbonates + acid $\rightarrow$ salt + water + carbon dioxide

• To make a soluble salt, we have to react an insoluble base with an acid

  • We gently heat an acid, and slowly and a base a bit at a time until it stopes dissolving

    • This shows the acid is neutralised and that the base is in excess

  • Then filter the excess base, and the product is dissolved soluble salt

  • To form crystals from this solution, we can gently heat it to evaporate any water (at less than 60 degrees)

    • When crystals begin to form, stop heating and leave to cool

    • This will cause more crystals to form, which can be filtered and dabbed dry

The reactivity series and displacement

• Elements below carbon can be reduced with carbon (in metal oxides)

• Reactivity is how easily something forms a positive ion

• Metals + acid $\rightarrow$ Salt + hydrogen

  • Reaction gets less violent down the reactivity series

  • Causes bubbles, explosions and temperature changes

• Metals + water $\rightarrow$ Metal hydroxide + hydrogen

  • Only in the most reactive metals

• Displacement example:   

  • Mg + FeSO₄ $\rightarrow$ MgSO₄ + Fe

    • Magnesium displaces iron, as it is more reactive

Separating metals from metal oxides

• Oxidisation is the process of losing electrons (gaining oxygen)

  • 2Mg + O₂ $\rightarrow$ 2MgO

• Reduction is the process of gaining electrons (losing oxygen0

  • 2MgO $\rightarrow$ 2Mg + O₂

• OILRIG - Oxidisation is loss, reduction is gain

• Most metals are fairly reactive

  • They reactive with oxygen (from the air) to form metal oxides

  • If they are really unreactive, they don’t react, like gold

• To get pure metals, we want to reduce them to remove the oxygen

• This is done by reacting metal oxides with carbon (displacement)

  • Oxygen will react with carbon and leave behind the pure metal

  • This only works with metals less reactive than carbon

  • It is cheap and quick

  • For example, 2Fe₂O₃ + 3C $\rightarrow$ 4Fe + 3CO₂, when mining for iron

• More reactive metals are extracted via electrolysis, which uses lots of energy and is expensive

The process of electrolysis

• Electrolysis involves using to break down electrolytes to form elements

• - ions are attracted to the anode (+ electrode)

  • They are then discharged, and become atoms

• + ions are attracted to the cathode (- electrode)

  • They are then discharged to become atoms, and produce either a solid or a gas

• At the cathode, Pb²⁺ + 2e- $\rightarrow$ Pb, reduction which forms solid lead

• At the anode, 2Br^- $\rightarrow$ Br₂ + 2e-, oxidisation which pairs up and forms gas (covalent bonds)

• Electrons at the anode that are loss are passed through the wire to the cathode, where they are given off to the positive ions to form atoms

Electrolysis of aqueous solutions

• Electrolysis is used to split compounds into their elements

• Soluble compounds can be dissolved in water to become aqueous, so charged particles are free to flow

• Because the compounds are dissolved in water, both the ions of the compound, and H and HO ions will be present

• For example:

• The cathode attracts positive ions - Na⁺ and H⁺

  • Only the less reactive will be discharged at the cathode

  • As hydrogen is less reactive than chlorine, hydrogen atoms will be formed, and given off as hydrogen gas

• The anode attracts negative ions - OH^- and Cl^-

  • If a halide is present, it will be discharged. If not, OH^- will be discharged

  • Halides are any of group 7 - F, Cl, Br, I and At

  • Chloride is a halide, so chloride is discharged, and forms chlorine gas at the anode

• This leaves NaOH in the beaker as a solution

Oxidation and reduction in terms of electrons

• Oxidation is loss, reduction is gain

• Reduction and oxidisation reactions both take at the same time - redox reactions

  • This is because the electrons that are loss also need to be gained

• Ionic equations are used to show displacement

  • They only show the particles that take part in the reaction and change

  • Ca + Fe²⁺SO₄^2- $\rightarrow$ Ca²⁺SO₄^2- + Fe

    • Ca + Fe²⁺ $\rightarrow$ Ca²⁺ + Fe

    • SO₄^2- is left out as it doesn’t change

• Half equations show the gain and loss of electrons for each element involved

  • Ca $\rightarrow$ Ca²⁺ + 2e-

  • Fe²⁺ + 2e- $\rightarrow$ Fe

  • The charged particles are always placed on the same side

Reactions of acids with metals

• The products of reactions of acids with metals are always a salt and hydrogen

• Acid + metal $\rightarrow$ salt + hydrogen

• They are redox reactions

  • 2H⁺ + Mg $\rightarrow$ Mg²⁺ + H²

  • Magnesium loses electrons - is oxidised

  • Hydrogen gains electrons - is reduced

Soluble salts

• Soluble salts can be prepared by reacting an acid with a suitable insoluble reactant

  • The reactant can be a metal, metal oxide, carbonate or a metal hydroxide

  • The reactant depends on the salt required

    • For example, copper doesn’t react with dilute acids, so can’t be used, and sodium is too reactive to use

• The reaction between an acid and a metal produces hydrogen

  • Hydrogen is flammable, so we usually use metal oxides or carbonates and an acid instead

• Making a salt

  • Add the powdered insoluble reactant to an acid in a beaker, 1 spatula at a time

    • The acid can be gently heated by a Bunsen burner to increase the reaction speed (particles have more energy)

  • Stir between each addition and continue until the powder is in excess and no longer reacts

  • Filter the mixture in the beaker with filter paper and a funnel to remove excess solid - this means the filtrate now only contains the salt and water

  • Heat the solution in a evaporating dish over a water bath

  • Stop when small crystals begin to appear - the solution is now saturated and most of the water has evaporated

  • Leave the solution for 1-2 days at room temperature to allow the rest of the water to evaporate and leave large crystals

  • Dry by gently dabbing with filter paper

Electrolysis of molten ionic compounds

• This is the process of separating elements in insoluble ionic compounds (splitting with electricity)

• In electrolysis, the electrolyte has to be a liquid or aqueous solution that contains an ionic compound - so the ions can move freely

  • CuSO₄ is soluble, and can be dissolve in water to create this

  • However, lead bromide is insoluble and has to be melted and made molten to allow ions to move freely

• PbBr₂ is heated to become molten

• It splits into it’s two ions - Br^- and Pb²⁺

• Br^- is attracted to the positive anode and is discharged as bromine gas

  • Br^- $\rightarrow$ Br + 2e-

  • Two bromines react covalently to form 2Br as a gas

• Pb²⁺ is attracted to the negative anode and is discharged as solid lead

  • Pb²⁺ + 2e- $\rightarrow$ Pb

  • This creates a layer of molten lead in the beaker

• Ions are being oxidised and reduced at the electrodes

  • Cathode = reduction

  • Anode = oxidisation

• Electrons that are lost when Br^- become Br are transported via the wire with current to the anode, where Pb²⁺ gains two electrons

Extracting metals with electrolysis

• Reactive metals are extracted from metal oxides by melting and making them molten compounds

• The cheapest way to reduced a metal from an oxide is with carbon, but it only works with metals that are less reactive than carbon

  • So we use electrolysis instead, which is expensive and needs lots of energy

• Electrolysis only works in ions can freely move through the solution

• For example, Al₂O₃ $\rightarrow$ Al + O₂

  • Aluminium is more reactive than carbon, so electrolysis is used

  • Aluminium oxide is found as a solid, and found mixed with bauxite when mined

    • This needs to become molten 🔥

• First, we need to purify Al₂O₃ from bauxite

• Then, we mix aluminium oxide with cryolite as it lowers the melting point

  • Aluminium oxide has a very high melting point (it was ~ 2000 degrees C), but is lowered by cryolite

  • We then melt this mixture to become molten - lots of energy is required as the melting point is still high

• The electrodes are made of graphite, and the cathode is found around the outside of the steel case

• O^2- is attracted to the anode in the centre

  • At the anode, it is discharge and becomes oxygen gas (makes a pair and covalently bonds once atoms0

  • 2 electrons are lost per oxygen, and they are transferred through the wire to the cathode for reduction

  • 2O^2- $\rightarrow$ O₂ + 4e-

• Al³⁺ is attracted to the cathode around the edge

  • It is discharged and becomes Al

    • 3 electrons are gained from the electrons lost at the anode

    • The aluminium formed pools at the bottom, and leaves via a channel at the bottom to be collected

    • Al³⁺ + 3e- $\rightarrow$ Al

• 2Al₂O₃ (l) $\rightarrow$ 4Al (l) + 3O₂ (g)

Atomic structure and the periodic table

Atoms, elements and compounds

• Atoms make up everything and have a radius of around 0.1 nanometres (1 × 10^-1m)

• Neutrons have a relative mass of 1, and no charge

• Protons have a relative mass of 1, and a positive 1 charge

• Electrons have a very small mass, and a charge of negative 1

• Elements are different types of atoms, and are shown on the periodic table

  

• An isotope is an element with the same number of protons, but different numbers of neutrons, so therefore different mass numbers

• Relative atomic mass of an element’s isotopes = the sum of(isotope abundance x mass) / sum of percentage of all isotopes

  • Should be divided by 100 if all isotopes are present

• Compounds are two or more different elements that are chemically bonded

  • H₂O or CO₂, but not O₂ OR Cl₂

  • Elements in compounds are always found in the same proportions

Molecules

• Molecules are two or more atoms chemically bonded

  • For example, O₂ and CO₂ are both molecules

Electronic structure

• The first shell of an atom has 2 electrons, then 8 electrons and then 8 electrons (2,8,8)

• The group number tells you the number of electrons an element has in it’s outer shell

• The number of shells an element has is the same as the number of shells it has

The size and mass of atoms

• Atoms have a radius of around 1 × 10^-10 m, or 0.1 nanometres

• Their mass is made of the total number of neutrons and protons

Development of the model of the atom

• Democritus - Atomic theory

  • In 500 B.C

  • Everything is made up of tiny indestructible particles, with empty space between

• John Dalton - In the 1800’s

  • Everything is made up of ‘solid spheres’

  • Different types of spheres are different elements

• JJ Thompson - The plum pudding model

  • In 1897

  • He found that atoms couldn’t be solid spheres and that most contained charged particles

  • He thought the atom was a ball of positive charge with discrete negative particles embedded in it

• Ernest Rutherford - Alpha scattering experiment

  • In 1909

  • He fired positive alpha particles at a thin sheet of gold

  • If the positive mass was spread out, like in the plum pudding model, all the particles should pass through as the positive charge is not dense enough to repel

    • However, not all of the particles did and some were deflected

  • This showed that there was a concentrated positive nucleus - the nuclear model

• Niels Bohr - In 1913

  • He suggested that electrons orbit the nucleus in shells, which prevents the atom from collapsing like in Rutherford’s model (1909)

• Rutherford - In 1913

  • He suggested that the nucleus in made up of small discrete particles, called protons, and wasn’t just a cloud of charge

• James Chadwick - In 1932

  • He found that the nucleus contained neutrons with mass

The periodic table

• The periodic table was created by Dimitri Mendeleev in the mid 19th century

• He left spaces to predict new, undiscovered elements

• Columns are called group and have similar chemical properties and have the same number of electrons in the outer shell

• Rows are called periods and shows the number of shells an element has

The development of the periodic table

• Before Dimitri Mendeleev, neutron and proton scientists tried to classify their discoveries in order of atomic weights

  • They were incomplete, and the groups were often inaccurate 

• Dimitri Mendeleev made the periodic table in 1869, and he predicted new elements and left spaces in his table

  • He included elements mass number, element symbol and atomic number in his table

  • He made groups with similar chemical properties, with the same number of electrons in the outer shell

Metals and non-metals

• Metals are found on the left side of the periodic table

  • They form positive ions - they lose electrons

    • The fewer electrons they have in their outer shells, the easier they are lost

  • They get more reactive the further down the table you go

    • This is because they have more shells, so the outer shell electrons have little attraction to the nucleus and are lost easily

  • Metals metallically bond so are very strong

    • They are malleable, conductors of electricity and heat

    • They have high melting and boiling points, and they are shiny and sonorous

• Non-metals are found on the right side of the periodic table

  • They form negative ions, or none at all

  • Non-metals have lower densities than metals, are brittle and dull in colour

    • They don’t conduct and have low melting and boiling points

• The transition metals have typical metal properties

  • They can form more than one ion

    • For example, iron (II) and iron (III)

  • They also form good catalysts

Group 1 - The alkali metals

• The group 1 elements are lithium, sodium, potassium, rubidium, caesium and francium

• Their properties include:

  • Soft, low densities and low melting points, which are different to other metals

  • They get more reactive down the group

  • Their melting and boiling points decrease down the group

• They have one electron in their outer shells, which can be lost mor easily the further down the group as the are more shell

  • They are further from the nucleus, so less attraction as there is a further distance between

• They are almost always found as ionic compounds

  • With electrostatic forces between non-metal and alkali metals

• Alkali metals and water react vigorously to form metal hydroxide and hydrogen gas

• Alkali metals and chlorine gas react to form metal chloride salt

  • Na + Cl $\rightarrow$ NaCl (salt)

• Alkali metals and oxygen react to form metal oxides

  • Lithium and oxygen react to form lithium oxide

  • Sodium and oxygen react to for sodium oxide OR sodium peroxide

  • Potassium and oxygen react to form potassium peroxide OR potassium superoxide

Group 7 - The halogens

• All of the halogens are a bit dangerous ☹

• They all exist as diatomic molecules

  • Make covalent bonds to gain a full outer shell, like chlorine

• They can form covalent bonds with other non-metals as well

• Properties and trends of the halogens:

  • Their melting and boiling points increase down the group

  • Their reactivity decreases down the group as the more shells there are, the weaker forces of attraction there is from the nucleus to gain more electrons

• They form negative 1 ions with metals

  • Their name changes to end in -ide, like chloride or fluoride

• In displacement reactions, a more reactive element can displace another element that is less reactive

  • Cl₂ + 2KBr $\rightarrow$ Br₂ + 2KCl - bromide is displaced by chlorine

  • The more reactive element will always displace the less reactive one

Group 0 - The noble gases

• The noble gases exist as colourless gases

• They are all unreactive and have a full outer shell

  • They exist as singular atoms

  • They are non-flammable

  • Their boiling points increase down the group

The transition metals

• The transition metals are found in the middle of the periodic table (between groups 2 and 3)

• Their properties in comparison to group 1

  • They both conduct electricity 

  • They are both shiny 

  • They have higher melting and boiling points 

  • They have higher densities and are stronger and harder than the alkali metals

• They have slow or non-existent reactions with oxygen, water and the halogens

  • They react vigorously with group 1

• They make good catalysts

Energy changes

Energy transfers during exothermic and endothermic reactions

• Different chemicals store different amounts of energy in their bonds

• CH₄ + 2O₂ $\rightarrow$ CO₂ + 2H₂0

  • Before this reaction, the bonds have more energy, and after energy has been released to the surroundings, through heat

• Exothermic reactions release energy to the surroundings

  • Energy is EXiting the reaction

  • The surroundings get hotter

  • Transfers energy to surroundings

• Endothermic reactions take in energy from the surroundings

  • Energy is ENtering the reaction

  • The surroundings get cooler

• Activation energy is the minimum energy the reactant particles need in order to collide with each other enough to cause a reaction

Energy changes of reactions

• Bond energy is the amount of energy required to break one mole of a particular bond

  • For example, a H-Cl bond requires 431 kJ/mol (energy needed to break one mole - 6.02 × 10²³ - of these bonds)

• Bond breaking is exothermic, and making bonds is endothermic

• To calculate:

  • Work out how many bonds are breaking in the reactants, and calculate the total energy required by multiplying the number of specific bonds by the kJ required

  • Work out the same for the products

  • Find the difference in the total energy required for the reactants and the products 🙂

Cells and batteries

• Electrochemical cells use chemical reactions to produce electricity

• The electrodes have to be made out of two different types of metal, and they conduct

• A cell can be made by connecting two different metal electrodes with wire and placing them in contact with an electrolyte solution

  • An electrolyte is a liquid through which charged particles can flow - creates a flow of charge, and therefore a cell

• Batteries are similar, but consist of two or more cells connected in series to provide a greater voltage

  

• Factors that affect the voltage of a cell or battery include:

  • The metals used - the greater the difference in their reactivities, the greater the voltage

  • The type and concentration of electrolyte

  • The conditions, such as temperature

• Rechargeable batteries work as the chemical reactions inside can be reversed when an external electrical current is supplied

  • They are used in phones and laptops

• Non-rechargeable batteries are where the reactions stop once one of the reactants has been used up

  • Used in smoke alarms, TV remotes, etc.

  • Also called Alkaline batteries

Fuel cells

• Fuels cells are electrochemical cells that converts energy between chemical and electrical

• We can convert the energy of oxygen and a fuel to release electrical energy we can use

• The most common type of fuel cell is the hydrogen-oxygen fuel cell

  • It forms water and creates lots of electrical energy

• H₂ enters through the left of the fuel cell, and is oxidised by the anode (-) to split into two positive hydrogen ions, and two electrons

  • The oxidisation of hydrogen - H₂ $\rightarrow$ 2H⁺ + 2e-

    • The electrons pass through the wire, creating a current and electrical energy, to the cathode

    • The hydrogen ions pass through the electrolyte to the cathode (+)

    • Oxygen enters from the right side

    • Hydrogen ions and electrons can react with oxygen to produce water

      • Two H₂ and one O₂

      • O₂ + 4H⁺ + 4e- $\rightarrow$ 2H₂O

    • The water leaves the fuel cell via the outlet, as well as heat (non-useful energy)

• The overall equation of this process is O₂ + 2H₂ $\rightarrow$ 2H₂O

• The electrical energy comes from the flow of electrons through the wire

• As the fuel enters the cell, it becomes oxidised, creating a potential difference across the cell

• PROS - Only requires oxygen and hydrogen

  • No waste is created

  • They last longer than batteries

  • Simple process

• CONS - H₂ is expensive to store as it takes up lots of space

  • H₂ is explosive in air ☹

  • To make hydrogen fuel, we need energy, often from fossil fuels

Bonding, structure and the states of matter

Ionic bonding

• An ion is a charged particle, which is formed when elements gain or lose electrons to gain a full outer shell

• Elements are more likely to form ions if there are not many electrons lost or gained

  • Takes less energy

  • Groups 1,2,6 and 7 are most likely

  • Groups 3,4 and 5 are rarely seen as ions

• Happens between a metal and non-metal

• For example, Na$\rightarrow$Na⁺ + e^- and Cl + e^- $\rightarrow$ Cl^-

• Electrons are transferred

• NaCl becomes an ionic compound and are attracted by strong electrostatic forces (have opposite charges)

Ionic compounds

• Metals and non-metals form ionic compounds with strong electrostatic forces and ionic bonds

• Usually group in large numbers

  • Regular lattice structures (3D)

  • Each ion is attracted to all those around it

• Properties - very high melting points as lots of energy is required to overcome strong ionic bonds

  • can conduct electricity when molten or aqueous as charged particles are free to move

• NEED TO KNOW - Hydroxide = OH^-

  • Sulphate = SO₄^-

  • Nitrate = NO₃^-

  • Carbonate = CO₃^-

  • Ammonium = NH₄⁺

Covalent bonding

• The sharing of electrons in the outer shell between non-metals to gain a full outer shell

• Cl - Cl

  

• They can become:

  • Simple molecular substances - Small molecules with strong covalent bonds between atoms and weak intermolecular forces between molecules

    • Like water, ammonia, chlorine or methane

  • Polymers - Long chains of repeating units (monomers)

  • Giant covalent structures - diamond, graphite + silica dioxide

Metallic bonding

• Happens between metal atoms

• Solid metals are in a giant structure arranged in a regular pattern with delocalised electrons

  • They give up their outer shell electrons and share them with the other metals

  • The atoms all become + ions

  • The lost electrons can freely move so are delocalised

  • There are strong forces of attraction between the ions and electrons, which hold everything together in a regular structure

• Metals are strong, so have high melting and boiling points

• They are good conductors of heat and electricity ⚡

  • Their delocalised electrons can feely move and carry electrical current through the structure

• Metals are malleable - their regular structure allows layers to slide

• Alloys are when 2 or more different metals or a metal and non-metal form metallic bonds, with different sized atoms

  • This disrupts the regular structure, so layers can no longer slide

• Alloys are stronger than pure metals

States of matter

• Solids - Strong forces of attraction (holds them close together)

  • Fixed, regular position

  • Definite shape and volume - can vibrate

• Liquids - Weak forces of attraction (particles are free to move and flow)

  • Can flow

  • Compact and definite volume, but not shape

  • Move to fit a container

• Gas - Very very weak forces of attraction

  • No definite shape of volume

  • Fill a container

  • Particles are free to move

  • Constantly moving with random motion - move in a straight line and are deflected when hit

State symbols

• Solid - s

• Liquid - l

• Gas - g

• Aqueous - aq

Properties of ionic compounds

• High melting and boiling point

  • Lots of energy is required to overcome strong electrostatic forces of attraction, and there are lots of forces

• Can conduct when aqueous or molten

  • Charged particles are free to flow through the structure

Properties of small molecules

• Low melting and boiling points

  • Weak intermolecular forces between molecules, that need little energy to break

  • COVALENT BONDS ARE NOT BROKEN

• Generally liquids or gases at room temperature

• Do not conduct

Properties of metals and alloys

• Metals are soft and malleable, shiny, good conductors of heat and electricity and have high melting and boiling points

• Alloys are hard (no layers), have high melting and boiling points and are good conductors

Giant covalent structures

• Simple molecular substances have low melting points, strong bonds between atoms and don’t conduct

• Giant covalent structures have huge numbers of non-metal atoms

  • Arranged in a regular repeating lattices

  • Have high melting and boiling points as there are a lot of covalent bonds

  • Very strong - lots of bonds

  • Generally don’t conduct (apart from graphite and graphene)

• Silica dioxide is made of silicon and oxygen in a ration of 1:2

  • Makes up sand 🏖

Diamond and graphite

• Allotropes of carbon

• Diamond is a giant covalent structure💎

  • Is very strong

  • Each carbon is bonded to 4 other carbons (max amount)

  • It doesn’t conduct as there are no delocalised electrons

• Graphite is a giant covalent structure

  • Is very strong

  • Each carbon makes 3 out of 4 covalent bonds possible

  • Is arranged in layers with weak intermolecular forces between them

    • This allows the layers to slide over one another and makes it soft

  • Has a high melting and boiling point

  • Can conduct electricity - only ¾ bonds are made, so there are left over electrons

    • Become delocalised (one per carbon atom) and are free to move through the structure and carry charge

• Graphene is a single layer of graphite

Graphene and fullerenes

• Are allotropes of carbon

• Graphene is a single layer of graphite and can conduct electricity as there are delocalise electrons

  • Useful in electronics (conducts and is small)

• Fullerenes are tubes and spheres made out of a single sheet of graphite

  • Spheres can be used to surround molecules (like drugs) and used to deliver to specific areas of the body

  • They have a large surface area : volume ratio, so make good industrial catalysts

  • Tubes can be used in nanotechnology as conductors, to strengthen tennis rackets (adds strength without weight as high length : diameter)

  • Buckminster fullerene is a hollow sphere that is made of 60 carbon atoms and is used for drug delivery

  • Using tiny structures is called nanotechnology

    • Medicine, fashion, batteries and food

Nanoparticles

• Nanoparticles are really really really tiny particles - 1nm - 100nm (0.00000001m)

• Nanoscience is produces new nanoparticle materials

• They have a large surface area : volume ratio

  • Good for catalysts as surface area increases its efficiency

  • Nanomedicine uses fullerenes to deliver drugs around the body

  • Electrical circuits use them to make tiny computer chips as some can conduct

  • Silver nanoparticles have antibacterial properties so can be infused into wound dressings and masks

• Issues with nanoparticles

  • They are relatively new so we are not aware of all risks (long term)

  • For example, sun cream with nanoparticles allows for better skin coverage but we are unaware if it can enter our body through the skin and potential damage cells

  • They are also possibly damaging to the environment

Sizes of particles and their properties

• Atoms and small molecules - 0.1 nm

• Nanoparticles - 1 to 100nm

• Fine particles - 100 to 2500 nm

• Coarse particles - 2500 to 10000 nm

• The smaller the particle, the higher the surface area : volume ratio, so increased reactivity

Polymers

• Polymers have very large molecules and their atoms are joined by strong covalent bonds in long chains

• Solid at room temperature

• Higher boiling points than strong intermolecular forces and lots of bonds to overcome