Year 11 Module 1 – The Chemical Earth

The Earth is made up of a large number of different substances: elements, compounds & mixtures

We can classify elements & compounds as pure substances. These are always HOMOGENEOUS. Mixtures are not pure substances and can be either homogeneous or HETEROGENEOUS.

Pure substances have a fixed compositioneg copper metal, sulfur, carbon dioxide (CO2), methane (CH4).

Mixtures have variable composition and can also be separated into its components relatively easily.Eg. Alloys – brass Cu: 65% Zn 35%

Elements are pure substances that cannot be broken down into any other substance. Compounds can be broken down, but this requires chemical means and is not as easy as separating mixtures.

The Spheres of the Earth

Atmosphere:Hydrosphere: define each, Lithosphere: and give an example of a mixtureBiosphere: found in each.

Separating Mixtures

Solids of different sizes - seiving

Solids & liquids – evaporation, filtration, distillation, decanting

Dissolved solids – crystallisation by evaporating

Solvents - distillation

Two liquids – immiscible liquids using a separating funnel- fractional distillation due to different boiling points

H/W – P.18 Qu. 11-18

Identifying Pure Substances

Colour – some substances have a characteristic colour, eg CuSO4 = blue. We can use this knowledge in a qualitative way to determine if a sample has impurities.

Melting & Boiling Point – The mp & bp of a substance is a good indicator of its purity. A pure substance will have a sharp melting point and its mp & bp will not change after further attempts at purification.

Impurities can lower or raise bp, but only lower mp.

Density = mass/volume – pure substances have characteristic densities – eg glass.

Gravimetric Analysis

This is a method used to determine the composition by mass of a substance.• Mining – determine the % of a mineral in an ore deposit• Quality control of products• Determining the composition of a product made by a rival company• Composition of a soil• Pollutant in air or water

Calculating % composition

% A in ABC = mass of A present x 100 Mass of total sample (ABC)

Eg % BaSO4 in Ore sample

3.61g of sampleMass of BaSO4 = 1.52gMass of MgSO4 = 2.07gTotal returned mass = 3.59g

% BaSO4 = 1.52/3.61 x 100 = 42%

% MgSO4 = 100 – 42 = 58 %

= 2.07/3.61 x 100 = 58%

P.25 Qu 26 & 27 for 26 → 1 ppm = 1 mg/L

Physical Properties & Uses of Elements

Metals & non-metals have differing properties. The properties of an element will determine what it is used for. Metals Shiny, ductile, malleable, good conductors of electricity, variable (but mostly high melting point)Eg – Gold – jewellery as it’s shiny

Aluminium, Iron – car parts as malleable

Copper – wire as it is ductile, a very good conductorTungsten – light bulb filaments due to v. High m.p

Non-metals: Dull, poor conductors, brittle, usually soft, variable mp (v.low – v.high)Eg. Carbon (graphite) can conduct, so used as electrodes

Carbon (diamond) used a toolsLiquid N2(l) used in cooling

Some elements have properties of both metals & non-metals, these are classified as semi-metals or metalloidsEg silicon.

Comparing Chemical & Physical Changes

Chemical changes are reactions – they are difficult to reverse, at least one new substance formed, mass will be conserved, often a large input/output of heat.

Physical changes – No new substances are formed, easily reversed and only small energy changes are required/produced.E.g 2H2O(l) + E 2H2(g) + O2(g) – chemical change

H2O(l) + E H2O(g) – physical change

Other physical changes are: freezing, melting, condensing, evaporating, subliming, dissolving & crystallisation.

In the electrolysis of H2O

Negative electrode: 4H+(aq) + 4e- 2H2(g)

Positive electrode: 4OH-(aq) O2(g) + 2H2O(l) + 4e-

Evidence of Chemical Reactions

• New products formed (eg precipitate)• Reactants consumed• Colour change• Smell• Gas evolved• Significant change in temperature

Different reaction classes include:

1. Decomposition A B + C eg CuCO3(s) CuO(s) + CO2(g)

2. Synthesis A + B C (combination reactions)2Mg(s) + O2(g) 2MgO(s)

3. Redox Reduction & oxidation

4. CombustionCxHy + (2x + 1/2y) O2 xCO2 + ½ y H2OOr CH4(g) + 2O2(g) CO2(g) + 2 H2O(g)

Reactions may be EXOTHERMIC – release energy or ENDOTHERMIC – absorb energy. All reactions require a small amount of E to get them started – activation energy.

Atoms, Molecules & Ions

The smallest particle that is still recognisable as an element is called an atom. They consist of protons, neutrons & electrons and have no overall charge.

Some elements, such as hydrogen, chlorine, oxygen etc do not exist as individual atoms. They are found in “molecules” H2, Cl2 & O2. These are diatomic molecules and the other ones that form are: F2, Br2, I2, N2. A molecule is the smallest unit of a substance that can exist separately.

Molecules are not always elements, eg: CO2, H2O etc. Some elements also exist as monatomic molecules: noble gases, whilst others have larger molecules: S4 & S8, P4,

C60.

Symbols – come from the first letter, first two letters, a combination of the 1st and other letter or Greek, Latin or German. Some are names after places, others after people.

Cf, Fr, Ge, Am, Pu, Bk, Hg, U, Ga, Sc, Np, Es, Cm,

P.39; Ques 4, 5, 6, 7 – (odd letters if easy!!)

Atomic Structure

All matter is made up of tiny indivisible particles called atoms. Each chemical element contains the same types of atom.

Atoms consist of three sub-atomic particles

Particle Symbol Mass Charge LocationProton P 1 +1 In the nucleus

Neutron N 1 0 In the nucleus

Electron E ~ 0 -1Orbiting the nucleus (in shells)

- Atomic Number = No of P = No of E- Mass Number = No of P + No of N

Number of Protons = atomic numberNumber of Electrons = atomic numberNumber of Neutrons = mass number – atomic number

Calculate P,N & E for Au, Fe, Ar & U235

Au P = E = N =

Fe P = E = N =

Ar P = E = N =

U P = E = N =

Isotopes

These have different numbers of neutrons. Some elements exist only as one isotope – eg Al 27, whereas others may have 2 or more. Eg Cl 35 & Cl 37.

Cl 35 – 75%Cl 37 - 25%

Relative Atomic Mass = the weighted mean of all naturally occurring isotopes.

R.A.M = (35 0.75) + (37 0.25) = 35.5

Calculate the RAM of potassium – K39 = 93.3% K40 = 0.01% K41 = 6.69%

Electron Structure/configuration

The electrons in atoms do not all group together in one shell. They are arranged in sub-shells or orbitals. There are many types:

S-sub-shells: can hold 2 electrons, these represent the lowest energy sub-shell in each complete shell.

P-sub-shells: each can hold 2 electrons and each complete shell has 3 p sub-shells (total = 6)

D-sub-shells: each can hold 2 electrons, each complete shell has 5 d sub-shells (total electrons = 10)

F-sub-shells: each can hold 2 e-, there are 7 in each complete shell (total = 14 electrons)

The sub-shells fill up in order of increasing energy:

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2

Eg H 1s1 1He 1s2 2Li 1s2 2s1 2,1Be 1s2 2s2 2,2B 1s2 2s2 2p1 2,3

Etc – Sc ZnKr

Ext- Y, Rh, I, Cs, Cu, Cr

Stable electron configurations

The number of electrons in the outer shell of an atom is the same as the group number (except transition metals). Eg all Group 1 (Li, Na, K have 1 electron in their outer shell).

The outer shell is sometimes called the valence shell, and knowing that the most stable atoms have a “full” outer shell of 8 electrons (2 for He), we can determine the reactivity and valency (combining power) for elements.

Group I II III IV V VI VII VIII

Valency

1 2 3 4 3 2 1 0

Charge +1 +2 +3 +4/-4 -3 -2 -1 0

Formation of Ions - Achieving Noble Gas Status

Having a full outer shell of electrons, like the noble gases is the most stable state to be in. Atoms will lose, gain or share electrons to get to this state.

(i) Forming ions

Metal atoms will tend to lose electrons, forming positive ions (cations) whose charge is determined by the number of electrons lost.

Examples – see written notes – diagrams eg Li → Li+ , Mg → Mg2+, Al → Al3+

Non-metal atoms tend to gain electrons, forming negative ions (anions), achieving a stable structure.

Examples – see written notes – diagrams eg Cl → Cl-, O → O2-, P → P3-

Ionic bonding

Metal atoms tend to “give” their electrons to non-metal atoms, forming a cation & anion that attract each other as an ionic bond. The + & - attract through electrostatic force.

Examples – see written notes – diagrams eg Na + Cl, Na + O, Mg + Cl, Al + O

Writing Ionic Formulae

If we know the valency of the two substances reacting together, we can determine the formula by ensuring that the number of electrons lost = no of electrons gained.

Eg. Na Cl both have a valency of 1 NaCl

1 2Li + O 2 Li ions are needed to balance O Li2O

Qu. Practice questions

Issues occur with transition metals, valencies are given as (II) etc.Eg. Cu(II), Fe(III)

If a polyatomic ion is used, brackets are required if more than one is needed.

Sulfate = SO42-

Carbonate = CO32-

Nitrate = NO3-

Phosphate = PO43-

Hydroxide = OH-

Ammonium = NH4+

Qu. Practice questions HW – P.79 Qu: 11-16

Covalent Bonding

This occurs between non-metal atoms that both need to gain electrons. They do this by SHARING electrons in their outer shell.

The number of electrons that they need to share is determined by their valency.

Carbon = share 4Nitrogen = 3Oxygen = 2Hydrogen, Chlorine, Fluorine = 1

Examples – see written notes – diagrams eg H2, H2O, CH4, - all single

Double – two pairs of shared electrons – O2, CO2

Triple 3 pairs of shared electrons N2, C2H2

Naming – uses mono-, di-, tri-, tetra-, etc as prefixes

Eg CO2 = carbon dioxide

Questions – P.92, 17-19

Balancing Chemical Equations

Due to the Law of Conservation of Mass, that states that matter cannot be created or destroyed, merely transformed from one form to another, we need to ensure that chemical equations have equal number of each type of atom on each side.

Eg. CH4(g) + 2O2(g) CO2(g) + 2H2O(g)

This has the following:Elemen

tReactan

tsProduct

sC 1 1H 4 4O 4 4

When balancing always use the following order:Metals, Non-Metals, C, H, O(Na, Fe), (S, Cl, N, P)

Balancing Equations – (rewrite in your book)

1. Na + Cl2 NaCl 10. H2 + Cl2 HCl

2. KClO3 KCl + O2 11. CO + O2 CO2

3. H2 + O2 H2O 12. KNO3 KNO2 + O2

4. P + O2 P4O10 13. Pb3O4 PbO + O2

5. Zn + H2SO4 ZnSO4 + H2 14. Fe + H2O Fe3O4 + H2

6. CH4 + O2 CO2 + H2O 15. Fe2O3 + C Fe + CO2

7. N2 + H2 NH3 16. NH3 + O2 N2 + H2O

8. Pb(NO3)2 PbO + NO2 + O2 17. NaOH + H2SO4 Na2SO4 + H2O

9. C7H6O2 + O2 CO2 + H2O 18. BCl3 + P4 + H2 BP + HCl

19. Ca(OH)2 + HCl CaCl2 + H2O

20. Na2CO3 + Ca(OH)2 NaOH + CaCO3

21. C2H2Cl4 + Ca(OH)2 C2HCl3 + CaCl2 + H2O

22. HCl + Na2CO3 CO2 + H2O + NaCl

23. Al2(SO4)3 + NaOH Al(OH)3 + Na2SO4

24. (NH4)2Cr2O7 N2 + Cr2O3 + H2O

25. Zn3Sb2 + H2O Zn(OH)2 + SbH3

Metallic Bonding

When metal atoms bond together, their valence electrons become ‘delocalised’ in a ‘sea of electrons’ across the whole structure.

Diagram of metallic bonding

Properties of Substances

1. Ionic substances

Ionic lattices are massive 3D structures arranged in an orderly manner with alternating + & - ions. Eg NaCl.

Insert diagramThe ions are held together by electrostatic force. This is an ionic bond. Because the lattice has millions of these strong bonds, ionic compounds are solid at room temp and have high melting & boiling points.The regular structure means that they are hard & brittle.As a solid all its ions & electrons are in a fixed position, so it cannot conduct electricity. However, when molten or dissolved in water, the ions are free to move and conduct.

2. Covalent molecules These are formed when 2 or more covalent atoms share electrons forming

covalent bonds. Eg; H2O, CO2, CH4, NH3, CH3COOH, C2H5OH. They all have low boiling & melting points because they have very weak

forces between the molecules (IMF = intermolecular forces). When solid, most are soft Pure covalent substances do not conduct electricity as solids or liquids as all

electrons are locked up in covalent bonds. When they dissolve to form aqueous solutions, they don’t conduct unless they react with the water forming ions.

3. Covalent networksEg, Carbon (graphite & diamond), SiO2

These networks have millions of covalent bonds They all have very high melting & boiling points Most do not conduct electricity (except graphite) Most are hard and brittle

Graphite can conduct as one of its outer electrons is ‘delocalised’.

4. Metals

Variable melting points from 25 °C for reactive metals and below 0°C for Hg to mid 3000 °C for WThey all conduct electricity due to their freely moving ‘delocalised’ electrons.Malleable & ductile as the ions can move over each other easily.