Chemistry - Module 4

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History of the Atom

John Dalton - solid spheres that make up different elements

J J Thompson - not solid spheres. Contain electrons with negative charge. 'plum pudding model'

Rutherford - fired +charged particles at thin sheet of gold. Expected most to be deflected by positive 'pudding' but most went through.
+ve nucleus at centre, surrounded by 'cloud' of negative electrons. Most of atom is empty space.

Bohr - Electrons exist in fixed orbits or shells.
Supported by many experiments and explains other scientists observations

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- At centre
- Contains protons and neutrons
- Has a positive charge bc of protons
- Almost whole mass of atom concentrated in nucleus

- Move around in shells
- Negatively charged
- Tiny but cover space
- Virtually no mass 

Protons - heavy (1) - positive
Neutrons - heavy (1) - neutral
Electrons - tiny (1/2000) - negative

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Atomic number and mass


Colums - Groups - No of electrons in outer shell
Rows - Periods - No of shells

Isotopes - different forms of the same element, same no of P, different no of N

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History of the Periodic Table and electron shells

History of the Atom

Dobereiner - arranged elements in order of atomic mass. Put into groups of three called triads.

Newlands - Every eighth element has similar properties. Pattern broke down on third row. Ignored bc he didnt leave gaps for undisovered elements

Mendeleev - Order of atomic mass, left gaps, similar properties in vertical groups

Electron Shells
- Electrons occupy shells (energy levels)
- Lowest energy levels filled first
- 1st shell, 2 electrons... rest of shells, 8 electrons

Electronic configuration
- Atomic no = no of electrons. (2, 8, 8...)

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Ionic bonding

A shell with one or two electrons is keen to lose electrons, so they are stable - become a positive ion

- A nearly full shell keen to gain electrons to become stable - become negative ions

- Ions are very reactive, be easily attracted to passing ions with opposite charge

Ionic bonds between metals and non-metals produce giant ionic structures

- closely packed in lattice. Ions are not free to move so do not conduct electricity when solid.

- Very strong chemical bonds between all ions

e.g MgO and NaCl

- High melting points bc of strong attraction

- MgO has higher MP than NaCl because its made of Mg^2+ and O^2- ions, so double charge

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Dot and Cross Diagrams

Metals and Non-Metals form Ionic Bonds

- when writing ionic compound equations, remember to balance the charges! -

Dot and cross diagrams - all atoms end with full outer shell


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Covalent bonding

Covalent bonds - sharing electrons - non-metals and non-metals


- The atoms within the molecules are held together by very strong covalent bonds

- Forces of attraction between molecules are weak

- Low melting and boiling points so molecules are easily parted from eac other

- Dont conduct electricity because there are no free electrons

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Group 1 - Alkali Metals

As you go down Group 1, the alkali metals become more reactive - outer electron is easily lost becuase its further from the nucleus so less energy is needed to remove it

- Low melting/boiling point (compared to other metals)
- Low density - lithium, sodium and potassium float on water
- Very soft - can be cut

Always form ionic compunds because they want to lose the outer electron
is the Loss of Electrons

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Alkali metals (2)

Reaction with cold water produces Hydrogen Gas
- Lithium, sodium or potassium react vigorously with water
- Move around surface, fizzing and producing hydrogen
- Sodium and potassium melt in heat
- An alkali forms which is the hydroxide of the metal

e.g Sodium + Water --> Sodium hydroxide + Hydrogen

Alkali metal compunds burn different colours
- dip wire loop in hydrochloric acid to clean and moisten
- put loop in powdered compund, place into a blue Bunsen flame 

Lithium: Red flame
Sodium: Yellow/orange flame
Potassium: Liliac flame

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Group 7 - Halogens

Halogens have 7 electrons in outer shell so react by gaining one electron to form a negative ion 

As you go down group 7, the halogens become less reactive - theres less inclination to gain the extra electron to fill the outer shell when its further away from the nucleus

As you go down group 7 the melting/boiling points increase
At room temp:
Chlorine: reactive, dense, green gas
Bromine: dense, poisonous, orange liquid
Iodine: dark grey crystalline solid

Reduction - the gain of electrons

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Halogens (2)

The halogens react with Alkali metals to form Salts called Metal Halides

Sodium + Chlorine --> Sodium Chloride
Potassium + Bromine --> Potassium Bromide

More reactive Halogens will Displace less reactive ones

Chlorine + Potassium Iodide --> Iodine + Potassium Chloride

Chlorine + Potassium Bromide --> Bromine + Potassium Chloride

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- Held together with metallic bonds
- Allows outer electron(s) to move freely

- Creates 'sea' of delocalised electrons throughout metal

Strong attraction between negative delocalised electrons and closely packed positive ions - high melting and boiling points

Good conductors because of the deleocalised electrons which carry the current


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Superconductors and Transition Metals

At very low temperatures some metals are Superconductors
- Normally metals have electrical resistance, so some energy is wasted as heat
- At very low temperatures (less than -265°c) the resistance in some metals disappears completely so no energy is wasted
- Current would flow forever

However -265°c is expensive and hard

Transition metals and their compunds
make good catalysts.

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Thermal Decomposition and Precipitation

Thermal decomposition - breaking down with heat

Transition metal carbonates (substances with CO3 at end) break down on heating into a metal oxide (e.g copper oxide, CuO) which usually results in a colour change

Copper(ll) carbonate CuCO3 --> copper oxide CuO + carbon dioxide CO2

Precipitation - two solutions react to form a solid
transition metal compund reacts with sodium chloride to form an insoluble hydroxide which precipitates out
CuSO4 + 2NaOH --> CU(OH)2 + Na2SO4
copper(ll) sulfate + sodium hydroxide --> copper(ll) hydroxide + sodium sulfate

Copper(II) hydroxide = blue solid
Iron(II) hydroxide = grey/green solid
Iron(III) hydroxide = orange/brown

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Water Purity

Surface water - Lakes, rivers and resevoirs which can start to run dry in Summer
Groundwater - aquifers (rocks that trap water underground)
Resources are limited depending on annual rainfall, and demand increases yearly.

Surface water needs a lot of treatment:
1. Filtration - wire mesh gets rid of large twigs etc and then gravel and sand beds remove any other solid bits
2. Sedimentation - iron sulfate or alluminium sulfate is added which makes fine particles clump together and sink to bottom
3. Chlorination - Chlorine gas is bubbled through to kill harmful bacteria and other microbes

Tap water can still contain impurities:
- Nitrate residues from fertiliser run-off
- Lead comounds from old lead pipes
- Pesticide residues from spraying near lakes and rivers

Sea water can be distilled to produce drinking water, but it is expensive and not practical

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Testing Water Purity

Test for Sulfate Ions using Barium Chloride
1. Add some dilute hydrochloric acid to test sample
2. Add 10 drops of Barium Chloride solution
3. If you see a white precipitate, there are sulfate ions in the sample

e.g barium ions + sulfate ions --> barium sulfate

Test for Halide Ions using Silver Nitrate
1. Add some dilute Nitric Acid to test sample
2. The 10 drops of Silver Nitrate solution
3. If Halide Ions are present, a precipitate will form

Chloride Ions = white precipitate    AgNO3 + NaCl --> AgCl + NaNO3
Bromide Ions = cream precipitate   AhNO3 + NaBr --> AgBr + NaNO3
Iodide Ions = yellow precipitate      AgNo3 + NaI --> AhI + NaNO3

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