Additional Science- C4

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  • Created by: Janviixo
  • Created on: 14-12-15 18:12

Atoms

Nucleus contains protons and neutrons

Nucleus is positively charged, almost whole mass is concentrated in nucleus

Electrons move around nucleus in shells

Electrons are negatively charged, they are tiny but cover lot of space

Volume of orbit determines size of atom, virtually have no mass (0.0005)

Neutral atoms have no charge, same number of protons and electrons

Mass number = number of protons and neutrons

Atomic number = number of protons. Atoms of same element have same no. of protons

Neutron number = mass number - atomic number

Periodic table: atomic number = mass number

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Elements & Isotopes

Periodic table is all known elements:

Reactive metals, transition elements, post-transition elements, non-metals and noble gases

Columns (down) = electrons on outer shell

Rows (across) = number of full shells

Isotopes are different forms of the same element, which have the number of protons but different number or neutrons

They have different atmoic mass but same atomic number otherwise it'd be different element

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Electron Shells

Electrons always occupy shells

Lowest engery levels filled first

Only certain number of electrons allowed in each shell: 1st = 2  2nd = 8  3rd = 8

Electron configuration = number of electrons on each shell e.g. Argon has 18 electrons, its electron configuration is 2,8,8

Electron configuration can be used to find period, group and atomic number of element

Argon is in group 8, 8 electrons on outer shell. In period 3, 3 shells. Atomic number is 18, add up no. of electrons.

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Ionic Bonding

Ionic bonding- to lose or gain electrons to form charged particles (ions)

Shell with 1/2 electrons on outer shell keen to lose it. Shell with 6/7 electrons find extra electron

Lose electrons = positive charge

Gain electrons = negative charge

Ions very reactive, attract to ion with opposite charge and sticks to it

Ionic bonds form between metals and non-metals, produce giant ionic structures. They form closely packed regular lattice arrangement. Ions not free to move, don't conduct electricity

MgO and NaCl are both giant ionic structures, have high boiling and melting points. When melted, ions free to move so conducts electricity

NaCl dissolves to form solution that conducts electricity.

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Ions and Ionic Compounds

Simple ions- groups 1&2 and 6&7

Ions try to get full outer shell

Metals lose electrons

Non-metals gain electrons

Metal + Non-metal = Ionic bond

To work out formula of ionic compound, you have to balance the +ve and -ve charges

'Dot and Cross' diagram shows what happens in ionic bonds

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Covalent Bonding

Covalent- sharing electrons

When non-metals combine they form covalent bonds. Both atoms get a full outer shell

Each covalent bond provides 1 extra shared electron

Hydrogen- form single covalent bond

Chlorine gas- form single covalent bond

Water- shares two electrons

Carbon dioxide- need 4 extra electrons, 2 double covalent bonds

Simple molecular substances ( CO2, H2O) held together by very strong covalent bonds

Force or attraction very weak, melting/ boiling points low

Molecular substance don't conduct electricity, no free electrons

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Group 1- Alkali Metals

Group 1 metals inlcude lithium, sodium, potassium, rubidum etc

As you go down becomes more reactive, outer electron more easily lost, further away from nucleus

Low melting/ boiling points, low density, very soft, forms ionic compounds

Alkali metals + water reacy vigorously , move around surface fizzing, produces hydorgen gas. kali forms the hyroxide of the metal

Sodium and Potassium melt in heat of reaction

Dip wire loop into hydorchloric acid to clean it, put loop in powdered sample and place in burner

Lithium = red

Sodium = yellow/ orange

Potassium = purple/ lilac

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Group 7- Halogens

Group 7 made up of flourine, chlorine, bromine, iodine and astatine

react by gaining 1 electron to form -ve charge

As you go down becomes less reactive, further away from nucleus less inclination to gain electron. Boiling/ Melting points increase

Room Temp= chlorine- fairly reactive, poisonous, dense green gas: bromine- dense, poisonous orange liquid: Iodine- dark grey crystalline solid

Halgens react with Alkali metals to form salts 'metal halides'

More reactive halogens displace less reactive ones

e.g Chlorine + Potassium Iodide ------> Iodine + Potassium Chloride

                Cl2    +    2Kl        ----->        Br2    +     2KCl

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Metals

Metals held together with metallic bonds, allow outer elctrons to move freely (delocalised = free)

Metals are hard, dense and shiny, strong attraction between delocalised electrons and the +ve ions

Strength and melting point decrease as atomic radius increases

Metals have high tensile strength- strong and hard to break but malleable

More delocalised electrons = good conducter of heat and electricity

Saucepans- conducts heat, doesn't rust easily: stainless steel, cheap

Electrical wiring- conducts electricity, easily bent: copper, best conductor

Aeroplanes- low density, strong, doesn't corrode: aluminium/ titanium (expensive)

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Superconductors and Transition Metals

All metals have electrical resistance, heats up when electricty flows through

When metals cold enough, lose electrical resistance completely- superconductors

Make power cables, really strong electromagnets, electronic circuits

Metals only superconduct when -265 C: getting that cold is expensive

Metals in middle of periodic table = transistion metals

Transition metals and their compounds make good catalysts

Iron catalyst- haber process

Nickel catalyst- hydrogenation of alkenes

They are colourful due to transition metal ion they contain.

Iron (II) = light green      Iron (III) = orange/brown       Copper = blue

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Thermal Decomposition and Precipitation

thermal decomposition- breaking down with heat

transition metals carbonates break down (iron carbonate): break down into metal oxide and carbon dioxide- results in colour change

Precipitation- a solid forms in solution

some transition metals react with sodium hyrdoxide to form insoluble hydorxide

copper (II) sulphate + sodium hydroxide -----> copper (II) hydroxide + sodium sulphate

Use precipitation to test for transition metals, they have distinctive colours

Add sodium hydroxide to salt:

copper (II) hydroxide- blue

iron (II) / (III) hydroxide- grey or green/ orange or brown

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Water Purity

We get water from lakes, rivers, reservoirs and aquifers

Resources are limited and demand increases every year

Water is purified in water treatment plants

filtration- wire of mesh screens out large twigs etc. gravel and sandbeds filter out othe solid bits

sedimentation- iron/aluminium sulphate added which makes fine particles clump together & settle

chlroination- chlorine gas bubbles through to kill bacteria and microbes

tap water can still contain impurities- must meet strict safety standards

pollutants still found, comes from nitrate residues, lead compounds, pesticides residues

you can get fresh water by distilling sea water, needs lots of energy so really expensive and not practical for large quantities

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Testing Water Purity

Water companies test water to check pollutant levels don't exceed strict limits. You can test dissolved ions using precipitation reactions

Test for sulpahte ions:

Add some dilute hydrochloric acid to test sample

Add 10 drops of barium chloride solution

Test for halide ions:

Add dilute nitric acid to test sample

Add 10 drops of silver nitrate solution

Iodide ions = pale yellown precipitate

Chloride = white

Bromide = cream

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