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Chemistry
Unit One

Definitions
• Relative atomic mass: the average mass of an atom when compared with 1/12
of carbon-12
...

• Relative molecular mass: the average mass of a molecules when compared
with 1/12 of carbon-12
...

• Base: a species that is a proton (H+) acceptor
...

• Salt: where a positive metal ion replaces a H+ ion from an acid
...


•

Covalent bond: the sharing of a pair of electrons between two atoms
...


•

Electronegativity: the ability of an atom to attract the bonded pair of electrons in a
covalent bond
...


•

First ionisation energy: the energy required to remove one electron from each atom
in one mole of gaseous atoms to form one mole of gaseous +1 ions
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•

Disproportionation: when a single element is simultaneously oxidised and reduced
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0005, Charge -1

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Most of the mass is concentrated in the nucleus
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Mass number: the total number of protons and neutrons in the nucleus
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Ions: different number of electrons
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Isotopes have the same chemical properties, but different physical properties
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Atomic Models
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Greeks: believed matter was made from indivisible particles
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Thomson (1897): concluded atoms weren’t solid or indivisible – his measurements showed an
atom must contain smaller, negative particles
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Rutherford (1909):
Gold foil experiment: alpha (positive) particles were fired at a thin sheet of gold – most passed
straight through the gold atoms, very few were deflected (plum model couldn’t be right)
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Moseley: found positive charge increased from one element to the next
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Nucleus of
atom was too heavy for just protons – must contain neutral particles as well (neutrons)
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Bohr: proposed that electrons existed in fixed orbitals, each orbital has a fixed energy,
radiation emitted or absorbed has fixed frequency, shells must only hold a fixed number of
electrons
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The quantum model is more accurate
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02 X 1023 (Avogadro's number)
• Molar mass: the mass of one mole of something
...
02 X 1023

• Mass = Moles X RAM
(g) (mol) (g/mol)
• Moles = Concentration X Volume
(mol)
(mol/dm3)
(dm3)
• Volume of a gas = Moles X 24 dm3
(dm3)
(mol)

Formulas
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•

Molecular formula: the actual amount of atoms of each element in a molecule
...


•

If you know the empirical formula and the molecular mass, you can calculate its
molecular formula:
1
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2
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3
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•
1
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3
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Calculating empirical formulas from data:
Use the “mass=moles X ram” to work out how many moles of each product
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Write down ratio
...
of moles
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4g of carbon dioxide and 1
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What is the empirical formula?
1
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3
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Moles of carbon dioxide: 4
...
1 moles
Moles of water: 1
...
1 moles
1 mole of carbon dioxide contains 1 mole of carbon atoms so you must have
started with 0
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1 mole of water contains two moles of hydrogen atoms (H20) so you must have
started with 0
...
1 x 2)
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1:0
...
1/0
...
2/0
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Assume you’ve got 100g of the compound, then turn percentages into masses
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Work out the moles by using the “mass=moles X ram” equation
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Divide each number of moles by the smallest number of moles to give a ratio
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Apply, and round, the numbers from the ratio to the formula
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2
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4
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2
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4
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Work out how many moles of the reactant you have
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Calculate the mass of that many moles of product (
If working out gas volumes instead of masses, use the “moles X 24” equation
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4Fe + 3O2  2Fe2O3
Moles of Fe: 28g/55
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5 moles
Molar ratio of Fe:Fe2O3 is 2:1
Every 2 moles of Fe will produce 1 mole of Fe2O3
0
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25 moles (moles of Fe2O3)
RAM of Fe2O3 is 159
...
25 moles X 159
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• Base: a species that is a proton (H+) acceptor
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• Acid + Base  Salt + Water
• The base could be a metal oxide or a metal hydroxide
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Anhydrous: doesn’t contain water
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All solid salts form a lattice of positive and negative ions
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Many hydrated salts lose their water of crystallisation when heated
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Mass of water = mass of hydrated salt – mass of anhydrous salt
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Find the number of moles of water lost
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Find the number of moles of anhydrous salt produced
...
Work out the molar ratio of anhydrous salt to the moles of water
...
Divide each by the smallest mole, and round the answer
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For example: CuSO4
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2
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4
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6
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Titration: how much of an acid is needed to neutralise an alkali quantity
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Phenolphthalein: turns pink to colourless when adding an acid to an alkali
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Add acid to the alkali using a burette – open the tap to run acid into the alkali a
little at a time
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Do a rough titration to get an idea of the end point (just enough to neutralise)
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Record the amount of acid used
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Oxidation States
• The oxidation state of an element tells you the total number of electrons
donated or accepted when it’s in a particular compound or ion
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• Oxidation state of a simple monatomic ion is the same as its charge
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• Sum of oxidation states for neutral compounds is 0
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• Combined hydrogen is nearly always +1, except in metal hydrides where it’s -1
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Redox Reactions
• Oxidation: loss of electrons
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• Redox reaction: something loses electrons and something else gains them
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• Reducing agent: donates electrons and gets oxidised
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Each orbital holds two electrons
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In an orbital, two electrons spin in opposite directions due to repulsion
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P-Subshell orbital: dumbbell-shaped
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Rule Two: electrons fill orbitals singly before sharing
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Energy has to be put in to ionise an atom or molecule
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Factors that affect ionisation energy:
Nuclear charge: the more protons in the nucleus, the more positively charged it is, meaning a
stronger attraction between the nucleus and the electrons
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Shielding: as the number of electrons between the outer shell and the nucleus increases, the
outer electrons feel less nuclear attraction
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•

Within each shell, successive ionisation
energies increase
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• Ions are formed when electrons are transferred from one atom to another
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• Ionic compounds: an electrostatic attraction holds oppositely charged ions
together
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• Electrical conductivity: ionic compounds conduct when molten or dissolved –
not when solid
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• Melting and boiling points: ionic compounds have high melting and boiling
points as it takes a lot of energy to overcome the ionic bond force
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Covalent Bonding
•

Covalent bond: the sharing of a pair of electrons between two atoms
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Exceptions:
Boron trifluoride: boron only has 6 electrons in the outer shell
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Some compounds can use the d orbital to expand the octet
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Covalent compounds have strong bonds within the molecules (covalent bond), but
weak forces between the molecules (Van Der Waals)
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Melting and boiling points: low – there are weak forces to overcome
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Giant Covalent Structures
Graphite
• Arranged in sheets of flat hexagons – the carbons are covalently bonded with
three bonds each
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• The sheets of hexagons are bonded together by weak Van Der Waals forces
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Strong covalent bond – very high melting point
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• Insoluble – bonds are too difficult to break
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Giant Covalent Structures
Diamond
• The carbons have four covalent bonds – arranged in tetrahedral shape to form a
crystal lattice structure
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• Extremely hard solid
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• Can’t conduct electricity – electrons held in
localised bonds
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Shapes of Molecules
• Linear Shape
• Two bonded electrons
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• Bond angle: 109
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• Bond angle: 120˚

• Trigonal Pyramidal Shape
• Three bonded electrons
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• Bond angle: 107˚
Example: NH3

Shapes of Molecules
• Non-linear Shape
• Two bonded electrons
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• Bond angle: 116˚
Example: O 3

• Non-linear Shape
• Two bonded electrons
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• Bond angle: 104
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• Bond angle: 120˚ and 90˚

• Octahedral Shape
• Six bonded electrons
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Fluorine is the most electronegative element
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Covalent bonds in diatomic gases are non-polar
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•

In a polar bond, the difference in electronegativity between two atoms causes a
dipole
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Metallic Bonding
•

Metal elements exist as giant metallic lattice structures
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This leaves a positive ion
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Melting and boiling points: number of delocalised electrons per atom affects
melting and boiling points – more there are higher the points
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Conductivity: delocalised electrons pass kinetic energy to each other – good
thermal conductors
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Solubility: insoluble – due to strength of the metallic bonds
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Weakest intermolecular force
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This causes a temporary dipole
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The two dipoles attract to each other – Van Der Waals
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Larger molecules have more electrons and a greater surface area – stronger force
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When you boil a liquid, you need to overcome the intermolecular forces so the
particles can escape – liquids with strong Van Der Waals have higher boiling points
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Intermolecular Forces
Permanent Dipole-Dipole
• Second strongest intermolecular force
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• Molecules have these forces due to the shift in electron density
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• Hydrogen bonding happens between molecules that have hydrogen covalently
bonded to fluorine, nitrogen or oxygen
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• The bond is so polarised that the hydrogen of one molecule forms a weak bond
with the lone pair of F, N and O of another molecule
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Higher boiling and freezing points than non-polar molecules of similar size
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• In ice, water molecules are held together in a simple molecular lattice by
hydrogen bonds – relatively long and less dense than liquid water
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1817: Döbereiner attempted to group similar elements into Döbereiner’s triads
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Pattern broke down when transition metals were
organised
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This predicted properties of unknown elements
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He also added noble gases to the table
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Periods: show the number of electron shells
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Group 1 & 2: S-block
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Groups 3-8: P-block
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Periodic Trends
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Periodicity: the trends in repeating physical and chemical properties of elements across the
periodic table
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Macromolecular structures: have strong covalent bonds to overcome
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Noble gases: lowest melting and boiling points – exist as single atoms
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Ionisation energy 
Across a period: ionisation energy increases – number of protons increase so stronger nuclear
attraction (and smaller atomic radius), extra electrons add to the same level so very little extra
shielding effect
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Group 2 – The Alkaline Earth Metals
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When group 2 elements react, they lose electrons to form positive ions
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Ionisation energy decreases down group – reactivity increases
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Go down the group – hydroxides are more soluble
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Reaction: Group 2 carbonate  metal oxide + carbon dioxide
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Calcium hydroxide is used to neutralise acid soils
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Group 7 – The Halogens
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Fluorine: pale yellow gas
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Bromine: red-brown liquid
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• Boiling and melting points:
• Increase down the group – increasing strength of Van Der Waals forces due to
more electrons
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• Halogens are reduced to form negative halide ions
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Halogens - Displacement Reactions
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Halogen’s relative oxidising strength can be seen in their displacement reactions with halide
ions
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Chlorine will displace Br- and IBromine will displace IIodine will not displace fluoride, chloride or bromide
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g
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The halogen will dissolve and the solvent
settles out as a solvent layer
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Chlorine: pale yellow/green
Bromine: orange/red
Iodine: purple

Testing for Halides
• Halogens are distinctive to look at, but halide solutions are colourless
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Add dilute nitric acid remove ions that might interfere
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Add the silver nitrate solution (AgNO3)
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A precipitate of silver halide is formed
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• Chloride: white precipitate
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• Iodide: yellow precipitate
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To be sure, ammonia solution is added – each halide has a different solubility:
Chloride: precipitate dissolves in dilute ammonia solution
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Iodide: precipitate is insoluble
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• This happens when halogens react with cold, dilute alkali
solutions
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• Chlorine (Cl2) with an oxidation state of 0
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Disproportionation and Water Treatment
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Bleach: can be made by mixing chlorine gas with dilute sodium hydroxide at r
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p to
get a sodium chlorate solution
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Chlorine gas is toxic and liquid chlorine can cause chemical burns
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Fluoridated water:
Fluoride ions are added to drinking water to prevent tooth decay
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