Group 7 - Halogens and Redox Flashcards
Define oxidation and an oxidising agent
The process of electron loss
Accepts electrons (itself is reduced) allowing other substances to be oxidised
Define reduction and a reducing agent
The process of electron gain
Donates electrons (itself is oxidised) allowing other substances to be reduced
Diatomic and uncombined elements have oxidation states of ?
0
Combined oxygen has an oxidation state of?
-2
Unless with peroxides where it is -1
Combined hydrogen has an oxidation state of?
+1
Except metal hydrides where its -1
Combined fluorine always has an oxidation state of?
-1
Increasing oxidation state numbers signify?
Oxidation
Decreasing oxidation states indicate?
Reduction
Define electronegativity?
The power of an atom to withdraw electron density to itself in a covalent bond
Describe and explain the trend in electronegativity down group 7?
Decreases down the group
Atoms get larger down the group meaning the outer electrons are further from the nucleus.
Increased shielding - Increased number of electron energy levels.
Weaker attraction between nucleus and outer electrons.
Describe and explain the trends in boiling points down G7
Melting and boiling points increase
Larger atoms have more electrons and therefore the VDW forces between molecules are stronger.
Colour and physical state of fluorine
Yellow
Gas
Colour and physical state of chlorine
Green
Gas
Colour and physical state of bromine
Red/brown
Liquid
Colour and physical state of iodine
Grey
Solid
Describe and explain the trend in oxidising ability of halogens down G7
Oxidising ability decreases
When halogens react they gain an electron so as the atoms become larger down the group. The force of attraction between the nucleus and outer electrons is weaker. Therefore, its harder for the halogen to gain an electrons and react.
Generally a halogen will displace a halide from solution if the halide is below it in the PT
What is disproportionation?
When an element is both oxidised and reduced
What are chlorate ions used for?
Pros and cons
Killing bacteria
Pros:
Kills pathogenic microorganisms
Prevents growth of algae which eliminates bad tastes and smells
Cons:
Chlorine gas irritates the respiratory system
Liquid chlorine can cause severe chemical burns
Can react with organic compounds in water to form carcinogenic chlorinated hydrocarbons
Describe the trend in the reducing ability of halide ions down group 7
Reducing power increases
The larger the ion the more easily it can lose and electron as the outer electrons are further away from the nucleus and there is a greater shielding affect.
How can the reactions of sodium halides with sulfuric acid reflect the reducing power of the halides?
Sodium chloride reacts with H2SO4 to form NaHSO4 and HCl gas.
This is not a redox reaction as the cl- is not a strong enough reducing agent to reduce the sulfur. (Oxidation state remains the same)
Sodium Bromide reacts with H2SO4 in 2 ways:
1. A similar acid base reaction to NaCl
2. -> SO2 + H2O + Br2 - Here the Br- ions are strong enough reducing agents to reduce the sulfuric acid to sulfur dioxide. Redox reaction
Sodium Iodide reacts with H2SO4 similarly as before in a acid base reaction but also can further reduce H2SO4 to H2S (smells of bad eggs)
How do you test for Halide Ions?
Add dilute nitric acid to remove other ions.
Then add silver nitrate solution
F- = no ppt
Cl- = white ppt
Br- = cream ppt
I- = yellow ppt
To be extra sure you can test by adding ammonia solution.
Cl- ppt will dissolve in dilute ammonia
Br- ppt will dissolve in conc ammonia
I- ppt will not dissolve in **conc ammonia
Reaction of chlorine with water?
And in sunlight?
Cl2(g) + 4H2O(l) ⇋ HClO(aq) + HCl(aq)
Disproportionation occurs
2Cl2(g) + 2H2O(l) -> 4HCl (aq) + O2(g)
Reaction of chlorine and sodium hydroxide
What is the household name for the product?
Cl2(g) + 2NaOH(aq) -> NaClO(aq) (bleach)+ NaCl(aq) +H2O(l)
Disproportionation occurs
