C6

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Magnesium and HCL react to produce H2 Gas

1. Add set volume of dilute hydrochloric acid to conical flask, & put on mass balance
2. Add magnesium ribbon to acid & quickly plug the flask with cotton wool
3. Start stop watch & record mass from balance, take readings at regular intervals
4. Plot results in a table and work out lost mass
5. Plot graph with time (x) and mass most (y)
6. Repeat previous steps with more concentrated acid solutions (keep everything else the same)

You should find the higher a concentration the faster the rate of reaction

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Sodium thiosulfate and HCL produce cloudy precipit

1. Add set volume of dilute sodium thiosulfate to conical flask
2. Put flask on a black cross
3. Add dilute HCL and start stopwatch
4. Time how long it takes for the black cross to disappear through the cloudy Sulfur
5. Repeat areaction with two more concentrations of HCL ( keep everything else the same)

Should find that as concentration of hydrochloric acid (HCL) increases time taken for cross to disappear is less time. So higher concentration quicker RoR as mark disappears quicker

This reaction doesn’t form a graph

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Reversible reaction

Reversible reactions reach equilibrium
A+B ⇌ C+D
As reactants react their concentrations fall so forward reaction slows but as more product is made their concentrations rises and so speeds up the backwards reaction
After a while the forward and backward reactions will be going at the same rate so it has reached equilibrium- where each reaction is still happening but there is no effect.
Equilibrium only takes place if the reaction takes place in a closed system -no products or reactants can escape or get in

They can be exo/endothermic.
If it’s endothermic in one direction it will be exothermique the other.
Energy transferred from surroundings by endothermic reactions is the same as energy to surroundings in exothermic reaction.
E.G. thermal decomposition of hydrated copper sulfate
Heating blue hydrated copper sulfate crystals removes water and leaves white anhydrous copper sulfate powder -ENDOTHERMIC
If u add water to white powder the blue crystals come back- EXOTHERMIC

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Position of equilibrium

A reaction at equilibrium doesn’t mean the amount of product/reactants are equal
If equilibrium lies to right- concentration of product is greater than reactant
If equilibrium lies to left- concentration of reactant is greater than product

Position of equilibrium depends on temperature, pressure, concentration

Temperature example:
Ammonia chloride ⇌ ammonia+hydrogen chloride

Heating moves equilibrium to the right (more ammonia and hydrogen chloride) and cooling it moves it to the left (more ammonia chloride)

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Le Chatelier’s principal

The principle: if you change conditions of a reversible reaction at equilibrium the system will try counteract the change.

Temperature:
If decrease temp equilibrium moves in direction of exothermic to produce more heat -more product for exothermic reaction-
If raise temp, moves to endothermic to lose heat-more product for endothermic reaction

Pressure: equilibrium involving gases only
Increase pressure, moves to direction of fewer molecules of gas, decrease pressure, direction of more molecules
USE BALANCED SYMBOL EQUATION TO FIND TTHE MOLECULE AMOUNTS
Concentration:
Increase either products/reactants it’s no longer at equilibrium

Higher concentration more products are made or of decrease products, reactants reduced

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