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    at dynamic equilibrium the rate of the forward reaction and backward reaction are equal
    ------------------------------------------------
    so if two reactants react and reached dynamic equilibrium how does the product form if the rate of the forward reaction and backward reaction are equal
    and does it happen in all the chemical reaction
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    Think of it like earning and spending money..... you are at dynamic equilibrium when you exactly spend what you make each week/month..... the overall amount of money you have in the bank account isn't changing, but money is constantly going in and out.

    Dynamic equilibrium requires an equilibrium reaction (not many are) and a closed system. Think about a bottle of coke, if it's sealed (closed system) it'll stay fizzy (dynamic equilibrium, concentration of dissolved CO2 remains constant), if it's open it'll go flat (not a closed system, CO2 gas escapes).
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    (Original post by gingerbreadman85)
    Think of it like earning and spending money..... you are at dynamic equilibrium when you exactly spend what you make each week/month..... the overall amount of money you have in the bank account isn't changing, but money is constantly going in and out.

    Dynamic equilibrium requires an equilibrium reaction (not many are) and a closed system. Think about a bottle of coke, if it's sealed (closed system) it'll stay fizzy (dynamic equilibrium, concentration of dissolved CO2 remains constant), if it's open it'll go flat (not a closed system, CO2 gas escapes).
    One small issue here..

    The vast majority of reactions (if not all) are equilibrium reactions, it's just that the position of equilibrium lies so far to one side (or the other) that the reaction seems to go to completion or not at all.

    The very fact that the activation energy for any reversible process is non-infinite means that the reverse reaction is also 'possible'.

    Use of ΔG to predict feasibility of reaction only appplies under standard conditions.
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    (Original post by charco)
    One small issue here..

    The vast majority of reactions (if not all) are equilibrium reactions, it's just that the position of equilibrium lies so far to one side (or the other) that the reaction seems to go to completion or not at all.

    The very fact that the activation energy for any reversible process is non-infinite means that the reverse reaction is also 'possible'.

    Use of ?G to predict feasibility of reaction only appplies under standard conditions.
    Pedant much? While what you say is strictly true, given that the OP is struggling to understand equilibrium, ?G, equilibrium constants etc is a little much to be going into in answering their questions. I'd rather not confuse them any more than they already are.
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    In industry, to get the most useful product, the forward conditions are amplified to get the best yield. They then quickly remove this useful product to prevent it from converting to reactants again. This way the concentration gradient is still steep, so more product will be made.

    Hope this was useful ;/
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    (Original post by reb0xx)
    at dynamic equilibrium the rate of the forward reaction and backward reaction are equal
    ------------------------------------------------
    so if two reactants react and reached dynamic equilibrium how does the product form if the rate of the forward reaction and backward reaction are equal
    and does it happen in all the chemical reaction
    Luffy avatar :five:

    Secondly a dynamic equilibrium is just an equilibrium that is not static i.e. it does not really stop occurring rather the forward and backward rates are constant as you said. This does not mean they are equal however, and in addition, if you put in something from one side of the equation, the dynamic equilibrium will react to counteract such a change so as to restore it to the equilibrium.

    So you asked how does it form if they are equal, they are most likely (if not always) not equal, they remain in the dynamic by keeping their constant rates so one side is dependant on the other. So one side could have a higher yield than the other, it merely means once the equilibrium point is reached it will stay like at this ratio of yields and any changes will be liable to counteraction by the reaction to restore equilibrium. And yes of course it happens in the chemical reaction, where else would it occur :p:
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    thx
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    (Original post by reb0xx)
    at dynamic equilibrium the rate of the forward reaction and backward reaction are equal
    ------------------------------------------------
    so if two reactants react and reached dynamic equilibrium how does the product form if the rate of the forward reaction and backward reaction are equal
    and does it happen in all the chemical reaction
    Let's say that the rate of forward reaction is given by:

    \text{rate}_\text{forward} = k_\text{forward} [\text{reactants}]

    And backward reaction is given by:

    \text{rate}_\text{backward} = k_\text{backward} [\text{products}]

    So, when they're equal (i.e. equilibrium), we have:

    k_\text{forward} [\text{reactants}] = k_\text{backward} [\text{products}]

    Rearrange for the concentration of the products:

    [\text{products}] = \frac{k_\text{forward}}{k_{ \text{backward} }} [\text{reactants}]

    So, when the forward rate constant is much bigger than the backward rate constant, the concentration of products is much bigger than the concentration of reactants.

    That's what happens when you carry out a chemical reaction until equilibrium. The products are just what you get if you follow the direction with the biggest rate constant.
 
 
 
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