Chemical Equilibrium Flashcards
What are they?
Many chemical reactions continue until one of the reactants is completely used up and the reaction stops (go to completion). Other reactions take place in both directions and are reversible: the conversion from reactants to products is incomplete no matter how long the reaction is allowed to continue (such reactions will eventually reach a state of equilibrium in which both reactants and products are present).
If you disrupt (w/ addition?), one will go faster, other slower, until they meet at a new equilibrium (?). Can go forward to products or back to reactants (will eventually reach equilibrium): always has some reactants and some products (never used up).
*Acids produce hydrogen ions; weak acids don’t completely dissociate (weakness/strength indirectly related to pH level).
Properties
Reaction is dynamic (still occurring in both directions [static is when nothing changes—dynamic, changing. In terms of direction?]).
Takes time for reaction to reach equilibrium (so, not spontaneous).
Can only be reached in a closed system (where species can’t escape [ex. w/ demo, splashing water, or gas being produced and we’re letting it escape).
The rate of the forward reaction equals the rate of the backward reaction (really, the definition [has to be—though others do too, of course]): the reaction appears to have stopped, but it is still occurring equally in both directions (like, if weights are matched, running up an escalator the wrong way: if we look from the side, looks like he’s in the same position, but he’s just running up at the same rate it’s moving down).
Concentrations of products and reactants do not appear to change (constant [even though products are forming into reactants, and reactants into products, the concentrations don’t change. This doesn’t, however, mean that they’re the same]).
Le Chatelier’s Principle
A + B ⇋ C + D: more A means a greater chance of colliding w/ B, which means more C + D, which means more A + B, which means more C + D, and so on…like a never-ending cycle. So, you’re also making more C + D (rate going up) b/c of concentration: if more A, reaction increases, and over time, b/c A + B are being used up, it goes down, and, when equal w/ other, at equilibrium.
*Ammonia is an equilibrium reaction: if you don’t make it right, it could reverse (you have to have the right conditions to make as much as possible).
Equilibrium can be disturbed (ex. adding more reactants, taking something out, changing pressure/temp, etc.: constraints are changes in concentration, temperature, and pressure) and the position of equilibrium can be moved to the left/right.
Obey Le Chatelier’s Principle: When a constraint is applied to a system in equilibrium, the system will move in such a way as to remove/cancel out the effect of the constraint (position of equilibrium changes to minimize the effect of any imposed change in conditions).
Reacting to form each other. Will do the opposite of what you do (if you add more A, it will remove A: it will do something to return to equilibrium).
Changing concentration
Equilibrium position will shift away from the increased concentration: if we lowered the conc of C, it would shift to the right to try to make more C.
If the concentration of any reactant or product is changed, then the equilibrium will move to undo the change in reactions: when the concentration of reactant A is increased, position of equilibrium shifts right (towards products), forward reaction rate increased and more product is formed, and, as the concentration of product has now increased, the backward rate also starts to increase until equilibrium is re-established.
Equilibrium means equal (so rates are equal): if disturbed, no longer in equilibrium.
If you increase conc of A, rate forward will increase, and as you start to use A + B up, rate decreases: as you start to make more C + D, backward will increase (eventually, it will shift to the right to undo A and help it reach equilibrium).
*We always talk about it in the way it’s written (left side is reactant side, right is product).
Change in pressure
Equilibrium position will shift away from the increased pressure to the side with fewer moles of gas (if you decrease, will shift to side more).
Pressure really only matters w/ gases (tiny bit w/ liquids [hydraulics is abt the fact that you can’t really compress liquids]). An increase in pressure for gases is equivalent to an increase in concentration: if you increase the pressure, it will want to decrease it, so it will shift to the side with less moles of gas.
The reaction will be faster and equilbrium will be established more quickly. Position of equilibrium is only affected if the reaction involves gases where there are different numbers of has molecules (moles) on either side of the equation (if same # of moles, no change).
If we increase the pressure, the equilbrium position moves to try to reduce the pressure: moving to the side with the least number of gas molecules will lower the pressure. Same container volume with more moles of gas means higher pressure. Add up moles of gas on each side (coefficients) and compare.
Increase in pressure means more ammonia (though you need a compressor, which is expensive): you shift the reaction to the product side, which is what you’re trying to make (really important industrial reaction: before process was discovered, crop output worldwide w[ammonia is in a lot of cleaning products, but its biggest use is in fertilizers: it’s a way to get nitrogen to plants, so it’s important to feed the world]). Bleach and ammonia make chlorine gas, which is poisonous.
*Also, an important part of the process to make sulfuric acid. Also, if you have bromide or iodide ions, if you add silver nitrate and it produces NO3, you can use it to measure reaction rate (AgCl precipitate [like disappearing cross]).
Change in temp
Changing the temp changes the rate of the forward and reverse reactions by different amounts. An increase in temp increases the rate of reaction (in both directions [not necessarily by the same amount, though]).
*Cl and CaCl4 are not the same: if you’re adding Cl, consider that.
W/ endo, write on opposite. For exothermic reactions, compromise temperatures are usually used which are high enough to increase the rate of reaction, but not too high that equilibrium is too far to the left
If you increased temp, you’d reach equilibrium faster b/c heat makes everything happen faster. An increase in temp increases the rate of reaction but does not necessarily increase the yield of product. The direction of the equilibrium shift depends on whether the forward reaction is exothermic or endothermic. If forward reaction if exothermic, the reverse will be endothermic.
If exo (negative change in enthalpy), heat is a product (surroundings get hot): if endo (positive change in enthalpy), heat is absorbed (it’s a reactant; it needs heat). If exothermic and we’re increasing the heat, it will go the other way to decrease it. Think of it like concentration. If endo, on left; if exo, on right.
Using up heat? Breaking the bonds between the water and CO2, so it’s endothermic. If you increased pressure there, it’d shift left (ex. opening a soda bottle: when you freeze soda, it can no longer hold the gas, so the gas comes out and the bottle could explode). With pressure, we’re making it more compact (if increased)?
Addition of a catalyst
Has no effect on position of equilibrium: enable equilibrium to be established more quickly and speed up both the forward and reverse reaction by the same amount).
In other words, it takes less time for equilibrium to be reached (still in equilibrium, so no position shift).
Haber process
Produces ammonia, which is used to make nitrogen-based fertilizers and polymers such as nylon (compromise temp and heated iron catalyst used).
We get hydrogen from steam reforming (we use methane), so it’s a hydrocarbon (something we’ll run out of?).
Very energy-intensive (liquify air [get it really cold to condense]; oxygen and nitrogen boil at different temps, so to separate them, they put them in a column and let one boil off [keep one as a liquid]).
Contact process
More sulfuric acid is manufactured in the world than any other chemical (about 110 million tonnes per year) using the Contact Process.
Starts with sulfur, which is burned in air to produce SO2
The SO2 is oxidized (burned) further to form SO3 (a equilibrium process). Sulphuric acid is then produced from the SO3. The production of the SO3 from SO2 is the bottleneck in the process, the reaction being reversible and slow. The conditions used for this reaction are 450°C, a V2O5 catalyst, and P = 1-2 atm.
But 1-2 atm is not very high… this is because even at relatively low pressures, there is a 99.5% conversion of sulfur dioxide into sulfur trioxide. So there is an expense for increasing the pressure and very little to be gained, which does not make sense economically. Even at lower pressures, if you find a decent compromise temp and add a catalyst, you’ll still get a really good conversion rate.
*Decreasing heat also lowers reaction rate, so compromise.
NOTE
- Do reviews…
- Consider rate
- Questions for AFTER you make change (?)
