Explain what is wrong with the following statements. (a) Once a reaction has reached equilibrium, all reaction stops. (b) If more reactant is used, the equilibrium constant will have a larger value.

In the world of chemistry, the equilibrium constant is a vital number that tells us a lot about a chemical reaction. It represents the balance point of a reaction when the rate of the forward reaction equals the rate of the reverse reaction. Let's imagine a simple seesaw with products on one side and reactants on the other. The equilibrium constant, denoted as K, is like the point where the seesaw doesn't tilt either way, it's perfectly horizontal.

Now, let's clear up a common mistake: some think that if you pile on more reactants onto our imaginary seesaw, K will change. But that's not quite right! K stays the same, no matter how much reactant you add. Why? Because it's determined by the ratio of the products to the reactants at equilibrium — think of it like a recipe that always needs the same proportions of ingredients to taste just right, no matter how much you make.

Here's the equation for the equilibrium constant:
\( K = \frac{[products]}{[reactants]} \)
Where the square brackets represent the concentration of substances and we raise them to the power of their coefficients from the balanced chemical equation. The key takeaway: K is consistent for a given reaction at a certain temperature, and does not change with the amount of reactants or products you start with.

Let's talk about reaction rates in a way that would help anyone understand, like comparing it to a dance-off between two teams. In chemistry, when we say 'reaction rates,' we are talking about the speed at which the reactant molecules dance (react) to form product molecules. When a reaction starts, the teams are all hyped up and the dance-off (reaction) is super-fast. But as time goes on, the teams get tired (the concentration of reactants decreases) and the dance-off slows down.

At equilibrium, something neat happens — the speed of the forward team (forward reaction) matches that of the reverse team (reverse reaction). It's like both teams are matching each other's moves perfectly. This doesn't mean they’ve stopped dancing! They're still grooving, but now the score stays even. So, even when it looks like things are still, there's a whole bunch of action going on at a microscopic level.

To measure the speed of this chemistry dance, scientists use rates. The faster the reaction, the higher the rate. But remember, at equilibrium, the rates of the forward and reverse reactions are the same, and that special point is what makes chemical equilibrium so fascinating and important in understanding how reactions work.

What if we think of a balanced chemical equation like a perfectly portioned recipe? In cooking, if you don't balance your ingredients, your recipe doesn't turn out right. It's similar in chemistry. A balanced chemical equation ensures that we have the correct amount of each molecule dancing in our reaction so that none are left out and no one crashes the party with too much of themselves.

The law of conservation of mass tells us that we can't create or destroy atoms in a chemical reaction, so we have to balance the equation to show that. In other words, we're arranging the number of atoms of each element to be the same on both sides of the equation. This gives us a clear blueprint for understanding how much of a reactant we’ll need to produce a certain amount of product.

Here’s a simple example:
\[ \ce{2H_2 + O_2 -> 2H_2O} \]
This equation tells us that two molecules of hydrogen gas react with one molecule of oxygen gas to make two molecules of water. Easy, right? Just remember, like a perfect recipe, the numbers in front of each molecule (called coefficients) ensure that atoms are conserved and the dance of a chemical reaction can happen smoothly.