![]() Notice that oxygen occurs twice on the left-hand side, so that is not a good element to start out with. Let’s illustrate this with an example by balancing this chemical equation:įirst, let’s look at the element that appears least often. In the end, make sure to count the number of atoms of each element on each side again, just to be sure.Įxplore Balancing Chemical Equations on Albert Example of Balancing a Chemical Equation Then, move on to the atom that shows up the second least number of times, and so on. It is best to start with the atom that shows up the least number of times on one side, and balance that first. We can break the house apart and build an airplane, but the color and shape of the actual blocks do not change.īut how do we go about balancing these equations? We know that the number of atoms of each element needs to be the same on both sides of the equation, so it is just a matter of finding the correct coefficients (numbers in front of each molecule) to make that happen. An easy way to understand this is to picture a house made of blocks. This means that chemical reactions do not change the actual building blocks of matter rather, they just change the arrangement of the blocks. So, if we start with ten atoms of oxygen before a reaction, we need to end up with ten atoms of oxygen after a reaction. This stems from the universal law of the conservation of mass, which states that matter can neither be created nor destroyed. Of molecules we have and that's what these coefficients in front of these molecular formulas represent.The ultimate goal for balancing chemical equations is to make both sides of the reaction, the reactants and the products, equal in the number of atoms per element. But notice, we didn't change the actual structure of the molecules. And here some of the oxygensĪre in the carbon dioxides, some of them are in the waters. We have four hydrogen atoms on both sides. In order to balance it, we're going to have six oxygens on the reactant side. On this side we have one, two, three, four, five, six oxygens. So O H, and actually let me write this right over here, so Of these water molecules? So, let's draw another Say, "okay, we have four hydrogens here, "we only have two here." Well, what if we had two So now we've balanced the carbons, two carbons on the reactant side. We're just going to have another molecule. ![]() Or if you want to do it, if you want to have it more visual, you could write it, okay, we're just going to The actual structure here, so let's write the two out front. We're definitely producing carbon dioxide, but how many carbon dioxides do we produce for each molecule of ethylene? And so that's where we say, "Okay, we have two carbons here, "we want two carbons here." We don't want to change We don't want to do, we want to say, we're definitely You would be changing what this molecule is. Somehow write a subscript of two right over there, you would somehow be changing the structure, Your water right over here, this is an oxygen bonded to two hydrogens. Finally, water, actually I'm going to do this in a slightly different color. Carbon dioxide is a carbon, double bonded, to two oxygens each. Notice you have twoĬarbons and four hydrogens. Then each carbon isīonded to two hydrogens. It a little bit differently, let's draw each of these molecules So ethylene looks like this. And that's why we do notĬhange these subscripts. You can't change the number of constituents within the molecule. So the only thing that you can change when you're changing these is the number of molecules. When it's unbalanced, it just doesn't have the numbers right in terms of number of molecules. When you're balancing chemical reactions, the reaction itself is, even before it's balanced is describing something that happens. You're actually changing the reaction when you're doing that. It's no longer carbon dioxide, it's now this bizarre thing that doesn't really exist in nature. " And the reason why you can't do that is that's actually changing the molecule. Subscript little twos look, "so why not put a two right over there. Might have been thinking, "Why put this big two out front "of the entire carbon dioxide, "I like the way these little Is let's just put a two out front here and so now we have two for every molecule of ethylene and we're not done balancing And right now, I only have oneĬarbon on the product side. And so I would want twoĬarbons on the product side. And we've seen that, okay, if let's say, we're trying to balance thisĮquation right over here and we started with the carbons. We've now seen a couple of examples of balancingĬhemical equations.
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