#### Related to [[Chemical nomenclature]] --- ## How to balance chemical formulas - When a chemical reaction takes place atoms can **never be created or destroyed** and so the act of *balancing equations* ensures that the **reactants and products** follow this rule. - Generally **start with the largest/most complex molecules**. - Progress to less complex elements leaving **lone elements** last. >[!example] ></br> > >![[Balancing equations example.png]] --- ## Symbols - A **(g) in parenthesis** indicates that a molecule *or element* is in its gaseous state. - **(l) indicates that it is a liquid.** - **(s) indicates that it is a solid** *or precipitate*. - **(⇌) indicates that the reaction can go in both directions**, which means the products may react back into the *reactants.* --- ## Finding out the mass of product - To find the mass of the product first find the limiting reactant. Then find the stoichiometric ratio between that reactant and the product you are solving for. **Multiply the moles of the reactant by the stoichiometric ratio**, the result of this equation is the number of moles of **the product**. - For more info on moles and conversion between moles and mass see [[Moles]]. --- ## Limiting reactant/reagent - The limiting reactant/reagent is the **reactant that will run out first** thus, stopping the entire reaction even if there is still enough of other reactants. - In the following example hydrogen would be the limiting reactant because there is not enough of it to react **completely with the nitrogen.** >[!example] ></br> > >![[Balancing equations reagent.png|800]] - When **calculating** the limiting reactant **always use/convert masses into [[Moles]]** since mass varies between elements *(some elements are more dense)* but the number of moles does not. - Using the ratio of each element given determine *how many X would I need to react with X2 fully?*. If the answer is less than the actual amount X2 is the limiting reactant, however if it is **more** then X is. >[!quote] > >##### These concepts in action: ></br> > >![[Balancing equations example-1.png|800]] --- ## Stoichiometric ratio - This is the ratio of the product per mol of reactant consumed. - When calculating the stoichiometric ratio always put the moles of **the product on the top of the fraction** and the moles of the reactant consumed on the bottom. - In situations where you are finding the ratio between **two reactants** the **KNOWN REACTANT** *(meaning we know the mass or moles)* always goes on the bottom of the fraction! >[!example] >![[Balancing equations.png|800]] --- ## Balancing redox reactions - When balancing a redox reaction the most important thing to keep in mind is that **no electrons will be gained or lost**. This means that both the reactants and products will have the same number of total electrons. The only difference is that the oxidizing agent will **accept electrons** and the reducing agent **will give them away**. >#### BASIC steps to follow: > - When balancing a redo reaction **first determine the oxidation state of each element within that reaction**. > - Then split the reaction up into two different formulas. One showing the **reduction and one showing the oxidation**. > - To balance these formulas first make sure the coefficients are correct *(the number before the elements/compounds)*. Then add any H+ ions OH- ions or water that is necessary to balance out excess hydrogen or oxygen. > - Then **add electrons on the reactant side of the equation if they are lost or add them on the product side if they are used**. Make sure to consider the H+ and OH- ions you added in the previous step! *The reduction reaction should have them on the reactant side and the oxidation reaction should have them on the product side.* > - The final step is to **change the coefficients of elements so that the electrons lost in the oxidation reaction *EQUALS* the electrons gained in the reduction reaction.** This means there will be the same number of electrons in **both** of the half-reactions! > - Then all that is left to do is combine the half reactions. - It is important to note that if their is excess oxygen in the **reactants** **water will be produced**. In order for this to be balanced make sure **to add enough H+ ions on the reactant side of the equation** so that all the oxygen can be "used up" creating water. - If their is excess hydrogen H+ ions **make sure to add the appropriate number of OH- molecules *also on the reactant side* so that all the hydrogen ions can be used up creating excess water**. >[!tip] >- Remember that if you have to add H+ ions to the equation to balance out other oxygen atoms then the redox reaction is taking place in an *acidic solution.* >- If you need to add OH- to balance out excess H+ ions then the redox reaction is happening in a **basic solution.** *For more: [[Acids and bases]].* >[!example] >![[Stoichiometry and balancing.png|800]] --- #mainpage