Balancing Chemical Equations

Balanced Equations Atoms can not be created or destroyed All atoms we start with we must end up with…and vice versa! A balanced equation has the same number of each element on both sides of the equation.

But what if an equation isn’t already balanced? C + O2  CO2 This equation is already balanced But what if an equation isn’t already balanced?

We need one more ________ in the products. Like this…. C + O2  CO We need one more ________ in the products. You can’t change the formula, because it describes what is being produced…CO (Carbon Monoxide.)

The other oxygen must be used to make another CO This gives a better/closer idea of what is happening… BUT…. The other oxygen must be used to make another CO But where does the other C come from?

Must have started with two C’s 2 C + O2  2 CO

Rules for Balancing Write the correct formulas for all the reactants and products Count the number of atoms of each type appearing on both sides. Balance the elements one at a time by adding coefficients (the numbers in front) 2 CO2 Check to see if it is balanced

Never Never change a subscript to balance an equation CO2 If you change the formula you are describing a different reaction. H2O is a different compound than H2O2 Never put a coefficient in the middle of a formula 2 NaCl is ok Na2Cl is not.

Example H2 + O2  H2O Make a table to keep track of atoms

Example H2 + O2  H2O R P 2 H O 1 Need twice as much O in the product

Example H2 + O2  2H2O R P 2 H O 1 Changes the O

Example H2 + O2  2H2O R P 2 H O 1 2 Also changes the H

Example H2 + O2  2H2O R P 2 H O 1 4 2 Now we need twice as much H in the reactant

Example 2H2 + O2  2H2O R P 2 H O 1 4 2 Recount to check

Example 2H2 + O2  2H2O Your answer R P 2 H O 1 4 4 2 Recount to check

Types of Reactions Millions of reactions Too many to remember They fall into several categories We will focus on Double Replacement in today’s lab

Double Replacement Two things replace each other Reactants must be two ionic compounds or acids. Usually in aqueous solution NaOH + FeCl3  The positive ions change place NaOH + FeCl3  Fe+3OH- +Na+1Cl-1 NaOH + FeCl3  Fe(OH)3 + NaCl

Double Replacement Will only happen if one of the products Doesn’t dissolve in water and forms a solid (look at solubility rules) Or is a gas that bubbles out Or is a covalent compound usually water After adding lead nitrate Potassium iodide 2KI(aq) + Pb(NO3)2 (aq)  2KNO3(aq) + PbI2 (s) PbI2 lead (II) iodide is insoluble

General Rules for the Water Solubilities of Common Ionic Compounds Compounds that are mostly soluble: All nitrates Alkali metal (group 1A) and ammonium compounds Chlorides, bromides, and iodides, except for those of Pb2+, Ag+, Hg2+ Sulfates except for those of Sr2+, Ba2+, Pb2+, and Hg2+ CaSO4 is slightly soluble

General Solubility Rules Compounds that are mostly insoluble: Carbonates, hydroxides, and sulfides, except for ammonium compounds and those of the group 1A metals. (The hydroxides and sulfides of Ca2+, Sr2+, and Ba2+ are slightly to moderately soluble.)