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Balancing Chemical Equations

Balancing Chemical Equations. Balanced Equation. Atoms can not be created or destroyed All atoms we start with we must end up with A balanced equation has the same number of each element on both sides of the equation. C + O 2  CO 2 This equation is already balanced

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Balancing Chemical Equations

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  1. Balancing Chemical Equations

  2. Balanced Equation • Atoms can not be created or destroyed • All atoms we start with we must end up with • A balanced equation has the same number of each element on both sides of the equation.

  3. C + O2 CO2 • This equation is already balanced • But what if it isn’t already?

  4. C + O2 CO • We need one more oxygen in the products. • Can’t change the formula, because it describes what is.

  5. The other oxygen must be used to make another CO • But where does the other C come from?

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

  7. 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

  8. Never • Never change a subscript to balance an equation CO2 • If you change the formula you are describing a different reaction. • H2Ois a different compound than H2O2 • Never put a coefficient in the middle of a formula • 2 NaCl is OK Na2Cl is not.

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

  10. Example H2 + O2 H2O Need twice as much O in the product

  11. Example H2 + O2 2H2O Changes the O

  12. Example H2 + O2 2H2O 2 Also changes the H

  13. Example H2 + O2 2H2O 4 2 Now we need twice as much H in the reactant

  14. Example 2H2 + O2 2H2O 4 2 Recount to check

  15. Example 2H2 + O2 2H2O Your answer 4 4 2 Recount to check

  16. 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

  17. 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

  18. 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

  19. 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

  20. 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.)

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