1. Balancing Chemical Equations and Stoichiometry
Chemistry 100 Chapter 3
Balancing Chemical Equations and Stoichiometry

2. The Chemical Reaction What happens in a chemical reaction?
Example
2 H2 + O2  2H2O
Starting materials (reactants) are converted into different chemical substance(s) (the product(s)).
Described by ‘standard shorthand’ - chemical equation

3. Balancing Equations
Balanced equations have the same number of atoms of a given element on the LHS and the RHS
Law of conservation of mass.
All the reactants and products must be identified!
Chemical equations report the results of experimentation!

4. A General Method of Balancing Chemical Equations.
Write the basic, or ‘skeletal’ equation, showing the formula of all reactants and products.
Balance the equation according to the law of conservation of mass.
Balance by adjusting the coefficients in front of the chemical formulas, never by adjusting the subscripts.
Best to start with an atoms that appears only once on the left and right hand side of an equation.

5. A Balanced Chemical Equation

6. Chemical Reaction Types
Combustion reaction  can be written for any compound containing C, H, O or C, H, S etc. O2 (g) is present as a reactant and is usually in excess.
We can also have combination and decomposition reactions
Combination reaction  two or more substances combine to form one product
2Mg(s) + O2(g)  2MgO(s)

7. Reactions Types (II)
Decomposition reaction  one substance breaks down to form two or more different substances
MgCO3(s)  MgO(s) + CO2(g)
NH4Cl(s)  NH3(g) + HCl(g)
Decomposition of sodium azide, NaN3
2NaN3(s)  2Na (s) + 3N2(g)

8. Molecular Masses
How do we use the atomic masses of elements to determine the molecular masses of molecules?
We simply add the masses of the constituent elements in the molecule.

9. The Mole and Avogadro’s Number
NOTE: 1 amu = x g; this is a very small mass!!
It is not a very convenient unit of measurement on the laboratory scale.
In chemistry, we use a special unit (the mole) when dealing with atoms, molecules, and ions.

10. The Definition of a Mole
1 mole (SI definition)
the amount of substance that contains as many elementary particles (atoms, molecules, ions) as there are atoms in 12 grams (exactly) of carbon-12.
This number is called Avogadro's number (it is an experimentally determined quantity).

11. The Molar Mass 1 mole = 6.022 x 1023 particles.
1 mole of carbon-12 has a mass of g exactly!
The molar mass of any element in g/mole is the same numerically as its atomic mass in amu’s!

12. Numbers in periodic table
The molar mass of a compound is the mass (in grams) of 1 mole of a compound.
Numbers in periodic table
e.g., Cl = g/mole
= amu/atom

13. Converting Between Moles and Molecules

14. Masses of Anions and Cations
Electron mass is small
mass of Na+  mass of Na
mass of Cl-  mass of Cl
mass of O2-  mass of O atom!!

15. Percent Composition
Percent composition is the percent by mass of each element in a compound.
Q: If we have the percent composition of the compound, can we calculate its empirical formula?

16. A Schematic for Doing Empirical Formula Calculations

17. Chemical Analysis
Use a variety of techniques to obtain the identity and the % by mass of each element in a sample.
Combustion analysis  a sample containing C, H, or C,H, and O is combusted (burned).
(C,H,O)  CO2 + H2O
all the C in original compound gets converted to CO2
all the H in original compound is converted to H2O

18. Combustion Analysis

19. Quantitative Information from Chemical Reactions
Example
A + B  C + D
Normally, we start out with certain quantities of reactants.
How much product can we expect?
How much reactant would we need to obtain a specified amount of product?

20. The Mole Method
Stoichiometric coefficients in the balanced chemical equation represent the number of moles of reactants and products.
certain mass of reactants  ? mass of products
initial volume of reactants  ? Volume of products

21. Example N2 (g) + 3 H2 (g)  2 NH3 (g)
1 mole N2  3 moles H2  2 moles NH3
the  symbol means stoichiometrically equivalent

22. A Schematic of the Mole Method

23. Limiting Reagent
Chemical equations  give the molecular or molar ratios of reactants needed and the products obtained.
N2 (g) + 3 H2 (g)  2 NH3 (g)
We assumed 100% complete conversion of of the N2 (g) and H2 (g) to NH3 (g).
Normally, the reactants will react in the proper ratio until one of them is consumed completely.

24. The Definition of the Limiting Reagent
The amount of product formed is limited by the reactant that is completely consumed  the limiting reagent.
The amount of product obtained assuming complete consumption of the limiting reagent  theoretical yield.
The reactant that is left over  the excess reagent.

25. Limiting Reagents

26. The Limiting Reagent Schematic
Use amount of Limiting Reagent
Mass of A and B
Use Molar Masses of A, B
Moles of A, B
Determine Limiting Reagent
Obtain Theor. Yield
Use Coefficients from balanced Equation