1. PRINCIPLES OF CHEMICAL REACTIVITY: CHEMICAL REACTIONS
Chapter 4

2. Chemical equations
A chemical equation is the representation of a chemical reaction using the symbols of the elements and the formulas of the compounds
The reacting species (starting materials or reactants) are listed on the left, and the products of the reaction are listed on the right

3. Four Things to Know About Equations
One: Reactants are on the left and products are on the right of the arrow.
Two: Coefficients are the numbers in front of each formula. If no number is shown, a one is understood.
Three: The coefficients tell us how many molecules (moles) of each reactant used and how many molecules (moles) of each product made.
Four: The phases of the reactants and products are indicated in parenthesis after the formula.
gas = (g)
liquid = (l)
solid (s)
aqueous (in a solution) =(aq)

4. Balanced Chemical Equation
C6H12O6(s) + 6O2(g)  6CO2(g) + 6H2O(l)

5. Balancing Chemical Equations
We balance chemical equations because of the Law of the Conservation of Mass
We use coefficients in front of the element or compound formula, which are whole numbers that adjust the amounts of the species
Note that we cannot balance an equation by changing the subscripts of the compound

6. Balancing coefficients
The subscripts of the compound are fixed: they cannot be changed in an equation
The coefficients used should be the smallest whole numbers possible
The coefficient multiplies every number in the formula (e.g. 2 MgCl2 means that there are 2 Mg atoms and 4 Cl atoms)
We will balance equations by inspection

7. Balance the following Si2H3 + O2  SiO2 + H2O
Al(OH) H2SO4  Al2(SO4) H2O

8. Types of chemical reactions
Combustion reactions - reaction of an element or a compound with elemental oxygen (O2)
These reactions are often accompanied by the release of large amounts of heat (and flame)
Consider: 2C2H6 (g) + 7O2 (g)  4CO2(g) + 6H2O(g)

9. Combination reactions
Often called synthesis reactions
Involves two or more substances combining to form a single substance
Consider:
2Mg(s) + O2 (g)  2MgO(s)

10. Decomposition reactions
The reverse of a combination reaction
One substance breaking apart into two or more substances
Consider:
H2CO3 (aq)  H2O(l) + CO2 (g)

11. Reactions in Aqueous Solution
Single Replacement
Double Replacement
Oxidation-Reduction
All require an understanding of aqueous solutions!

12. Formation of ions in solution
When we consider a substance dissolving in another substance, we define the solute as the substance in lesser amount, and the solvent as the substance in greater amount
Solubility is the relative degree to which a substance dissolves in a solvent
Soluble substances dissolve to a considerable degree
Insoluble substances do not dissolve at all or only to a small degree

13. Dissolution of ionic compounds
When ionic compounds dissolve in water, they dissociate into separate anions and cations: NaCl(s)  Na+(aq) + Cl-(aq)
Acids are so named because when they dissolve, the ionize to yield H+ and the associated anion HCl(aq)  H+(aq) + Cl-(aq)

14. Solubility Rules Rule Exception Soluble Group 1 elements, NH4+
NO3 , ClO3, ClO4 
Chlorides, bromides, iodides
Ag+, Pb2+, Hg22+
Acetates
Ag+, Hg22+
Sulfates
Sr2+, Ba2+, Pb2+, Ca2+
Insoluble Compounds
Carbonates, phosphates, oxalates, chromates, sulfides
Group 1, NH4+
Hydroxides, oxides
Group 1, Ba2+

15. Strong Acids
Strong acids completely dissociate into ions upon dissolution in water
With strong acids the major components in solution are the hydrogen ion and the anion.
Examples are HCl, H2SO4

16. Weak Acids
Weak acids dissolve, but do not completely dissociate into H+ and anions.
With weak acids the major component in solution is the weak acid.
Examples are CH3CO2H, HF

17. Single replacement reactions
An element replaces another in a compound
Zn(s) + CuCl2 (aq)  ZnCl2 (aq) + Cu(s)
This particular form of the equation is termed the molecular equation, where all reactants and products are shown as neutral compounds

18. Ionic equations
The total ionic equation writes all of the anions and cations separately Zn(s) + Cu2+(aq) + 2Cl-  Zn2+(aq) + 2Cl- + Cu(s)
Ions (such as the Cl- ions) that are in the same state on both sides of the equation are called Spectator ions.

19. Net ionic equation
The net ionic equation leaves out the spectator ions
The complete ionic equation Zn(s) + Cu2+(aq) + 2Cl-  Zn2+(aq) + 2Cl- + Cu(s)
The net ionic equation Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s)

20. Double replacement reactions
Where the anions and cations in an ionic equation swap places.
AgNO3 (aq) + NaCl(aq)  AgCl(s) + NaNO3 (aq)

21. What causes double replacements?
There has to be some process to cause the reaction to take place
Precipitation of a solid
Evolution of a gas
Formation of a Molecular Substance
Neutralization reaction between an acid and a base
If one of these processes does not occur, no reaction takes place

22. Double replacement reactions
Precipitation
AgNO3 (aq) + NaCl(aq)  AgCl(s) NaNO3 (aq)
Gas evolution
2HCl(aq) + Na2CO3 (aq)  2NaCl(aq) CO2 (g)

23. Double replacement reactions
Neutralization
an acid is a proton donor
a base is a hydroxide donor
proton plus hydroxide yields water
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)

24. Net Ionic Equations
In a net ionic equation the spectator ions are omitted.
Spectator ions are those that do not change in the reaction.
Net ionic equations are only used for reactions occurring in solution between ionic compounds – single and double replacement reactions – as well as oxidation-reduction

25. Write the Net Ionic Equation
AgNO3 (aq) + NaCl(aq)  AgCl(s) + NaNO3 (aq)
2HCl(aq) + Na2CO3 (aq)  2NaCl(aq) + CO2 (g)
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)

26. Oxidation – Reduction Oxidation
loss of at least one electron during a reaction..
Ni(s) + H+(aq)  Ni2+(aq) + H2(g)
Reduction
gain of at least one electron during a reaction.
In above example, H+ gains an electron to become reduced.

27. Oxidation-Reduction Cont.
Every reaction must have an oxidation and reduction.
Metals react with acids to form salts and hydrogen gas.
Cu(s) + 2HNO3(aq)  Cu(NO3)2(aq) + H2(g)
Metals also oxidized with salts
Fe(s) + Cu(NO3)2(aq)  Fe(NO3)2(aq) + Cu(s)

28. Oxidation Number Oxidation number (state):
The charge that an atom in a substance or monatomic ion has or would have if all the bonding electrons were assigned to the most electronegative element.
Ca in CaO +2
Ca2+(aq)
Cl(aq) 1
Cr in CrO3 +6
Fe in Fe2O3 +3
Cr in K2Cr2O7

29. Assigning Oxidation Numbers
Elemental form: 0
Monatomic ions: charge of ion
Oxygen: 2, except in H2O2 and other peroxides.
Hydrogen: +1, except with metal hydrides when it is 1.
Halogens: 1 (except when bound to oxygen or a halide above it)
Alkali and alkaline earth metal ions have a charge of +1 and +2, respectively.
Compounds and ions: sum of the charges on the atoms in a compound add up to 0 and to the ion charge in the ion.

30. Displacement Reactions: Activity series of the elementsLi
Reacts vigorously with acids to give H2
Reacts with H2O to give H2 K Ba Ca Na Mg
Reacts with acids to give H2
Reacts slowly with H2O to give H2; more vigorous with steam Al Zn Cr Fe Cd
Etc.
A relative reactivity scale allows us to predict if reaction will occur when two substances are mixed together.
E.g. Copper ions in solution are reduced to the metal when an iron nail is placed in the solution.
Cu2+(aq) + Fe(s)  Fe2+(aq) + Cu(s)  Iron displaces copper.
Fe2+(aq) + Cu(s) NR  copper will not displace iron.
Iron more reactive than copper.
E.g. Predict which reaction will occur when:
Li is mixed with K+ and
Li+ is mixed with K.
E.g. In which of the following mixtures will reaction occur:
Li+ + Mg
Al + Mn2+
Fe + Cd2+
Cr + Zn2+

31. Balancing: Oxidation-Number Method
Determine oxidation # for each atom- both sides of equation.
Determine change in oxidation state for each atom.
Left side: make loss of electrons = gain.
Balance other side.
Insert coefficients for atoms that don't change oxidation state.

32. Example
FeS(s) + CaC2(s) + CaO(s)  Fe(s) + CO(g) + CaS(s)

33. Balancing: Half-Reaction Method
Write unbalanced half reactions for the oxidation and the reduction
Balance the number of elements except O and H for each.
Balance O's with H2O to the deficient side.
Balance H's with H+ to the hydrogen deficient side
Acidic: add H+
Basic: add H2O to the deficient side and OH to the other side.
Balance charge by adding e to the side that needs it.
Multiply each half-reaction by integers to make electrons cancel.
Add the two half-reactions and simplify.

34. Half-Reaction Acidic
Zn(s) + VO2+(aq)  Zn2+(aq) + V3+(aq).

35. Half Reaction Basic
Ag(s) + HS(aq) + CrO42(aq)  Ag2S(s) + Cr(OH)3(s).