1. Chemistry 068, Chapter 7

2. Chemical Equations and Stoichiometry
Chemical equations are representations of chemical reactions using formulas rather than words.
Correctly written chemical equations must follow a number of rules.
Stoichiometry is the quantitative study of the relationships between the amounts of products and reactants in a chemical reaction.
Stoichiometry also involves calculations using chemical equations.

3. Evidence of Chemical Reactions
Many, but not all, chemical reactions result in visible, or at least detectable, changes which indicate that a reaction is taking place.
Some examples are:
Color change.
Formation of a solid.
Formation of a gas.
Heat absorption or emission.
Light emission.

4. Writing and Balancing Chemical Equations
We will look at the reaction of methane and oxygen to form carbon dioxide and water as we go through the rules.
When writing chemical reactions it is necessary to follow a number of rules.
Reactants are always written on the left, products are always written on the right.
CH4 O2 CO2 H2O
Plus signs separate formulas on each side.
CH4 + O2 CO2 + H2O

5. Writing and Balancing Chemical Equations (Cont’d)
An arrow () or an equilibrium sign (⇌) separate the products and reactants.
CH4 + O2  CO2 + H2O
Reactions must be consistent with experimental results (correct phase, molecular formulas rather than empirical formulas).
Optionally, phase can be represented by writing (s) for solid, (l) for liquid, (g) for gas, and (aq) for aqueous (dissolved in water) after each formula.
CH4 (g) + O2 (g)  CO2 (g) + H2O (l)

6. Writing and Balancing Chemical Equations (Cont’d)
Reactions must be balanced (follow conservation of mass) for number and type of atoms but not phase.
There must be the same number of each type of atom on each side of the equation.
Coefficients in front of a formula tell you how many there are of that atom/molecule. Ones are excluded.
Coefficients must be the smallest whole number values possible.
CH4 + 2O2  CO2 + 2H2O
Reactions must be charge balanced (same overall ionic charge). We will look at this much later.

7. Writing and Balancing Chemical Equations from Pictures
Chemists sometimes like to represent chemical information with images of the product and reactant molecules.
It is possible to write chemical equations from these representations.

8. Writing and Balancing Chemical Equations from Pictures (Cont’d)
For example, consider a reaction between A (red) and B (blue) to make products.
3A2 and 3B molecules react to form 3A2B molecules. So, we write:
3A2 + 3B  3A2B
Since we always want the smallest possible coefficients, we finally write:
A2 + B  A2B

9. Balancing Chemical Equations from Pictures Problem
Write a balanced chemical equation for the reaction of A (red), B (blue), and C (green).

10. Strategies for Balancing Chemical Equations
Remember that number and type of atoms is the main thing that must be balanced.
Phase and number of molecules does not need to be balanced.
You can only change coefficients (the numbers in front).
You cannot do anything to change the molecules in the reaction.
You can’t change subscripts.

11. Strategies for Balancing Chemical Equations (Cont’d)
Try to work with one element at a time – it’s usually much easier.
It is often simpler to treat polyatomic ions as single units rather than individual atoms.
This only really works when the polyatomic ions remain as single units in both reactants and products.

12. Strategies for Balancing Chemical Equations (Cont’d)
Check your coefficients.
They must all be whole numbers (no fractions).
They should be the smallest possible whole numbers.
Once you have finished, double check everything.
Equal number of each element, no fractional coefficients, smallest whole number coefficients.

13. Balancing Chemical Equation Problems
Balance the following reactions:
PbO + NH3  Pb + N2 + H2O
NH3 + HNO3  NH4NO3
NBr3 + NaOH  N2 + NaBr +HBrO
C5H10 + O2  CO2 + H2O
Al + Sn(NO3)2  Al(NO3)3 + Sn

14. Balancing Chemical Equation Problems (Cont’d)
Balance the following reactions:
FeO + HCl  Fe + Cl2 + H2O
NaOH + H2CO3  Na2CO3 + H2O
Ba(C2H3O2)2 + (NH4)3PO4  Ba3(PO4)2 + NH4C2H3O2
C2H6 + O2  CO2 + H2O
NO + CH4  HCN + H2O + H2

15. Ions and Solutions
The basic theory of ionic solutions was established by Svante Arrhenius in 1884.
He proposed that certain substances split into freely moving ions when in solution and that those ions allow for electrical conduction.

16. Electrolytes and Nonelectrolytes
Nonelectrolytes – substances that dissolve in water to produce a poorly or non conducting solution.
Electrolytes – substances that dissolve in water to produce a conducting solution. They are usually ionic compounds.
Strong electrolytes completely dissociate into ions.
Weak electrolytes only partially dissociate into ions.
Same idea as strong acids/bases or weak acids/bases.

17. Solubility Rules
The solubility rules are a set of qualitative rules used to determine weather or not ions are soluble in water.
Soluble ions are normally strong electrolytes.
They arise from the “like dissolves like” rule.
There are still quantitative limits to these rules.

18. Solubility Rules (Cont’d)
Soluble
All group IA and ammonium – Li+, Na+, K+, NH4+.
All acetates and nitrates – C2H3O2-, NO3-.
Most chlorides, bromides, and iodides – Cl-, Br-, I-.
Exceptions: Silver [AgCl, Ag Br, Ag I], Mercury [Hg2Cl2, HgBr2, Hg2Br2, HgI2, Hg2I2], Lead [PbCl2, PbBr2, PbI2].
Most sulfates – SO42-.
Exceptions: Calcium [CaSO4], Strontium [SrSO4], Barium [BaSO4], Silver [AgSO4], Lead [PbSO4].
Insoluble
Most carbonates – CO32-.
Exceptions: Group IA and ammonium carbonates.
Most phosphates – PO43-.
Exceptions: Group IA and ammonium phosphates.
Most sulfides – S2-.
Exceptions: Group IA, IIA, and ammonium sulfides.
Most hydroxides – OH-.
Exceptions: Group IA hydroxides, Calcium [Ca(OH)2], Strontium [Sr(OH) 2], Barium [Ba(OH) 2].

19. Precipitation Reactions
Uses the solubility rules to predict weather or not any precipitate(s) form.
The insoluble compound(s) will form solid precipitate(s).
Simplest form is a double replacement reaction where ions from the two mixed solutions switch places.

20. Precipitation Reaction Problems
Predict the identity of any precipitate(s) that forms during the following reactions.
NaBr + AgNO3  ?
NaOH + CaCl2  ?
CaSO4 + BaS  ?

21. Molecular, Complete Ionic, and Net Ionic Equations
Three different kinds of equations can be written to describe ionic solution phase reactions.
Molecular – no ions shown.
This is the way we have been writing equations up until now.
Complete Ionic – all aqueous molecules written as ions, including spectator ions.
Spectator ions – ions in the chemical equation that don’t take part in the reaction – they are both products and reactants.
Net Ionic – spectator ions removed from the equation.

22. Molecular, Complete Ionic, and Net Ionic Equation Examples
Mg(s) + 2HBr(aq)  MgBr2(aq) + H2 (g)
Molecular
Mg(s) + 2H+(aq) + 2Br-(aq)Mg2+(aq) + 2Br-(aq) + H2(g)
Complete Ionic
Mg(s) + 2H+(aq)  Mg2+(aq) + H2(g)
Net Ionic
Note that all of these equations can be used to represent the same reaction.

23. Writing and Balancing Complete Ionic and Net Ionic Equations
Complete ionic and net ionic equations are written and balanced in much the same way as molecular equations.
The key difference is in how charge must be balanced as well as number/type of atoms.
The total charge of all ions added up must be the same on both the reactants and products side of the equation.
This usually, but not always, means a total charge of zero on both sides.

24. Writing and Balancing Complete Ionic and Net Ionic Equations (Cont’d)
It is also important to establish rules for what compounds do, and do not, split up into ions in water.
The following kinds of compounds split up:
All soluble salts.
All strong acids.
All strong bases.
All other compounds are written as molecules rather than ions.

25. Writing and Balancing Complete Ionic and Net Ionic Equations (Cont’d)
To write a complete ionic or net ion equation you should follow the following steps:
Write and balance the molecular equations as normal.
HCl + NaOH  NaCl + H2O
Split compounds into ions as appropriate, and double check charge.
At this point you are done if all you are writing is a complete ionic equation.
H+ + Cl- + Na+ + OH-  Na+ + Cl- + H2O
If writing a net ionic equation, cancel out any spectator ions.
H+ + OH-  H2O

26. Writing and Balancing Complete Ionic and Net Ionic Equation Problems
Write complete and net ionic equations for each of the following reactions:
MgCl2 + AgNO3  Mg(NO3)2 + AgCl
H2SO4 + KOH  H2O + K2SO4
NaNO3 + KCl  KNO3 + NaCl

27. Acid-Base Reactions
Arrhenius Acid – A substance that produces hydrogen ions (H+) when dissolved in water.
Ex: HCl(aq)H+(aq)+Cl-(aq)
Arrhenius Base – A substance that produces hydroxide ions (OH-) when dissolved in water.
Ex: NaOH(aq)Na+(aq)+OH-(aq)
Bronsted-Lowry Acid – A proton donor in a proton transfer reaction.
Bronsted-Lowry Base – The proton acceptor in a proton transfer reaction.
Ex:
NH3(aq)+H2O(aq)NH4+(aq)+OH-(aq)

28. Acid-Base Reactions (Cont’d)
Strong Acid/Base – Completely ionizes in water.
Weak Acid/Base – Only partially ionizes in water.
Strong/Weak only refers to ionization, not how harmful or how reactive it is.
Concentration is a much better indicator of how dangerous an acid/base is.
Strong Acids and Bases
Acids – HClO4, H2SO4, HI, HBr, HCl, HNO3.
Bases – LiOH, NaOH, KOH, Ca(OH)2, Sr(OH)2, Ba(OH)2.

29. Acid-Base Reactions (Cont’d)
Acid-Base Indicator – A dye used to distinguish between acidic and basic solutions by color.
Neutralization – A reaction of an acid and a base that produces a salt and possibly water.
HA(aq) + BOH(aq)  H2O(l) + AB(aq)
Ex: HCl(aq) + NaOH(aq)  H2O(l) + NaCl(aq)
Polyprotic Acid – An acid that contains more than one acidic hydrogen.

30. Acid-Base Problems
Write molecular, ionic, and net ionic equations for the following neutralization of Barium hydroxide by hydrobromic acid.

31. Acid-Base Problems (Cont’d)
Write molecular, ionic, and net ionic reaction equations for the neutralization of each acidic hydrogen in sulfuric acid by sodium hydroxide; and for the overall (1 step) reaction.

32. Acid-Base Gas Formation Reactions
Reactions with gas formation – certain salts give off gas when they react with acids.
Most notably carbonates (CO32-), sulfites (SO32-), sulfides (S2-), and ammoniums (NH4+).
They produce CO2, SO2, H2S, and NH3 respectively.

33. Acid-Base Gas Formation Reactions (Cont’d)
The reactions of sulfides are a simple one step metathesis reaction resulting in a salt and a gaseous product.
Example:
Na2S(aq)+2Hl(aq)  2NaI(aq)+H2S(g)

34. Acid-Base Gas Formation Reactions (Cont’d)
The reactions for carbonates and sulfites are very similar.
The reaction takes place in two steps. The first step similar to a metathesis reaction, which produces a salt and an unstable intermediate product.
The intermediate then decomposes into water and a gas.
Examples:
K2CO3(aq)+2HBr(aq)  2KBr(aq)+H2CO3(aq)
2KBr(aq)+H2CO3(aq)2KBr(aq)+H2O(l)+CO2(g)
Li2SO3(aq)+2HCl(aq)  2LiCl(aq)+H2SO3(aq)
2LiCl(aq)+H2SO3(aq)2LiCl(aq)+H2O(l)+SO2(g)

35. Acid-Base Gas Formation Reactions (Cont’d)
Like the reactions for carbonates and sulfites, ammonium reactions are two step..
The reaction takes place in two steps. The first step similar to a metathesis reaction, which produces a salt and an unstable intermediate product.
The intermediate then decomposes into water and a gas.
Example:
NH4Cl(aq)+KOH(aq)  KCl(aq)+NH4OH(aq)
NH4OH(aq)+KCl(aq) H2O(l)+KCl(aq)+NH3(g)

36. Acid-Base Gas Formation Reactions Problems
Write a reaction for the production of carbon dioxide when calcium carbonate reacts with hydrochloric acid.
If 12.0g of calcium carbonate are reacted, what mass of carbon dioxide is produced?

37. Oxidation – Reduction Reactions
Oxidation-reduction reactions (also called redox) involve electron transfer from one species to another, or in which atoms change oxidation number (state).
For example, consider these molecular and net ionic equations:
Fe(s) + CuSO4(aq)  FeSO4(aq) + Cu(s)
Fe(s) + Cu2+(aq)  Fe2+(aq) + Cu(s)
Iron and copper switch charged states. Iron goes from neutral to 2+ while copper goes from 2+ to neutral. Sulfate does not change state.

38. Combustion Reactions
Combustion reactions are reactions of a substance with oxygen which include a release of heat and flame.
Typically, oxygen is from the air.
Example reactions:
2C2H2 + 5O2  4CO2 + 2H2O
CS2 + 3O2  CO2 + 2SO2

39. Predicting Products of a Combustion
Carbon and hydrogen atoms commonly form the following compounds as products during a combustion:
C forms CO2
H forms H2O
O forms either CO2 or H2O
It is more difficult to predict products formed from other elements.
S forms SO, SO2, SO3
N forms NO, NO2, NO3
It is only possible to determine which is formed experimentally.

40. Predicting Products of a Combustion (Cont’d)
For the moment we will only deal with compounds of carbon, hydrogen, and oxygen.
Each reactant carbon forms one CO2.
Every two reactant hydrogens for one H2O.
Ignore oxygens – we assume they are all used up to make water or carbon dioxide.
You can then determine the number of oxygens needed as reactants by chemical balancing.

41. Predicting Products of a Combustion Problems
Predict the products of the combustion of ethanol C2H6O and write the balanced equation for the reaction.

42. Predicting Products of a Combustion Problems (Cont’d)
Predict the products of the combustion of benzene C6H6 and write the balanced equation for the reaction.

43. Predicting Products of a Combustion Problems (Cont’d)
Predict the products of the combustion of methane CH4 and write the balanced equation for the reaction.

44. Types of Chemical Reactions
There are many types of chemical reactions and a single course could not possible look at all of them.
The following one method to classify chemical reactions.
We will look at 4 classes of chemical reactions in this chapter.
They are:
Synthesis reactions.
Decomposition reactions.
Single replacement reactions.
Double replacement reactions.

45. Synthesis Reactions
Synthesis, sometimes called combination, reactions are reactions which form a single product from two or more reactants.
Reverse of decomposition reactions.
Example reactions:
2NO2 + H2O2  2HNO3
Ni + S  NiS

46. Decomposition Reactions
Decomposition reactions are reactions where a single reactant breaks down into two or more products.
Reverse of synthesis reactions.
Example reactions:
2H2O  2H2 + O2
Al2(CO3)3 Al2O3 + 3CO2

47. Single Replacement Reactions
Single replacement reactions are reactions where one element in a compound is replaced by another (single) element.
Single elements but can be multiple atoms of the same element.
Example reactions:
Fe + CuSO4  Cu + FeSO4
Mg + Ni(NO3)2  Ni + Mg(NO3)2

48. Double Replacement Reactions
Double replacement reactions are reactions where two compounds switch parts with one another to create two new compounds.
Again, can be multiples of the same part.
Example reactions:
NaF + HCl  NaCl + HF
AgNO3 + HCl  AgCl + HNO3

49. Types of Reactions Problems
Determine the type of chemical reaction for each of the following:
SO3 + H2O  H2SO4
Na2CO3 + Ca(OH)2  2NaOH + CaCO3
Fe + 2CuNO3  2Cu + Fe(NO3)2
2SO2 + O2  2SO3
K2CO3  K2O + CO2