1. Chemical Reactions
SC2. Obtain, evaluate and communicate information about how the Law of Conservation of Matter is used to determine chemical composition in compounds and chemical reactions.
a. Use mathematics and computational thinking to balance chemical reactions (i.e., synthesis, decomposition, single replacement, double replacement, and combustion) and construct an explanation for the outcome of a simple chemical reaction based on the outermost electron states of atoms, trends in the periodic table, and knowledge of the patterns of chemical properties.
b. Plan and carry out an investigation to determine that a new chemical has been formed by identifying indicators of a chemical reaction (e.g., precipitate formation, gas evolution, color change, water production, and changes in energy to the system).

2. Chemical Reaction
A chemical reaction occurs when one or more substances change into one or more new substances.
Indicators of a chemical reaction include:
Energy change (gain or loss of heat)
Gas production (bubbles or odor)
Precipitate production (liquids make a solid)
Color change
Water production

3. Chemical Equations REACTANTS  PRODUCTS
A chemical reaction is represented in writing by a chemical equation.
REACTANTS  PRODUCTS
Multiple reactants and products are separated from each other by + signs
Reactants are separated from products by a yield sign ()
Other notations include Δ, ↔, s, l, g, aq

4. Writing Chemical Equations
Iron metal combines with oxygen gas in the presence of water to produce solid iron(III) oxide.
Aqueous sodium hydroxide combines with aqueous copper(II) nitrate to form aqueous sodium nitrate and solid copper(II) hydroxide.

5. Reading Chemical Equations
ZnS(s) + O2(g)  ZnO(s) + SO2(g)
KClO3 (s)  KCl(s) + O2(g)

6. It’s the Law!

7. The Law Says… Law of Conservation of Matter:
In a chemical reaction, matter can be neither created nor destroyed.
In a chemical reaction, the amount of reactants equal the amount of products.

8. It’s the Law! How it applies to us now: Law of Conservation of Atoms –
The number of atoms of each type of element must be the same on each side of the equation.
This makes the equation BALANCED!

9. Matter Is Conserved Mg + Cl2 MgCl2 Total atoms = Total atoms
Total Mass = Total Mass
1(24.3) (35.5) ( )
95.3 g = g

10. hydrogen + oxygen water
Balancing Equations
hydrogen + oxygen water
H2 + O H2O
FORMULAS MUST BE CORRECT!
FORMULAS CANNOT BE CHANGED when you start balancing

11. Balancing Equations H2 + O2 H2O
If the formulas cannot be changed, how can the atoms be made equal?
Adjust the number of molecules by changing the coefficients.

12. Balancing Equations
H2 + O H2O

13. Learning Check Fe3O4 + 4 H2 3 Fe + 4 H2O 1. Number of H atoms in 4 H2O
A) 2 B) 4 C) 8
2. Number of O atoms in 4 H2O
3. Number of Fe atoms in Fe3O4
A) 1 B) 3 C) 4

14. Balancing Equations
N2 + H NH3

15. Matter Is Conserved N2 + 3 H2 2 NH3 Total atoms = Total atoms
Total Mass = Total Mass
2(14.0) (1.0) [ (3x1.0)]
34.0 g = g

16. Learning Check
Balance each equation. The coefficients for each equation are read from left to right.
4. Mg N Mg3N2
A) 1, 3, B) 3, 1, C) 3, 1, 1
5. Al Cl AlCl3
A) 3, 3, 2 B) 1, 3, 1 C) 2, 3, 2

17. Balancing Equations Balancing Hints for Complex Eqns:
Balance tallies with “1s” first.
Balance the metals next.
Balance the ion groups third.
Save oxygen and hydrogen until last.

18. Balancing Equations
Ca3(PO4)2 + H2SO4 CaSO4 + H3PO4

19. Balancing Equations
NH O2 NO + H2O

20. Balancing Equations
Cu + H2SO CuSO H2O + SO2

21. Learning Check 6. Fe2O3 + C Fe + CO2
A) 2, 3, 2, 3 B) 2, 3, 4, 3 C) 1, 1, 2, 3
7. Al FeO Fe Al2O3
A) 2, 3, 3, 1 B) 2, 1, 1, 1 C) 3, 3, 3, 1
8. Al H2SO Al2(SO4) H2
A) 3, 2, 1, B) 2, 3, 1, 3 C) 2, 3, 2, 3

22. Types of Reactions There are six general types of reaction:
Synthesis (also called combination)
Decomposition
Single replacement
Double replacement
Combustion
Oxidation-reduction (REDOX)

23. Synthesis
These reactions have two or more reactants that combine into a single product.
A + B  AB
Mg + O2  MgO
Fe + S  Fe2S3

24. Decomposition
These reactions have one reactant that breaks down into two or more products.
AB  A + B
HgO  Hg + O2
H2O  H2 + O2

25. Single Replacement A + BX  B + AX
These reactions involve an elemental reactant that takes the place of another element in a compound.
A + BX  B + AX
K + H2O  KOH + H2
Zn + Cu(NO3)2  Cu + Zn(NO3)2

26. Single Replacement Y + BX  X + BY
These reactions involve an elemental reactant that takes the place of another element in a compound.
Y + BX  X + BY
NaI + Br2  NaBr + I2
F2 + KCl  KF + Cl2

27. Single Replacement Not all combinations result in chemical reactions
Successful replacement depends on where the exchanging elements are placed on the reactivity series

28. Reactivity Series
More reactive elements will displace less reactive elements from compounds
Hydrogen
(below = unreactive with acids)

29. Double Replacement AY + BX  AX + BY
These reactions involve aqueous reactant compounds that swap ions to produce other compounds.
AY + BX  AX + BY
Na2S + Cd(NO3)2  CdS + NaNO3
FeS + HCl  H2S + FeCl2

30. Double Replacement
Neutralization and precipitation reactions are specific examples of double replacement reactions that can occur:
Neutralization – when acids and bases combine to produce a salt and water
NaOH + HCl  NaCl and H2O

31. Double Replacement SOLUBILITY RULES
Precipitation – when solutions combine and a precipitate forms
How can you predict if one will form?
SOLUBILITY RULES

32. Solubility Rules NAG SAG PMS CastroBar Soluble: Nitrates Acetates
Group 1 Exceptions:
Sulfates ** * Lead, Mercury, Silver
Ammonium * Calcium, Strontium, Barium
Group 17 *
NAG SAG PMS CastroBar

33. Combustion CH4 + O2  CO2 + H2O
These reactions involve a fuel (typically a hydrocarbon) burning in oxygen gas to produce water and carbon dioxide.
CH4 + O2  CO2 + H2O

34. Oxidation-Reduction
If the oxidation states of any elements in a reaction change, the reaction is an oxidation-reduction reaction.
Oxidation state refers to the charge on an atom if it were bonded ionically
Oxidation states are identified for each element, per atom, in a compound with numbers like charges, but with sign first

35. Oxidation-Reduction Rules for Assigning Oxidation State:
1. Elements in a free, uncombined state, or as diatomic molecules, have an oxidation state of zero
2. Monatomic ion charge = oxidation state
3. Sum of all oxidation states in a neutral compound is zero; sum in a polyatomic ion is charge on ion

36. Oxidation-Reduction Rules for Assigning Oxidation State:
4. Oxidation state of elements in Group 1 = +1, Group 2 = +2
5. Oxidation state of oxygen is -2, unless as peroxide (-1) or with fluorine (+1)
6. Oxidation state of hydrogen is +1, unless as hydride (-1)

37. Oxidation-Reduction Rules for Assigning Oxidation State:
7. Oxidation state of fluorine is always -1; other halogens are usually -1, unless with oxygen or fluorine
MnO2 + HCl → MnCl2 + Cl2 + H2O
Al + Fe2O3 → Fe + Al2O3

38. Oxidation-Reduction
Oxidation is the process by which an atom’s oxidation state increases
For an atom’s “charge” to increase, what has to happen?
So, loss of electrons = oxidation (LEO)
Or, oxidation = loss of electrons (OIL)
The atom that undergoes oxidation is called the reducing agent

39. Oxidation-Reduction
Reduction is the process by which an atom’s oxidation state decreases
For an atom’s “charge” to decrease, what has to happen?
So, gain of electrons = reduction (GER)
Or, reduction = gain of electrons (RIG)
The atom that undergoes reduction is called the oxidizing agent

40. Oxidation-Reduction
MnO2 + HCl → MnCl2 + Cl2 + H2O Al + Fe2O3 → Fe + Al2O3 H2O2 → O2 + H2O

41. Net Ionic Equations
Net ionic equations represent what actually “happens” in a chemical reaction
Only the species directly involved in the reaction are included
Spectator ions are not included
Any substance that is a solid, liquid, or gas remains “as is” in the net ionic equation

42. Net Ionic Equations To write a net ionic equation:
Write the balanced molecular equation.
Write the balanced complete ionic equation.
Cross out the spectator ions that are present.
Write the "leftovers" as the net ionic equation.

43. Net Ionic Equations
(NH4)2CO3 (aq) + Al(NO3)3 (aq)  NH4NO3 (aq) + Al2(CO3)3 (s)
NaOH (aq) + H2SO4 (aq)  Na2SO4 (aq) + H2O (l)
Mg (s) + HCl (aq)  MgCl2 (aq) + H2 (g)
Zn (s) + CuSO4 (aq)  ZnSO4 (aq) + Cu (s)
Na2CO3 (aq) + HNO3 (aq)  NaNO3 (aq) + H2O (l) + CO2 (g)