1. Equations & Reactions

2. 8.1 Describing Chemical Reactions
A. Chemical Changes and Reactions
1. New substances are produced.
2. Chemical reaction – chemical bonds between atoms or ions break, and new bonds form between atoms or ions.
B. Evidence of a Chemical Reaction
1. color change
2. formation of a precipitate
3. temperature change
4. formation of a gas

3. C. Mechanics of a Chemical Reaction
1. Starting Materials – reactants
2. Ending Materials - products
3. reactants → products
Arrow yields or produces
4. Many reactions occur to complete a set of valence electrons.
5. Symbols above the yield sign represent conditions necessary for a reaction to proceed.
Ex)
= delta = heat
= electrolysis

4. Liquid = l Gas = g Solid = s Crystal = cr
6. Some reactions occur spontaneously.
7. Symbols represent the state of the reactants and products.
Liquid = l Gas = g Solid = s Crystal = cr
Aqueous = aq (solids in water solution)
DEMO
Ex) 2Al(s) + 3CuCl2(aq) → 2AlCl3(aq) + 3Cu(s)
silver blue gray red

5. 8. Complete chemical equations include the subscript to indicate the physical state of each substance.
9. Diatomic molecules – certain elements exist in nature as diatomic molecules (X2)
List them: N2 O2 F2 Cl2 Br2 I2 H2

6. Natural States of the Elements
Diatomic Molecules
Nitrogen gas contains
N2 molecules.
Oxygen gas contains
O2 molecules.

7. Highlight your Periodic Table

8. Natural States of the Diatomic Elements

9. 8.2 Balancing Equations A. Equations in Chemistry
1. Chemical equation: an expression that uses
symbols and formulas to describe a chemical
reaction.
2. + means “reacts with.”
3. → means produces (called the yield sign.)

10. B. Balancing Chemical Equations
1. Conservation of mass leads to balancing equations – the number of atoms of each element must be the same before & after the reaction.
2. The Law of Conservation of Mass also states that the total mass before and after the reaction must be the same. You cannot lose or gain mass.
3. Therefore the MASS OF THE PRODUCTS = MASS OF REACTANTS
4. Subscript – indicates number of atoms of an element present in a compound.
5. Coefficient – indicates the number of atoms or molecules involved in the reaction.

11. 6. Steps to Balance Equations:
A. Write equation with symbols.
B. Count # of atoms on each side of the
reaction.
C. Balance atoms using coefficients.
D. General Rule: Balance all elements first.
Then, balance C, H, and O.
E. NEVER CHANGE SUBSCRIPTS!!!!

12. 3 H N2 → 2 NH3
3 Na2SO Ca3(PO4)2 → 2 CaSO Na3PO4
2 NaNO3 → 2 NaNO O2
2 C8H O2 → 16 CO H2O

13. 8.3 Classifying Chemical Reactions
A. Synthesis Reactions (direct combination)
1. Two or more elements or compounds combine
to form a more complex product.
A + B → AB
2. Ex. Fe + S → FeS
CaO + H2O → Ca(OH)2
sodium reacts with chlorine
2 Na Cl2 → 2 NaCl

14. Synthesis Reaction

16. B. Decomposition Reactions (analysis)
A single reactant breaks down into simpler compounds or elements.
AB → A + B
2. The opposite of a synthesis reaction.
3. Carbonates (compounds ending in CO3) break down into metal oxides and carbon dioxide.
4. Ex HgO → Hg + O2
CaCO3 → CaO + CO2

17. Decomposition Reaction

19. CuCO3(s)  CO2(g) + CuO(s)
flame goes out
copper (II) carbonate
metal oxide

20. What is one of the products?
2HgO → 2Hg + O2
mercury (II) oxide
mercury (II) oxide

21. C. Single Replacement Reactions
1. Atoms of an uncombined element replace atoms of another element in a compound.
A + BX → AX + B
2. A more active element will replace a less active element. (See activity series.)

23. 3CuCl2 + 2Al  2AlCl3 + 3Cu

25. 2 Al + 3 CuSO4 → 3 Cu + Al2(SO4)3 Fe + MgCl2 → No Reaction
3. An Activity Series is a way of ranking elements (usually metals) in order from greatest to least reactivity. It can be used to predict whether a reaction will occur or not.
2 Al CuSO4 → 3 Cu + Al2(SO4)3
Fe MgCl2 → No Reaction
PbSO Au → No Reaction
AgCl2 + Cu → CuCl Ag

26. D. Double-Replacement Reactions
1. Atoms or ions from two different compounds replace each other.
AX + BY → AY + BX
2. These types of reactions will
(A) form precipitates (↓)
(B) form gases (↑)
(C) are acid-base neutralizations

27. Double Replacement Examples:
Pb(NO3) KI → PbI2 ↓ KNO3
B. CaCO HCl → CaCl H2CO3
C. NaOH + HCl → NaCl + HOH
3. In letter B above, carbonic acid, H2CO3, is unstable and will immediately decompose into carbon dioxide and water.
CaCO HCl → CaCl CO2↑ + H2O

31. E. Combustion Reactions
1. One substance reacts with oxygen, O2 to produce oxide compounds.
2. Occurs during burning or oxidation (rusting.)
3. The reactions that only add oxygen are classified as synthesis reactions.
Ex) S + O2 → SO2

32. Hydrogen Burning Video
H O2  2 H2O
Synthesis Reaction

33. 4. Combustion reactions are exothermic, releasing a large amount of energy as light, heat, or sound.
5. A true Combustion reaction occurs when a hydrocarbon (compound containing H & C ) reacts to form carbon dioxide and water that are always the products.
CxHx O2 → CO H2O
6. Ex. __CH4 + __O2 → __CO2 + __H2O kJ
C6H12 O O2 → CO H2O heat 2 2 6 6 6

34. Combustion Reaction

35. Combustion Reaction

36. 5 Types of Chemical Reactions Video

37. Chapter 8 Equations Test Objectives
1. Identify the parts of a chemical equation (reactants VS products).
2. Identify notation abbreviations used in equations (state symbols).
3. Properly assign subscripts in formulas (nomenclature!!!).
4. Use coefficients to correctly balance equations.
5. Calculate the number of atoms, molecules and formula units
6. Classify Reactions as Synthesis, Decomposition, Single/Double Replacement or Combustion.
7. Predict the products and balance all types of reactions (using state symbols).
8. Use the Activity Series to predict the products and then balance Single Replacement reactions.