CHAPTER 8 NOTE GUIDE
CHEMICAL REACTIONS
Chapter objectives:
1. Write a chemical reaction given the names of products and reactants.
2. Understand the various symbols that are used in a chemical reaction.
3. Balance a chemical reaction using coefficients.
4. Predict the products of the following types of reactions: combustion, synthesis, decomposition, single displacement, double displacement, and redox.
5. Write a net ionic equation for reactions in solution.
CHEMICAL EQUATIONS and WRITING REACTIONS FROM NAMES
In order to write a chemical reaction, it is necessary to review writing chemical formulas from the compound name.
1. Ionic
2. Covalent
3. Acids
4. Organic compounds
NAMING ACIDS and WRITING ACID FORMULAS
· Recall that acids have a hydrogen in the front of their formula.
· Acids are always aqueous.
· The two types of acids are binary and oxyacids
BINARY ACIDS
· Only contain 2 elements.
· Examples: HCl, HF, H2S
· Name using a prefix, a root, and a suffix
o Prefix – hydro
o Root – based on the 2nd element’s name
o Suffix - ic
· To write the formula of a binary acid:
o Hydro – indicates that it is binary.
o Write H first with a +1 charge, followed by the 2nd element with its most common charge, then balance charges.
o Ex. Hydrobromic, hydroiodic
OXYACIDS
· Include oxygen in their formula.
· Named based on a polyatomic ion in the formula.
· Examples: H2CO3, HNO3, HNO2, H3PO4
Polyatomic ion ending / Acid name ending-ate / -ic
-ite / -ous
· To write the formula for an oxyacid:
o Look at the acid name ending, if –ic, find the polyatomic ion ending in –ate, write it with its charge and then balance the charge by adding H+ to the front of the formula.
o Ex. Sulfuric acid, sulfurous acid,
ORGANIC NAMES and FORMULA WRITING
· Organic compounds include carbon.
o The naming rules for organics are quite complex, so we will focus on 3 groups of organic compounds: Alkanes, Alkenes, Alkynes
o All three are named based on the number of carbons in their formula.
Number of carbons / prefix1 / Meth-
2 / Eth-
3 / Prop-
4 / But-
5 / Pent-
6 / Hex-
7 / Hept-
8 / Oct-
9 / Non-
10 / Dec-
ALKANES
· Straight chains of carbon and hydrogen with single bonds throughout.
· Have the general formula: (CnH2n+2)
· Write the formula for an alkane and draw its Lewis structure.
· Names end in -ane
ALKENES
· Contain one double bond between 2 carbons, so the number of hydrogens is reduced.
· Have the general formula: (CnH2n)
· Name ends in – ene
· Ex. C5H10
ALKYNES
· Contain a triple bond between 2 carbons, so the number of hydrogens is reduced again.
· Have the general formula: (CnH2n-2)
· Name ends in –yne
· Ex. C4H6
DAY 2: BALANCING CHEMICAL REACTIONS
The law of conservation of mass states that matter cannot be created or destroyed, so the number of atoms of each element must remain constant during the reaction.
Equations are balanced by trial and error.
Three steps will ensure success:
1. balance polyatomic ions that appear on both sides of reaction as single units
2. save elements that are listed as atoms on one side of the reaction for last
3. balance H and O after all other elements have been balanced
Write and balance the following reactions:
1. hydrochloric acid (aq) and sodium hydroxide (aq) react to form sodium chloride (aq) and water
2. Iron (s) reacts with sulfur (s) to form iron (III) sulfide (s)
3. hydrogen, carbon monoxide and oxygen react in water to form carbonic acid
4. silicon dioxide and hydrofluoric acid react to form silicon tetrafluoride and water
5. Magnesium phosphate reacts with aluminum sulfate to form magnesium sulfate and aluminum phosphate.
DAY 3: TYPES of REACTIONS
1. combustion
2. synthesis
3. decomposition
4. single displacement
5. double displacement
6. oxidation reduction
COMBUSTION
· During complete combustion a hydrocarbon burns in oxygen to produce water and carbon dioxide.
· Incomplete combustion produces CO.
· General form:
CxHy(maybe Oz) +O2 à CO2 (g) + H2O(g)
Ex. Butane burns in oxygen
Ex. Pentyne burns in oxygen
Ex. sucrose burns in oxygen
HINTS:
· Burned in oxygen
· Hydrocarbon
Check point:
Write the formulas for the following:
1. Ethane
2. Hexyne
3. Octene
4. Hydrobromic acid
5. Hypobromous acid
6. Perbromic acid
7. Bromic acid
8. Bromous acid
TYPES of REACTIONS (CONT.)
2. SYNTHESIS
· One compound is formed in a reaction
General forms:
A. Element + element = compound
Ex. Sodium (s) reacts with liquid bromine
Ex. Carbon and oxygen form carbon dioxide
B. Compound + compound = compound
a. Metal oxide + water produces a basic solution
Ex. Sodium oxide reacts with water
b. Non-metal oxide produces an acidic solution
Ex. Sulfur trioxide + water
c. Metal oxide + non-metal oxide produces a salt
Ex. Calcium oxide reacts with sulfur trioxide
HINTS:
· Look for 2 elements combining
· Compounds with the same anion (usually oxygen)
· Form one product
Self check: list the reactions from your packet that are synthesis here.
3. DECOMPOSITION
· one compound is broken down into 2 or more pure substances
General forms:
A. binary compound decomposition à yields elements
a. Heating
Ex. Calcium oxide is heated
b. Electrolysis
Ex. Water has a current passed through it
B. metal hydroxide à metal oxide and water
ex. Magnesium hydroxide decomposes
C. oxyacid à non-metal oxide and water
ex. sulfuric acid decomposes
D. metal carbonate à metal oxide and carbon dioxide
Ex. magnesium carbonate decomposes
E. metal chlorate à metal chloride + oxygen
HINTS:
· Decomposes
· Heated strongly
· 1 reactant
Self check: List the reactions that are decomposition here.
4. SINGLE DISPLACEMENT
· A more reactive metal or non-metal replaces a less reactive metal or non-metal.
· Consult the activity series to determine if the reaction will work.
General forms:
A. metal replacement by a more reactive metal
ex. Na (s) + MgO à
B. hydrogen replacement by a more reactive metal
ex. Mg (s) + HCl (aq) à
ex. Na (s) + H2O (l) à
C. non-metal replacement by a more reactive non-metal
ex. Cl2 (g) + KI (aq) à
HINTS:
· a single element + a compound react
· check the activity series to see if the reaction takes place
Self check: List the reactions that are single displacement here.
5. DOUBLE DISPLACEMENT
· The cations and anions of two compounds exchange places.
· A precipitate, gas, or H2O must form for the reaction to proceed.
General forms:
A. formation of an INSOLUBLE solid (PPT.) – see solubility table.
ex. CaCl2 (aq) + Na2CO3 (aq)à
ex. Fe(NO3)3 (aq) + Ba(OH)2 (aq) à
B. formation of a gas
ex. CaCO3 (aq) + HCl (aq) à
Evidence of a reaction?
C. Formation of water
Ex. HBr (aq) + NaOH (aq) à
Evidence of a reaction?
HINTS:
· 2 compounds with different cations and anions are mixed
Self check: List the reactions that are double displacement here.
NET IONIC REACTIONS
· when reactions occur in aqueous solution it is necessary to show the ions that are present.
· Break all (aq) substances into their cation and anions with charges.
o Ex. H2SO4(aq) à
o Ex. Ba(OH)2 (aq) à
· Leave (s), (g) and (l) together.
Steps for writing a net ionic equation:
· Write and balance the reaction.
Ex. Sodium chloride reacts with silver nitrate
· Write the total ionic reaction
· Cancel spectator ions - ions that appear on both sides of the equation that do not participate in the reaction
· Write the net ionic reaction with charges and phases on all ions.
6. OXIDATION REDUCTION REACTIONS
· Causes a change in the oxidation state (charge) of an element
Ex. Find the oxidation state for Sulfur in each of the elements in the compounds or ions below:
SO3
SO3-2
SO4-2
S
S4O6-2
S2O3-2
Na2SO4
KMnO4
· Reduction – gain of electrons, causes a decrease in the ion/elemental charge
· Oxidation – loss of electrons, causes an increase in the ion/ elemental charge
· Electrons are transferred from one element/ ion to another. The number of electrons lost = number of electrons gained.
· Balance using the ½ reaction method.
o Reduction ½ reaction
Electrons are gained as a reactant. Add enough to balance the charges.
o Oxidation ½ reaction
Electrons are lost as a product. Remove enough to balance charges.
· Acidic solution
o Uses H+ and H2O to balance the reaction
Example: NO2 à NO3 -1 + NO
1. Break into ½ reactions
Reduction reaction
Oxidation reaction
2. Use H2O to balance extra oxygen
3. Use H+ to balance excess hydrogen
4. Add electrons to balance charge
5. Balance electrons by multiplying ½ reactions
6. Add reactions together, cancel whenever possible.
BALANCING IN BASIC SOLUTION
1. Follow steps 1-3 as if it were in basic solution
2. Add OH-1 to the H+ make H2O, add an = number of OH-1 to the other side of the reaction.
3. Pick up with step 4 above.
Example:
Zn à Zn(OH)4 2- + H2