Chemical Equations and Reactions. I. Chemical Equations.

Chemical Equationsand Reactions

I. Chemical Equations

II. Chemical EquationsA. The equation must represent facts.

B. The equation must contain the correct formulas for the reactants (on the left of the arrow) and the products (on the right of the arrow).

C. The law of conservation of mass and energy must be satisfied. Therefore the same number of atoms of each element must appear on each side of a correct chemical equation.

Symbol Explanation of symbol

+ separates 2 or more reactants or products “yield”, separates reactants from products. ↔ indicates a reversible reaction(s) solid state. Placed after the formula of a

substance↓ Alternative to (s) but used ONLY for a solid

PRODUCT, not reactants(l) indicates a liquid reactant or product(aq) indicates an aqueous solution (where some

solute has been dissolved in water)(g) indicates a gaseous reactant or product↑ alternative to (g), but used ONLY for a gaseous

PRODUCTΔ indicates that heat is supplied to the reaction

A formula written above or below the sign indicates that it is used as a catalyst (something

that speeds up the reaction)

C. Diatomic and Polyatomic Molecules:

Element Formula State

Hydrogen H2 gas

Nitrogen N2 gas

Oxygen O2 gas

Fluorine F2 gas

Chlorine Cl2 gas

Bromine Br2 liquid

Iodine I2 solid

Sulfur S8 solid

PhosphorusP4 solid

III. Writing and Balancing Equations

ExampleWrite a balanced equation for the following reaction:

Na + Cl2 NaCl first write an atom inventory for the total

number of atoms of each element on each side of the equation.

Na + Cl2 NaCl ReactantsProducts# Na # Na# Cl # Cl

Atom Inventory or Counting Atoms

• you must be able to count atoms in order to balance an equation. There are two ways to designate numbers in a formula:

• subscripts – small numbers within a formula of a compound. Tells the number of atoms in that compound

• MgCl2 – 1 atom of Mg and 2 atoms of chlorine• Sn3N2 – 3 atoms of tin and 2 atoms of nitrogen• Coefficient – the large number in front of the formula

of a compound. Tells the number of molecules (in a molecular compound) or formula units (in an ionic compound) or atoms of an element.

• Remember that atoms cannot be created or destroyed; we must balance an equation using coefficients. Never change a subscript to balance an equation!!

Algebraic Method

Write the skeleton equationWrite a,b,c… on each substanceList the atoms presentDetermine equalities for each atomAssign a value 6 for AUse A to determine other valuesReduce, if possiblePlug #s in as coefficients

Practice

__H2O __H2 + __O2

__Pb(NO3)2 + __Na __NaNO3 + __Pb

__C4H10 + __O2 __CO2 + __H2O

Examples

Zn + H2O Zn(OH)2 + H2

C2H6 + O2 CO2 + H2O

Na2SO4 + Ba(NO3)2 BaSO4 + NaNO3

IV. Types of Chemical Reactions

A. Combination or Synthesis

• where 2 or more simple substances (elements or compounds) combine to form ONE complex substance

• 8Fe + S8 8FeS

• 2Sr + O2 2SrO

A. Combination or Synthesis

Li + P4

N2 + Al

Cl2 + Ca

Na + N2

Special Combination or Synthesis Reactions (Pre-AP

Only)the metals that has a variable charge:

If one of these metals reacts with fluorine, oxygen, or nitrogen (F, 0, N), these nonmetals will

pull the metal to its HIGHEST charge or oxidation number. Otherwise, when these metals react in a

combination reaction, use their LOWEST charge or oxidation number when forming a new compound

Practice:Fe + O2

Pb + N2

Sn + S8

Cu + P4

Fe + Br2

Cu + F2

B. Decomposition

a complex substance (compound) decomposes into 2 or more simple substances. Heat or electricity is usually required.

Ex:

2NaCl 2 Na + Cl2

8MgS 8Mg + 2S8

Special decomposition reactions to know (Pre-AP only):

• 2KClO3 2KCl + 3O2 -

• all metal chlorates decompose into metal chloride + O2

• CaCO3 CaO + CO2

• metal carbonates decompose into a metal oxide + CO2

• 2KOH K2O + H2O

• metal hydroxides decompose into a metal oxide + H2O

C. Combustion Reactions

• where oxygen reacts with another substance, usually a hydrocarbon, resulting in the release of energy, usually heat or light.

• CH4 + 202 CO2 + 2H20

• Hydrocarbons always produce carbon dioxide and water

Name Molecular Formula

Methane CH4

Ethane C2H6

Propane C3H8

Butane C4H10

Pentane C5H12

Hexane C6H14

Heptane C7H16

Octane C8H18

Nonane C9H20

Decane C10H22

Common Hydrocarbons

Examples

C3H8 + O2

C2H2 + O2

Ca + O2

D. Single-Replacement

• occurs when one element displaces another element in a compound. You must check the “Activity Series of Metals” to see if the “lone” element is active or “strong” enough to displace the element in the compound

Activity Series of MetalsLiK Ba CaNaMg Al Zn FeNiSnPb HCuHgAg Au

Practice:

Li + KCl

Sn + ZnCl2

Sn + HCl

Ni + HOH

the halogens

F2

Cl2Br2

I2

Decreasing strength

E. Double-Replacement reactions

• occur when the cations (positive ions) “switch” places. You do NOT need the “activity series of metals” list in these reactions. When you switch places, be sure to correctly write the formula of the new compound!!!!!

• Ex:

• 2 NaCl + Mg0 MgCl2 + Na20

Practice

CuS04 + Al(OH)3

Ca3(P04)2 + ZnCr04

Rules for Predicting Double Replacement Reactions:

• 1. Predict the products of the double-replacement reaction and indicate the solubility of both of the products

• Use the “Solubility Rules” handout (at end of notes) to determine the solubility.

• If the compound is soluble that means that it will remain as ions in the solution, if it is insoluble then the compound precipitated out of the reaction (it became the precipitate or solid).

• 2. If at least one INSOLUBLE product is formed (which means a precipitate will form) the reaction will occur!

• 3. If only SOUBLE products are formed then the reaction will NOT occur (because no precipitate is formed)!

• 4. If water is produced the reaction will occur!• 5. If the reaction occurs and one of the compounds formed is

soluble then that compound is written as ions and not as a compound.

SOLUBILITY RULES

1. All salts whose cation is in Group 1 or is NH4+ are soluble – no matter what the anion is.

2. All nitrates and nitrites are soluble.3. All acetates are soluble.4. All chlorates and perchlorates are soluble – no matter what the

cation is.5. All chlorides are soluble except silver, lead (II), mercury6. All bromides are soluble except silver, lead (II), mercury7. All iodides are soluble except silver, lead (II), mercury8. All flourides are INSOLUBLE.9. All sulfates are soluble except silver, lead (II), mercury, calcium,

strontium and barium10. All sulfides are INSOLUBLE except Group 1 and 2.11. All phosphates, phosphites, carbonates, chromates, and

dichromates are INSOLUBLE unless the cation is in Group 1 or is NH4+.

12. All hydroxides are INSOLUBLE except Group 1, calcium,. Strontium, barium and ammonium

Net Ionic Equations

(Pre-AP Only)

• F. Net Ionic Equations (Pre-AP Only) – shows only the compounds and ions that undergo a chemical change in a double replacement reaction

• Example: Na2S + Cd(NO3)2 Na+ + NO3 + CdS(s)

• Step 1: Convert the chemical equation to an overall ionic equation. All reactants are shown as ions. For the products, all soluble ionic compounds are shown as dissociated ions and the precipitates are shown as solids.

• Na+ + S 2 + Cd 2+ + NO3 Na+ + NO3 + CdS(s)

• Step 2: All spectator ions (ions that do not take part in a chemical reaction and are found as ions both before and after the reaction) are removed from the equation.

• S2 + Cd2+ CdS(s)

Examples

BaCO3 + CuSO4 BaSO4(s) + CuCO3 (s)

K3PO4 + NaOH no reaction (no ppt)

Na2S + Cd(NO3)2 Na+ + NO3 + CdS(s)

Types of Reactions SummaryCombination (synthesis)

A + B AB

Decomposition

AB A + B

Combustion

CxHy + O2 CO2 + H2O

Single Replacement

A + BC AC + B

Double Replacement

AB + CD AD + CB

Practice Predicting Products l. AlCl3

2. C2H4 + 02

3. Zn + AgNO3

4. H20

Practice Predicting Products

5. Al + P4

6. NaI + MgS

7. Cl2 + NaBr

8. C6H1206 + O2

Practice Predicting Products

1. AlCl3 + Na2CO3

2. Ni + MgSO4

3. Cl2 + K

4. C5H12 + 02

Practice Predicting Products (Pre-AP)

1. sodium metal is placed into water

2. methane gas is burned in the presence of oxygen

3. potassium bromide solution is mixed with chlorine gas

4. a solution of aluminum dichromate is mixed with a solution of lithium oxalate

V. Oxidation Reduction Reactions (Pre-AP only)

What is REDOX?

• Oxidation-Reduction (Redox) – involves a transfer of electrons

• One specie is losing electrons

OIL – Oxidation is losing

Mg0 Mg+2 + 2 e-

• One specie is gaining electrons

RIG – Reduction is gaining

Mg+2 + 2 e- Mg0

• The Species that is oxidized is the reducing agent

• The Species that is reduced is the oxidizing agent.

• Mg0 + O20 Mg+2O-2

Ox

Red

Red ag

Ox ag

REDOX reactions MUST:

• 1 Have a species that is oxidized and one reduced – YOU cannot have one without the other

• The number of electrons gained and lost MUST be the SAME

• The number of atoms of each element must be the same on both sides of the equation

Balancing REDOX Reactions

• 1) Assign oxidation numbers to each atom in the equation

• 2) Determine the substances oxidized, reduced, oxidizing agent, reducing agent,

3) Write balanced half-reactions for the oxidation and reduction reactions.

• Mg0 + O20 Mg+2O-2

Mg0 Mg+2 + 2 e-

O20 + 4 e- 2O-2

• 4) Multiply each equation so that the number of electrons lost equals the number of electrons gained.

• Mg0 + O20 Mg+2O-2

• 2[Mg0 Mg+2 + 2 e-]= 2Mg0 2Mg+2 + 4 e-

O20 + 4 e- 2O-2

• 5) Add the two half-reactions. Place the coefficients into the original equation.

• Mg0 + O20 Mg+2O-2

2Mg0 2Mg+2 + 4 e-

O20 + 4 e- 2O-2

2Mg0 + O20 2Mg+2O-2

• 6) Adjust other ions if necessary. Check all atoms for conservation. Check hydrogen’s and oxygen’s last.

2Mg + O2 2Mg O