Chapter 11 Chemical Reactions

Physical Changes Physical Change – A change during which some properties of a material change, but the composition of the material does not change. Examples: - Cutting, breaking, grinding, tearing - phase changes (melting, boiling, freezing) - dissolving

Chemical Changes Chemical change – A change in which a substance undergoes a change in identity (composition). Examples: burning, rusting, decomposing Things that are NOT chemical changes: Any phase change (evaporation, freezing…) Dissolving

Signs of Chemical Changes Several signs that a chemical change has occurred: 1) Heat, light or gas given off 2) Change in color or odor 3) A precipitate (solid product) is formed (It will look cloudy.) 4) Sound 5) Bubbles (except boiling)

When signs of a chemical change can be deceiving. Boiling - phase change from liquid to gas. It is NOT a reaction. But you will see bubbles. Diluting – adding water to lower the concentration. It will alter the color, but it is NOT a reaction. Diluting is NOT a chemical change. Two clear liquids making a yellow precipitate IS a chemical reaction

Chemical Reactions Chemical Reaction – A chemical change in which new substances are formed. Chemical Equation – A symbolic representation of a chemical reaction.

Reaction Terms Reactants: Starting substances in a chemical reaction. A + B  C + D Products: Substances produced in a chemical reaction.

Reaction Terms Coefficient – Big number in front of a substance. Tells the number of moles, molecules, or units of a substance. 3H2O Subscript – Small number written below. Tells the numbers of moles or atoms of a particular element.

Reaction Symbols  “yields” (produces, results in)  reversible reaction, “equilibrium” (s) solid (l) liquid (g) gas (aq) aqueous – dissolved in water reaction occurred by heating

Collision Theory 1) Molecules must collide in order to react. 2) When they collide, they have to have a) The right orientation b) Enough energy

Law of Conservation of Mass Law of Conservation of Mass – Matter cannot be created or destroyed. Every chemical equation must satisfy this law. Therefore equations must be balanced so that: total mass of reactants = total mass of products.

Balancing Equations Only COEFFICIENTS can be changed ! (NEVER the subscripts!) Example: Al(s) + CuSO4(aq)  Cu(s) + Al2(SO4)3(aq)

Balancing Examples H2 + O2  H2O Zn(OH)2 + H3PO4  Zn3(PO4)2 + H2O Ag2S + Al  Al2S3 + Ag _____ Na + _____ I2 _____ NaI

More Balancing Examples ___ Ca3(PO4)2 + ___ H2SO4 ___ CaSO4 + ___ H3PO4 ___ KClO3  ___ KCl + ___ O2 SO2 + O2  SO3 C3H6 + O2  CO2 + H2O

Types of Reactions Synthesis – 2 or more reactants, 1 product Example: A + B  AB

Types of Reactions Decomposition: 1 reactant, 2 or more products Example: AB  A + B

Types of Reactions Single Replacement - a single element “switches places” with an element in a compound. D + BC  C + BD (nonmetal) (nonmetal) A + BC  B + AC (metal) (metal)

Types of Reactions Double Replacement – Two compounds “switch partners”. Don’t forget that in a compound the charges must be (+)(-). Example: AB + CD  + - + - + - + - AD + CB

Types of Reactions Combustion – Special type of reaction CxHy + O2  CO2 + H2O Watch for CO2 and H2O as products!

Predicting the Products – Single Replacement To predict if a single replacement will react, use the ACTIVITY SERIES. The free element MUST BE more reactive than the element in the compound.

Single Rep. If the free element is MORE ACTIVE than the element in the compound, the reaction WILL OCCUR. Example: Cl2 + 2HBr  If the free element is LESS ACTIVE than the element in the compound, the reaction WILL NOT OCCUR. Example: Br2 + 2HCl  (DNR = Does Not React) 2HCl + Br2 DNR

Note **Note – Don’t EVER bring a subscript across the arrow UNLESS it’s part of a polyatomic ion!!**

Activity Series Examples Ag + ZnCl2  Zn + 2AgNO3  Zn + CuSO4  NaCl + Li  NaCl + I2  LiOH + Na  2NaI + Br2  DNR 2Ag + Zn(NO3)2 Cu + ZnSO4 LiCl + Na DNR DNR 2NaBr + I2

Predicting the Products – Double Replacement In order for a double replacement reaction to occur, one product MUST BE: A gas A liquid (water) A precipitate (a solid) If all products are aqueous (soluble), it WILL NOT occur. Use the SOLUBILITY RULES: Soluble = Aqueous Insoluble = Solid precipitate

Solubility Rules

Double Replacement Examples Ba(OH)2 + H3PO4  K2SO4 + CaCl2  FeBr2 + AlCl3 KOH + H2SO4  (NH4)2CO3 + CaCl2 

Net Ionic Equations Rules: Aqueous ionic compounds can be split into ions. (Don’t forget charges!) Strong acids can be separated into ions. Substances that are solids, liquids, or gases cannot be separated. Spectator ions are removed from the ionic equation, leaving the net ionic equation.

If Aqueous… For all aqueous compounds: Step 1 – Split it up Step 2 – Write charges for each thing Step 3 – Write how many you have as a coefficient

Net Ionic Eqn. Example Na2SO4 (aq) + BaCl2 (aq)  2NaCl (aq) + BaSO4 (s) Ionic equation: Spectator ions: Net Ionic Equation:

Lab Tests – Burning splint A burning splint can be used to test for: Hydrogen (squeaky “pop” sound) Oxygen (blow out the splint and it will reignite) Because fire needs O2 to burn Carbon dioxide (flame will go out) Because CO2 smothers it

Lab test - Limewater and CO2 Clear, colorless limewater will turn a cloudy white if CO2 is added.