Understanding Chemical Reactions

• Chemical Reactions

Writing Reactions

• Physical changes are changes in a substance that do not alter chemical identity.
• Examples of physical changes include boiling, melting, vaporizing, and grinding.
• Chemical change is a change of a substance into another substance; when a chemical reaction occurs.
• Examples of chemical changes include burning, oxidizing, rotting, and fermenting.
• A chemical reaction is a process by which substances collide with enough energy that new bonds are formed between atoms, thus creating a new substance.
• Chemical reactions do not change matter.
• Chemical reactions change energy.
• Collision theory states that reacting particles must collide with enough force for a chemical reaction to occur.
• A reaction depends on the frequency of collisions.
• When a chemical change occurs, a chemical reaction occurs.
• Energy change in the system indicates a chemical reaction.
• The release of light or heat is a sign of chemical reaction.
• Sudden color change is a sign of chemical reaction.
• Odor release is a sign of chemical reaction.
• Gas release is a sign of chemical reaction.
• Sudden appearance of a solid (precipitate) is a sign of chemical reaction.
• The solid that forms is known as a precipitate.
• The Law of Conservation of Mass says that matter cannot be created or destroyed, it can only change forms.
• This is true even in a chemical change/reaction.
• The mass of the reactants = the mass of the products.
• If there are 25 g of reactants, there must be 25 g of product.
• If there are two oxygen atoms in the reactants, there must be two oxygen atoms in the products.
• Balancing equations is critical.
• Equations represent chemical reactions.
• Equations can be written in word or formula form.
• Hydrogen gas and Oxygen gas react to form liquid water is an example in word form.
• 2H2(g) + O2(g) -> 2H2O(l) and 2Na(s) + Cl2(g) -> 2NaCl(s) are examples in formula form.
• Reactants = starting substances, ingredients.
• Products = ending substances; what is made.
• "+" separates one or more reactants or products.
• "→" yields, produces.
• Coefficient tells us the number of each substance needed.
• Subscripts cannot be changed.
• If only 1 is needed, the 1 is not written (Ex. Cl2).
• How to read this the ratio of sodium to chlorine is a 2:1 this reaction.
• (s) = solid.
• (l) = liquid.
• (g) = gas.
• (aq) = aqueous; dissolved in water.
• Equations must be balanced to keep from violating the Law of Conservation of Mass.
• Coefficients are used to balance equations.
• Coefficients can be changed, but subscripts cannot!
• Steps to Balance Equations:
  • Write the equation.
  • Put the symbols of each element in a column under the reactants and products by keeping polyatomic ions together, if possible.
  • Count the number of each atom on both sides of the equation.
  • Change the coefficients only on either side of the equation to balance the number of each atom.
  • Reduce coefficients, as needed.
  • Coefficients must be written in the lowest whole number ratio.

Classifying Reactions

• Chemical reactions can generally be classified into five categories: synthesis, decomposition, combustion, single replacement, and double replacement.
• Reactions are classified so that it is easier to predict the products for a reaction.
• Synthesis:
  • One or more reactants combine to form a new compound.
  • A + B → AB.
  • For example, 2Na(s) + Cl2(g) -> 2NaCl(s).
  • When an element combines with oxygen, it produces an oxide.
  • An oxide is a binary compound with at least one atom of oxygen.
  • For example, C(s) + O2(g) -> CO2(g).
  • Certain metal oxides combine with nonmetal oxides to make salts.
  • For example, CaO(s) + SO2(g) -> CaSO3(s).
• Decomposition:
  • One reactant produces two or more simpler compounds.
  • AB -> A + B.
  • For example, Ca(OH)2(s) -> CaO(s) + H2O(g).
  • Most only take place when energy (in the form of heat or electricity) is added.
  • Electrolysis is the decomposition of a substance by an electric current.
  • An example of electrolysis is 2H2O(l) -> 2H2 + O2(g).
• The decomposition of:
  • Metal carbonates transforms into metal oxide + CO2(g).
    • For example, CaCO3(s) -> CaO(s) + CO2(g).
  • Metal hydroxides transforms into metal oxide + H2O(g).
    • Group I metals are an exception.
    • For example, Ca(OH)2(s) -> CaO(s) + H2O(g).
  • Metal chlorates transforms into metal chloride + O2(g).
    • For example, 2KClO3(s) -> 2KCl(s) + 3O2(g).
• Combustion occurs when a substance reacts with oxygen, often producing a lot of energy, heat, and light.
  • reactant + O2 -> product.
    • For example, 2H2(g) + O2(g) -> 2H2O(g).
  • Many combustion reactions occur with a hydrocarbon and result in carbon dioxide and water.
• a hydrocarbon is a compound made of carbon and hydrogen.
  • Fuel + O2 -> CO2 + H2O; CH4 (g) + 2O2(g) -> CO2(g) + 2H2O (g).
• Single Replacement:
  • When one element replaces an element like it in a compound.
  • A + BC -> B + AC.
  • For example, 2Al(s) + 3Pb(NO3)2(aq) -> 3Pb(s) + 2Al(NO3)3(aq) and Cl2(g) + 2KBr(aq) -> Br2(l) + 2KCl (aq).
  • Many single replacement reactions take place in an aqueous solution.
  • The most reactive metals can react with water to produce metal hydroxides and hydrogen.
    • For example, 2Na(s) + 2H2O(l) -> 2NaOH(aq) + H2(g).
• Double Replacement:
  • When the ions of two compounds exchange places in an aqueous solution to form two new compounds.
  • AB + CD -> AD + CB.
  • One of the resulting products is usually a precipitate.
  • A precipitate is an insoluble solid that forms from the ions of two aqueous compounds.
    • For example, NaCl (aq) + AgNO3 (aq) -> AgCl (s) + NaNO3 (aq).
  • Insoluble gas is released as bubbles.
    • For example, FeS (s) + 2HCl (aq) -> H2S (g) + FeCl2 (aq).
  • Molecular compound like water.
    • For example, HCl (aq) + NaOH (aq) -> NaCl (aq) + H2O (l).
• Neutralization reaction (aka, acid-base reaction) is when an acid and a base react to form water and a salt.
  • Acid: a substance that makes H+ ions when dissolved in water.
  • Base: a substance that makes OH- ions when dissolved in water.
  • Salt: an ionic compound made from the cation of a base and an anion from an acid.
    • For example, HCl (aq) + NaOH (aq) -> NaCl (aq) + H2O (l).
• Predicting Products:
  • Some elements react more readily than others.
  • Activity series: a list of elements in the order in which they will easily undergo certain chemical reactions.
  • Metals that have a greater activity activity means they lose electrons more easily.
  • Nonmetals that have a greater activity activity means they gain electrons more easily.
  • Elements can replace any element below them but not above them.

Chemical Equilibrium

• Reactions can be reversible:
  • 2SO2(g) + O2(g) ⇆ 2SO3(g).
  • 2SO3(g) ⇆ 2SO2(g) + O2(g).
• When the forward and reverse processes are happening at the same rate, we say that the system is in equilibrium.
  • Indicated by a double arrow.
  • 2SO2(g) + O2 (g) ⇆ 2SO3 (g).
• Chemical Equilibrium: A dynamic process when there is no net change occurring in the amount of reactants and products in the system, thus no visible change.
  • When products and reactants are forming at the same rate.
  • When reversible processes are occurring at the same rates.
  • Consider a reaction at equilibrium: A + B ⇆ AB.
  • A + B -> AB is happening at the same rate as AB -> A + B.
• Le Chatelier's Principle: Le Chatelier's Principle states that when conditions change for a system at equilibrium, the system responds by reducing the effect of the change.
• Consider: A + B ⇆AB:
  • Forward reaction = A + B -> AB:
    • If this reaction/direction is favored, over time, [A+B] decreases, [AB] increases.
  • Reverse reaction = AB -> A + B:
    • If this reaction/direction is favored, over time [A+B] increases, [AB] decreases.
• Changes (or stresses) to a system: change in concentration, temperature and pressure.
• Blood pH and exercise example:
  • Stress in exercise: produces CO2, shifting the equilibrium to reactants.
  • High CO2 builds up in the bloodstream, increasing the acidity of the blood.
  • The body responds with exercise by removing CO2 from the bloodstream.
  • Stress from breathing faster removes CO2, shifting the equilibrium to products.
• Factors That Affect Equilibrium:
  • Concentration change: adding or removing a reactant or product.
    • Adding a reactant favors the forward reaction.
    • Adding a product favors the reverse reaction.
    • Removing a reactant favors the reverse reaction.
    • Removing a product favors the forward reaction.
  • Temperature change:
    • Exothermic reactions release heat.
      • For example: N2(g) + 3H2(g) ⇆ 2NH3(g) + 91 kJ.
      • Adding heat would favor the reverse reaction.
    • Endothermic reactions absorb heat.
      • For example, N2O4(g) + heat ⇆ 2NO2(g).
      • Adding heat would favor the forward reaction.
  • Pressure change (in gases):
    • If pressure is increased, the reaction will favor whichever direction produces fewer gas molecules.
    • If pressure is decreased, the reaction will favor the direction that produces more gas molecules.
      • Reactions with no gases are unaffected by pressure.
      • For example, CaCO3 (s) ⇆ CaO (s) + CO2(g)
      • Reactions with an equal number of gas moles on each side are unaffected by pressure.
        • For example, H2(g) + Cl2 (g) ⇆ 2HCl (g).
      • If there are disproportionate number of molecules on each side of the equation, increase in pressure favors forward, decrease in pressure favors reverse.
        • For example, N2 (g) + 3H2(g) 2NH3 (g)