High School Chemistry/Stiochiometry

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Solube (aq)
  1. Soluble: Nitrates (NO3), Acetates (CH3COO), Ammonium (NH4), and Group 1 salts.
  2. Soluble: Chlorides (Cl), Bromides (Br), and Iodides (I) except with Silver (Ag), Lead (Pb), and Mercury (I) - Hg.
  3. Soluble: Flourides (Fl) except with Group 2 elements, Lead (II), and Iron (III).
  4. All sulfates SO42- except with Calcium, Strontium, Barium, Mercury, Lead (II), and Silver.
Insoluble (s)
  1. All carbonates and phosphates except with Group I and Ammonium.
  2. All hydroxides (OH) except with Group I, Strontium, Barium, and Ammonium.
  3. All sulfides except Group I and II and Ammonium.
  4. All oxides (O2) except Group I.

Reactions[edit | edit source]

  1. Synthesis/Combination
  2. Decomposition
  3. Single Replacement
  4. Double Replacement, also known as double-displacement reactions
  5. Combustion
  6. Neutralization
  7. Oxidation-Reduction

Synthesis/Combination Reactions[edit | edit source]

  1. Formation of a binary compound: A + B --> AB
  2. Metal oxide-water reactions go like this: [Metal]O + H2O --> [base]
  3. Nonmetal oxide-water reactions go like this: [Nonmetal]O + H2O --> [acid]
  • Put them together! Two reactants come together to form one product.

Decomposition Reactions[edit | edit source]

  1. Binary compounds form: AB --> A + B
  2. Metallic carbonates: [Metal]CO3 --> [Metal]O + CO2
  3. Metallic hydrogen carbonates: [Metal]HCO3 --> [Metal]O + H2O
  4. Metallic hydroxide: [Metal]OH --> [Metal]O + H2O
  5. Metallic chlorate: MClO3 --> [Metal]Cl + O2
  6. Oxyacids decompose to nonmetal oxides and water: [Acid] --> [Nonmetal]O + H2O
  • A product is broken down into two reactants.

Single Replacement Reactions[edit | edit source]

  1. Metal replaces metal: A + BC --> AC + B
  2. Active metal replaces hydrogen from H2O: [Metal] + H2O --> [Metal]OH + H2
  3. Active metal replaces hydrogen from acid: [Metal] + H[acid] --> [Metal][Acid] + H2
  4. Halide-Halide replacement: [Halide A] + B[Halide B] --> [Halide B] + B[Halide A]
  • 1 single + 1 compound
  • "Like replace like" - cations/anions try to kick each other out

Metals can only replace another metal if they are more reactive than the metal they are positioned to replace. See an activity series graph (right), which showcases the reactivities of several metals. For the main trend, reactivity increases as you go down the periodic table while reactivity decreases as you go from left to right on the periodic table. If a metal cannot replace another metal because the latter metal is more reactive than the former metal, then no reaction occurs.

Double Replacement Reactions[edit | edit source]

  1. AB + CD --> AD + CB (Outer-Outer, Inner-Inner)
  2. Formation of a precipitate (reaction) from a solution
  3. Acid-Base neutralization reaction
  • 2 compounds turn into 2 new compounds (must include states of matter)

Combustion Reactions[edit | edit source]

  1. Hydrocarbon + oxygen --> carbon dioxide + water
  • "Combustion" translates to "burning", so: burning a hydrocarbon, which is any CH combo, and O2 turns into carbon dioxide and water. Answer is always CO2 + H2O

Acid-Base Neutralization Reactions[edit | edit source]

Very similar to double replacement reactions:
[Acid] (H) + [Base] (OH) --> salt (ionic compounds) + H2O

Oxidation-Reducation Reactions[edit | edit source]

Type of chemical reaction that consists of a transfer of electrons between two compounds.

Chemical Kinetics[edit | edit source]

A reaction rate is a change in concentration over a period of time. The collision theory states that atoms must collide in order to react. Effective collisions, collisions resulting in a chemical reaction, will cause a chemical reaction only if reactions collide when colliding molecules run into one another with the appropriate amount of energy and proper orientation.

The minimum amount of energy needed for a reaction to occur is called the activation energy.

Reaction Coordinate Diagram.png

In the above reaction coordinate diagram, the reactants, A, produce the products, C (as displayed at the bottom of the graph). The graph is evidently exothermic, as energy is being released. The △G≠ represents the activation energy while B represents the activation complex. The △G° represents the ending energy. Since this graph is exothermic, the ending energy will be a negative value (as energy has been lost). The formula for △G° (heat of reaction) is (products minus reactants = ending energy).

Fast vs. Spontaneous

A spontaneous reaction is a reaction that will occur under shown conditions. The rate addresses how fast it will occur. This, to a certain extent, depends on the activation energy. A reaction with a large activation energy will only occur at a slow pace/will not occur at all without an initial input of energy to overcome the activation energy barrier.

Factors affecting reaction rate[edit | edit source]

  1. Temperature
    • As temperature increases, so does the reaction rate. More reactants have enough kinetic energy to overcome its activation energy.
  2. Concentration
    • Increased concentration causes the reaction rate to increase. Due to the increase in the number of particles present, more collisions will occur.
  3. Surface Area
    • As surface area increases, the reaction rate increases as well. Greater surface area allows more particles to collide with each other. Surface area can be increased by either crushing, grinding, or dissolving.
  4. Catalyst
    • Decreasing the activation energy, it increases the reaction rate.
  5. Volume/Pressure
    • Volume and Pressure are opposites
    • If the volume of a reaction container decreases, the number of successful collisions and reaction rate will increase.

Chemical Equilibrium[edit | edit source]

Reactions that can happen in both directions are called reversible reactions. At equilibrium, the forward and reverse reactions occur at the same rate while concentration might not be generally equal. If the forward reaction is favored, this indicates that more products exist than reactants. Vice versa applies.

Le Chatelier's Principle[edit | edit source]

Le Chatelier's Principle is a system where equilibrium remains during the equilibrium state. When a reaction is disturbed, the system will return the reaction back to its equilibrium state.

Concentration[edit | edit source]

  • When something is added, it shifts away from that thing
  • When something is removed, it shifts towards that thing
  • This only happens with amounts of (aq - aqueous) and (g - gases) count. Solids and liquids do not have any effect. Adding a substance not in the reaction also has no effect.

Volume/Pressure[edit | edit source]

  • If the volume is reduced/pressure is increased, it shifts to the side with fewer gas particles.
  • If the volume is increased/pressure is reduced, it shifts to the side with more gas particles.

Temperature[edit | edit source]

  • If temperature is increased, it shifts to the side without heat
  • If temperature is decreased, it shifts to the side with heat

Equilibrium Expressions[edit | edit source]

Keq =

You only include aqueous and gaseous reactants and products. You keep the subscripts inside the paranthesis and add the coefficients as exponents. If:

  • K>1: Products are favored
  • K<1: Reactants are favored

Stoichiometry[edit | edit source]

Stoichiometry is scientists' explanation that gives a description of the quantitative relationship between different quantities of the reactants and products. You must know how to do dimensional analysis and molehill conversions (see lesson 1 and lesson 5). This is the way you figure out how much of a reaction/product you need in a chemical reaction.

We're going to analyze the chemical equation below:

  • 1N2 + 3H3 --> 2NH3

The coefficients here represent the number of moles. So, this equation means it takes 1 mole of nitrogen and 3 moles of hydrogen to produce 2 moles of ammonia. The mole ratio of this equation is 1:3:2.

Mole → Mole[edit | edit source]

  1. How many moles of H2O can be obtained from the reaction of 6.0 moles of oxygen in the following equation?
2H2 + 1O2 --> 2H2O
6.0mol O2 • 2mol H2O/1mol O2 = 12mol H2O

Mass → Mole[edit | edit source]

Mole → Mass[edit | edit source]

Mass → Mass[edit | edit source]