Unit 6: Chemical Reactions

Chemical reactions

Can be represented using chemical equations and word equations.

Coefficient

From the balanced chemical equation.

Subscript

Tell us the nature of the chemical (e.g. mono-, di-, tetra- or octatomic, etc.).

State

Solid (s), liquid (l), gas (g), and aqueous (aq).

System

The part of the universe, being studied, open or closed.

Open system

A system where things (mass, energy) can enter or leave (e.g. a beaker of water on a hotplate).

Closed system

A system where nothing can enter or leave (e.g. a thermos)

Gases (g)

• H₂, O₂, N₂, Cl₂, F₂, I₂ are diatomic gases at room temperature.

• Br₂ is a diatomic liquid at room temperature.

• “Have No Fear Of Ice Cold (Cl) Beer (Br)”

• During a chemical reaction, gases may evolve:

  • Hydrogen (H₂): a burning splint POP.

  • Carbon dioxide (CO₂): a burning splint goes out.

  • Oxygen (O₂): a distinguished splint glows.

Liquid (l)

• Br, Hg, and H₂O are liquid at room temperature.

• Pure substances.

Aqueous (aq)

• In order to react, we dissolve solutes in water to form solutions.

• Bases and acids are always aqueous.

• Nitrate salts (“X”NO₃) and Alkali salts (1^st column on P. Table) are always aqueous.

  → Ex: Mg(NO₃)₂ and K₂S.

Solids (s)

• A solid forming from a chemical reaction is called a precipitate (ppt.).

• Elemental metals may also form during a reaction, it will usually be more dense than a precipitate.

• Most elements on the periodic table are present in solid form at room temperature.

Law of conservation of mass (LOCOM)

Total mass of all reactants = Total mass of all products (mass is conserved)

Law of conservation of atoms

The total number and type of atoms does not change during a chemical reaction.

Law of conservation of energy

Total energy does not change during a chemical reaction (may change form).

Synthesis

• Combine 2 or more substances; usually results in 1 product.

• May need a catalyst to speed up the process.

  • Catalyst: a substance that speeds up a reaction without being used up.

  • Catalysts are NOT included in the chemical reaction.

• # reactants > # products.

Decomposition

• One substance breaks down into 2 or more substances.

• May need heat, light, or a catalyst for a reaction.

• # products > # reactants.

Single replacement

• One element is replaced by another in a compound.

  • Metal replaces metal

  • Non-metal replaces non-metal

• Look for single switch.

Double replacement

• Metals and non-metals switch partners to form new products.

• Look for double switch.

Neutralization

• A special type of double replacement reaction.

• Acid + Base → Salt (ionic compound) + Water.

  → Ex: HBr + KOH → H₂O + KBr

• Base has metal + OH^- (e.g. KOH, NaOH, etc.)

• Acid has H⁺ + non-metals/ polyatomic anions (except O₂) (e.g. H₂SO₄, HNO₃, etc.)

Combustion

• Burning, explosive; the reaction of a substance with O₂ that produces energy as heat and light.

• Look for a compound reacting with oxygen (O₂).

• Common combustion reaction:
  C_xH_y + O₂ → CO₂ + H₂O (x and y are whole numbers)

• C_xH_yS_z + O₂ → CO₂ + H₂O + SO₂

• C_xH_yN_z + O₂ → CO₂ + H₂O + NO₂

Prediction of chemical formulas (Single replacement reactions)

• Metal replacement

• Hydrogen replacement

• Non-metal replacement

Chemical reactivity

• The tendency of a substance to undergo chemical change.

• Related to periodic trends in electronegativity and ionization energy.

• If the reactants do not react, we write “NR”.

Energy exchange during chemical reactions

• Bonds breaking: Energy from the surroundings enters the system.

• Bonds forming: Energy leaves the system to enter the surroundings.

• Energy is transferred as heat (average kinetic energy).

• Energy stored in bonds is potential energy.

• Energy exchange can be represented using
  Δ delta H notation (ΔH = H of products - H of reactants) and a thermochemical equation.

Enthalpy

• The total energy contained in a system.

• Chemical potential energy.

• Symbol for enthalpy: “H”.

• Change in enthalpy during a chemical reaction: ΔH (Δ - delta means “change in”).

  → Ex: Gasoline contains more stored chemical potential energy = enthalpy than water.

Exothermic reactions

• Convert the enthalpy stored in a substance into heat.

• ΔH is negative.

• Products are lower than Reactants in the Enthalpy Diagram.

• Heat (kinetic energy) is released to the surroundings. Surroundings feel warmer relative to the system.

Endothermic reactions

• Heat is absorbed from the surroundings and converted into enthalpy.
  → Ex: NaHCO₂ + HCl + heat → H₂O + CO₂ + NaCl

• ΔH is positive

• Products are higher than Reactants in the Enthalpy Diagram.

• Surroundings feel cooler relative to the system.