Chemical Reactions Overview

Chemical Reactions

• Chemical reactions involve the transformation of one or more substances into different ones.
• This change alters the arrangement of atoms and molecules.
• Chemical reactions are evidenced by various observations, including color change, solid formation, gas evolution, absorption/emission of heat, or light emission.
• However, these are not always definitive.
• Sometimes, a reaction occurs without apparent signs.

Evidence of Chemical Reactions

• Color change
• Formation of a solid
• Formation of a gas
• Absorption/emission of heat
• Light emission

Lab Evidence in Chemical Reactions

• Wooden splint test: A "lit" splint is used to test for hydrogen (flammable), carbon dioxide (causes the splint to extinguish), or oxygen (keeps the splint burning).
• pH indicator (phenolphthalein): This indicator turns pink in the presence of hydroxide ions, which signals a basic solution. It is colorless in neutral or acidic solutions.

Law of Conservation of Matter and Mass

• Matter and mass are neither created nor destroyed in a chemical reaction.
• The total mass of reactants equals the total mass of products.
  • Example: Water decomposing into hydrogen and oxygen, the mass of water equals the mass of hydrogen and oxygen combined.
  • Example: Butane reacting with oxygen. Mass of butane and oxygen reacted must equal the mass of carbon dioxide and water formed.

Chemical Equations

• Chemical equations represent chemical reactions.
• Reactants (on the left side of the equation) give way to products (on the right side).
• Use reaction arrows to connect reactants to products.
• States of matter (s, l, g, aq) are written in parentheses.

Diatomic Elements

• Some elements naturally exist as diatomic molecules (two atoms bonded together).
• These elements are: hydrogen (H₂), oxygen (O₂), fluorine (F₂), chlorine (Cl₂), bromine (Br₂), iodine (I₂), and nitrogen (N₂).
  • Remember these diatomic elements (HOFBrINCl) when writing equations.

Chemical Equations (continued)

• Chemical equations must be balanced (conservation of matter).
• Coefficients (numbers in front of compounds) determine the number of molecules
• Subscripts(numbers below elements in formulas) do not change when balancing.

Balancing Guidelines

• Balance elements that appear in only one compound on either side first.
• Balance metal atoms before nonmetal atoms.
• Treat polyatomic ions as units when balancing.

Writing/Balancing Reactions

• Write the skeleton equation.
• Place correct formulas on the correct sides.
• Use coefficients to balance the equation.

Correct Formulas and States of Matter

• Ionic compounds must be neutral.
• Molecular compounds have prefixes.
• States of matter are important for correctness.
• Consider common experiences (state of matter at room temp, etc.)
• Acids are aqueous (aq).
• Solubility rules are essential for aqueous compounds (are they soluble or insoluble? e.g. (aq) or (s)).

Aqueous Solutions

• Aqueous solutions are homogeneous mixtures where a substance is dissolved in water.
• Water molecules pull ions apart (NaCl) due to stronger attractions than between ions, creating solutions.
• Ionic compounds dissociate into their separate ions when dissolved in water (NaCl → Na+ + Cl–).
• Polyatomic ions remain intact when dissolved in water (e.g., NaNO3 → Na+ + NO3−).

Ionic Compound Solubility

• Insoluble ionic compounds form precipitates in solutions.
• Solubility rules dictate which ionic compounds are soluble or insoluble.

Classifying Reactions (continued)

• Synthesis: Two or more reactants combine to form a single product.
• Decomposition: A single reactant breaks down into two or more products.
• Single Displacement: One element replaces another element in a compound.
• Double Displacement: Two elements in different compounds exchange places to form new compounds.

Acid-Base Reactions

• Acid-base reactions, often called neutralization, involve the formation of water (H₂O) and a salt.
• Acids contain hydrogen (H⁺) ions. Bases contain hydroxide (OH⁻) ions.
• Acid-base reactions are double displacement.

Precipitation Reactions

• Reactions where two aqueous solutions mix, creating an insoluble solid (precipitate).
• Solubility rules are used to predict precipitates.

Redox Reactions

• Redox reactions involve electron transfer between reactants.
• One substance loses electrons (oxidation) while another gains them (reduction).

Combustion Reactions

• Combustion reactions are redox reactions where a substance reacts with oxygen to produce energy in the form of heat or light.
• When carbon and hydrogen compounds react with oxygen, they always yield carbon dioxide and water.

Predicting Products

• Use the type of reaction and reactants involved to predict the products.
• Consider solubility rules, the activity series, and the chemical nature of reactants to predict products.

Writing Equations and Predicting Products

• Use the type of reaction and the correct and/or neutral formulas for reactants and products.
• Use states of matter.
• Balance the equation using coefficients.

Predicting Products (continued)

• Determining the reaction type helps predict products.
• Consult the activity series for single-replacement reactions.

Limiting Reactions

• The limiting reactant determines the maximum amount of product.
• Determine the limiting reactant to determine how much product can be formed.
• Identify which reactant is limiting by comparing the amount of reactants available to their stoichiometric ratios, calculate how much product will be formed for each ratio.

Limiting Reactant (continued)

• Calculate the amount of product based on the complete consumption for each reactant.
• Determine which reactant is limiting using stoichiometry.
• Determine which reactant produces the smaller amount of product.

Yield

• Theoretical Yield: The maximum amount of product possible, calculated from the limiting reactant.
• Actual Yield: The amount of product experimentally obtained.
• Percent Yield: The ratio of the actual yield to the theoretical yield, expressed as a percentage.