Lecture 7: Chemical Reactions Study Notes Study Notes on Chemical Reactions and Equations

Fundamental Principles of Chemical Reactions

• Definition: Chemical reactions involve chemical changes in matter that result in the formation of new substances.
• Mechanism: Reactions occur through the rearrangement and exchange of atoms to produce new molecules.
• Transmutation: It is important to note that elements are not transmuted (changed into different elements) during a chemical reaction.
• Basic Notation: Reactions are represented as:     $\text{Reactants} \rightarrow \text{Products}$

Diatomic Molecules

• There are 7 nonmetals that naturally occur as diatomic molecules in their elemental state.
• These elements must always be written as diatomic molecules whenever they appear in chemical reactions:
  • Hydrogen: $H_2$
  • Nitrogen: $N_2$
  • Oxygen: $O_2$
  • Fluorine: $F_2 elements$
  • Chlorine: $Cl_2$
  • Bromine: $Br_2$
  • Iodine: $I_2$

Predicting and Identifying Chemical Reactions

Driving Forces of Reactions

Several "forces" or thermodynamic drivers can indicate that a reaction will occur:

• Formation of a solid: The creation of an insoluble precipitate.
• Formation of water: Common in acid-base neutralization reactions.
• Formation of a gas: Observed as bubbles or effervescence.
• Transfer of electrons: The basis for oxidation-reduction (redox) reactions.

Visible Evidence of Reactions

Clues that suggest a chemical reaction has taken place include:

1. Color Changes: The solution or substance changes hue.
2. Formation of a Solid: A precipitate appears in a previously clear solution.
3. Formation of Bubbles: Indicates the evolution of a gas.
4. Energy Changes: Heat and/or a flame is produced (exothermic), or heat is absorbed (endothermic).

Chemical Equations and Symbols

• Definition: A chemical equation is a shorthand method for describing a reaction.
• Information Provided:
  • The specific formulas of reactants and products.
  • The physical states of all substances involved.
  • The relative numbers of molecules or moles required to satisfy mass balance.

Standard Symbols in Equations

• $\rightarrow$: Produces, yields, or gives (separates reactants from products).
• $+$: Reacts with, added to, or plus (separates two or more reactants/products).
• $\xrightarrow{\Delta}$: Heat is used as a catalyst for the reaction.
• $\xrightarrow{Fe}$: Iron (or another specified element/compound) is used as a catalyst.
• NR: Indicates no reaction occurs.

States of Matter Notations

• $(s)$: Solid substance or precipitate.
• $(l)$: Liquid substance.
• $(g)$: Gaseous substance.
• $(aq)$: Aqueous solution (substance dissolved in water).

Catalysts

• A catalyst is a substance written above the reaction arrow (e.g., $A + B \xrightarrow{\text{Catalyst}} C + D$).
• Function: It speeds up the rate of the reaction without being consumed in the process.

The Law of Conservation of Mass

• Core Principle: Matter cannot be created or destroyed.
• Atomic Persistence: In any chemical reaction, every atom present at the beginning (reactants) must still be present at the end (products).
• Mass Constancy: The total mass of the reactants must equal the total mass of the products.
• Example: Combustion of Methane:     $CH_4(g) + 2 O_2(g) \rightarrow CO_2(g) + 2 H_2O(l)$
  • Reactant side: $1\,C, 4\,H, 4\,O$
  • Product side: $1\,C, 4\,H, 4\,O$

Balancing Chemical Equations

• Balanced Equation: An equation where the number of atoms of each element is identical on both sides of the arrow.
• Coefficients: Whole numbers placed in front of chemical formulas to balance the equation.
• Coefficient Multiplication: A coefficient multiplies all subscripts in the formula.
  • Example: $3 H_2O$ contains $6$ hydrogen atoms ($3 \times 2$) and $3$ oxygen atoms ($3 \times 1$).

Guidelines for Balancing

1. Verify that all chemical formulas are written correctly first.
2. Never change the subscripts in a formula to balance the equation; change only the coefficients.
3. Start balancing with the element found in the most complex formula.
4. If a polyatomic ion appears unchanged on both sides of the equation, balance it as a single unit.

Examples of Writing and Balancing Equations

Example 1: Burning Magnesium Metal

• Word Equation: Magnesium(s) + Oxygen(g) $\rightarrow$ Magnesium oxide(s)
• Formula Identification: Metals are solids (except liquid mercury, $Hg$). "Burning in air" implies reaction with $O_2$.
• Unbalanced Formula: $Mg(s) + O_2(g) \rightarrow MgO(s)$
• Balanced Equation: $2 Mg(s) + O_2(g) \rightarrow 2 MgO(s)$

Example 2: Ammonia and Oxygen at $1000^{\circ}C$

• Word Equation: Ammonia(g) + Oxygen(g) $\rightarrow$ Nitrogen monoxide(g) + water(g)
• Unbalanced Formula: $NH_3(g) + O_2(g) \rightarrow NO(g) + H_2O(g)$
• Balanced Equation: $4 NH_3(g) + 5 O_2(g) \rightarrow 4 NO(g) + 6 H_2O(g)$

Learning Check: Practice Balancing

• Problem 1: $O_2 + PCl_3 \rightarrow POCl_3$
  • Solution: $O_2 + 2 PCl_3 \rightarrow 2 POCl_3$
• Problem 2: $P_4 + N_2O \rightarrow P_4O_6 + N_2$
  • Solution: $P_4 + 6 N_2O \rightarrow P_4O_6 + 6 N_2$
• Problem 3: $C_6H_5CH_3 + O_2 \rightarrow H_2O + CO_2$
  • Solution: $C_6H_5CH_3 + 9 O_2 \rightarrow 4 H_2O + 7 CO_2$
• Problem 4: $CO_2 + NH_3 \rightarrow CO(NH_2)_2 + H_2O$
  • Solution: $CO_2 + 2 NH_3 \rightarrow CO(NH_2)_2 + H_2O$
• Problem 5: $UO_2 + HF \rightarrow UF_4 + H_2O$
  • Solution: $UO_2 + 4 HF \rightarrow UF_4 + 2 H_2O$
• Problem 6: $Ag + H_2SO_4 \rightarrow Ag_2SO_4 + SO_2 + H_2O$
  • Solution: $2 Ag + 2 H_2SO_4 \rightarrow Ag_2SO_4 + SO_2 + 2 H_2O$

Types of Chemical Reactions

1. Combustion Reactions

• A subclass of Oxidation-Reduction reactions.
• Involves the reaction of a compound (usually carbon-based) with $O_2(g)$.
• Products:
  • Carbon compounds produce $CO_2(g)$.
  • Hydrogen-containing compounds produce $H_2O(g)$.
• Example: $C_3H_8(g) + 5 O_2(g) \rightarrow 3 CO_2(g) + 4 H_2O(g)$

2. Synthesis (Combination) Reactions

• Reactions where two or more chemicals combine to form a single, more complex product.
• General Format: $A + B \rightarrow AB$
• Example 1: $2 Na(s) + Cl_2(g) \rightarrow 2 NaCl(s)$
• Example 2: $N_2(g) + O_2(g) \rightarrow 2 NO(g)$

3. Decomposition Reactions

• Reactions where a single reactant breaks down into two or more simpler molecules or elements.
• General Format: $AB \rightarrow A + B$
• Example (using electric current): $2 NaCl(l) \rightarrow 2 Na(l) + Cl_2(g)$

4. Single-Replacement Reactions

• A reaction in which a more active metal displaces a less active metal from a compound.
• General Format: $A + BZ \rightarrow AZ + B$
• Example: $Fe(s) + CuSO_4(aq) \rightarrow FeSO_4(aq) + Cu(s)$

5. Double-Displacement (Double-Replacement) Reactions

• Reactions involving the exchange of ions between two compounds.
• General Format: $AB + CD \rightarrow AD + CB$
• Example: $AgNO_3(aq) + KCl(aq) \rightarrow AgCl(s) + KNO_3(aq)$
• Subtypes:
  1. Precipitation reactions: Produce an insoluble solid.
  2. Acid-Base (Neutralization) reactions: Produce water and a salt. ($HX + BOH \rightarrow BX + HOH$)

The Activity Series

Concept of Activity

• The activity of a metal measures its ability to compete in a replacement reaction.
• Metals are arranged based on how easily they lose electrons.
• Reactivity Order: The most reactive metals appear first; the least reactive appear last.
• Substitution Rule: A metal can only replace another metal in a compound if it precedes it in the activity series.

Specific Displacements

• Acid Displacements: Metals that precede Hydrogen $(H)$ in the activity series react with acids to produce $H_2(g)$ and an ionic compound.
  • Example: $Fe(s) + 2 HCl(aq) \rightarrow FeCl_2(aq) + H_2(g)$
  • Example (No Reaction): $Au(s) + H_2SO_4(aq) \rightarrow NR$ (Gold is less active than hydrogen).
• Active Metals and Water: A few highly active metals (e.g., alkali metals) react directly with liquid water to produce a metal hydroxide and $H_2(g)$.
  • Example 1: $2 Na(s) + 2 H_2O(l) \rightarrow 2 NaOH(aq) + H_2(g)$
  • Example 2: $Ca(s) + 2 H_2O(l) \rightarrow Ca(OH)_2(aq) + H_2(g)$

Solubility Rules for Precipitation Reactions

To predict if a precipitate will form in a double-displacement reaction, use the following rules for aqueous solutions:

Soluble Compounds$(aq)$

1. All salts containing $NO_3^{-}$ are soluble.
2. All salts containing $Na^{+}$, $K^{+}$, or $NH_4^{+}$ are soluble.
3. Most salts containing $Cl^{-}$, $Br^{-}$, or $I^{-}$ are soluble, except those containing $Ag^{+}$, $Hg_2^{2+}$, or $Pb^{2+}$.
4. Most salts containing $SO_4^{2-}$ are soluble, except those containing $Ba^{2+}$, $Pb^{2+}$, or $Ca^{2+}$.

Insoluble Compounds$(s)$

1. Most salts containing $S^{2-}$, $CO_3^{2-}$, or $PO_4^{3-}$ are insoluble.
2. Most salts containing $OH^{-}$ are insoluble, except $NaOH$ and $KOH$ which are soluble, and $Ba(OH)_2$ and $Ca(OH)_2$ which are moderately soluble.

Ionic and Net Ionic Equations

Dissociation

• When ionic compounds dissolve in water, the anions and cations separate from each other.
• Example 1: $KCl(aq) \rightarrow K^{+}(aq) + Cl^{-}(aq)$
• Example 2: $K_2SO_4(aq) \rightarrow 2 K^{+}(aq) + SO_4^{2-}(aq)$

Types of Chemical Equations

1. Molecular Equation: Shows the complete formulas of all reactants and products as if they were molecules.
   • Example: $KCl(aq) + AgNO_3(aq) \rightarrow KNO_3(aq) + AgCl(s)$
2. Ionic Equation: Shows all strong electrolytes (dissolved ionic compounds) as free ions.
   • Example: $K^{+}(aq) + Cl^{-}(aq) + Ag^{+}(aq) + NO_3^{-}(aq) \rightarrow K^{+}(aq) + NO_3^{-}(aq) + AgCl(s)$
3. Net Ionic Equation: Shows only the components directly involved in the chemical change. Spectator ions are removed.
   • Example: $Ag^{+}(aq) + Cl^{-}(aq) \rightarrow AgCl(s)$

Spectator Ions

• Definition: Ions that appear on both the reactant and product sides of an ionic equation.
• They do not participate in the actual chemical reaction.

Learning Check: Comprehensive Practice

Case 1: Copper(II) nitrate + Potassium chromate

• Molecular Equation: $Cu(NO_3)_2(aq) + K_2CrO_4(aq) \rightarrow CuCrO_4(s) + 2 KNO_3(aq)$
• Question: Does reaction occur? Yes, because solid $CuCrO_4$ (Copper(II) chromate) forms.
• Product Names: Copper(II) chromate and Potassium nitrate.

Case 2: Lead(II) nitrate + Cesium sulfate

• Formulas: $Pb(NO_3)_2$ and $Cs_2SO_4$
• Molecular Equation: $Pb(NO_3)_2(aq) + Cs_2SO_4(aq) \rightarrow PbSO_4(s) + 2 CsNO_3(aq)$
• Product Names: Lead(II) sulfate and Cesium nitrate.

Case 3: Sodium sulfide + Copper(II) chloride

• Reaction: $Na_2S(aq) + CuCl_2(aq) \rightarrow 2 NaCl(aq) + CuS(s)$
• Conclusion: Only $CuS$ will precipitate from the solution; $NaCl$ remains aqueous.

Case 4: Net Ionic Equation Problem

• Scenario: Lead(II) nitrate + Sodium chloride
• Molecular: $Pb(NO_3)_2(aq) + 2 NaCl(aq) \rightarrow PbCl_2(s) + 2 NaNO_3(aq)$
• Ionic: $Pb^{2+}(aq) + 2 NO_3^{-}(aq) + 2 Na^{+}(aq) + 2 Cl^{-}(aq) \rightarrow PbCl_2(s) + 2 Na^{+}(aq) + 2 NO_3^{-}(aq)$
• Net Ionic: $Pb^{2+}(aq) + 2 Cl^{-}(aq) \rightarrow PbCl_2(s)$

Case 5: Complex Beaker Analysis

• Beaker contains: $Na^{+}(aq), C_2H_3O_2^{-}(aq)$, and solid $SrSO_4$
• Molecular: $Na_2SO_4(aq) + Sr(C_2H_3O_2)_2(aq) \rightarrow SrSO_4(s) + 2 NaC_2H_3O_2(aq)$
• Net Ionic: $Sr^{2+}(aq) + SO_4^{2-}(aq) \rightarrow SrSO_4(s)$