Type of Reactions

This document explores the major types of chemical reactions, including their definitions, examples, and applications. The key categories covered are combustion, precipitation, acid-base (neutralization), and oxidation-reduction reactions.

1. Types of Chemical Reactions

Chemical reactions can be categorized based on how reactants interact to form products.

(a) Combination (Synthesis) Reactions

• Definition: Two or more reactants combine to form a single product.

• General Form: A+B→ABA + B \rightarrow ABA+B→AB

• Examples: 2CO+O2→2CO22 CO + O_2 \rightarrow 2 CO_22CO+O2​→2CO2​ 4Al+3O2→2Al2O34 Al + 3 O_2 \rightarrow 2 Al_2O_34Al+3O2​→2Al2​O3​

(b) Decomposition Reactions

• Definition: A compound breaks down into simpler compounds or elements.

• General Form: AB→A+BAB \rightarrow A + BAB→A+B

• Example: CaCO3→CaO+CO2CaCO_3 \rightarrow CaO + CO_2CaCO3​→CaO+CO2​

  • Electrolysis of Water: 2H2O→2H2+O22 H_2O \rightarrow 2 H_2 + O_22H2​O→2H2​+O2​

(c) Single Replacement Reactions

• Definition: One element replaces another in a compound.

• General Form: A+BC→AC+BA + BC \rightarrow AC + BA+BC→AC+B

• Examples: Zn+2HCl→ZnCl2+H2Zn + 2 HCl \rightarrow ZnCl_2 + H_2Zn+2HCl→ZnCl2​+H2​ 2AgNO3+Zn→2Ag+Zn(NO3)22 AgNO_3 + Zn \rightarrow 2 Ag + Zn(NO_3)_22AgNO3​+Zn→2Ag+Zn(NO3​)2​

(d) Double Replacement Reactions

• Definition: Two compounds exchange ions to form new compounds.

• General Form: AB+CD→AD+CBAB + CD \rightarrow AD + CBAB+CD→AD+CB

• Examples: NaOH+HCl→H2O+NaClNaOH + HCl \rightarrow H_2O + NaClNaOH+HCl→H2​O+NaCl Ba(OH)2+2HNO3→Ba(NO3)2+2H2OBa(OH)_2 + 2 HNO_3 \rightarrow Ba(NO_3)_2 + 2 H_2OBa(OH)2​+2HNO3​→Ba(NO3​)2​+2H2​O

2. Combustion Reactions

• Definition: A reaction where oxygen (O₂) reacts with another substance, releasing heat and energy.

• Characteristics:

  • Produces CO₂ and H₂O when hydrocarbons react with O₂.

  • Always exothermic.

• Examples: C6H12O6+6O2→6CO2+6H2OC_6H_{12}O_6 + 6 O_2 \rightarrow 6 CO_2 + 6 H_2OC6​H12​O6​+6O2​→6CO2​+6H2​O (Glucose combustion in respiration)

3. Precipitation Reactions

• Definition: A reaction in which an insoluble solid (precipitate) forms when two aqueous solutions mix.

• General Form: AB(aq)+CD(aq)→AD(s)+CB(aq)AB (aq) + CD (aq) \rightarrow AD (s) + CB (aq)AB(aq)+CD(aq)→AD(s)+CB(aq)

• Example: 2KI(aq)+Pb(NO3)2(aq)→PbI2(s)+2KNO3(aq)2 KI (aq) + Pb(NO_3)_2 (aq) \rightarrow PbI_2 (s) + 2 KNO_3 (aq)2KI(aq)+Pb(NO3​)2​(aq)→PbI2​(s)+2KNO3​(aq)

  • PbI₂ is the precipitate.

• Solubility Rules (Key Points):

  • Nitrate (NO3−NO_3^-NO3−​) and Acetate (C2H3O2−C_2H_3O_2^-C2​H3​O2−​) salts are always soluble.

  • Chlorides (Cl−Cl^-Cl−), Bromides (Br−Br^-Br−), and Iodides (I−I^-I−) are soluble EXCEPT with Ag⁺, Pb²⁺, and Hg₂²⁺.

  • Sulfates (SO42−SO_4^{2-}SO42−​) are soluble EXCEPT with Sr²⁺, Ba²⁺, Hg₂²⁺, and Pb²⁺.

Net Ionic Equations for Precipitation Reactions

• Molecular Equation: AgNO3(aq)+NaCl(aq)→AgCl(s)+NaNO3(aq)AgNO_3 (aq) + NaCl (aq) \rightarrow AgCl (s) + NaNO_3 (aq)AgNO3​(aq)+NaCl(aq)→AgCl(s)+NaNO3​(aq)

• Complete Ionic Equation: Ag+(aq)+NO3−(aq)+Na+(aq)+Cl−(aq)→AgCl(s)+Na+(aq)+NO3−(aq)Ag^+ (aq) + NO_3^- (aq) + Na^+ (aq) + Cl^- (aq) \rightarrow AgCl (s) + Na^+ (aq) + NO_3^- (aq)Ag+(aq)+NO3−​(aq)+Na+(aq)+Cl−(aq)→AgCl(s)+Na+(aq)+NO3−​(aq)

• Net Ionic Equation (removing spectator ions): Ag+(aq)+Cl−(aq)→AgCl(s)Ag^+ (aq) + Cl^- (aq) \rightarrow AgCl (s)Ag+(aq)+Cl−(aq)→AgCl(s)

4. Acid-Base (Neutralization) Reactions

• Definition: A reaction between an acid (H⁺ donor) and a base (OH⁻ donor) to form salt and water.

• General Form: HA+BOH→BA+H2OHA + BOH \rightarrow BA + H_2OHA+BOH→BA+H2​O

• Examples: HCl(aq)+NaOH(aq)→NaCl(aq)+H2O(l)HCl (aq) + NaOH (aq) \rightarrow NaCl (aq) + H_2O (l)HCl(aq)+NaOH(aq)→NaCl(aq)+H2​O(l) H2SO4(aq)+2NaOH(aq)→Na2SO4(aq)+2H2O(l)H_2SO_4 (aq) + 2 NaOH (aq) \rightarrow Na_2SO_4 (aq) + 2 H_2O (l)H2​SO4​(aq)+2NaOH(aq)→Na2​SO4​(aq)+2H2​O(l)

Strong vs. Weak Acids & Bases

5. Oxidation-Reduction (Redox) Reactions

• Definition: A reaction involving the transfer of electrons, leading to changes in oxidation states.

• Rules for Oxidation States:

  • Elements: 0 (e.g., O₂, N₂, Na).

  • Alkali metals: +1 (e.g., NaCl → Na⁺).

  • Alkaline earth metals: +2 (e.g., CaO → Ca²⁺).

  • Oxygen: -2 (except peroxides, where it’s -1).

  • Hydrogen: +1 (except metal hydrides, where it’s -1).

Redox Terms

• Oxidation = Loss of electrons (increase in oxidation state).

• Reduction = Gain of electrons (decrease in oxidation state).

• Oxidizing Agent = Causes oxidation (gets reduced).

• Reducing Agent = Causes reduction (gets oxidized).

Examples of Redox Reactions

1. Rusting of Iron:

   4Fe+3O2→2Fe2O34 Fe + 3 O_2 \rightarrow 2 Fe_2O_34Fe+3O2​→2Fe2​O3​

   • Fe oxidized (0 → +3).

   • O₂ reduced (0 → -2).

2. Combustion of Carbon:

   C+O2→CO2C + O_2 \rightarrow CO_2C+O2​→CO2​

   • C oxidized (0 → +4).

   • O₂ reduced (0 → -2).

3. Reaction of Magnesium and Hydrochloric Acid:

   Mg+2HCl→MgCl2+H2Mg + 2 HCl \rightarrow MgCl_2 + H_2Mg+2HCl→MgCl2​+H2​

   • Mg oxidized (0 → +2).

   • H⁺ reduced (+1 → 0).

Key Takeaways

• Chemical reactions are classified into synthesis, decomposition, single/double replacement, combustion, precipitation, acid-base, and redox reactions.

• Solubility rules help predict precipitates in double replacement reactions.

• Acid-base reactions always form salt + water.

• Redox reactions involve electron transfer and changes in oxidation states.

• "LEO goes GER" (Loss of Electrons is Oxidation, Gain of Electrons is Reduction).

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Comprehensive Summary of "Types of Reactions"

This document categorizes and explains the fundamental types of chemical reactions, their mechanisms, examples, and applications. It includes precipitation, acid-base (neutralization), combustion, and oxidation-reduction reactions with a focus on how substances interact and transform.

1. Types of Chemical Reactions

Chemical reactions are classified based on how reactants rearrange to form products.

(a) Combination (Synthesis) Reactions

• Definition: Two or more reactants combine to form a single product.

• General Equation: A+B→ABA + B \rightarrow ABA+B→AB

• Characteristics:

  • Simpler substances combine to form more complex compounds.

  • Usually exothermic (releases energy).

• Examples: 2CO+O2→2CO22 CO + O_2 \rightarrow 2 CO_22CO+O2​→2CO2​ 4Al+3O2→2Al2O34 Al + 3 O_2 \rightarrow 2 Al_2O_34Al+3O2​→2Al2​O3​ NH3+HCl→NH4ClNH_3 + HCl \rightarrow NH_4ClNH3​+HCl→NH4​Cl

(b) Decomposition Reactions

• Definition: A single compound breaks down into two or more simpler substances.

• General Equation: AB→A+BAB \rightarrow A + BAB→A+B

• Characteristics:

  • Requires energy input (heat, light, or electricity).

  • Opposite of combination reactions.

• Examples:

  • Thermal decomposition of limestone: CaCO3→CaO+CO2CaCO_3 \rightarrow CaO + CO_2CaCO3​→CaO+CO2​

  • Electrolysis of water: 2H2O→2H2+O22 H_2O \rightarrow 2 H_2 + O_22H2​O→2H2​+O2​

  • Decomposition of ammonium dichromate: (NH4)2Cr2O7→Cr2O3+N2+4H2O(NH_4)_2Cr_2O_7 \rightarrow Cr_2O_3 + N_2 + 4 H_2O(NH4​)2​Cr2​O7​→Cr2​O3​+N2​+4H2​O

(c) Single Replacement Reactions

• Definition: An element replaces another less reactive element in a compound.

• General Equation: A+BC→AC+BA + BC \rightarrow AC + BA+BC→AC+B

• Characteristics:

  • Occurs in aqueous solutions.

  • Reactivity series predicts whether a reaction occurs.

• Examples:

  • Zinc reacting with hydrochloric acid: Zn+2HCl→ZnCl2+H2Zn + 2 HCl \rightarrow ZnCl_2 + H_2Zn+2HCl→ZnCl2​+H2​

  • Iron displacing copper in copper sulfate: Fe+CuSO4→FeSO4+CuFe + CuSO_4 \rightarrow FeSO_4 + CuFe+CuSO4​→FeSO4​+Cu

(d) Double Replacement Reactions

• Definition: Two compounds exchange ions, forming two new compounds.

• General Equation: AB+CD→AD+CBAB + CD \rightarrow AD + CBAB+CD→AD+CB

• Characteristics:

  • Often leads to the formation of a precipitate, gas, or water.

• Examples:

  • Formation of barium sulfate precipitate: BaCl2+Na2SO4→BaSO4(s)+2NaClBaCl_2 + Na_2SO_4 \rightarrow BaSO_4 (s) + 2 NaClBaCl2​+Na2​SO4​→BaSO4​(s)+2NaCl

  • Neutralization of NaOH with HCl: NaOH+HCl→NaCl+H2ONaOH + HCl \rightarrow NaCl + H_2ONaOH+HCl→NaCl+H2​O

2. Combustion Reactions

• Definition: A reaction where a substance reacts with oxygen (O₂), releasing energy.

• General Equation: CxHy+O2→CO2+H2OC_xH_y + O_2 \rightarrow CO_2 + H_2OCx​Hy​+O2​→CO2​+H2​O

• Characteristics:

  • Highly exothermic.

  • Involves hydrocarbons, organic compounds, or metals.

• Examples:

  • Combustion of methane (natural gas): CH4+2O2→CO2+2H2OCH_4 + 2 O_2 \rightarrow CO_2 + 2 H_2OCH4​+2O2​→CO2​+2H2​O

  • Glucose combustion in respiration: C6H12O6+6O2→6CO2+6H2OC_6H_{12}O_6 + 6 O_2 \rightarrow 6 CO_2 + 6 H_2OC6​H12​O6​+6O2​→6CO2​+6H2​O

3. Precipitation Reactions

• Definition: A solid (precipitate) forms when two aqueous ionic solutions mix.

• General Equation: AB(aq)+CD(aq)→AD(s)+CB(aq)AB (aq) + CD (aq) \rightarrow AD (s) + CB (aq)AB(aq)+CD(aq)→AD(s)+CB(aq)

• Characteristics:

  • Uses solubility rules to predict product formation.

• Example: 2KI(aq)+Pb(NO3)2(aq)→PbI2(s)+2KNO3(aq)2 KI (aq) + Pb(NO_3)_2 (aq) \rightarrow PbI_2 (s) + 2 KNO_3 (aq)2KI(aq)+Pb(NO3​)2​(aq)→PbI2​(s)+2KNO3​(aq)

  • PbI₂ is the precipitate.

Solubility Rules

1. Soluble: Nitrate (NO3−NO_3^-NO3−​) and Acetate (C2H3O2−C_2H_3O_2^-C2​H3​O2−​) salts.

2. Insoluble: Carbonates (CO32−CO_3^{2-}CO32−​), phosphates (PO43−PO_4^{3-}PO43−​) except with Na⁺, K⁺, NH₄⁺.

Net Ionic Equations

• Complete Ionic Equation: Ag+(aq)+NO3−(aq)+Na+(aq)+Cl−(aq)→AgCl(s)+Na+(aq)+NO3−(aq)Ag^+ (aq) + NO_3^- (aq) + Na^+ (aq) + Cl^- (aq) \rightarrow AgCl (s) + Na^+ (aq) + NO_3^- (aq)Ag+(aq)+NO3−​(aq)+Na+(aq)+Cl−(aq)→AgCl(s)+Na+(aq)+NO3−​(aq)

• Net Ionic Equation (removing spectator ions): Ag+(aq)+Cl−(aq)→AgCl(s)Ag^+ (aq) + Cl^- (aq) \rightarrow AgCl (s)Ag+(aq)+Cl−(aq)→AgCl(s)

4. Acid-Base (Neutralization) Reactions

• Definition: An acid and a base react to form salt and water.

• General Equation: HA+BOH→BA+H2OHA + BOH \rightarrow BA + H_2OHA+BOH→BA+H2​O

• Examples: HCl+NaOH→NaCl+H2OHCl + NaOH \rightarrow NaCl + H_2OHCl+NaOH→NaCl+H2​O H2SO4+2KOH→K2SO4+2H2OH_2SO_4 + 2 KOH \rightarrow K_2SO_4 + 2 H_2OH2​SO4​+2KOH→K2​SO4​+2H2​O

5. Oxidation-Reduction (Redox) Reactions

• Definition: Reactions involving the transfer of electrons.

• Key Terms:

  • Oxidation: Loss of electrons.

  • Reduction: Gain of electrons.

  • Oxidizing Agent: Causes oxidation (gets reduced).

  • Reducing Agent: Causes reduction (gets oxidized).

Rules for Assigning Oxidation Numbers

• Elements = 0.

• Group 1 metals = +1.

• Oxygen = -2 (except peroxides, -1).

• Hydrogen = +1 (except with metals, -1).

Examples

1. Rusting of Iron:

   4Fe+3O2→2Fe2O34 Fe + 3 O_2 \rightarrow 2 Fe_2O_34Fe+3O2​→2Fe2​O3​

   • Fe oxidized (0 → +3).

   • O₂ reduced (0 → -2).

2. Combustion of Carbon:

   C+O2→CO2C + O_2 \rightarrow CO_2C+O2​→CO2​

   • C oxidized (0 → +4).

   • O₂ reduced (0 → -2).

3. Reaction of Magnesium and Hydrochloric Acid:

   Mg+2HCl→MgCl2+H2Mg + 2 HCl \rightarrow MgCl_2 + H_2Mg+2HCl→MgCl2​+H2​

   • Mg oxidized (0 → +2).

   • H⁺ reduced (+1 → 0).

Key Takeaways

• Reactions are classified into combination, decomposition, single/double replacement, combustion, precipitation, acid-base, and redox reactions.

• Solubility rules help predict precipitates.

• Acid-base reactions always form salt + water.

• Redox reactions involve electron transfer (LEO = Lose Electrons Oxidation, GER = Gain Electrons Reduction).