Unit 4: Chemical Reactions

Types of Chemical Reactions

• Synthesis (Combination) Reactions occur when two or more reactants combine to form a single product. These reactions are typically exothermic because bond formation releases energy, and they are often used in industrial processes like ammonia production via the Haber process. The general form is \( A + B \rightarrow AB \), and the reaction can involve elements or simpler compounds combining to create more complex substances.
• Decomposition Reactions involve a single compound breaking down into two or more simpler products, usually requiring energy input in the form of heat, light, or electricity. These reactions are often endothermic because bonds in the original compound must be broken. The general form is \( AB \rightarrow A + B \), and common examples include the thermal decomposition of calcium carbonate into calcium oxide and carbon dioxide.
• Single Replacement (Displacement) Reactions occur when an element reacts with a compound and replaces one of its components. The reactivity of the elements determines whether the reaction will proceed, and activity series charts help predict outcomes. For example, a more reactive metal can replace a less reactive metal from its compound, as in \( Zn + CuSO_4 \rightarrow ZnSO_4 + Cu \).
• Double Replacement (Metathesis) Reactions involve the exchange of ions between two compounds to form two new compounds. These often occur in aqueous solutions and result in the formation of a precipitate, gas, or weak electrolyte like water. The general form is \( AB + CD \rightarrow AD + CB \), and examples include precipitation reactions and acid–base neutralizations.
• Combustion Reactions involve the rapid reaction of a substance with oxygen to produce energy in the form of heat and light. Hydrocarbon combustion typically yields carbon dioxide and water, and complete combustion requires excess oxygen. Incomplete combustion, due to limited oxygen, produces carbon monoxide or soot, which is less efficient and can be hazardous.
• While these five categories help organize reaction types, many real-world reactions do not fit perfectly into a single category and may involve overlapping characteristics. Understanding the driving forces—such as the formation of a precipitate, gas evolution, or energy changes—helps predict and classify reactions accurately.

Balancing Chemical Equations

• Balancing chemical equations is based on the Law of Conservation of Mass, which states that matter cannot be created or destroyed in a chemical reaction. This means the number of atoms of each element must be the same on both sides of the equation. A balanced equation ensures that the stoichiometric relationships used in calculations reflect reality.
• To balance an equation, adjust the coefficients (the numbers in front of compounds or elements) rather than changing the subscripts in chemical formulas. Changing subscripts alters the identity of the substance, which is chemically incorrect. Coefficients represent the relative number of moles of each reactant and product.
• A common strategy for balancing involves starting with the most complex molecule, typically the one containing the most elements. Balance atoms of elements that appear in only one reactant and one product first, then proceed to elements that appear in multiple compounds. Save balancing hydrogen and oxygen for last, especially if they appear in multiple species.
• For reactions with polyatomic ions that remain unchanged on both sides, it is often easier to treat the polyatomic ion as a single unit. This reduces the complexity of balancing by minimizing the number of species to track. For example, in a reaction involving sulfate (\( \text{SO}_4^{2-} \)), treat the sulfate as a single entity if it appears intact on both sides.
• It is important to verify your final equation by counting the number of each atom type on both sides to ensure perfect equality. If the equation includes fractions during balancing, multiply all coefficients by the denominator to obtain whole numbers. The smallest set of whole-number coefficients is considered the correctly balanced form.
• Balancing equations is not only necessary for accurate representation of reactions but also crucial for stoichiometric calculations in determining reactant amounts, product yields, and energy changes. Without a balanced equation, mole ratios are incorrect, and any further chemical calculations will be invalid.

Stoichiometry and Mole Ratios

• Stoichiometry is the quantitative relationship between reactants and products in a balanced chemical equation. It allows chemists to predict the amounts of substances consumed and produced in a reaction. These relationships are based entirely on the mole ratios given by the coefficients in the balanced equation.
• Mole ratios are derived directly from the balanced equation and are essential for converting between moles of different substances. For example, in the equation \( 2H_2 + O_2 \rightarrow 2H_2O \), the mole ratio of hydrogen to water is \( 2:2 \) or \( 1:1 \). These ratios enable calculations from one known quantity to another through dimensional analysis.
• Stoichiometry problems often require multiple conversions: mass → moles → moles → mass. First, convert the known substance’s mass to moles using molar mass, then use the mole ratio to find moles of the unknown substance, and finally convert back to mass if needed. The same principle applies for volume relationships in gases at standard conditions.
• Reactions are rarely perfect, so percent yield is used to compare the actual yield (measured experimentally) to the theoretical yield (calculated stoichiometrically). Percent yield is calculated as \( \frac{\text{actual yield}}{\text{theoretical yield}} \times 100\% \), and values less than 100% usually result from incomplete reactions or loss of product.
• When reactants are not present in exact stoichiometric proportions, the limiting reactant is the one completely consumed first. The limiting reactant determines the maximum amount of product that can be formed, while the excess reactant is left over. Identifying the limiting reactant requires calculating the product yield from each reactant and selecting the smaller result.
• Mastering stoichiometry is critical for laboratory efficiency, industrial production, and environmental impact calculations. It ensures that resources are not wasted and helps control the proportions of chemicals in large-scale manufacturing to minimize costs and hazards.

Net Ionic Equations

• Net ionic equations show only the species that actually participate in the chemical reaction, removing the ions that remain unchanged in solution. These unchanged ions are called spectator ions, and while they are present in the reaction mixture, they do not affect the chemical change taking place.
• To write a net ionic equation, start with a correctly balanced molecular equation, which lists all reactants and products as complete compounds. Next, rewrite it as a complete ionic equation by breaking all strong electrolytes (strong acids, strong bases, and soluble ionic salts) into their constituent ions while leaving solids, liquids, and gases intact.
• After writing the complete ionic equation, identify and cancel the spectator ions that appear on both sides. The remaining species make up the net ionic equation, which highlights the actual chemical transformation taking place. For example, \( AgNO_3(aq) + NaCl(aq) \rightarrow AgCl(s) + NaNO_3(aq) \) becomes \( Ag^+(aq) + Cl^-(aq) \rightarrow AgCl(s) \).
• It is important to know which compounds dissociate in water, which is determined by solubility rules. For example, nitrates (\( NO_3^- \)) are always soluble, whereas most chlorides are soluble except those of silver, lead(II), and mercury(I). Solubility knowledge is critical for predicting which products will precipitate and thus appear in the net ionic equation.
• Net ionic equations are valuable because they provide a clear, simplified picture of the core chemical event. They are especially useful in analyzing precipitation reactions, acid–base neutralizations, and certain redox processes, all of which involve ionic species in aqueous environments.
• By focusing only on the reacting species, net ionic equations make it easier for students to understand reaction mechanisms, stoichiometry in ionic systems, and the role of specific ions in driving chemical changes.

Predicting Products of Chemical Reactions

• Predicting products begins with identifying the type of reaction, as each reaction class follows general patterns. Common types include synthesis (two or more reactants form one product), decomposition (one compound breaks into simpler substances), single replacement (an element replaces another in a compound), double replacement (ions exchange partners between compounds), and combustion (a hydrocarbon reacts with oxygen to produce CO₂ and H₂O).
• In synthesis reactions, the products can often be predicted by combining reactants in the simplest ratio. For example, metal oxides reacting with water generally form bases (\( CaO + H_2O \rightarrow Ca(OH)_2 \)), while nonmetal oxides with water often form acids (\( SO_3 + H_2O \rightarrow H_2SO_4 \)). Recognizing these patterns helps with quick product identification.
• In decomposition reactions, products are usually more stable substances. Metal carbonates often decompose into metal oxides and CO₂, and metal chlorates decompose into metal chlorides and O₂. Heat or electricity is often required to break the bonds in these compounds, making decomposition less common without energy input.
• Single replacement reactions depend on the activity series, which ranks elements by reactivity. A more reactive element will replace a less reactive element from its compound. For example, \( Zn + CuSO_4 \rightarrow ZnSO_4 + Cu \) occurs because zinc is more reactive than copper, but reversing the reaction would not occur spontaneously.
• In double replacement reactions, predicting products requires knowing solubility rules to determine if a precipitate, gas, or weak electrolyte forms. For instance, when \( BaCl_2 \) and \( Na_2SO_4 \) react, \( BaSO_4 \) precipitates because it is insoluble, leaving \( NaCl \) in aqueous form. This product prediction step is key for identifying actual chemical changes.
• Combustion reactions involving hydrocarbons almost always produce CO₂ and H₂O if oxygen is sufficient. If oxygen is limited, incomplete combustion occurs, producing CO and soot instead. Balancing combustion equations often involves adjusting oxygen last to simplify the process.
• Understanding these patterns allows chemists to not only write correct product formulas but also predict reaction feasibility. These skills are essential in laboratory planning, industrial chemical production, and environmental chemistry, where knowing possible reaction outcomes helps prevent hazardous byproducts.

Redox Reactions

Definition and Concept of Redox

• Redox reactions involve the simultaneous processes of oxidation (loss of electrons) and reduction (gain of electrons). In any redox process, one species donates electrons while another accepts them, making electron transfer the central event. These reactions are essential to understanding energy flow in chemistry, as they underpin processes like combustion, corrosion, and electrochemical cell operation.
• The terms “oxidizing agent” and “reducing agent” describe the roles of reactants: the oxidizing agent gains electrons (and is reduced), while the reducing agent loses electrons (and is oxidized). This dual nature means every oxidation must be accompanied by a reduction, which can be tracked using oxidation numbers or half-reaction methods. Recognizing this relationship is key to balancing and predicting redox chemistry outcomes.

Oxidation Numbers and Their Rules

• Oxidation numbers are a bookkeeping tool that help identify which atoms undergo oxidation or reduction. Common rules include: elements in their pure form have an oxidation number of 0, oxygen is usually −2, hydrogen is usually +1, and the sum of oxidation numbers in a neutral compound is 0. Applying these systematically allows chemists to detect changes in electron ownership during reactions.
• Changes in oxidation numbers from reactants to products indicate electron transfer. If the oxidation number increases, the element is oxidized; if it decreases, the element is reduced. This method works even in complex ionic compounds and polyatomic ions, making it a reliable way to classify reactions as redox or non-redox.

Half-Reaction Method

• The half-reaction method separates a redox reaction into its oxidation and reduction components, allowing each to be balanced individually for mass and charge. First, assign oxidation numbers, then write separate half-reactions for oxidation and reduction. Each half-reaction is balanced for atoms other than oxygen and hydrogen first, then for oxygen by adding water, and hydrogen by adding H⁺ (in acidic solution) or OH⁻ (in basic solution).
• After balancing atoms, balance charges by adding electrons to the more positive side of each half-reaction. The number of electrons lost in oxidation must equal the number gained in reduction, so multiply each half-reaction by the necessary factor before adding them together. This method ensures both mass and charge conservation, which are fundamental to all chemical equations.

Electrochemical Cells and Redox

• Electrochemical cells use redox reactions to generate or consume electrical energy. In a galvanic cell, a spontaneous redox reaction drives electron flow from the anode (site of oxidation) to the cathode (site of reduction) through an external circuit. A salt bridge maintains charge neutrality by allowing ion migration without mixing the solutions.
• In electrolytic cells, an external voltage source forces a nonspontaneous redox reaction to occur, such as electrolysis of water or metal plating. Understanding cell notation, standard reduction potentials, and the direction of electron flow is crucial for predicting cell behavior and calculating cell voltages. Both types of cells illustrate the direct conversion between chemical and electrical energy.

Balancing Redox Reactions

Importance of Balancing Redox Reactions

• Balancing redox reactions ensures that both mass and charge are conserved, which is a fundamental law of chemistry. Since redox processes involve electron transfer, it is essential to verify that the number of electrons lost in oxidation equals the number gained in reduction. This guarantees that the reaction reflects the true stoichiometry occurring in nature or in the laboratory.
• Unbalanced redox equations can lead to incorrect predictions about reactant and product quantities, misinterpretations of reaction mechanisms, and errors in related calculations such as cell potentials or thermodynamic values. Therefore, mastering the balancing process is crucial for accuracy in both academic problem-solving and industrial applications.

Half-Reaction Method in Acidic Solutions

• Begin by assigning oxidation numbers to all atoms in the reaction to identify which species are oxidized and reduced. Separate the equation into two half-reactions: one for oxidation and one for reduction. This separation allows for individual balancing of mass and charge in each half.
• Balance all elements other than oxygen and hydrogen first. Then, balance oxygen atoms by adding H₂O molecules and hydrogen atoms by adding H⁺ ions. Finally, balance the charge of each half-reaction by adding electrons to the more positive side. Multiply the half-reactions by the necessary factors so that the electrons cancel when the two halves are combined.

Half-Reaction Method in Basic Solutions

• Follow the same initial steps as in acidic conditions: assign oxidation numbers, separate into half-reactions, and balance non-oxygen/non-hydrogen atoms first. In basic solutions, oxygen is still balanced with H₂O, but hydrogen atoms are balanced by adding OH⁻ ions instead of H⁺.
• After balancing hydrogen with OH⁻, combine any excess OH⁻ and H⁺ (if present) to form water, simplifying the equation. Check that the electrons lost and gained are equal, then combine the half-reactions to produce the overall balanced equation. This process ensures the reaction correctly reflects the alkaline environment.

Common Errors to Avoid

• Students often forget to multiply the entire half-reaction (including water, H⁺/OH⁻, and electrons) when equalizing electron transfer between oxidation and reduction halves. This omission leads to mismatched electron counts and an unbalanced overall equation.
• Another common mistake is neglecting to double-check both mass and charge after combining the half-reactions. Even if atoms appear balanced, a charge imbalance indicates a step was skipped or an electron count was miscalculated, requiring careful review of each stage.

Solubility Rules

Purpose of Solubility Rules

• Solubility rules are guidelines used to predict whether an ionic compound will dissolve in water or form a precipitate. Understanding these rules allows chemists to determine the products of double-displacement reactions and design precipitation experiments. These rules are especially useful for quickly predicting outcomes without requiring experimental testing.
• While solubility is often treated as a yes-or-no property, in reality it exists on a spectrum. Many compounds labeled as “insoluble” may dissolve slightly, producing a small but measurable ion concentration in solution. Recognizing this nuance is important when dealing with equilibrium problems and ionic strength calculations.

General Soluble Compounds

• All salts containing alkali metal cations (Li⁺, Na⁺, K⁺, Rb⁺, Cs⁺) and the ammonium ion (NH₄⁺) are soluble with no significant exceptions. These ions do not form insoluble salts because their hydration energy outweighs lattice energy, making dissolution favorable.
• Nitrates (NO₃⁻), acetates (CH₃COO⁻), and most perchlorates (ClO₄⁻) are soluble in water regardless of the cations involved. Their large, delocalized anions interact weakly with cations, making it easy for water molecules to separate and stabilize them in solution.
• Most chlorides (Cl⁻), bromides (Br⁻), and iodides (I⁻) are soluble, except when paired with Ag⁺, Pb²⁺, or Hg₂²⁺. These exceptions arise because the lattice energies of these salts are high enough to outweigh the hydration energy, preventing dissolution.

General Insoluble Compounds

• Most hydroxides (OH⁻) and oxides (O²⁻) are insoluble, except those containing alkali metals or the larger alkaline earth metals (Ca²⁺, Sr²⁺, Ba²⁺), which have lower lattice energies. Insolubility occurs because the strong ionic bonds in these compounds resist disruption by water molecules.
• Carbonates (CO₃²⁻), phosphates (PO₄³⁻), and sulfides (S²⁻) are generally insoluble unless paired with alkali metals or ammonium. These polyatomic anions form rigid crystal lattices that are difficult for water to break apart, leading to precipitation in most conditions.
• Sulfates (SO₄²⁻) are generally soluble, but exceptions include BaSO₄, PbSO₄, SrSO₄, and CaSO₄. These compounds have particularly strong lattice energies that make dissolution energetically unfavorable.

Applications in Chemical Reactions

• Predicting precipitates in double replacement reactions relies on applying solubility rules. If both potential products are soluble, no visible reaction occurs, but if one is insoluble, a precipitate forms, indicating a successful precipitation reaction.
• Solubility rules also guide laboratory separations, such as isolating specific ions from mixtures via selective precipitation. By carefully controlling pH, temperature, and added reagents, chemists can exploit solubility differences to purify substances or identify ions in qualitative analysis.

Acid–Base Reactions

Definition and Nature of Acid–Base Reactions

• Acid–base reactions involve the transfer of protons (H⁺ ions) between reactants, following the Brønsted–Lowry definition. An acid donates a proton, while a base accepts it, and these processes always occur in pairs—every acid has a conjugate base, and every base has a conjugate acid.
• In aqueous solutions, strong acids (such as HCl, HNO₃, and H₂SO₄) dissociate completely to release H⁺, while strong bases (such as NaOH and KOH) dissociate completely to release OH⁻. Weak acids and bases only partially dissociate, leading to equilibrium systems where both undissociated and dissociated species coexist.

Neutralization and Ionic Representation

• Neutralization occurs when an acid and base react to form water and a salt, such as HCl + NaOH → NaCl + H₂O. In ionic form, this simplifies to H⁺ + OH⁻ → H₂O, showing the essential proton–hydroxide combination that defines acid–base neutralization.
• Strong acid–strong base reactions typically go to completion, while weak acid–strong base or weak base–strong acid reactions require equilibrium analysis. The net ionic equation helps strip away spectator ions, revealing the actual reacting species.

Applications and Significance

• Acid–base reactions are essential in titrations, where a solution of known concentration is used to determine the unknown concentration of an acid or base through stoichiometric relationships. Indicators are chosen based on their pH transition ranges relative to the expected equivalence point.
• These reactions are also crucial in biological systems, industrial processes, and environmental chemistry. For example, the neutralization of stomach acid by antacids or the treatment of acidic industrial waste streams relies on acid–base principles.

Gas-Forming Reactions

Definition and Types

• Gas-forming reactions produce a gaseous product that escapes from the reaction mixture, driving the reaction forward due to the decrease in pressure and increase in entropy. This gas evolution often results from the decomposition of unstable intermediates in aqueous reactions.
• Common examples include the reaction of acids with carbonates or bicarbonates to produce CO₂, acids with sulfides to produce H₂S, and acids with sulfites to produce SO₂. Each gas evolves due to the instability of the intermediate acid form.

Examples of Gas Evolution

• Acid + Carbonate/Bicarbonate: \( \text{CaCO}_3 + 2\text{HCl} \rightarrow \text{CaCl}_2 + \text{H}_2\text{O} + \text{CO}_2 \uparrow \) — carbon dioxide escapes as bubbles, making this a classic laboratory demonstration.
• Acid + Sulfide: \( \text{FeS} + 2\text{HCl} \rightarrow \text{FeCl}_2 + \text{H}_2\text{S} \uparrow \) — hydrogen sulfide gas is produced, known for its distinct rotten egg smell.
• Acid + Sulfite: \( \text{Na}_2\text{SO}_3 + 2\text{HCl} \rightarrow 2\text{NaCl} + \text{H}_2\text{O} + \text{SO}_2 \uparrow \) — sulfur dioxide gas evolves, often with a pungent odor.

Significance and Safety Considerations

• Gas-forming reactions are used in applications ranging from baking (CO₂ from baking soda and acid) to industrial gas generation. In analytical chemistry, gas evolution can confirm the presence of certain ions, such as carbonates or sulfites.
• Some evolved gases, such as SO₂ and H₂S, are toxic and require proper ventilation or fume hood use in laboratories. Safety protocols ensure that chemists can perform these reactions without exposure to harmful concentrations of gas.

Energy Changes in Reactions

Thermochemistry Basics

• Chemical reactions involve changes in energy due to the breaking and forming of chemical bonds. Breaking bonds requires energy input (endothermic), while forming bonds releases energy (exothermic). The net change in energy determines whether the reaction absorbs or releases heat overall.
• Enthalpy change (\( \Delta H \)) is a measure of heat transfer at constant pressure. A negative \( \Delta H \) indicates an exothermic process, while a positive \( \Delta H \) indicates an endothermic process. This energy change reflects the balance between bond energies in reactants and products.

Potential Energy Diagrams

• Potential energy diagrams illustrate the energy of reactants, products, and the transition state. The activation energy (\( E_a \)) is the energy barrier that must be overcome for a reaction to proceed, and it determines the reaction rate.
• Catalysts lower the activation energy without affecting \( \Delta H \), enabling reactions to occur more quickly by providing an alternative pathway for the reaction.

Connection to Reaction Spontaneity

• While enthalpy is important, spontaneity also depends on entropy (\( \Delta S \)) and temperature, as described by the Gibbs free energy equation: \( \Delta G = \Delta H - T\Delta S \). A negative \( \Delta G \) means the reaction is thermodynamically spontaneous under the given conditions.
• Exothermic reactions often, but not always, are spontaneous; endothermic reactions can also be spontaneous if accompanied by a large increase in entropy.

Representations of Reactions

Balanced Chemical Equations

• A balanced chemical equation shows the relative number of moles of each reactant and product, ensuring conservation of mass and atoms. Coefficients are adjusted to balance each atom on both sides without changing subscripts, which would alter the compound identity.
• Equations can be written in molecular form, complete ionic form, or net ionic form, depending on the level of detail needed. Net ionic equations show only species that participate directly in the chemical change, omitting spectator ions.

State Symbols and Energy Information

• State symbols (s, l, g, aq) indicate the physical state of each substance and provide insight into reaction conditions. For example, (aq) denotes species dissolved in water, often important in acid–base and precipitation reactions.
• Energy terms such as \( \Delta H \) values or "heat" above the arrow can be included to show thermodynamic information, while catalysts or reaction conditions (e.g., \( \Delta \), hv, Pt) can be written above the reaction arrow to clarify how the reaction is driven.

Reaction Mechanisms

• In some cases, a reaction is represented by a series of elementary steps in a mechanism, each with its own molecularity and rate law. The sum of these steps equals the overall reaction.
• Mechanistic representations are especially important in organic and biochemical contexts, where multiple intermediates and transition states occur before reaching the final products.

Types of Chemical Reactions

Synthesis (Combination) Reactions

• Two or more reactants combine to form a single product, such as \( 2H_2 + O_2 \rightarrow 2H_2O \). These reactions often release energy, though not always, and are common in material formation and biological processes.
• Synthesis reactions can involve elements or compounds and may require catalysts or specific conditions to proceed efficiently.

Decomposition Reactions

• A single compound breaks down into two or more simpler substances, such as \( 2HgO \rightarrow 2Hg + O_2 \). These reactions usually require energy input in the form of heat, light, or electricity.
• Decomposition is important in processes like thermal breakdown of carbonates or electrolysis of water.

Single Replacement Reactions

• An element replaces a similar element in a compound, such as \( Zn + CuSO_4 \rightarrow ZnSO_4 + Cu \). The reactivity of the replacing element must be greater than that of the displaced element, as determined by the activity series.
• These reactions are often redox in nature and are common in metal corrosion and halogen displacement.

Double Replacement Reactions

• Two ionic compounds exchange ions to form new compounds, such as \( AgNO_3 + NaCl \rightarrow AgCl \downarrow + NaNO_3 \). These reactions typically occur in aqueous solutions and may produce a precipitate, gas, or weak electrolyte.
• They are important in analytical chemistry and in processes like water treatment, where unwanted ions are precipitated out of solution.

Combustion Reactions

• A substance reacts rapidly with oxygen, releasing heat and light. Hydrocarbon combustion produces CO₂ and H₂O, such as \( CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O \). Incomplete combustion produces CO or soot.
• Combustion is the basis of energy production in engines, power plants, and biological respiration, though incomplete combustion can pose environmental and health hazards.

Common Misconceptions in Chemical Reactions

Misunderstanding Conservation Laws

• Many students incorrectly believe that mass or atoms can be lost or created during a chemical reaction. In reality, the Law of Conservation of Mass ensures that the total number of atoms for each element remains constant; only their arrangement changes. Coefficients in balanced equations adjust quantities without altering the identity of substances.
• A related misconception is thinking that electrons, energy, or matter disappear entirely in a reaction. In truth, electrons may be transferred or shared, energy may change form (e.g., chemical to thermal), and matter is simply rearranged into different molecules or ions.

Confusing Coefficients and Subscripts

• Some learners mistakenly change subscripts in a chemical formula when trying to balance an equation, which changes the compound’s identity. For example, turning H₂O into H₂O₂ to balance oxygen atoms is incorrect because it alters the chemical substance.
• Coefficients adjust the quantity of a compound without changing its chemical nature, while subscripts indicate the fixed number of atoms in a molecule as determined by chemical bonding.

Misinterpreting Reaction Types

• Students sometimes classify reactions incorrectly by focusing only on surface appearances, such as assuming all reactions producing a gas are combustion reactions. A careful analysis of reactants, products, and electron transfer is necessary for proper classification.
• For example, not all double replacement reactions form precipitates; some may produce gases or weak electrolytes. Similarly, not all exothermic reactions are combustion reactions.

Overgeneralizing Energy Changes

• It is often assumed that all exothermic reactions are “fast” and all endothermic reactions are “slow.” In reality, reaction rate is determined by activation energy and reaction mechanism, not directly by enthalpy change.
• Endothermic reactions can proceed rapidly if they have low activation energies, and some exothermic reactions can be very slow without a catalyst.

Neglecting States of Matter and Conditions

• Many students overlook the importance of state symbols and reaction conditions in chemical equations, leading to misinterpretations of reaction feasibility. For instance, a reaction may occur in aqueous solution but not in the solid state due to differences in ion mobility.
• Pressure, temperature, and concentration also strongly influence whether a reaction occurs as written, especially in gas-producing and equilibrium processes.