Fiche de révision : Fundamentals of Chemical Reactions and Energy

📋 Course Outline

1. Chemical Reaction Balancing
2. Redox Half Equations
3. Avogadro’s Constant
4. Moles Calculation
5. Solution Concentration
6. Mass Conservation
7. Endothermic vs Exothermic
8. Energy in Reactions
9. Oxidation and Reduction
10. pH and Ion Types

📖 1. Chemical Reaction Balancing

🔑 Key Concepts & Definitions

• Matter is conserved: The principle that in a chemical reaction, the total number of each type of atom remains the same on both sides of the equation, ensuring mass is neither created nor destroyed.

• Adjust coefficients (numbers in front), not subscripts: When balancing chemical equations, only change the coefficients (the numbers in front of molecules or atoms) to achieve balance; subscripts (the small numbers indicating atom counts within molecules) must remain unchanged.

• Oxygen is diatomic (O₂): Oxygen naturally exists as a diatomic molecule, meaning two oxygen atoms bonded together (O₂). This must be carefully balanced in equations, especially when balancing oxygen atoms on both sides of the reaction.

📝 Essential Points

• To balance a chemical equation, ensure the same number of each atom appears on both sides, respecting the conservation of matter.

• When balancing oxygen, always consider its diatomic form (O₂). For example, in the reaction 2Fe + O₂ → 2FeO, the oxygen atoms are balanced by adjusting the coefficient in front of O₂.

• Do not alter subscripts within chemical formulas; only modify coefficients to balance the equation.

• Example of balanced equation: 2Fe + O₂ → 2FeO (iron reacts with oxygen to form iron oxide, with balanced atoms and molecules).

💡 Key Takeaway

Balancing chemical equations involves adjusting coefficients to ensure the same number of each atom appears on both sides, with special attention to diatomic oxygen (O₂), while maintaining the integrity of chemical formulas by not changing subscripts.

📖 2. Redox Half Equations

🔑 Key Concepts & Definitions

• Oxidation (OIL): The loss of electrons during a chemical reaction, as described by OXIDATION = LOSS OF ELECTRONS (OIL).
• Reduction (RIG): The gain of electrons during a chemical reaction, as described by REDUCTION = GAIN OF ELECTRONS (RIG).
• Oxidation Half Equation: Represents the process where a substance loses electrons; for example, Fe → Fe²⁺ + 2e⁻.
• Reduction Half Equation: Represents the process where a substance gains electrons; for example, O₂ + 4e⁻ → 2O²⁻.
• Iron is oxidised: In redox reactions involving iron, iron undergoes oxidation, losing electrons.
• Oxygen is reduced: In redox reactions involving oxygen, oxygen gains electrons and is reduced.

📝 Essential Points

• Redox reactions involve simultaneous oxidation and reduction processes, which can be represented by their respective half equations.
• The oxidation half equation shows the species losing electrons, while the reduction half equation shows the species gaining electrons.
• The electrons lost in oxidation are equal to those gained in reduction, maintaining charge balance.
• The concepts of oxidation and reduction are fundamental to understanding electron transfer in chemical reactions, especially in processes like corrosion, electrolysis, and metal extraction.
• The half equations are used to balance complex redox reactions and to analyze electron flow.

💡 Key Takeaway

Redox half equations clearly depict the electron transfer process, with oxidation involving electron loss (OIL) and reduction involving electron gain (RIG); iron is oxidised, and oxygen is reduced in these reactions.

📖 3. Avogadro’s Constant

🔑 Key Concepts & Definitions

• Avogadro’s constant (6.02 × 10²³): The number of particles (atoms, molecules, or ions) in one mole of a substance, as defined by Avogadro (early 19th century).
• A mole: A fixed number of particles (6.02 × 10²³), used as a counting unit in chemistry to quantify atoms, molecules, or ions.
• Particles per mole: The total number of individual particles contained in one mole of a substance, directly given by Avogadro’s constant.

📝 Essential Points

• Avogadro’s constant (6.02 × 10²³) provides a bridge between the microscopic world of atoms/molecules and the macroscopic world of measurable quantities like mass and volume.
• A mole is a fundamental unit in chemistry, allowing chemists to count particles indirectly through measurable quantities such as mass or volume.
• The concept of a mole simplifies calculations involving large numbers of particles, making it easier to relate mass, number of particles, and volume in chemical reactions.

💡 Key Takeaway

Avogadro’s constant (6.02 × 10²³) defines the number of particles in one mole, enabling scientists to count and relate microscopic particles to macroscopic measurements efficiently.

📖 4. Moles Calculation

🔑 Key Concepts & Definitions

• Moles: A measure of amount of substance, defined as the number of particles (atoms, molecules, ions) in a sample.
• Molar mass: The mass of one mole of a substance, calculated as the sum of atomic masses of all atoms in a compound.
• Moles = mass / molar mass: The fundamental formula to calculate the number of moles in a given mass of a substance.

📝 Essential Points

• To find the number of moles in a sample, divide the mass of the sample by its molar mass.
• Molar mass is calculated by adding the atomic masses of all elements in the compound, e.g., FeO = 56 + 16 = 72 g/mol.
• Example: For 45 g of FeO, the moles are calculated as 45 / 72 = 0.625 mol.
• This calculation is essential for stoichiometry, enabling conversion between mass and number of particles, which is crucial in chemical reactions and balancing equations.

💡 Key Takeaway

Moles provide a bridge between the mass of a substance and the number of particles it contains, using the relationship: moles = mass / molar mass.

📖 5. Solution Concentration

🔑 Key Concepts & Definitions

Concentration | The amount of a substance (mass) present in a given volume of solution.
Definition: CONCENTRATION = MASS / VOLUME (see example: concentration of NaCl solution = 16 g / 250 cm³ = 0.064 g/cm³)

Mass | The amount of matter in a substance, measured in grams (g).
Definition: The quantity of matter contained in an object or substance, used in calculating concentration.

Volume | The space occupied by a solution, measured in cubic centimeters (cm³) or milliliters (mL).
Definition: The amount of space that a solution takes up, used in the concentration formula.

📝 Essential Points

• Concentration quantifies how much solute is dissolved in a specific volume of solution, directly calculated via mass / volume.
• The units of concentration can vary (e.g., g/cm³, g/mL), but the formula remains consistent.
• Example: To find the concentration of NaCl solution with 16 g in 250 cm³, divide the mass by the volume:
  16 g / 250 cm³ = 0.064 g/cm³.
• Accurate measurement of mass and volume is crucial for precise concentration calculations.
• This concept is fundamental in preparing solutions with desired strengths for chemical reactions and laboratory experiments.

💡 Key Takeaway

Concentration is a measure of how much solute is dissolved in a given volume of solution, calculated by dividing the mass of the solute by the volume of the solution.

📖 6. Mass Conservation

🔑 Key Concepts & Definitions

Mass is neither created nor destroyed in a reaction
(source content): During a chemical reaction, the total mass of the reactants remains the same as the total mass of the products, meaning no mass is lost or gained.

Total mass of reactants equals total mass of products
(source content): The sum of the masses of all reactants before a reaction is equal to the sum of the masses of all products after the reaction, illustrating the principle of conservation of mass.

Atoms are rearranged, not lost
(source content): In a chemical reaction, atoms are simply reorganized to form new substances; no atoms are created or destroyed, only redistributed.

📝 Essential Points

• The principle of mass conservation is fundamental to understanding chemical reactions, exemplified by the balanced equation for iron reacting with oxygen: 2Fe + O₂ → 2FeO.
• When balancing equations, coefficients are adjusted to ensure the same number of each atom appears on both sides, reflecting that atoms are conserved.
• This concept applies universally, regardless of the reaction type, emphasizing that the total mass remains constant throughout the process.
• The idea that atoms are rearranged, not lost, underpins the law of conservation of mass, which was established through careful experimental observations and is a cornerstone of modern chemistry.

💡 Key Takeaway

Mass conservation states that in a chemical reaction, the total mass of reactants equals the total mass of products because atoms are only rearranged, ensuring no atoms are lost or gained during the process.

📖 7. Endothermic vs Exothermic

🔑 Key Concepts & Definitions

• Exothermic: releases energy during a chemical reaction because bond making releases more energy than bond breaking absorbs (source content).
• Endothermic: absorbs energy in a reaction because bond breaking requires more energy than is released during bond making (source content).
• Breaking bonds: requires energy input (energy in) because energy must be supplied to overcome atomic forces (source content).
• Making bonds: releases energy (energy out) as atoms settle into more stable arrangements, releasing excess energy (source content).

📝 Essential Points

• The energy change in a reaction depends on the balance between energy required to break bonds and energy released when bonds form.
• In exothermic reactions, the energy released from making bonds exceeds the energy needed to break bonds, resulting in a net release of energy.
• Conversely, in endothermic reactions, more energy is absorbed to break bonds than is released during bond formation, leading to a net energy absorption.
• The concepts of energy in and energy out are fundamental to understanding whether a reaction is exothermic or endothermic, based on the energy change calculation: energy change = bonds broken - bonds formed.
• These energy dynamics influence reaction profiles and are critical in predicting reaction behavior, such as temperature change and energy flow (source content).

💡 Key Takeaway

Exothermic reactions release energy because bond making releases more energy than is required to break bonds, whereas endothermic reactions absorb energy since breaking bonds requires more energy than is released during bond formation.

📖 8. Energy in Reactions

🔑 Key Concepts & Definitions

• Energy change: The difference between the energy required to break bonds and the energy released when new bonds form during a chemical reaction.
• Bonds broken: The energy input needed to break chemical bonds in the reactants.
• Bonds formed: The energy released when new bonds are created in the products.
• Example reaction: Li + H₂O → LiOH + H₂, illustrating how energy change can be calculated based on bonds broken and formed.
• More energy released than absorbed: Indicates an exothermic reaction, where the overall energy change is negative, releasing heat to the surroundings.

📝 Essential Points

• The energy change in a reaction is calculated as bonds broken - bonds formed.
• In the example of lithium reacting with water, bonds are broken (O–H, Li–Li) and formed (Li–O, H–H).
• If the total energy released from forming new bonds exceeds the energy needed to break the original bonds, the reaction is exothermic.
• Exothermic reactions release heat, making the surroundings warmer, while endothermic reactions absorb heat.
• This concept helps explain why some reactions, such as combustion, are highly exothermic, and others, like melting ice, are endothermic.

💡 Key Takeaway

Energy change in a reaction is determined by subtracting the energy needed to break bonds from the energy released when new bonds form; a greater release indicates an exothermic process.

📖 9. Oxidation and Reduction

🔑 Key Concepts & Definitions

• Oxidation (see section 10): Loss of electrons during a chemical reaction, often associated with an increase in oxidation state.
• Reduction (see section 10): Gain of electrons during a chemical reaction, often associated with a decrease in oxidation state.
• Oxidation (alternative definition): Gain of oxygen atoms in a substance, indicating a chemical change where oxygen is added.
• Reduction (alternative definition): Loss of oxygen atoms from a substance, indicating a chemical change where oxygen is removed.

📝 Essential Points

• Oxidation involves the loss of electrons, which results in an increase in the oxidation state of the element involved.
• Reduction involves the gain of electrons, leading to a decrease in the oxidation state.
• In redox reactions, oxidation and reduction occur simultaneously; one substance is oxidised while another is reduced.
• Oxidation can also be described as gain of oxygen, and reduction as loss of oxygen, which is particularly relevant in processes like combustion and corrosion.
• The concepts of oxidation and reduction are fundamental to understanding electron transfer in chemical reactions, as exemplified by the half equations for iron and oxygen.

💡 Key Takeaway

Oxidation and reduction are complementary processes involving the transfer of electrons and changes in oxygen content, essential for understanding redox reactions in chemistry.

📖 10. pH and Ion Types

🔑 Key Concepts & Definitions

Acidic ions: H⁺

• Ions that characterize acids, responsible for acidity in solutions.
• H⁺ ions are protons, which increase the hydrogen ion concentration, making a solution acidic.

Alkaline ions: OH⁻

• Ions that characterize alkalis, responsible for alkalinity in solutions.
• OH⁻ ions are hydroxide ions, which increase the hydroxide ion concentration, making a solution alkaline.

pH value ranges for acids and alkalis

• The pH scale measures the acidity or alkalinity of a solution, from 0 to 14.
• Acids: pH 0–6 (strong acids: 0–3, weak acids: 4–6)
• Neutral: pH 7
• Alkalis: pH 8–14 (weak alkalis: 8–10, strong alkalis: 11–14)

Methods to find pH: pH probe and indicators

• pH probe: An electronic device providing precise, numerical pH readings.
• Indicators: Substances like universal indicator or pH paper that change color depending on the pH, providing a visual estimate.

📝 Essential Points

• H⁺ ions are the defining feature of acids, increasing hydrogen ion concentration and lowering pH.
• OH⁻ ions are characteristic of alkalis, increasing hydroxide ion concentration and raising pH.
• The pH scale helps determine the strength of acids and alkalis, with lower pH indicating stronger acids and higher pH indicating stronger alkalis.
• pH measurement methods include the use of a pH probe for accuracy and indicators for quick, visual assessment.
• Accurate pH measurement is essential in many chemical processes, including neutralisation and titration.

💡 Key Takeaway

The concentration of H⁺ and OH⁻ ions determines whether a solution is acidic or alkaline, which can be accurately measured using pH probes or indicators to assess the solution’s pH value.

📊 Synthesis Tables

⚠️ Common Pitfalls & Confusions

1. Changing subscripts when balancing equations instead of coefficients.
2. Forgetting that oxygen exists as diatomic O₂, leading to incorrect balancing.
3. Confusing oxidation (loss of electrons) with reduction (gain of electrons); mixing up OIL and RIG.
4. Miscalculating moles by using incorrect molar mass or unit conversions.
5. Using incorrect units for concentration (e.g., mixing g/mL and g/cm³) without proper conversion.
6. Assuming mass is lost or gained during reactions, violating conservation of mass.
7. Overlooking the importance of balancing electrons in redox half equations.
8. Forgetting to include the diatomic nature of oxygen in redox reactions involving O₂.

✅ Exam Checklist

• Know the principle of matter conservation and how to balance chemical equations by adjusting coefficients, not subscripts.
• Understand that oxygen exists as diatomic O₂ and how to balance oxygen atoms accordingly.
• Be able to write and interpret oxidation and reduction half equations, identifying species that are oxidized and reduced.
• Recall Avogadro’s constant (6.02 × 10²³) and its role in relating particles to moles.
• Calculate moles from mass and molar mass, and vice versa, for various substances.
• Determine solution concentration using the formula: concentration = mass / volume.
• Explain the law of conservation of mass, emphasizing that atoms are rearranged, not destroyed or created, during reactions.
• Differentiate between endothermic and exothermic reactions, including energy flow and temperature change.
• Describe how energy is involved in reactions, including activation energy and energy profiles.
• Understand oxidation and reduction processes, including electron transfer and half equations.
• Master pH calculations, including the ion types involved (H⁺, OH⁻), and how to interpret pH values.
• Know key authors and references such as SMITH’s definition of the invisible hand (if applicable), and fundamental principles in chemistry.