Physical vs Chemical Changes

Physical Change

• Definition: Alters the physical state or appearance of a substance without changing its chemical composition.
• No new substance is formed.
• Substances remain chemically the same.
• Easily reversible (e.g., melting, freezing, dissolving).
• Examples: Ice, snow, and water (all H₂O), dissolving sugar or ethanol in water, melting (endothermic, takes in heat), and freezing/condensation (exothermic, gives out heat).

Chemical Change (Chemical Reaction)

• Definition: A change in which a new substance is formed.
• Results from a chemical reaction.
• Often irreversible.
• Energy is often given out or taken in.
• Example: Burning magnesium in oxygen forms magnesium oxide (new substance) and produces a brilliant white flame, it is an exothermic reaction (gives out heat and light).

Key Differences Between Physical and Chemical Changes

• Physical change: alters the physical state or appearance without a chemical change.
• Chemical reaction: a change where a new substance is formed.

Exothermic Change

• Definition: A chemical reaction that releases energy to the surroundings, usually as heat or light.
• Examples: Burning methane in a domestic cooker, fireworks, respiration, and combustion reactions.
• Effects: Temperature of surroundings increases, reactions can produce intense heat and flames.

Endothermic Change

• Definition: A chemical reaction that absorbs energy from the surroundings.
• Examples: Photosynthesis (plants absorb light energy to make food), reaction between nitrogen and oxygen to form nitrogen monoxide (happens during lightning strikes).
• Effects: Temperature of surroundings decreases, less common than exothermic reactions.

Chemical Change - Key Points Recap

• New substances form during chemical reactions.
• Reactions can be hard to reverse.
• Energy is either given out (exothermic) or taken in (endothermic) with chemical reactions.
• Exothermic reactions are more common than endothermic ones.

Identifying Physical vs. Chemical Changes

• Melting of ice is a physical change.
• Burning of magnesium is a chemical change.
• Evaporation of ethanol is a physical change.
• Dissolving of sugar in water is a physical change.

Identifying Exothermic vs. Endothermic Reactions

• Condensation of steam to water is exothermic.
• Burning of magnesium is exothermic.
• Addition of concentrated sulfuric acid to water is exothermic.
• Evaporation of a volatile liquid is endothermic.

Exothermic Process

• Definition: Releases heat energy to the surroundings.
• Indicators: Rise in temperature of the reaction mixture, air around the test tube, the test tube itself, or the thermometer.
• Examples: Dissolving magnesium chloride in water (exothermic physical process); adding zinc to copper(II) sulfate solution (exothermic chemical process).
• Memory tip: "EXothermic = Heat EXits".

Endothermic Process

• Definition: Takes in heat energy from the surroundings.
• Indicators: Drop in temperature of surroundings due to energy being absorbed by reactants.
• Examples: Photosynthesis, cooling reactions like dissolving certain salts.
• Memory tip: "ENdothermic = Heat ENters".

Energy Diagrams: Reaction Pathway Diagrams

• Used to show energy levels in exothermic and endothermic reactions.
• Y-axis represents energy of reactants and products.
• X-axis represents the progress of the reaction (time/pathway).
• A downward arrow indicates an exothermic reaction (energy released), while an upward arrow indicates an endothermic reaction (energy absorbed).

Reaction Pathway Diagram

• Illustrates the energy levels of reactants and products and indicates whether a reaction is exothermic or endothermic.

Exothermic Reactions - Combustion of Methane

• Equation: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g)
• Diagram Notes: Energy of reactants is higher than that of products, arrow points downward indicating energy is given out, heat is released to surroundings.

Endothermic Reactions - Nitrogen + Oxygen

Equation: N₂(g) + O₂(g) → 2NO(g)

• Diagram Notes: Energy of reactants is lower than that of products, arrow points upward indicating energy is taken in. Heat absorbed from surroundings.

Energy Diagrams

• Exothermic Reaction: Reactants have higher energy than products.
• Endothermic Reaction: Products have higher energy than reactants.
• Exothermic Reaction: Arrow direction is downwards.
• Endothermic Reaction: Arrow direction is upwards.
• Exothermic Reaction: Heat flows released to the surroundings.
• Endothermic Reaction: Heat flows absorbed from the surroundings.

Tips for Exam Questions on Energy Diagrams

• Downward arrow: Heat given out = Exothermic
• Upward arrow: Heat taken in = Endothermic

Enthalpy (H)

• Is the total thermal (heat) energy content of a system.

Enthalpy Change (ΔH)

• Represent change in heat during a chemical reaction.
• It is measured in kilojoules per mole (kJ/mol).
• It can be negative or positive:
  • Exothermic reaction: Heat released which means ΔH is negative.
  • Endothermic reaction: Heat absorbed which means ΔH is positive.

Methane Combustion

• Equation: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g)
• Enthalpy change: ΔH = -728 kJ/mol (strongly exothermic)

Bond Breaking and Making

• Bond Breaking: Requires energy input, always an endothermic process (e.g., breaking C–H and O=O bonds before combustion).
• Bond Making: Releases energy, always an exothermic process (e.g., forming C=O and O–H bonds in CO₂ and H₂O).

Why Combustion is Exothermic

• More energy is released from bond making (in products) than is absorbed during bond breaking (in reactants).
• A net energy release results in an exothermic reaction.

Visual Representation of Bond Energy

• Energy goes up during bond breaking and then drops down during bond making.

Key Definitions

• Enthalpy (H) represents the heat content of a system.
• ΔH (Enthalpy change) represents the change in heat during a reaction (negative for exothermic, positive for endothermic).

Bond Energies

• Bond Energy defines the energy needed to break one mole of a covalent bond (in kJ/mol).
• Bond Energy is used to calculate ΔH (enthalpy change) of a reaction using the formula: ΔH=Total energy needed to break bonds−Total energy released when making bonds

Bond Energies - Values

• H–H: 436 kJ/mol (in hydrogen)
• C–H: 435 kJ/mol (average in methane)
• O–H: 464 kJ/mol (in water)
• C–C: 347 kJ/mol (average in many compounds)
• O=O: 498 kJ/mol (in oxygen)
• C=O: 803 kJ/mol (in carbon dioxide)
• N≡N: 945 kJ/mol (in nitrogen)

Example Calculation - Methane Combustion

• Equation: CH4+2O2→CO2+2H2O
• Bonds Broken (Reactants):
  • 4 × C–H = 4 × 435 = 1740 kJ
  • 2 × O=O = 2 × 498 = 996 kJ
  • Total = 2736 kJ
• Bonds Formed (Products):
  • 2 × C=O = 1606 kJ
  • 4 × O–H = 1856 kJ
  • Total = 3462 kJ ΔH=2736−3462=−726 kJ/mol, which is an exothermic reaction.

Mnemonics

• MEXO: Making bonds = EXOthermic
• BENDO: Breaking bonds = ENDOthermic

Activation Energy (Ea)

• Represents minimum energy required to start a chemical reaction.

Activation Energy - Key Points

• Most reactions need extra energy to start, even if exothermic.
• This energy is needed to break bonds in the reactants before new bonds can form in products.
• This initial energy input is called activation energy (Ea).

Burning of Methane

• Reaction: CH4(g)+2O2(g)→CO2(g)+2H2O(g)
• Needs a spark or flame to start, which provides activation energy.

Reaction Pathway Diagrams

• Exothermic Reaction: Ea is required to start the reaction, products have lower energy than reactants, and energy is released (ΔH is negative).
• Endothermic Reaction: Ea is also required, products have higher energy than reactants, and energy is absorbed (ΔH is positive).

Importance of Activation Energy

• If particles don’t have enough energy, they will bounce off each other instead of reacting.
• The reaction only happens if their energy meets or exceeds the activation energy.
• Some reactions, like sodium in water, have a very low Ea, so they happen spontaneously.