THERMODYNAMIC

Breaking Bonds and Making Attractions:

• First Laws of Thermodynamic 

  • The System: This refers to the specific chemical currently reacting, such as the dry salt in a chemical pack.

  • The Surroundings: This encompasses everything else around the chemical, including the water inside the pouch, the plastic bag, and the climbers' skin using the pack.

• The Microscopic Process of Dissolving: When a solid salt (e.g., a $NaCl$ crystal) dissolves in water, it occurs in two distinct microscopic steps:

  • Step 1: Breaking the Salt Apart (Lattice Energy):
    Solid salts consist of positive ions (e.g., Sodium ion $Na^+$) and negative ions (e.g., Chloride ion $Cl^-$) tightly locked in a 3D grid known as an ionic crystal structure. To dissolve, these strong ionic bonds must be physically torn apart.

    • Energy Requirement: Breaking bonds always REQUIRES energy (effort), defined as Lattice Energy will ABSORB in the system in order to break the bonds.

  • Step 2: Forming Attractions with Water (Hydration Energy):
    Water ($H_2O$) molecules are polar, operating like tiny magnets with a slightly positive end ($\delta^+$) and a slightly negative end ($\delta^-$). Once ions are separated, polar water molecules rush in to surround the floating ions (forming hydrated sodium and chloride ions).

    • Energy Release: Making new connections always RELEASES energy (relaxation), termed Hydration Energy.

Real-Life Example: When you add table salt (sodium chloride) to water to make brine, you observe the cooling sensation as the salt dissolves. This is because breaking apart the salt and hydrating the ions requires more energy than is released when the ions are surrounded by water, hence it absorbs heat from the surroundings, making the solution feel colder.

Potential Energy Diagrams:

• Key Components of the Graph:

  • Reactants: Starting chemicals before the reaction begins, represented by a flat line on the left (e.g., Solid salt + liquid water).

  • Products: Final chemicals after the process is complete, shown as a flat line on the right (e.g., Dissolved floating ions).

  • Activation Energy ($E_a$): The initial energy required to break bonds, measured from the Reactants line to the top of the energy hill.

  • Heat of Reaction ($\Delta H$): Overall net change in energy from start to finish, calculated as $\text{Products} - \text{Reactants}$, representing the difference between Lattice Energy and Hydration Energy.

• Specific Graph Examples:

  • Graph A (Exothermic Reaction):
    Reactants start at $100\,\text{kJ}$, peak energy at $120\,\text{kJ}$, dropping to $50\,\text{kJ}$ for Products. Energy is released.

    • Exothermic reactions release heat, light, or energy to the surroundings, causing temperatures to rise, often feeling hot    

      

  • Graph B (Endothermic Reaction):
    Reactants start at $50\,\text{kJ}$, peak at $150\,\text{kJ}$, dropping to $100\,\text{kJ}$ for Products. Energy is absorbed.

    • Endothermic reactions absorb energy, causing surroundings to cool down, often feeling cold

Real-Life Example: In an exothermic reaction like the combustion of gasoline in an engine, energy is released as the fuel reacts with oxygen, driving the pistons and powering the vehicle. In an endothermic reaction, such as photosynthesis, plants absorb energy from sunlight to convert carbon dioxide and water into glucose and oxygen.

Calorimetry Math:

• The Calorimetry Equation: $q = mc\Delta T$

  • $q$ : Heat energy transferred (J).

  • $m$: Mass of the surroundings, typically water (g).

  • $c$: Specific Heat Capacity of water, $4.18\,\text{J/g}^{\circ}\text{C}$ (AWAYS)

  • $\Delta T$: Change in temperature, calculated as $\text{Final Temp} (T<em>f) - \text{Initial Temp} (T</em>i)$.

  • Solving for Final Temperature: For calorimetry specific problems, the equation is rearranged to find final temperature based on heat gained or lost.

  • Calorimetry with a Different Substance: Adjust calculations based on the specific heat of the substance involved, ensuring correct understanding of $m$ and $c$ values.

• Finding Experimental Molar Heat of Solution and Percent Error:

  • Experimental Molar Heat of Solution ($\Delta H$):
    $\text{Heat per mole} = \frac{q_{\text{salt}}}{\text{moles}}$

  • Percent Error Calculation:
    $\text{Percent Error} = \frac{|\text{Theoretical} - \text{Experimental}|}{|\text{Theoretical}|} \times 100$

Real-Life Example: In a laboratory setting, when dissolving ammonium nitrate in water to create a cold pack, the resultant temperature drop is measured using calorimetry, allowing for determination of the molar heat of solution and validation against theoretical values to check accuracy.