HS Chemistry - Heat in Reactions

Unit Summary

Enthalpy

• Atoms are less reactive when they're bonded together in a molecule than when they're separate, as their valence shells are full when they're bonded together. They're more energetically stable, and therefore have less energy, when bonded together in a molecule.

• The energies of all the separate parts of the system combined will give the system's enthalpy.

• When a system undergoes a reaction, its enthalpy changes, as energy is either released to or absorbed from the surroundings.

• Enthalpy change is shown with a ∆H value.

  • A positive ∆H value shows that the enthalpy of the system has increased in the reaction.
  • A negative ∆H value shows that the enthalpy of the system has decreased in the reaction.

• The enthalpy change of a reaction can be shown using a potential energy diagram.

• One way to calculate enthalpy change $∆H$ is:

  • New Enthalpy - Old Enthalpy
  • This is the same as Enthalpy of Products - Enthalpy of Reactants

Exothermic & Endothermic Reactions

• In an exothermic reaction, there is a net decrease in the enthalpy of the system. In an endothermic reaction, there is a net increase in the enthalpy of the system.

• When atoms bond together and form a molecule, they release heat, and the enthalpy of the system decreases. When a molecule breaks apart into separate atoms, the atoms absorb heat in order to break apart, and as a result the enthalpy of the system increases.

• Most reactions involve the atoms of a molecule being broken apart and then joined together in different ways to form different molecules $for example, 2H and 2O might form H~2~O~2~ instead of H~2~ \+ O~2~$.

  • When this happens, it's unlikely that the heat absorbed in order to break apart the molecules is equivalent to the heat released in forming new molecules. There is thus a net enthalpy change.
  • If the net enthalpy change is positive, the reaction is exothermic. If the net enthalpy change is negative, the reaction is endothermic.

• In order for a reaction to occur, the reactant molecules first need to be broken up into individual atoms. The energy required for this to happen is called the activation energy.

  • If this activation energy is not provided, then the reaction won't occur, even if the products have a lower energy level than the reactants. Such a system is kinetically stable, but energetically unstable.
  • The point in the reaction where the atoms are all separated is called the activated complex. At this moment, the system's enthalpy is highest.

• Once the atoms are separated in the activated complex, they bond together in ways that achieve the lowest energy state.

  • So then how are endothermic reactions possible? They happen because of something called entropy, but that won't be covered here. For now, just remember that they happen sometimes.
  • In an exothermic reaction, the energy released after one set of molecules form back together is more than the activation energy, so part of the released heat acts as the activation energy for the next set of molecules.
  • In an endothermic reaction, it's not enough, and therefore endothermic reactions need a continuous supply of heat to continue (until the supply of one of the reactants is exhausted).

Hess's Law

• The enthalpy change of a reaction is always the same for that reaction, even if a different set of molecules/atoms/ions is used $given that they are the same type of molecules/atoms/ions. That is, it doesn't matter which Hydrogen atom or which Fluorine atom is used in H + F → HF, as long as a Hydrogen atom and a Fluorine atom are used as the reactants$.

• When a reaction is reversed, its enthalpy change is also reversed $the enthalpy change becomes negative if it was previously positive, and positive if it was previously negative, but the number doesn't change$.

• Hess's Law allows for reactions to be multiplied and added/subtracted, provided that the same manipulation is applied to both the reactants and the products and the enthalpy changes of the reactions. This is particularly helpful in finding enthalpy changes of reactions.

Calorimetry

• Q = mC∆T can be used when thermal energy is added to a substance that remains in a certain state of matter.

• Q = mL can be used when thermal energy is added to a substance, but instead of the substance's temperature increasing, its state of matter changes.

  • If the substance is going from a lower-energy state to a higher-energy state, then thermal energy is being added to the substance, and L is positive.
  • If the substance is going from a higher-energy state to a lower-energy state, then thermal energy is being removed from the substance, and L is negative.

• When using a calorimeter, at the end of the experiment, the thermal energies of the two substances are equal. Q_s1 = Q_s2, and this can be broken down further depending on the experiment that was done:

  • If the experiment doesn't include phase changes $like placing a substance in water$, then it can be expanded to:

    m_s1 × C_s1 × ∆T_s1 = m_s2 × C_s2 × ∆T_s2

    Which further expands out to:

    m_s1 × C_s1 × (T_s1f − T_s1i) = m_s2 × C_s2 × (T_s2f − T_s2i)

  • If one of the substances changes phases once in the experiment, then the equation can be expanded to:

    m_s1 × C_s1 × ∆T_s1 = m_s2 × C_s2 × ∆T_s2 + m_s2 × L_s2

    Which further expands out to:

    m_s1 × C_s1 × (T_s1f − T_s1i) = m_s2 × C_s2 × (T_s2f − T_s2i) + m_s2 × L_s2