Enthalpy

• A positive value means that heat has been absorbed by the surroundings, so it's an endothermic process.
• • Hess’s Law: the energy change of any chemical reaction is the same regardless of the route taken.
• • Exothermic reaction: gives out heat. More chemical energy present in the reactants than the products.
• • Endothermic reaction: takes in heat. More chemical energy in the products than the reactants.
• • Enthalpy change of reaction, Δf​HΘ :the enthalpy change when the number of moles of the substances in the equation as written react under standard conditions of 298K and 101 kPa.
• Energy cycles: Hess’s Law

  ‘Hess’s Law’ states that the total enthalpy change for a chemical reaction is independent of the route where the reaction takes place, provided the initial and final conditions are the same.
• Standard enthalpies of combustion can be calculated using the following formula:
• The enthalpy change of formation (∆f) is defined as the enthalpy change when one mole of substance is formed from its elements at their reference state.
• Standard enthalpy of combustion = -394kJ/mol
• If you’re asked to calculate the enthalpy of a reaction and are given the enthalpies of formation only, then there’s an easier way of working out the enthalpy change. It’s by using the following equation:

  Enthalpy change = Δf​HΘ(products) - Δf​HΘ(reactants)

  Remember that if you have numbers before substances in an equation, you must include these numbers in your calculation or the enthalpies will not balance.
• There are several different types of enthalpy changes, Whatever the type of enthalpy change, they must all be quoted under standard conditions.

  These standard conditions are:

  • temperature of 298K (25°C)
  • concentration of 1 mol dm-3 for solutions
  • pressure of 1 atmosphere (atm) or 101 kPa for gases.
• The standard state of a substance is its physical state under standard conditions, e.g. carbon dioxide gas, water or lithium chloride solid.
• Standard enthalpy change of formation,Δf​HΘ

  This is the enthalpy change that occurs when one mole of compound is formed from its constituent elements in their standard states.

  Na(s)+21​Cl2​(g)→NaCl(s)Δf​HΘ=–411.15kJmol−1
• Standard enthalpy change of atomisation, Δat​HΘ

  This is the enthalpy change that occurs when one mole of gaseous atoms is formed from the elements in its standard state.

  Na(s)→Na(g)Δat​HΘ=+107.32kJmol−1
• Ionisation energy, ΔIE​HΘ

  1st ionisation energy, ΔIE​HΘ : The energy required to remove one mole of electrons from one mole of atoms in the gaseous state.

  Na(g)→Na+(g)ΔIE​HΘ=+498.3kJmol−1
• Electron affinity, ΔEA​HΘ

  1st electron affinity, ΔEA​HΘ : The enthalpy change that occurs when one mole of gaseous 1- ions are formed from gaseous atoms.

  Cl(g)→Cl−(g)ΔEA​HΘ=−346kJmol−1
• Enthalpy change of solution, Δsol​HΘ

  The standard enthalpy change of solution, Δsol​HΘ, is the enthalpy change when one mole of a compound dissolves in water under standard conditions.

  e.g.

  NaCl(s)→Na+(aq)+Cl−(aq)Δsol​HΘ=+4kJmol−1
• Enthalpy of hydration, Δhyd​HΘ

  The standard enthalpy of hydration, Δhyd​HΘ, is the enthalpy change that occurs when one mole of aqueous ions are formed from their gaseous ions under standard conditions.

  Hydration is an exothermic process:

  Na+(g)→Na+(aq)Δhyd​HΘ=–420kJmol−1

  Cl−(g)→Cl−(aq)Δhyd​HΘ=–363kJmol−1
• Lattice Enthalpy

  When ionic compounds form, they form giant ionic 3D lattices. These are regularly arranged with alternating positive and negative ions, held together by the electrostatic attraction of the oppositely charged ions. When these ions come together to form a lattice, they are forming something very stable and as a result energy is released. This energy is called the lattice enthalpy.
• Lattice enthalpy, Δlatt​HΘ (sometimes seen as ΔLE​HΘ ) is the energy released when one mole of ionic compound is formed from its constituent ions in the gaseous state, under standard conditions.

  In general:

  xAy+(g)+yBx−(g)→Ax​By​(s)

  Specifically:

  Na+(g)+Cl−(g)→NaCl(s)

  Δlatt​HΘ=−786kJmol−1
• The larger the negative value, the more energy is released when the ions come together, thus the ionic bond is stronger and will require more energy to break it.
• Born-Haber cycles

  These are energy cycles with several steps that are used to calculate the energy of formation of ionic compounds from elements in their standard states. The steps include many of the energy transfers
• Route 1 shows the energy of the formation of LiF from its elements in their standard states.

  Route 2 shows a series of steps leading to the formation of LiF.

  
• Step 1: formation of one mole of gaseous atoms – enthalpy of atomisation.

  Energy is required to form one mole of gaseous atoms of lithium from standard states:

  Li(s)→Li(g)

  
• Step 2: formation of lithium cation – ionisation energy

  Energy is required for one mole of lithium atoms to remove one mole of electrons:

  Li(g)→Li+(g)

  
• Step 3: conversion of fluorine molecules to atoms – enthalpy of atomisation

  21​F2​(g)→F(g)

  
• Step 4: formation of fluoride anion – electron affinity

  Energy is usually given out when an atom gains an electron to form an anion. If the anion requires more than one electron in its formation, energy is usually required.

  F(g)→F−(g)

  
• Step 5: formation of LiF solid from its ions – lattice formation

  Energy is released when the ions join to form the solid ionic compound. The more exothermic this process is, the more stable the ionic compound.

  Li+(g)+F−(g)→LiF(s)

  As for the other energy cycles we've studied:ROUTE 1 = ROUTE 2

  
• Stability of ionic compounds

  we can determine the stability of an ionic compound by the value of the lattice enthalpy. We can also identify the stability by the value of the enthalpy of formation, Δf​HΘ :
• • If the value Δf​HΘ of is negative, energy is given out in the formation of the compound from its elements, it means that the ionic compound is more stable than the elements from which it is formed. The more negative the value of Δf​HΘ , the more stable the ionic compound is.
• • If the value of Δf​HΘ is positive, it means that the ionic compound is unstable compared with the elements from which it is formed. It is possible for these compounds to exist but energy is needed to convert the elements into the compound. Those that do have a positive Δf​HΘ tend not to decompose once formed as the process is too slow.
• How to find the lattice enthalpy for sodium chloride:

  Na+(g)+Cl−(g)→NaCl(s)
• Step 1: Place the ionic product on a line at the base of the diagram, and place a line with the elements above. The energy change between the two is Δf​HΘ

  
• Step 2: The ionic solid can also be made from the constituent ions in the gaseous state ( ΔLE​HΘ ). This can be added to the diagram:

  
• Step 3: The elements need to be converted into the gaseous ions. This can be done by adding stages to represent Δat​HΘ of sodium and chlorine, ΔIE​HΘ, Δ

  
• On a Born-Haber cycle, all arrows pointing upwards are endothermic values, all arrows pointing downwards are exothermic values.

  For any Born-Haber cycle:

  Sum of all clockwise enthalpies = Sum of all anti-clockwise enthalpies

  Δf​HΘ= sum of all Δat​HΘ+ sum of all ΔIE​HΘ+ sum of all ΔEA​HΘ+ΔLE​HΘ
• an enthalpy change also occurs when an ionic compound dissolves. This is called the standard enthalpy change of solution, Δsol​HΘ.
• The standard enthalpy change of solution, Δsol​HΘ, is the enthalpy change when one mole of a compound dissolves in water under standard conditions.

  e.g.

  NaCl(s)→Na+(aq)+Cl (aq)Δsol​HΘ=+4kJmol−1
• The breakdown of the ionic lattice requires energy. This is the opposite of the lattice energy, so will be endothermic to the value of the lattice enthalpy of that compound.
• Hydration involves the formation of solvation shells around the gaseous ions, with the δ+ hydrogen side of water surrounding the cation and the δ- oxygen side of water surrounding the anion.