Lattice enthalpy Flashcards
What is bond enthalpy?
the enthalpy change required to break 1 mole of bonds between 2 atoms in the gaseous state
What is the first ionisation energy of an element?
the energy needed to remove 1 electron from every atom in mole of gaseous atoms of an element
What is the second ionisation energy of an element?
the energy needed to remove 1 electron from every ion in 1 mole of gaseous +1 ions of an element
Are all ionisations endothermic or exothermic? and why?
endothermic = energy is needed to overcome the attraction between the outer electrons and the nucleus
Why is the second ionisation energy more endothermic than the first?
there is less repulsion between the remaining electrons as 1 has already been removed, allowing the remaining electrons to be more attracted to the nucleus
What is the enthalpy change of atomisation of an element?
the enthalpy change when 1 mole of gaseous atoms of an element are formed from the element in its standard state
Why does the DHatomisation increase as metallic bonding increases?
Na –> Mg –> Al
the number of outer shell electrons increases so there are more electrons contributing to metallic bonding and stronger attracttion b/w e- and metal cation so more energy is needed
What is the equation for the DHatomisation when the element is already a gas in its standard state? use Cl2 as an example
1/2 CL2 (g) —> Cl (g)
produces 1 mol of gaseous Cl atoms
Which group in the periodic table is the only group whose DHatomisation is not positive?
group 0 noble gases = they all exist as individual gas atoms in their standard state
What is the first electron affinity of an element?
the enthalpy change when one electron is added to each atom in a mole of gaseous atoms of an element
Are first electron affinities endothermic or exothermic and why?
exothermic = the electron that is added is attracted to the nucleus more strongly than it is repelled by the other electrons
Is the second electron affinity endothermic or exothermic and why?
endothermic = energy must be put in to overcome the repulsion between the negatively charged X- ion and the e- that is added
What is the lattice enthalpy of an ionic compound?
the enthalpy change when 1 mole of a solid ionic compound is formed from its gaseous ions
What is lattice enthalpy a measure of?
a measure of the strength of ionic bonding in a compound
Are lattice enthalpies exothermic or endothermic and why?
exothermic = when ions come together to form ionic bonds, heat is always released
How can you tell how strong the ionic bonding is from its lattice enthalpy?
the more exothermic (more negative) the stronger the ionic bonding
What are the 2 factors affecting lattice enthalpy?
the size of the ions
the charge on the ions
How does the size of the ion affect the lattice enthalpy?
the smaller the ionic radii of the ions, the greater their charge density, hence the stronger the attraction b/w them and the more exothermic the lattice enthalpy
How does the charge on the ions affect the lattice enthalpy?
the greater the charge on the ions, the greater their charge density, so the stronger the attraction b/w them and the more exothermic the lattice enthalpy
What is the trend of lattice enthalpy down a group?
charge on the ion stays the same, but the ionic radius increases
decreases charge density so the attraction to oppositley charged ions becomes weaker and lattice enthalpies become less exothermic
What is the trend of lattice enthalpy across a period?
the charge on the ion increases and the radius of the ion decreases
theses effects work together to increase the attraction to other ions so lattice enthalpies become more exothermic
Why do lattice enthalpies and melting points roughly correlate?
more exothermic lattice enthalpy, typically the higher the MP
both depend on how much energy is needed to disrupt the ionic lattie
How are lattice enthalpies measured?
using born-haber cycles, they cannot be measure directly
What do you need to do when dealing with a 2+ ion when calculating DHionisation energy?
multiply the ionisation energy by 2 or use the 1st and 2nd ionisation energies
