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| Question | Answer |
| Define first ionisation energy | The energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions |
| TRUE OR FALSE: Ionisation is an endothermic process | TRUE You have to put energy in to ionise an atom or molecule, so it's an endothermic process |
| What are the 3 factors affecting ionisation energy? | Nuclear charge Atomic radius Shielding |
| How does nuclear charge affect ionisation energy? | The more protons there are in the nucleus, the more positively charged the nucleus is and the stronger the attraction for the electrons |
| What happens to the ionisation energy when nuclear charge increases? | Ionisation energy increases |
| How does atomic radius affect ionisation energy? | Attraction falls off very rapidly with distance. An electron close to the nucleus will be much more strongly attracted than one further away |
| What happens to ionisation energy when atomic radius increases? | Ionisation energy decreases |
| How does shielding affect ionisation energy? | As the number of shells between the outer electrons and the nucleus increases, the outer electrons feel less attraction towards the nucleus |
| What happens to ionisation energy when shielding increases? | Ionisation energy decreases |
| Describe and explain the trend in ionisation energy down the group | As you go down a group, ionisation energies generally fall (less energy is needed to remove outer electrons). This is because: Shielding increases Atomic radius increases Nuclear charge increases (BUT! Effect is overridden by the effect of the extra shells) |
| Describe and explain the trend in ionisation energy across a period | As you move across a period, the ionisation energies generally increase. This is because: Shielding stays the same Atomic radius decreases Nuclear charge increases HOWEVER! There are small drops between Groups 2 and 3, and 5 and 6 |
| The ionisation energies across a period generally increase. However, there is a small drop between Groups 2 and 3. Why? | The drop between Groups 2 and 3 are due to the outer electrons in Group 3 elements being in a p-orbital rather than an s-orbital. The p-orbital has a slightly higher energy than the s-orbital, so the electron is found (on average) further from the nucleus. The p-orbital also has additional shielding provided by the s-orbital electrons. Both of these factors together are strong enough to override the effect of the increased nuclear charge, so IE drops slightly. |
| The ionisation energies across a period generally increase. However, there is a small drop between Groups 5 and 6. Why? | The drop between Groups 5 and 6 is due to electron repulsion. For example: P: [Ne] 3s² 3p³ S: [Ne] 3s² 3p⁴ The shielding is identical and the electrons are being removed from the 3p sub-shell. In P's case, the electron is being removed from a singly-occupied orbital. But in S, the electron is being removed from an orbital containing two electrons. The repulsion between two electrons in an orbital means that electrons are easier to remove from shared orbitals (full>half>partial) |
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