Chapter 7 - Periodicity Flashcards
Who created the modern periodic table?
Dmitri Mendeleev
How were the elements ordered by Mendeleev?
Atomic mass
What did Mendeleev do that no one before him did?
- Put elements into groups with similar properties. Swapping some elements around so they fit trends.
- Left gaps for undiscovered elements so the patterns would fit. He therefore predicted properties of these elements.
What are the elements now ordered in?
Atomic number
What are the vertical columns that elements are arranged in?
What are the horizontal columns that elements are arranged in?
Groups
Periods
What is periodicity?
Repeating trend across a period
How does electron configuration change across a period?
The S and P sub-shells are filled the same way.
S subshell fills with 2 electrons, then p subshell to to 6. - up n=4 shell
Periodic pattern
How does electron configuration change down the group?
- Same number of electrons in outer shell.
- Same number of electrons in each sub-shell.
Give the definition of the first ionisation energy?
Energy required to remove one electron from each atom in one mole of gaseous atoms of an element toform one mole of 1+ ions.
Give the first ionisation half equation of Sodium?
Na(g) -> Na+(g) +e-
Give 3 factors that affect ionisation energy
- Atomic radius - Greater the distance between nucleus and the outer electrons the less the nuclear attraction. Inversely proportional to nuclear attraction.
- Nuclear charge - More protons = greater attraction between nucleus and outer electrons.
- Electron shielding - More electrons shells = more shielding effect, reducing attraction between nucleus and outer electrons.
How many ionisation energies does Magnesium have?
12 - as magnesium has 12 electrons to lose
Define second ionisation energy
Energy required to remove one electron from each ion in one mole of 1+ ions of an element to form one mole of gaseous 2+ ions.
What does large jump in ionisation energy between ionisation energy number 7 and 8 signify in analysis of successive ionisation energy graph?
8th electron must be removed from a different from a different shell, closer to the nucleus with less shielding - and therefore takes more energy to remove.
How does ionisation energy change down the group? And why?
Decreases down the group. Atomic radius increases, due to more electron shells meaning more shielding - so decreases nuclear attraction on outer electrons, even though nuclear charge increases - this is outweighed by shielding. And so less energy is required to remove an electron from its outer shell.
How does the first ionisation energy change across a period? And why?
Increases across a group. This is because the atomic radius decreases due to same or similar shielding in all elements but an increase in nuclear charge across the group and so increased nuclear attraction. And so increased first ionisation energy.
When is the rule of increasing ionisation energies across a period wrong?
- When a new sub-shell starts e.g. when first adding electrons to P sub-shell. e.g. Be to B
- When starting to pair electrons as paired electrons repel and less energy required to remove e.g. N to O
What is a metallic bond?
Strong electrostatic attraction between cations (positive ions and delocalised electrons.
Why do metals conduct electricity?
Delocalised electrons (given out by cations) in sea of delocalised electrons carry charge through the structure.
Why do metals have high boiling and melting points?
High temperatures/ large amounts of energy required to overcome strong metallic bonds holding the atoms together in the giant metallic lattice.
List the properties of a giant covalent molecule
- High melting + boiling point
- Insoluble
- Non-conductors of electricity (except hexagonal layers of carbon e.g. graphene and graphite)
How does the melting point change across a period?
Melting point increases from group 1 to group 4 (giant structures). Then a sharp decrease between group 4 and 5 and then the melting points are low from group 5 to 8 as these are simple molecules.
Why does melting point increase from group 1 to group 4?
Ions have greater charges (greater charge density) and also more electrons - so metallic bond is stronger due to increased attraction between ions and electrons.
