periodicity Flashcards
(22 cards)
how are elements ordered in periodic table?
order of atomic number
define periodicity
a repeating trend in properties of the elements acorss each period of the periodic table.
what do elements in a group have in common and why?
each element in a group has similar chemical properties because they have same number of outer shell electrons
Describe and explain the periodic trend in electron configuration across periods 2 and 3
each period starts with an electron in a new highest energy shell.
Across period 2, the 2s sub shell fills with 2 electrons, followed by 2p sub shell with six electrons
Across period 3, the same pattern of filling is repeated for 3s and 3p sub shells
What is the link between the block an element is in and its electron configuration
The highest energy occupied orbital indicates which block the element is in. i.e. if the electron configuration is 1s,2s2,2p3, then the highest energy occupied orbital is the 2p orbital and the element is in the p block
What is the definition of first ionisation energy? Give an equation for first ionisation energy for sodium.
-the energy required to remove 1 mole of electrons from 1 mole of gaseous atoms
Na(g) > Na+ (g) + e-
Describe and explain the overall trend in first ionisation energies across Periods 2 and 3, using the terms of attraction, nuclear charge and atomic radius?
Ionisation energy increases across a period. This is because the nuclear charge increases, electrons are in the same shell so there is a similar amount of shielding, nuclear attraction (the attraction of the nucleus for the outermost electrons) increases, the atomic radius decreases and so more energy is required to remove the outermost electron
Describe and explain the anomalies in first ionisation energy at group 3 and group 6
-there is a fall in first ionisation energy from group 2 to group 3. this is because the outermost electron in group 3 is in a p sub shell which is higher in energy and further from the nucleus than the s sub shell.
-there is a fall in first ionisation energy from group 5 to group 6 because the electrons begin to pair in the p sub shell in group 6. there is repulsion between the paired electrons and this makes the 6th electron slightly easier to remove
Describe and explain the trend in first ionisation energies down a group, using the terms attraction, nuclear charge and atomic radius?
Ionisation energy decreases down a group. the nuclear charge increases, however nuclear attraction (the attraction of the nucleus for the outermost electrons) decreases because the atomic radius increases and so less energy is required to remove the outermost electron.
Describe how you could predict from successive ionisation energies of the number of electrons in each shell of an atom and the group of an element?
look for large jumps in ionisation energy. this marks the change from one shell to another.
define metallic bonding
the electrostatic attraction of positive ions for delocalised electrons
What is meant by a solid giant covalent lattice?
atoms are held together by a network of strong covalent bonds in a giant lattice structure.
which elements form solid covalent lattices? which compounds?
Boron, carbon and silicon form giant covalent lattices.
Examples are diamond, graphite, graphene, silicon, silicon dioxide.
features of metals
-high melting points
-not soluble
-conduct electricity when solid or liquid
why do metals have high melting points?
there is a strong electrostatic attraction between positive metal ions and delocalised electrons which requires lots of energy to overcome.
why aren’t metals soluble?
any interactions between the metal ions and water would lead to a reaction rather than dissolving.
why can metals conduct electricity
electrons are mobile.
features of giant covalent elements and compounds
-have high melting points
-not soluble
-cant conduct electricity
why do giant covalent elements and compounds have high melting points?
giant covalent lattices are held together by many strong covalent bonds, these require a lot of energy to overcome.
why are giant covalent elements not soluble?
covalent bonds holding the atoms together are far too strong to be broken by interaction with solvents
why can’t giant covalent elements and compound conduct electricity
no mobile electrons all used in bonding.
describe and explain the variation in melting points across period 2 and 3 in terms of structure and bonding
-going across periods 2 and 3, metallic points increase from the group 1elements to the group 3 elements
-metallic bonding increases in strength as the group number increases
-this is because the cation has a bigger positive charge and there are more delocalised electrons per atom
-the metallic bond inc in strength and requires more energy to overcome
-the elements in group 4 have giant covalent structures which require even more energy to overcome and so these elements have higher melting points.
-the melting points of the remaining elements are much lower as they are simple covalent molecules with weak London forces between molecules which require little energy to overcome.
-the bigger the molecule the higher the melting point.
