Module 3 - Periodicity Flashcards
how is the periodic table arranged
-by atomic (proton) number
define periodicity
-the regular repeating patterns in the physical and chemical properties of elements
define ionisation
-when an atom loses an electron from its outer shell
define first ionisation energy
-the energy required to remove 1 mole of electrons from one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions
describe how to read successive ionisation energy graphs
-where the first big jump is shows the number of electrons in the outer shell
-this determines the group
describe how to draw a successive ionisation graph for an element
-ionisation energy on y axis
-number of electrons on x axis
-ionisation energy slowly increases for electrons in the same shell
-big jump when moving into the next shell
name the factors affecting ionisation energy
-shielding
-nuclear charge
-atomic radius
define shielding
-the effect of the inner electrons protecting the outer electrons from the effect of the nuclear charge
describe the trend in shielding across periods (eg period 3)
-the shielding stays the same
describe the trend in nuclear charge across periods (eg period 3)
-nuclear charge increases
define nuclear charge
-the positive charge on the nucleus
define atomic radius
-the distance between two nuclei of atoms, then halving the distance (to measure the radius of an atom)
describe the trend in atomic radius across a period (eg period 3)
-decreases
explain the trend in atomic radius across a period (eg period 3)
-nuclear charge increases
-shielding stays the same
-nuclear attraction increases
-so atomic radius decreases
describe the trend in first ionisation energy across a period (eg period 3)
-increases
explain the trend in first ionisation energy across a period (eg period 3)
-nuclear charge increases
-atomic radius decreases
-nuclear attraction increases
-shielding stays the same
-it takes more energy to remove the first electron
equation for ionisation energy (eg magnesium)
1st : Mg(g) -> Mg+(g) + e-
2nd : Mg+(g) -> Mg2+(g) + e-
explain why the first ionisation energy goes down between magnesium and aluminium
-aluminium has one electron in a higher sub shell (3p)
-that one electron is removed more easily as it is further away from the nucleus
-therefore the first ionisation energy is lower than magnesium
explain why the first ionisation energy goes down between phosphorous and sulphur
-sulphur has one 3p orbital that contains a pair of electrons
-these paired electrons repel each other, so one of these electrons is easier to remove
-therefore sulphur has a lower first ionisation energy than phosphorous
describe the trend in shielding down a group
-increases
describe the trend in nuclear charge down a group
-increases
describe the trend in atomic radius down a group
-increases
explain the trend in atomic radius down a group
-shielding increases
-nuclear charge increases but it is cancelled out by the extra shielding
-so nuclear attraction decreases
-so atomic radius increases
describe the trend in first ionisation energy down a group
-decreases
explain why the first ionisation energy decreases down a group
-more shielding
-nuclear charge increases
-nuclear attraction of electrons to the nucleus decreases (as the increased nuclear attraction is cancelled out by the increased shielding)
-atomic radius increases
-so less energy is needed to remove the outer electron
describe the trend in melting point in period 3
-increases from Na (group 1) to Si (Group 4)
-then it decreases
describe the types of bonding across period 3
- metallic bonding - Na, Mg, Al
- giant covalent - Si
- simple covalent - P, S, Cl, Ar
explain why the melting and boiling points rise across the three metals in period 3
-the charge on the metal ions increases from +1 to +3
-so the number of delocalised electrons increases
-so the strength of the metallic bonding increases
-so the melting/boiling points increase as more energy is needed to break stronger metallic bonds
explain why silicon has the high melting/boiling point in period 3
-silicon is a metalloid
-it has a giant covalent structure exactly the same as carbon in diamond (each silicon is bonded to four other silicons)
-silicon has a high melting point
-you have to break strong covalent bonds to melt it
-so a lot of energy is needed to break
what does phosphorous exist as
P4
what does sulphur exist as
S8
what does chlorine exist as
Cl2
what does argon exist as
Ar atoms
explain why the simple molecular elements in period 3 have low melting points
-when these 4 substances melt/boil, the London force are broken
-these are very weak bonds so little energy is needed to overcome them
-the strength of the London forces increases with the number of electrons
describe the order of melting point decrease in simple molecular elements in period 3
S8 > P4 > Cl2 > Ar
define metallic bonding
-the electrostatic force of attraction between the positive metal ions and the delocalised electrons
name the factors that affect the strength of metallic bonding
- number of protons/ strength of nuclear attraction
-the more protons the stronger the bond - number of delocalised electrons per atom (the outer shell electrons are delocalised)
-the more delocalised electrons the stronger the bond - size of ion
-the smaller the ion, the stronger the bond
structure of giant metallic structures
-giant metallic lattice
-lattice of positive metal ions
-electrons
boiling and melting points of giant metallic structures
-high: strong electrostatic forces between positive
ions and sea of delocalised electrons
solubility in water of giant metallic structures
-insoluble
conductivity when solid of giant metallic structures
-good: delocalised electrons can move through
structure
conductivity when molten of giant metallic structures
-good
general description of giant metallic structures
-shiny metal
-malleable as the positive ions in the lattice are all identical
-so the planes of ions can slide easily over one another
-attractive forces in the lattice are the same whichever ions are adjacent
