Module 3 (chapter 7 and 8) - periodicity, group 2 and 17

• arranged the 60 elements known at the time into the ‘periodic table’ and left gaps for undiscovered elements. (e.g. predicted ekasilicon) • arranged them in columns of similar chemical properties • if an element appeared to be in the wrong place due to its atomic weight he moved it to where it fitted with the pattern he had discovered (arranged by mass number) • discovered the Nobel gases in 1890s which fitted in his final group

• arranged the periodic table by atomic number • using an X-ray gun fired at elements he measured the wavelength of X-rays given, using it to calculate the frequency • when square rooted and plotted against atomic number, the graph showed a perfect straight line

• the vertical columns - each element in a group has the same number of outer shell electrons and therefore similar chemical properties

• horizontal rows • the number of the highest energy electron shells in an element’s atom (e.g. period 2, the 2s sub shell fills with two electrons followed by the 2p sub shell with six electrons) • for each period, the S- and P-s sub shells are filled in the same ways

• the elements in the periodic table can be divided into blocks corresponding to their highest energy sub shell • this gives four distinct blocks (s, p, d, f)

measures how easy an atom loses an electron to form positive ions

the energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions -(e.g. Na (g) –> Na+ (g) + e-)

• atomic radius = the greater the distance between the nucleus and the outer electrons, the less the nuclear attraction. the force of attraction falls of sharply with increasing distance so atomic radius has a large effect • nuclear charge = the more protons there are in the nucleus of an atom, the larger the attractive force on the outer electrons • electron shielding = electrons are negatively charged and so inner shell electrons repel outer shell electrons. this repulsion is called the shielding effect, reducing the attraction between the nucleus and outer shell electrons

• an element has many ionisation energies as there are electrons • second ionisation energy is the energy required to remove one electron from each ion in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions

• evidence for the different electron energy level in an atom (shells) as by analysing jumps scientists can deduce: • they allow predictions to be made about the number of electrons in the outer shell • the groups of the element in the periodic table • the identity of an element • there is a general increase in ionisation energy across each period but a sharp decrease between the end of one period and the start of another.

• as each electron is removed there is less repulsion between the remaining electrons and so each shell will be drawn in slightly closer to the nucleus • the positive nuclear charge will outweigh the negative charge every time an electron is removed • as the distance of each electron from the nucleus decreases slightly, the nuclear attraction increases and so more energy is needed to remove an electron.

• atomic radius increases • more inner shells so shielding increases • nuclear attraction on outer shell electrons decreases • first ionisation energy decreases

• nuclear charge increases due to more protons • same shell so similar levels of shielding • nuclear attraction increases • atomic radius decreases • first ionisation energy increases

• indicate a new sub shell which is slightly further away form the nucleus (e.g. boron less than beryllium) • the even smaller drops reflect that despite an increased nuclear charge there are two electrons in the same orbital compared to an unpaired electron. these two repel each other rathe than spinning at right angles where there is equal repulsion in paired, making it easier to remove (e.g. oxygen lower than nitrogen)

plasma displays where gas is electrically ionised to form a mix of positive ions.

• increases across a group as there are more shells so shielding increases. despite the nucleus getting mote protons the shielding negates this increased nuclear charge • decreases across a period due to more protons and therefore a higher nuclear charge. there is the same number of electron shells and therefor shielding so there is a greater pull of electrons to the nucleus

• top of group 13 to bottom of group 17 - elements near this divide (e.g. Si) are often called semi metals or metalloids

• strong electrostatic force of attraction between positive metal ions and delocalised electrons • regularly arranged rows of metal cations held together by a sea of delocalised electrons • the delocalised electrons are there because each metal atom donates it outer shell electrons to a shared pool of electrons which are delocalised throughout the whole structure • for 12 cation each with 1+ charge there are 12 electrons each with a 1- charge. • for metals containing 2+ cations, twice as many electrons are present to balance the charge

• in solid and liquid states • when a potential differences applied, electrons are delocalised and can therefor remove through the whole structure, carrying charge

• strong electrostatic forces of attraction between cations and electrons • these require a large amount of energy to overcome

• it is expected that there would be some interaction between polar solvents and the charges in the metallic lattice • but any interactions would lead to a reaction rather than dissolving • giant lattice structures have which are far too strong

• ductile - malleable as electrons can move and be bent into shapes without shattering the structure and breaking bonds

• more delocalised electrons so stronger electrostatic force of attraction between oppositely charged ions • there is also a greater charge density (ionic charge is greater and ionic radius decreases) so electrons are more attracted to the nucleus

a repeating pattern in either chemical or physical properties across different periods