Module 3.1 The Periodic Table

Aristotle believed the world was made up of 4 elements (earth, water, air and fire) - similar to solid, liquid and gas - “fire” represented weird things like plasma Materials were extracted (some from ores) - deepening knowledge of how substances behaved

1789 - produced first modern chemical textbook - contained a list of elements “substances that could not be broken down further” He devised a theory about the formation of compounds His list distinguished between metals and non-metals His list included mistakes e.g. compounds, mixtures, light and heat

1828 - published a table of atomic weights He determined the composition by mass of many compounds He introduced letter based symbols for elements (previously signs were used)

Noticed certain groups of 3 elements (“triads”) ordered by atomic weight would have a middle element with the weight and properties that were roughly the average of the other 2 elements

Created a periodic table in order of relative atomic mass 1865 - suggested elements show similar properties to the element 8 places after - “law of Octaves”

1869 - published his periodic table Elements ordered by atomic masses and also periodically Gaps helped predict new elements

1913 - Moseley determined the atomic number of all known elements Moseley modified Mendeleev’s table so the elements were organised by atomic no. He corrected the order in some instances Seaborg discovered the transuranic elements and placed the actinides below the lanthanides

The trend in properties that is repeated across each period

s-block d-block p-block f-block

The energy required to remove one electron from each atom in one mole of the gaseous element to form one mole of gaseous 1+ ions

Atomic radius Nuclear charge Electron shielding

As each electron is removed, there is less repulsion between the remaining electrons and each shell will be drawn slightly closer to the nucleus The positive nuclear charge will outweigh the negative charge each time an electron is removed As the distance of each electron from the nucleus decreases slightly, the nuclear attraction increases. More energy is needed to remove each successive electron

Decrease in atomic radius - bc of increased nuclear charge The number of protons increase - higher attraction on the electrons Electrons are added to the same shell - outer shell drawn inwards slightly Same no. of inner shells - shielding barely changes Attraction between the nucleus and the outer electrons increases - more energy needed to remove an electron

Group 13 elements have their outermost electrons in a p-orbital Group 2 elements have their in an s-orbital p-orbitals have slightly higher energy than s-orbitals - marginally further from the nucleus Electrons in these orbitals are slightly easier to remove = lower I.Es

As you move from group 13 towards group 18, outer electrons are found in p-orbitals In groups 13,14 and 15, each of the p-orbitals contains only a single electron In group 16, however, the outermost electron is now spin-paired in the Px orbital Electrons that are spin paired experience some repulsion - making the first outer electron slightly easier to remove, so the 1st I.E is slightly lower

The no. of shells inc. - distance of the outer shell electrons from the nucleus inc. - weaker force of attraction on the outer electrons More inner shells - inc. shielding - weaker attraction No. of protons inc. - the resulting inc. attraction is far outweighed by the inc. in distance and shielding Attraction dec. as you move down a group - less energy needed to remove an electron

Cations fixed in the lattice Delocalised electrons - can move

Strong electrostatic attraction between the positively charged metal ions and the negative delocalised electrons A high temperature is needed to overcome the metallic bonds

Delocalised electrons can move freely within the lattice This allows metals to conduct electricity, even when in the solid state

Delocalised electrons can move - the metallic structure has a degree of “give” - allows atoms or layers to slide past each other

Group 1 to 14: Melting points increase steadily Giant structures For each successive group: If metallic lattice - nuclear charge increases, no. of electrons in outer shell inc., stronger attraction If giant covalent lattice - each successive group has more electrons with which to form covalent bonds Group 14-15: Sharp dec. in melting point Simple molecular substances Relatively weak intermolecular forces Group 15-18: Relatively low melting points Simple molecular structures

2D giant lattice One carbon atom thick Interlocking hexagonal carbon rings V strong V light Can conduct electricity Many uses in the field of nanotechnology

Reasonably high melting points and boiling points Low densities Form white compounds

Reactive metals Strong reducing agents Group 2 elements are oxidised in their reactions Form 2+ ions