The Periodic Table The how and why

Atomic Size First problem where do you start measuring. The electron cloud doesn’t have a definite edge. Chemists get around this by measuring more than 1 atom at a time.

Atomic Size } Radius Atomic Radius = half the distance between two nuclei of a diatomic molecule (homo-nuclear molecule).

Trends in Atomic Size Influenced by two factors. Energy Level Higher energy level is further away. The effective charge from the nucleus The greater the nuclear charge reaching the valence electrons the closer these electrons are pulled in.

Group trends H Li Na K Rb As we go down a group Each atom has another energy level, so the atoms get bigger. There are more levels in the kernel and therefore greater shielding of valence electrons (weaker attraction). Li Na K Rb

Periodic Trends Na Mg Al Si P S Cl Ar As you go across a period the radius gets smaller. Valence electrons are in the same energy level. Greater nuclear charge gets to the valence electrons – same shielding e- but greater nuclear charge. Outermost electrons are closer. Row 3 elements all have 10 shielding electrons; +11 (nuclear charges) +18 Na Mg Al Si P S Cl Ar

Rb K Overall Na Li Atomic Radius (nm) Kr Ar Ne H 10 Atomic Number

Ionization Energy The amount of energy required to completely remove an electron from a gaseous atom. Removing one electron makes a +1 ion. The energy required is called the first ionization energy.

Ionization Energy The second ionization energy is the energy required to remove the second electron. Always greater than first IE. The third IE is the energy required to remove a third electron. Greater than 1st of 2nd IE.

Symbol First Second Third 11749 14840 3569 4619 4577 5301 6045 6276 HHeLiBeBCNO F Ne 1312 2731 520 900 800 1086 1402 1314 1681 2080 5247 7297 1757 2430 2352 2857 3391 3375 3963

Symbol First Second Third 11749 14840 3569 4619 4577 5301 6045 6276 HHeLiBeBCNO F Ne 1312 2731 520 900 800 1086 1402 1314 1681 2080 5247 7297 1757 2430 2352 2857 3391 3375 3963

What determines IE The greater the effective nuclear charge the greater IE. Less distance from nucleus increases IE Filled and half filled sublevels have lower energy, so removing them raises the IE. Shielding: effective blocking of nuclear charge weakens the attraction of valence electrons.

Shielding The electron on the outside energy level has to look through all the other energy levels to see the nucleus

Shielding The electron on the outside energy level has to look through all the other energy levels to see the nucleus. A second electron has the same shielding.

Group trends As you go down a group first IE decreases because The electron is further away. More effective shielding.

Periodic trends All the atoms in the same period have the same energy level. Same shielding. Increasing nuclear charge So IE generally increases from left to right. Exceptions at full and 1/2 fill sublevels.

He has a greater IE than H. no shielding greater nuclear charge First Ionization energy Atomic number

further away (outweighs greater nuclear charge) Li has lower IE than H more shielding further away (outweighs greater nuclear charge) H First Ionization energy Li Atomic number

greater nuclear charge removing an electron from a full s sublevel He Be has higher IE than Li same shielding greater nuclear charge removing an electron from a full s sublevel First Ionization energy H Be Li Atomic number

greater nuclear charge He B has lower IE than Be same shielding greater nuclear charge removing an electron from a partially filled p sublevel First Ionization energy H Be B Li Atomic number

First Ionization energy He First Ionization energy H C Be B Li Atomic number

First Ionization energy He N First Ionization energy H C Be B Li Atomic number

First Ionization energy He N First Ionization energy H C O Be Breaks the pattern because in N the electron gets removed from a ½ filled p sublevel B Li Atomic number

First Ionization energy He F N First Ionization energy H C O Be B Li Atomic number

Both have full valence shells Ne has more shielding Greater distance Ne has a lower IE than He Both have full valence shells Ne has more shielding Greater distance F N First Ionization energy H C O Be B Li Atomic number

Na has a lower IE than Li Both are s1 Na has more shielding He Ne Na has a lower IE than Li Both are s1 Na has more shielding Greater distance F N First Ionization energy H C O Be B Li Na Atomic number

First Ionization energy Atomic number

Driving Force Full Energy Levels are very low energy. Noble Gases have full energy levels. Atoms behave in ways to achieve noble gas configuration (become isoelectronic).

2nd Ionization Energy For elements that reach a filled or half filled sublevel by removing 2 electrons 2nd IE is lower than expected. True for s2 Alkaline earth metals form +2 ions.

3rd IE Using the same logic s2p1 atoms have an low 3rd IE. Atoms in the aluminum family form +3 ions. 2nd IE and 3rd IE are always higher than 1st IE!!!

Electron Affinity The energy change associated with adding an electron to a gaseous atom. Easiest to add to group 17 (7A). Gets them to full energy level becomes stable – releases a large amount of energy. Increase from left to right atoms –metals are losers so to gain an electron will only bring about a small increase in stability (if any) – small energy change. Decrease as we go down a group.

Ionic Size Cations form by losing electrons. Cations are smaller than the atom they come from. Metals form cations. Cations of representative elements have noble gas configuration.

Ionic size Anions form by gaining electrons. Anions are bigger than the atom they come from. Nonmetals form anions. Anions of representative elements have noble gas configuration.

Configuration of Ions Ions always have noble gas configuration. Na is 1s22s22p63s1 Forms a +1 ion - 1s22s22p6 Same configuration as neon. Metals form ions with the configuration of the noble gas before them - they lose electrons.

Configuration of Ions Non-metals form ions by gaining electrons to achieve noble gas configuration. They end up with the configuration of the noble gas after them.

Group trends Li+1 Na+1 K+1 Rb+1 Cs+1 Adding energy level Ions get bigger as you go down. Li+1 Na+1 K+1 Rb+1 Cs+1

Periodic Trends N-3 O-2 F-1 B+3 Li+1 C+4 Be+2 Across the period nuclear charge increases so they get smaller. Energy level changes between anions and cations. N-3 O-2 F-1 B+3 Li+1 C+4 Be+2

Size of Isoelectronic ions Iso - same Iso electronic ions have the same # of electrons (same configuration) Al+3 Mg+2 Na+1 (Ne) F-1 O-2 and N-3 all have 10 electrons all have the configuration 1s12s22p6

Size of Isoelectronic ions Positive ions have more protons than electrons so they are smaller (strong attraction for electrons). N-3 O-2 F-1 Ne Na+1 Al+3 Mg+2

Electronegativity

Electronegativity The tendency for an atom to attract electrons to itself when it is chemically combined with another element. How fair it shares. Big electronegativity means it pulls the electron toward it. Atoms with large negative electron affinity have larger electronegativity.

Group Trend The further down a group the farther the electron is away and the more electrons an atom has – the lower the electronegativity value.

Periodic Trend Metals are at the left end. They let their electrons go easily Low electronegativity At the right end are the nonmetals. They want more electrons. Try to take them away. High electronegativity.

Ionization energy, electronegativity Electron affinity INCREASE

Atomic size decreases, shielding constant Atomic size increases Shielding increases Ionic size decreases