Periodic Trends

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

Atomic Size u Atomic Radius = half the distance between two nuclei of a diatomic molecule. } Radius

Trends in Atomic Size Influenced by two factors. 1) Energy Level  Higher energy level is further away. 2) Charge on nucleus (#protons)  More charge pulls electrons in closer.

Group trends u As we go down a group, each atom has another energy level so the atoms get bigger. H Li Na K Rb

Periodic Trends u As you go across a period the radius gets smaller. u Same energy level. u More nuclear charge. u Outermost electrons are closer. NaMgAlSiPSClAr

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

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

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

SymbolFirstSecond Third H He Li Be B C N O F Ne

SymbolFirstSecond Third H He Li Be B C N O F Ne

What determines IE o The greater the nuclear charge the greater IE. o The greater the distance from nucleus, the lower IE o Filled and half-filled orbitals have lower energy, so achieving them is easier, lower IE o Shielding - lowers the IE

Shielding u This outermost electron is shielded from the attractive nuclear forces by the inner electrons u When a new orbital is started, every orbital of lower energy shields these electrons from feeling the full nuclear charge.

Group trends u As you go down a group first IE decreases because the electron is further away (more energy levels) and there is more shielding.

Periodic Trends All the atoms in the same period have the same energy level and same shielding, BUT there is increasing nuclear charge SO IE generally increases from left to right. IE decreases to make full and half filled orbitals.

First Ionization energy Atomic number He u He has a greater IE than H  same shielding BUT greater nuclear charge H

First Ionization energy Atomic number H He l Li has lower IE than H  more shielding & further away which outweighs greater nuclear charge Li

First Ionization energy Atomic number H He l Be has higher IE than Li  same shielding BUT greater nuclear charge Li Be

First Ionization energy Atomic number H He l B has lower IE than Be  same shielding, greater nuclear charge BUT the 2p electron in B is easier to remove than the 2s electron in Be Li Be B

First Ionization energy Atomic number H He Li Be B C

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

First Ionization energy Atomic number H He Li Be B C N O u Oxygen breaks the pattern because removing an electron gets to 1/2 filled p orbital (1s 2 2s 2 2p 3 )

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

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

First Ionization energy Atomic number H He Li Be B C N O F Ne l Na has a lower IE than Li l Both are s 1 l Na has more shielding l Greater distance Na

First Ionization energy Atomic number

Driving Force u Full Energy Levels are very low energy. u Noble Gases have full orbitals. u Atoms behave in ways to achieve noble gas configuration.

2nd Ionization Energy u For elements that reach a filled or half filled orbital by removing 2 electrons, the 2nd IE is lower than expected. u True for s 2 u Alkali earth metals form +2 ions.

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

Electron Affinity The energy released when 1 mole of gaseous atoms each acquire an electron to form 1 mole of gaseous 1 - ions. X(g) + e- → X - (g) + energy  In other words, the neutral atom's likelihood of gaining an electron.

Electron Affinity u Easiest to add to group 7A….gets them to full energy level. u EA -Increase from left to right B/C atoms become smaller, with greater nuclear charge. u EA-Decrease as we go down a group.

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

Configuration of Ions u Na is 1s 2 2s 2 2p 6 3s 1 u Forms a +1 ion - 1s 2 2s 2 2p 6 u Same configuration as neon. u Metals form ions with the configuration of the noble gas before them - they lose electrons.

Ionic size: Anions u Anions form by gaining electrons u Anions are bigger that the atom they come from. u Non-metals form anions u Anions of representative elements have configuration of noble gas after them.

Sizes of Ions: Explanation 1) Cations are smaller than their parent ions. Electrons have been removed from the most spatially extended orbital. The effective nuclear charge has increased. Therefore, the cation is smaller than the parent. 2) Anions are larger than their parent ions. Electrons have been added to the most spatially extended orbital. This means total e - -e - repulsion has increased. The nuclear charge has remained the same, but the number of screening electrons has increased. Therefore, anions are larger than their parents.

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

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

Size of Isoelectronic ions u Isoelectronic ions have the same # of electrons (same electron config) Eg. Al +3 Mg +2 Na +1 Ne F -1 O -2 and N -3 u all have 10 electrons u all have the configuration 1s 2 2s 2 2p 6

Size of Isoelectronic ions u Positive ions have more protons so they are smaller. Al +3 Mg +2 Na +1 Ne F -1 O -2 N -3

Sizes of Ions: Summary For ions of the same charge, ion size increases down a group. All the members of an isoelectronic series have the same number of electrons, BUT as nuclear charge increases in an isoelectronic series the ions become smaller: O 2- > F - > Na + > Mg 2+ > Al 3+

Electronegativity A chemical property that describes the ability of an atom to attract electrons towards itself in a covalent bond.

Electronegativity u The tendency for an atom to attract electrons in a bond to itself u How fairly it shares electrons. u Big electronegativity means it pulls the electron toward it.

Group Trend u The further down a group the farther the electron is away and the more electrons an atom has. u More willing to share. u Low electronegativity.

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

Ionization energy, electronegativity Electron affinity INCREASE

Atomic size increases, Ionic size increases - see p Heath