TRENDS IN THE PERIODIC TABLE

ATOMIC SIZE - Estimated by measuring the distance between 2 neighboring atoms

FACTORS AFFECTING ATOMIC SIZE PRINCIPAL QUANTUM NUMBER (n) OF THE VALENCE ELECTRON – the farther the valence electron is from the nucleus, the greater is the atomic size. EFFECTIVE NUCLEAR CHARGE (Zeff ) – charge “felt” by the valence electron due to the attractive force of the protons (Z) and the shielding effect (S) by the core electrons Zeff = Z - S

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

Periodic Trends As you go across a period, the radius gets smaller. Electrons are in same energy level. More nuclear charge. Outermost electrons are closer. Na Mg Al Si P S Cl Ar

Periodic Trends As you go across a period, the radius gets smaller. Electrons are in same energy level. More nuclear charge. Outermost electrons are closer. Na Mg Al Si P S Cl Ar

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

+1-0=+1 +2-0=+2 +4 -2=+2 +3 -2=+1

EFFECT OF CHARGE ON SIZE OF IONS

IONIZATION ENERGY (IE) - Amount of energy needed to remove an electron from an atom.

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

Symbol First Second Third HHeLiBeBCNO F Ne 1312 2731 520 900 800 1086 1402 1314 1681 2080 5247 7297 1757 2430 2352 2857 3391 3375 3963 11810 14840 3569 4619 4577 5301 6045 6276

Symbol First Second Third 11810 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 nuclear charge, the greater IE. Greater distance from nucleus decreases IE Filled and half-filled orbitals have lower energy, so achieving them is easier, lower IE. Shielding effect

Shielding The electron in the outermost energy level experiences more inter-electron repulsion (shielding). Second electron has same shielding, if it is in the same period

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

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

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

Outer electron further away outweighs greater nuclear charge Li has lower IE than H Outer electron further away outweighs greater nuclear charge H First Ionization energy Li Atomic number

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

greater nuclear charge He B has lower IE than Be same shielding greater nuclear charge p orbital is slightly more diffuse and its electron easier to remove 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 Breaks the pattern, because the outer electron is paired in a p orbital and experiences inter-electron repulsion. N First Ionization energy H C O Be B Li Atomic number

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

Ne has a lower IE than He Both are full, 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 require lots of energy to remove their electrons. Noble Gases have full orbitals. Atoms behave in ways to achieve noble gas configuration.

ELECTRON AFFINITY (EA) – amount of energy needed to add electrons to an atom

Trends in Electron Affinity The energy change associated with adding an electron (mole of electrons) to a (mole of) gaseous atom(s). Easiest to add to group 7A. Gets them to full energy level. Increase from left to right: atoms become smaller, with greater nuclear charge. Decrease as we go down a group.

ELECTRONEGATIVITY – tendency of an atom to attract electrons; guided by the octet rule.