Ch. 9 - Periodicity and Ionic Bonding

periodicity

the tendency to recur at intervals

electrostatic

pertaining to the forces of attraction or repulsion in charged species

effective nuclear charge (Z_eff)

the net positive charge from the nucleus that a valence electrons experiences.

effective nuclear charge trend

• Z_eff increases across a row 

  • Z (proton #) increases by one for each element, and S (shielding e-) increase by less than 1.

  • same # of shielding e- across a period but the number of protons increases.

• Z_eff slightly decreases going down a group

  • even though Z and S both increase, the addition of a new shell of core e- causes S to increase at a slightly higher rate than Z

  • addition of a new shell increases shielding from inner e- 

atomic radii decreases from left to right across a period

• the nuclei of atoms become more strongly positive, causing Z_eff to increase

• valence e- experiencing higher Z_eff are pulled more tightly to the nucleus, making the atomic radii smaller.

atomic radii increases from down a group on the periodic table

• each successive element down a group has an additional electron shell, increasing the distance between the nucleus and the valance e- 

cation radii

• removing an e- decreases the total negative charge of an atom while retaining the positive charge of the nucleus.

• this increases the attraction between the valence e- and the nucleus

• all cations have smaller radii than their corresponding neutral (parents) atoms

anion radii

• adding an e- to the valence shell increases the negative charge of an atom and retains the positive charge of the nucleus

• this decreases the attraction between the valence e- and the nucleus and increases the repulsions between valance e-

• all anions have larger radii than their corresponding neutral atoms 

ionization energy (IE)

the energy required to remove an e- from a gaseous atom to produce a gaseous cation

IE increases across a period

• as atoms get smaller and the nuclear charge increases (more protons)m the nucleus pulls the valance e- closer making them harder to remove (requiring more energy) 

IE decreases down a group

• as atoms get larger (smaller Z_eff), the valence e- are further from the nucleus and are less attracted to the nucleus, making them easier to remove (requiring less energy). 

exceptions to IE trend: nitrogen to oxygen

• half-filled sub shells are stable

• you would expect O has a higher IE than N, however O < N

• it takes less energy to remove one pair of electrons

exceptions to IE trend: beryllium to boron

• each time a new subshell begins, the IE of that element is lower (easier to ionize) than the previous one

• you would expect that B has a higher IE than Be, but B < Be

• the 2s subshell is lower in energy than the 2p subshell, making it harder to remove an electron

electron affinity (EA) trend

the energy change that occurs when an e- is added to the gaseous anion

lattice energy

the energy released when gas lands ions combine to form a solid ionic compound

born-haber cycle

a thermodynamic cycle that uses hess’s law, ionization energy, electrons affinity, and the energy of their processes to calculate the lattice energy of an ionic compound

negative EA

• represent a decrease in energy when an e- is added to a gas-phase atom (exothermic)

positive EA

• reflect endothermic process: energy is required to add an electron to the gaseous atom. noble gases and some transition metals

EA trend

• become more negative from left to right across a period