Presentation on theme: "The Periodic Law Electron Configuration and Periodic Properties."— Presentation transcript:
The Periodic Law Electron Configuration and Periodic Properties
Objectives: 1.Define atomic radius and ionization energy. 2.Compare the periodic trends of atomic radii and ionization energy.
Atomic Radius Atomic Radius – distance from the nucleus to the outermost electron. can be measured as one- half the distance between nuclei of identical atoms bonded together
What are the trends on the periodic table? 1.Atomic radius increases going down a group on the P.T. 2.Atomic radius decreases moving left to right across a period.
Atomic Radius How can these trends be explained? Group Trend 1.Outer electrons occupy higher energy levels; electrons are farther from the nucleus Period Trend 2.Increasing nuclear charge causes atoms to be smaller (left to right); increased pull on outer electrons due to nuclear charge.
Ionization Energy Ionization energy – energy required to remove one electron from a neutral atom (first ionization energy) X + energy X + + e - Measured using isolated atoms in the gas phase.
Ionization Energy What are the trends on the periodic table? 1.I.E. increases going left to right across a period. 2.I.E increases from bottom to the top of a group
Ionization Energy How can these trends be explained? Period Trend I.E. increases left to right as a result of increasing nuclear charge; a larger nuclear charge results in stronger attraction for electrons Group Trend I.E decreases going down a group due to the fact that electrons are in higher energy levels and farther from the nucleus; also due to the shielding effects of inner electrons.
Successive Ionization Energies
Electron Affinity The energy change that occurs when a neutral atom acquires an electron. X + e - X - + energy (- energy change) or X + e - + energy X - (+ energy change)
Electron Affinity Trends Halogens have the highest electron affinities. Fluorine - 1s 2 2s 2 2p 5 Electron affinity increases moving to the right on the periodic table. (not always & not including noble gases) Carbon – 1s 2 2s 2 2p 2 Nitrogen - 1s 2 2s 2 2p 3 Electron affinity increases moving up a group on the p.t. This trend is not as clear as it is for ionization energy. E.A is affected by nuclear charge and atomic radius
Ionic Radii When atoms become cations, they become smaller. When atoms become anions, they become larger.
Valence Electrons Valence Electrons – the electrons located in s and p orbitals for main group elements Atoms gain, lose, or share electrons to in order to have a complete set of s and p electrons. N 1s 2 2s 2 2p 3 P 1s 2 2s 2 2p 6 3s 2 3p 3
Electronegativity Electronegativity – the ability of an atom to attract electrons in a chemical bond.
Electronegativity Trends N, O, and the halogens have the highest electronegativities Electronegativity increases going left to right across the p.t. up until the halogens Electronegativity increases going up groups on the p.t.
Periodic Properties 1.Which element (Cs, Hf, Au) has the smallest atomic radius? Au 2.Arrange the following elements in order of decreasing electron affinity: C, O, Li, Na, Rb, F F, O, C, Li, Na, Rb
Periodic Properties 3.Arrange the following elements in order of decreasing first ionization energy: Li, O, C, K, Ne, F Ne, F, O, C, Li, K 4.Which element in #3 would have the highest second ionization energy and why? Li
Periodic Properties 5.Which element is the most electronegative among C, N,O, and S? O 6.Which ion has the smallest radius, K + or Ca 2+ ? Ca 2+