Unit 5 Notes p. 3-4 January 6

Jan 6 - Objectives You will be able to define – Atomic radius – Electronegativity – Ionization Energy – Electron Affinity You will be able to identify trends in each of the above categories

Periodic Trends Ch. 5 - The Periodic Table

Let’s look at some data… In groups of four Each of you will have a different set of data on a periodic table Look at your table until I call time (~5 min) Refer to instructions on packet

Identify the element with the LOWEST value. Circle it. Identify the element with the HIGHEST value. Circle it. Identify any elements that have NO values. Cross them out. Draw two arrows on your table showing: In which direction numbers increase in a period (horizontal row) In which direction numbers increase in a group (vertical column)

What did you see: Atomic Radius

Electronegativity

First Ionization Energy

Electron Affinity

Periodic Law (Periodicity) When elements are arranged in order of increasing atomic number, elements with similar properties appear at regular intervals.

Shielding Decrease in attraction between an electron and the nucleus in an atom with more than one electron shell

Atomic Radius The size of an atom Li

Atomic Radius Atomic Radius: the size of atom Increases to the LEFT and DOWN

Atomic Radius Which atom has the larger radius? Why? – Be or Ba

Atomic Radius Which atom has the larger radius? Why? – Be or Ba Ba: Farther down the periodic table, A.R. increases farther down (it has higher energy levels (larger shells), plus shielding) – Ca or Br Ca: To the left of the table, A.R. increases to the left (Both are in the same energy level, but there is less nuclear charge for Ca, so electrons are not held as tightly)

Electronegativity  Ability of a bonded atom to attract electrons Note: Atoms must be BONDED for electronegativity to be measured  Fluorine the most electronegative atom

Electronegativity Increases UP and to the RIGHT, Does NOT include noble gases

Electronegativity Which atom has the greater electronegativity? Why? – P or Cl? Cl: Farther to the right (nearly a full shell, smaller atom attracts electrons more strongly) – Rb or F? F: Fluorine is THE most electronegative atom; small atom and nearly a full shell

First Ionization Energy First Ionization Energy – energy required to remove one electron from a neutral atom He

First Ionization Energy Increases UP and to the RIGHT

First Ionization Energy Which atom has the higher first ionization energy? Why? – N or Bi N: Higher up the periodic table (it’s a smaller atom so electrons are held more tightly) – Ba or Ne Ne: smaller atom, full shell so electrons are stable and hard to remove

Successive Ionization Energies Successive Ionization Energies – Energy required to remove more than one electron – 2 nd IE – Energy needed to remove 2 electrons – 3 rd IE – Energy needed to remove 3 electrons – Etc.

Successive Ionization Energies Successive IEs increase because it gets harder to separate charges as the charges get bigger (Coulomb’s Law) Large jump in I.E. occurs when a CORE e - is removed.

Successive Ionization Energies Mg1st I.E.736 kJ 2nd I.E.1,445 kJ Core e - 3rd I.E.7,730 kJ Large jump in I.E. occurs when a CORE e - is removed.

Successive Ionization Energies Large jump in I.E. occurs when a CORE e - is removed.

Al1st I.E.577 kJ 2nd I.E.1,815 kJ 3rd I.E.2,740 kJ Core e - 4th I.E.11,600 kJ Successive Ionization Energies Large jump in I.E. occurs when a CORE e - is removed.

Successive Ionization Energies Which atom has the higher second ionization energy? Why? – K or Ca K: removing a second electron would require removing a core electron – S or Cl Cl: has almost a full shell, does not lose electrons easily

Electron Affinity Ability of an atom to attract electrons

Electron Affinity Electron affinity Increases UP and to the RIGHT, Does NOT include noble gases

Electron Affinity Which atom has the greater electron affinity? Why? – Na or F? F: It’s a smaller atom and holds electrons more strongly; nearly a full shell – O or Se O: smaller atom, attracts electrons more strongly