Unit 8: The Periodic Table Trends

Notes goals Understand basic history and arrangement of periodic table Be able to use the periodic table to predict chemical trends Be able to use periodic trends to determine the identity of elements

History Dmitri Mendeleev (1869, Russian) Organized elements by increasing atomic mass. Elements with similar properties were grouped together. There were some discrepancies.

History Henry Mosely (1913, British) Organized elements by increasing atomic number. Resolved discrepancies in Mendeleev’s arrangement.

Arrangement Period = a single row going across left to right Numbered from 1 to 7 Group/Family = a column going down top to bottom Different numbering systems (1-18) or (1A – 8A, etc.) Elements in the same group have similar physical and chemical characteristics

Periodic Trends Effective Nuclear Charge Atomic Radius Ionic Radius Ionization Energy Electron affinity Electronegativity Overall Reactivity

Effective Nuclear Charge Valence electrons = electrons in outer level Core electrons = electrons in levels underneath outer level Valence electrons are: Attracted to protons in nucleus Repelled by core electrons which are in the way (shielding)

Effective Nuclear Charge Effective Nuclear Charge = the net attraction of the valence electrons to the nucleus Greater effective nuclear charge means valence electrons are closer to the nucleus

Effective Nuclear Charge (Zeff ) Felt by Shielding electrons (S) Effective Nuclear Charge (Zeff) Atomic Number (Z = #p+) Zeff = Z - S Inner Core of Shielding e- = screening constant Attractive Charge felt by valence electrons Atomic Number (#p+) {Eff.Nuclear.Charge*}

Effective Nuclear Charge (Zeff ) depends on both the fact that valence electrons are both: attracted to the nucleus repelled by the other (shielding) electrons. Effective Nuclear Charge (Zeff.) Increasing +1 Shielding electrons +8 (0e-) +2 +3 +4 +5 +6 +7 (2e-) (10e-)

Effective Nuclear Charge Trend Increases to the RIGHT

6.3 Atomic Radius The atomic radius is one half of the distance between the nuclei of two atoms of the same element when the atoms are joined. This diagram lists the atomic radii of seven nonmetals. An atomic radius is half the distance between the nuclei of two atoms of the same element when the atoms are joined.

Atomic Radius Why larger going down? Why smaller to the right? Higher energy levels have larger orbitals Why smaller to the right? Increased nuclear charge without additional shielding pulls electrons in tighter

Atomic Radius 14

Atomic Radius 15

Atomic Radius Trend Atomic Radius Increases to the LEFT and DOWN

Ionic Radius Cations (positive ions) Anions (negative ions) Lose electron Become smaller than the neutral atom Anions (negative ions) Gain electron Become larger than the neutral atom Trend is the same as the atomic radius

Summary of Ionic Radius Cations are always smaller than the original atom. The entire outer EL is removed during ionization. Conversely, anions are always larger than the original atom. Electrons are added to the outer EL.

Cation Formation Effective nuclear charge on remaining electrons increases. Na atom 1 valence electron Remaining e- are pulled in closer to the nucleus. Ionic size decreases. 11p+ Valence e- lost in ion formation Result: a smaller sodium cation, Na+

Anion Formation A chloride ion is produced. It is larger than the original atom. Chlorine atom with 7 valence e- 17p+ One e- is added to the outer shell. Effective nuclear charge is reduced and the e- cloud expands.

Ionic Radius

Ionic Radius Trend Ionic Radius Increases to the LEFT and DOWN within the cations and the anions Cations (+) Anions (-)

Ionization Energy Ionization Energy = Energy required to remove an electron from a neutral atom Smaller atoms have electrons closer to the nucleus so they are held tighter. This means a higher ionization energy is required to take an electron away. There are a few exceptions where where s, p, or d orbitals are filled or half-filled which creates a more stable structure

Ionization Energy

Ionization Energy Trend First Ionization Energy Increases UP and to the RIGHT

Electron Affinity What does the word ‘affinity’ mean? Electron affinity is the energy change that occurs when an atom gains an electron (also measured in kJ). Where ionization energy is always endothermic, electron affinity is usually exothermic, but not always.

Electron Affinity Increases UP and to the RIGHT

Electronegativity Electronegativity is a measure of an atom’s attraction for another atom’s electrons. ranges from 0 to 4. Generally, metals are electron givers and have low electronegativities. Nonmetals are are electron takers and have high electronegativities. What about the noble gases?

Electronegativity Increases UP and to the RIGHT

Overall Reactivity This ties all the previous trends together in one package. However, we must treat metals and nonmetals separately. The most reactive metals are the largest since they are the best electron givers. The most reactive nonmetals are the smallest ones, the best electron takers.

Overall Reactivity Your help sheet will look like this:

The Octet Rule The “goal” of most atoms (except H, Li and Be) is to have an octet or group of 8 electrons in their valence energy level. They may accomplish this by either giving electrons away or taking them. Metals generally give electrons, nonmetals take them from other atoms. Atoms that have gained or lost electrons are called ions.