The Periodic Table and Trends Please have a periodic table out. SONG SONG

Use this one.

Dmitri Mendeleev 1834 – 1907 Russian chemist and teacher given the elements he knew about, he organized a “Periodic Table” based on increasing atomic mass (it’s now atomic #) he even left empty spaces to be filled in later

At the time the elements gallium and germanium were not known At the time the elements gallium and germanium were not known. These are the blank spaces in his periodic table. He predicted their discovery and estimated their properties.

Henry Moseley 1887 – 1915 arranged the elements in increasing atomic number (Z) properties now recurred periodically better than before

Design of the Table Groups (18 total) are the vertical columns. elements have similar, but not identical, properties most important property is that they have the same # of valence electrons

Know the names of these groups

valence electrons- electrons in the highest occupied energy level all elements have 1,2,3,4,5,6,7, or 8 valence electrons These are a lower level. Therefore the d sub-level is never included for valence electrons

The highest level is 4.

Lewis Dot-Diagrams/Structures a short cut wherevalence electrons are represented as dots around the chemical symbol for the element Na Cl

2 1 3 5 8 2

What two blocks will always be the highest occupied level?

Look, they are following my rule!

Br is [Ar] 4s2 3d10 4p5 How many valence electrons are present? B is 1s2 2s2 2p1; 2 is the outermost energy level it contains 3 valence electrons, 2 in the 2s and 1 in the 2p Br is [Ar] 4s2 3d10 4p5 How many valence electrons are present?

Periods (7) are the horizontal rows do NOT have similar properties however, there is a pattern to their properties as you move across the table that is visible when they react with other elements

Trends in the table

many trends are easier to understand if you comprehend the following the ability of a nucleus to “hang on to” or attract its valence electrons is the result of two opposing forces the attraction between the electron (-) and the nucleus (+) the repulsions between the electron (-) in question and all the other electrons (-) in the atom (this is called the shielding effect) the net resulting force of these two is referred to effective nuclear charge

This is a simple, yet very good picture. Do you understand it?

Atomic radii the distance from the nucleus to the outermost electron cannot measure the same way as a simple circle due to electrons are not in a fixed location therefore measure distance between two nuclei and divide by two

periods across the periodic table groups increases downwards as more levels are added more shielding pushes outer levels out periods across the periodic table radii decreases the number of protons in the nucleus increases increases the strength of the positive nucleus and pulls electrons in the given level closer to it added electrons are not contributing to the shielding effect because they are still in the same level H Li Na K Rb McGraw Hill video

Looking at ions compared to their parent atoms Ionic radii Looking at ions compared to their parent atoms atoms tend to gain or loose electrons in order to have the electron configuration of a noble gas

cations (+ ions) are smaller than the parent atom have lost an electron (actually, lost an entire level!) therefore have fewer electrons than protons less no shielding Li+ .078nm + Li forming a cation Li 0.152 nm

anions (- ions) are larger than parent atom have gained an electron(s) to achieve noble gas configuration effective nuclear charge has decreased since same nucleus now holding on to more electrons plus, the added electron repels the existing electrons farther apart (kind of “puffs it out”) F - 0.133 nm 10 e- and 9 p+ F 0.064 nm 9e- and 9p+

Ionization energy the minimum energy (kJ/mol) needed to remove an electron from a neutral gaseous atom in its ground state, leaving behind a gaseous ion X(g)  X+(g) + e- first ionization energy- energy to remove first electron second ionization energy- energy to remove second electron third ionization energy- and so on…

don’t forget-- gaseous

decreases down a group outer electrons are farther from the nucleus and therefore easier to remove inner core electrons “shield” the valence electrons from the pull of the positive nucleus and therefore easier to remove

increases across a period the nucleus is becoming stronger (effective nuclear charge) and therefore the levels, especially valence electrons, are pulled closer atomic radii is decreasing this makes it harder to remove a valence electron since it is closer to the nucleus or another way to look at it… a stronger nuclear charge acting on more contracted orbitals

Electronegativity measures the attraction for a shared pair of electrons in a bond Linus Pauling (1901 to 1994) came up with a scale where a value of 4.0 is arbitrarily given to the most electronegative element, fluorine, and the other electronegativities are scaled relative to this value.

trends (same as ionization energy and for the same reasons) as you go down a group electronegativity decreases the size of the atom increases the bonding pair of electrons (-) is increasingly distant from the attraction of the nucleus (+) the bonding pair of electrons (-) are shielded because of core electrons (-) interfering with the nucleus’ (+) hold on valence electrons H Li Na K Rb

as you go across a period electronegativity increases the atoms become smaller as the effective nuclear charge increases easier to attract the shared pair of electrons as they will be in an orbital closer to the nucleus moving from L to R on the table