Module 3.03
Periodic Trends
Periodic Table
To Review:
•Periods – rows of the periodic table
•Groups – columns of the period table
Periodic Table
Trends
•Trends are useful in making predictions
•Periodic table is arranged based on theproperties of the elements
•Reading across each period and down eachgroup you will see repeated trends inproperties
Effective Nuclear Charge
Effective Nuclear Charge – the charge (fromthe nucleus) felt by the valence electronsafter the number of shielding electrons aretaken into account.
Subtract the number of shielding (core)electrons from the total nuclear charge(number of protons) gives the effectivenuclear charge.
Effective Nuclear Charge
Example:
Nitrogen
7 protons (atomic number of 7)
-2 core electrons (1s2 electrons)
--------
5 is the effective nuclear charge
Each of the valence electrons in nitrogen feelsan effective nuclear charge of about +5.
Effective Nuclear Charge
Across a Period:
The effective nuclear charge felt by anatom’s valence electrons increases by onefor each element from left to right in aperiod
Down a Group:
The effective nuclear charge felt by anatom’s valence electrons stays constant foreach element going down a group
Atomic Radius
Atomic Radius – half the distance betweenthe centers of two atoms that are bondedtogether.
Atomic Radius
Across a Period:
There is a gradual decrease in atomicradii from left to right across a period.
Down a Group:
There is a general increase in atomic radiigoing down each group on the periodictable.
Atomic Radius
Ionization Energy
Ionization energy – energy required to removeon electron from an element, resulting in apositive ion.
Elements with a lower effective nuclear chargefelt by their electrons will give up an electroneasier than other elements.
Ionization Energy
Across a Period:
Ionization energy has a general increasefor elements across a period from left toright.
Down a Group:
Ionization energy of elements decreasesgoing down a group because the atomicradius of the atoms increases.
Ionic radius
Ionic radius – ½ the diameter of an ion.
Note: Metals naturally form cations (positiveions) by losing one or more electrons.Nonmetals naturally form anions (negativeions) by gaining one or more electrons.
Atomic Radius vs Ionic Radius
Ionic Radius
Across a Period:
Within each period, the metals at the left formcations and the nonmetals at the right formanions. There is a decrease in the ionic radii of thecations from left to right and a decrease in theionic radii of the anions from left to right.
Down a Group:
Ionic radii increase down a group, following thesame trend as atomic radii.
Electronegativity
Electronegativity – measure of the attraction ofan atom for the electrons in a chemical bond.
Electronegativity
Across a Period:
Electronegativity increases from left to rightacross a period because of the effective nuclearcharge.
Down a Group:
Electronegativity decreases down a group as aresult of a greater atomic radius.
Module 3.03
Periodic Trends