CH 115 Fall 2014
Worksheet 5
1. Describe how and why the common monatomic ions form. How can you use the
periodic table to decide what types of ions will form for a particular element?
Common monoatomic ions form because those elements want to have a
completely full valence shell (with 8 valence shell electrons). An element will
either lose or gain electrons (which ever is easiest for them) to get to a full
valence shell or a noble gas electron configuration. You can tell which ion an
element will form by the group number on the periodic table (this doesn’t
work for transition metals!). Anything in group 1 will lose one electron to
become a 1+ cation (e.g. sodium), etc.
2. Give the ground-state electron configurations for each of the following
elements/ions in noble gas notation. Indicate how many valence electrons each
element has and the total number of unpaired electrons present.
a). Ca [Ar]4s2; 2 valence electrons, 0 unpaired electrons
b). Au [Xe]6s15d10; 1 valence electron, 1 unpaired electron
c). Se2- [Ar]4s23d104p6 or [Kr] (I am checking with Dr. March to see which way
he would prefer you write it!}; 8 valence electrons, 0 unpaired electrons
d). Mo [Kr]5s14d5; 1 valence electron, 6 unpaired electrons
3. Define atomic radii. What is the periodic trend involving atomic radii? What causes
this trend?
If you stick two atoms side by side, atomic radii is half of the distance between
the nuclei of those two atoms. Atomic radii increases as you go down a group
and decreases as you move from left to right across a period. The group trend
is caused by placing electrons in higher energy levels  causes the atom to
expand. The period trend is caused by a mix between effective nuclear charge
and shielding (see below).
4. Define effective nuclear charge. How are effective nuclear charge and atomic radii
related?
Effective nuclear charge is a measure of how strongly the nucleus pulls on the
electrons surrounding it. As you go across a period, effective nuclear charge
increases as you keep adding protons (more protons means more positive
charge, which means a stronger pull on the electrons).
Shielding is also important for this story. Shielding occurs when an inner shell
electron blocks a valence electron from being pulled towards the nucleus – so
essentially, shielding and effective nuclear charge are opposing forces. As you
go across a period, shielding is not effected. Even though you are adding more
electrons, you’re adding them all to the valence shell, so there is no increase in
how many inner shell electrons participate in shielding.
So as you move from left to right across the periodic table, you have the same
amount of shielding but an increased effective nuclear charge, which causes
the nucleus to pull more on the electrons and shrinks the atom.