Download Lecture 10 - Unit 2 Part 6 SLIDES

Periodicity
Unit 2: Ch. 6
Coral Gables Senior High
Ms. Kiely
Pre-IB Chemistry I
Bell-Ringer #13
a)
b)
c)
d)
A
B
C
D
Bell-Ringer #13
Answer:
b) B
TEST
Oct 10 and 11
Quiz Review
Periodicity: Atomic Radii
Atomic Radius is not as simple as defining the radius in geometry class!
Remember, the electron cloud doesn’t have a definite edge.
Energy levels and orbitals are not at fixed distances from the nucleus.
Atomic Radius: half the distance between the nuclei of two bonded identical
atoms.
Explanations for trends:
1. Nuclear charge and the,
therefore, effective nuclear
charge experienced by
valence electrons (explains
why trend increases across
period but decreases down
group)
2. Shielding of the valence
electrons (explains
specifically why trend
decreases down a group but
increases across a period)
Across a Period: Atomic Radii Trends
-Atomic radii decreases left to right across a period because the attraction
between the nucleus and the valence electrons increases as the nuclear
charge increases, causing the atom to “shrink” inwards. Remember that
within a period, all elements have the same amount of energy levels.
Down a Group: Atomic Radii Trends
-Atomic radii increases down a group as the
number of occupied energy levels increases.
-As you go down a group, each element has an
additional energy level than the one above it.
Indeed the atoms get bigger because of more
protons and neutrons in the nucleus just like
they do across a period, but the additional
energy level explains why “shrinking” isn’t a
factor here- the shielding effect is at play
here.
Nuclear Charge: the amount of positive charge in an atom’s nucleus; given by
the number of protons in the nucleus; atomic number. It increases by one
between successive elements in the periodic table, as a proton is added to the
nucleus.
Effective Nuclear Charge: how “effective” the nuclear charge is at attracting or
pulling its own valence electrons.
Shielding effect: inner electrons, electrons that are positioned between the
nucleus of an atom and the valence electrons of that atom; inner electrons
repel the valence electrons away from the nucleus. The shielding effect
causes the nuclear charge of an atom to not be as effective as it could be
in attracting its own electrons.
These concepts explain why the radii of atoms of elements down a group
increase while across a period the radii decrease.
Does the effective nuclear charge increase across a period?
Why does the shielding effect remain constant across a period?
Why does the shielding effect increase down a group?
Periodicity: Ionic Radii
Ionic Radius: 1) Positive ions (cations) are smaller than their parent
atoms. 2) Negative ions (anions) are larger than their parent atoms.
Trends in radii size are the same as for regular atomic radii.
Periodicity: Ionic Radii
CATIONS are SMALLER than the
atoms from which they come.
The electron/proton attraction has
gone UP and so the radius
DECREASES.
ANIONS are LARGER than the atoms
from which they come.
The electron/proton attraction has
gone DOWN and so size INCREASES.
Answers