Download 5-3 Electron Configurations and Periodic Properties

5-3
Electron Configurations and
Periodic Properties
PreClass: What does the numerical
value of the period (ie. Period 4)
represent in addition to just physically
locating a particular row on the
Periodic Table?
• The period = n value of the outermost
electrons
• Ie. If n=4, than all elements on the 4th
period have their outermost electrons in the
4 energy level !!
Beaker Breaker
1. Mendeleev’s periodic table had the
elements ordered in increasing ________
2. What is the element in the same Period that
follows potassium, K ?
3. What is the name of the group on the far
left of the Periodic table?
4. What element ends with 3p1 ?
Atomic Radius
• One-half the distance between the nuclei of
identical atoms that are bonded together
Group Trend for Atomic Radii
• Down a group the radius of an atom gets
larger
• The principal quantum number increases
and the cloud grows by one shell
Period Trend for Atomic Radii
• Across a period the radius of an atom gets
smaller
• You’re adding electrons to approximately the
SAME region (same “n”) while the electrons
are being pulled in more tightly as nuclear
charge increases (more protons)
• Exception: Noble Gas Family – atoms don’t
interact and pull together like in other atoms
because atoms already have full outer level
Questions
• Of the elements Li, O, C, and F, identify the
one with the largest atomic radius and the one
with the smallest atomic radius
• Of the elements Br, At, F, I and Cl, identify the
one with the largest atomic radius and the one
with the smallest atomic radius
Answers
• Li = largest
• F = smallest
F=smallest
At = largest
Why does the atomic radii not
decrease very much as you move
across the d-sublevel?
• Increasing # inner electrons shield
outer level electrons from nucleus
Why would the atomic radius of
hafnium (#72) be LESS than that
of zirconium (#40)?
• Hf has such a greater nuclear charge
Beaker Breaker
1. Which has the largest and which has the
smallest atomic radius of the following:
C, Ge, Sn, Si
2. Which has the largest and which has the
smallest atomic radius of the following:
K, Cu, As, Br
Ionization energy, IE
• Energy required to remove an electron from a
neutral atom (or first ionization energy, IE1)
• Ion: charged particle
• Ionization: process of an electron being lost
or gained from an atom which results in the
formation of an ion (Na+ and Cl-)
Group Trend for Ionization
Energies
• Down a group the ionization energy of an
atom generally gets smaller
• Electrons are at a greater distance from the
nucleus
• Outer electrons are shielded from the
nucleus by inner electrons
Period Trend for Ionization
Energies
• Across a period the ionization energy
increases
• Nuclear charge gets greater while atomic
size decreases
Examples
• Which of the following has the largest
ionization energy?
Na, Mg, P, Cl
Which of the following has the smallest
ionization energy?
Be, Mg, Ca ,Sr
Homework
• Pg 156-157
# 28 (all), 29 (all), 30 (all), 31 (a &b only)
Beaker Breaker
• Which alkaline earth metal has the lowest
ionization energy?
• Which halogen has the largest atomic
radius?
What will determine whether an
electron is easily lost or not?
(4 things)
4 Factors Affecting Ionization
Energy
1. Nuclear charge (greater the charge, the greater the
IE)
2. Shielding Effect (greater the shielding, the lower
the IE)
3. Radius (the greater the radius, the less the IE)
4. Sublevel configuration (an electron from a halffull or full sublevel requires more IE)
nd
2
Ionization energy, IE2 is the energy
required to take a 2nd electron away
from an atom.
Why is IE2 always greater than IE1?
Why is the IE2 of Na so much greater
that the greater of Mg ??
(Hint: look at the electron config.)
Multiple Ionization Energies
• Energy required to remove the 2nd, 3rd, etc.
electron from an atom
• IE3 > IE2 > IE1 because remaining
electrons will be held more tightly as the
electron repulsion decreases and the cloud
is pulled in more tightly
Why does the IE go UP as you
go down a d-sublevel group but
go down in an s-sublevel?
• “f” sublevel has minimal shielding
effect while the nuclear charge
continues to grow (“Lanthanide
Contraction”)
Electron Affinity p. 147
• Energy change that occurs when an electron is
acquired by a neutral atom
• Most atoms RELEASE energy when they acquire an
electron: A + e-  A- + energy
– energy has negative sign
• Some atoms gain energy when they acquire an
electron: A + e- + energy  A– energy has positive sign; atom is unstable &
loses electron spontaneously
Group Trend for Electron
Affinities
• Down a group the electron affinity of an atom
tends to get smaller
• Although there is an increasing nuclear charge,
there are more levels so the size is greater
• Adding an energy level usually dominates!
Period Trend for Electron
Affinity
• Across the p-sublevel, the energy change
increases (becomes more negative)
–Electron config is close to being full
and the size is smaller
Why is the electron affinity of
nitrogen so low when compared
to carbon or oxygen?
• Adding an electron to carbon half
fills the 2p sublevel
• Adding an electron to nitrogen
forces the config to go from stable
(half-filled) to less stable (no spec
arrangement)
Multiple Electron Affinities
• It is always more difficult to add a
2nd electron to an already
negatively charged ion..therefore,
all 2nd electron affinities are
positive
Beaker Breaker
1. Which alkaline earth metal has the largest
electron affinity?
2. Which of the following has the smallest
electron affinity Na, Mg, P or Cl
Ionic Radii of Cations
• Cation: positive ion
• Formed by an atom losing electron(s)
• Always smaller because electron cloud is
smaller (less repulsion) & sometimes
even one less energy level!
Ionic Radii of Anions
• Anion: negative ion
• Formed by an atom gaining electron(s)
• Always larger because electron cloud
is greater (more repulsion among
electrons)
Group Trend for Ionic Radii
• Down a group the ionic radius of an
atom generally gets larger
• Electrons are at a greater distance from
the nucleus (higher E level) and have
more shielding
Period Trend for Ionic Radii
• Metals (left side):form cations
– Cationic Radius: Decreasing ionic radius
as nuclear charge increases without adding
an energy level
• Non-metals: form anions
– Anionic Radius:Decreasing ionic radius
as nuclear charge increases without adding
an energy level
Practice
• Pg 157 # 41
• Pg 157 # 46
PC: What are VALENCE
electrons?
Valence Electrons
• Electrons available to be lost,
gained, or shared in the formation
of chemical compounds
• Often the outermost electrons
because they are held most loosely
What would be the # of valence
electrons in……..
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Calcium
Lithium
Chlorine
carbon
What would be the # of valence
electrons in……..
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Calcium – 2 : 4s2
Lithium – 1: 2s1
Chlorine – 7: 3s23p5
Carbon – 4: 2s22p2
Beaker Breaker
1. How many valence electrons does Te
have?
2. Which has a smaller ionic radius Na+1 or
Ca+1 ?
Electronegativity
• Measure of the ability of an atom in a
chemical compound to attract electrons
• Fluorine is the MOST electronegative
element – assigned an arbitrary value of
4.0
• All other values are relative to F
• 3 highest values: F – O - N
Group Trend for Electronegativity
• Tend to decrease down a group (or
stay the same) as the atoms gets
larger
Period Trend for Electronegativity
• Tend to increase across the period
as the atoms gets smaller, the nuclear
charge becomes greater, and the
atom is getting closer to a noble gas
configuration
Determine the likely charge for the
following elements: Ca, O, Al
1. Write the noble gas configuration of the
element
2. Determine if electrons will be LOST or
GAINED to make the element stable
3. ID the noble gas whose electron
configuration by losing/gaining these
electrons
4. Write the formula for the ion
5. ID it as a cation OR anion
Determine the likely charge for the
following elements: Ca, O, Al
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Ca: [Ar]4s2
Ca will LOSE 2 eCa now has the Ar config
Ca+2
cation
O: [He]2s22p4
O will GAIN 2 eO now has a Ne config
O-2 , an anion
Al
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Al: [Ne]3s23p1
Al will LOSE 3 eAl now has a Ne config
Al+3 , a cation
How do d-block elements form
ions?
• Electrons in the highest occupied
sublevel are always removed first
Why does zinc become a +2 ion?
Which electrons are lost when
titanium becomes a +2, a +3 and a
+4 ion?
Beaker Breaker
Which one is larger?
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Na or K
Na or Mg
Na or Na+
O or FO or O-2
Which one is larger?
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Na or K
K
Na or Mg
Na
Na or Na+
Na
O or FFO or O-2
O-2