Electron Configurations and the Periodic Table

1

• Electron Configurations and the Periodic Table
• Mendeleevs Periodic Table (page 133-135)
• 1. In 1870 a Russian chemist named Dmitri
  Mendeleev proposed a periodic table on the belief
  that the properties of elements were a function
  of atomic mass.
• 2. Mendeleev stated that elements of similar
  properties repeated in rows (now we call them
  periods) of varying ___________.
• 3. Mendeleev placed the 67 elements in order of
  increasing atomic mass. He placed 7 elements in
  each of the first two _____________ (horizontal
  rows) much like periods 2 and 3 of the modern
  periodic table except for the noble gases which
  were not yet known.
• 4. Mendeleev placed ____ elements in each of the
  next two periods much like periods 4 and 5 of the
  modern periodic table.
• 5. For elements of similar properties to be in
  the same group (vertical column) Mendeleev had to
  leave some spaces in his table. He stated that
  the blanks in his table corresponded to _________
  elements.
• 6. From the blank spaces in his table Mendeleev
  predicted the existence and properties of ___
  unknown elements. An example can be seen in an
  unknown element that he called _____________ now
  known to be Germanium. Fig. 2 pg. 134 shows how
  close Mendeleevs predicted of the existence of
  unknown elements.
• 7. There was a problem with Mendeleevs periodic
  table. If the elements are arranged in order of
  increasing atomic mass as Mendeleev did, then Te
  and __ seemed to be in the wrong column although
  they seemed to be in the correct column when
  observing their properties. Mendeleev believed
  that the problem was that their atomic mass had
  been measured __________. Other pairs of
  elements were seen to have that same problem such
  as Co and __, Ar and ___, as well as others.
• Moseley and the Periodic Law (pg. 135)
• 1. In 1913 a British scientist named Henry
  Moseley solved Mendeleevs dilemma. Through the
  use of x-ray experiments he determined the number
  of protons in the nucleus which is called the
  atomic ___________.

length
periods
17
unknown
3
ekasilicon
I
wrong
Ni
K
number
2
repeat

• 2. When the elements are placed in order of
  increasing atomic number, their properties
  __________ in a periodic manner. Stating that
  chemical properties are a function of atomic
  number is referred to as the periodic law.
• 3. The modern periodic table lists the elements
  in order of increasing atomic ___________.
• Electron Configurations and the Periodic Table
  (pg. 138-149)
• 1. An elements properties are determined by the
  elements electron configuration (namely the
  valence shell electron configuration).
• 2. We also know that elements with the same
  properties appear in the same __________ on the
  periodic table. Therefore, elements in that same
  group must have similar valence shell electron
  configurations. Note that elements in group IA
  all have an ____ in their outer shell and
  elements in groups VIIA (group 17) all have an
  _________ configuration in their outer shells.
  We can see that all the Group A elements or
  representative elements have the same number of
  __ and or __ electrons in their outer most shell
  for a given group.
• Groups III-II B (groups 3-12) (the transition
  metals) have ___ electrons in the outer shell as
  well as ____ electrons in the shell __ energy
  level in from the outer shell for a given group.
  For group III B this would be ____ s electrons in
  the outermost shell and ____ d electron in the
  shell 1 energy level closer to the nucleus. Group
  IA (group 1) elements are called the ____________
  metals. Group IIA (group 2) elements are called
  the ________________ metals. Group VIA (group
  16) elements are called the _______________.
  Group VIIA (group 17) are called
  the_____________. Group VIIIA (group 18) are
  called the __________ gases. f block elements are
  called ______________ transition metals.
• Exceptional Electron Configurations (page 119 of
  text)
• 1. Some elements do not have the electron
  configurations that you would expect. Look at
  ___ and ___. For each of these elements you
  would expect there to be __ 4s electrons but they
  only have __ 4s electron. This is due to the fact
  that a sublevel is stabilized if it is half or
  completely full.

number
group
s1
s2p5
p
s
s
1
d
2
1
alkali
alkaline earth
chalcogens
halogens
noble
inner
Cr
Cu
2
1
2. This is due to the fact that when sublevels of
nearly equivalent energy (here the 4s and 3d) are
next to each other, the 3d can ______ an electron
from the 4s. Note, d and f orbitals can steal
electrons from an s orbital for d__ or d__ and
f__ or f __ (however, p2 will not take an s
electron due to being too large of a difference
in energy)
steal
6
13
4
9
3
3d4

• Cr

Ar
4s2
4s1
3d5
?
?
?
?
?
?
?
___
___
___
___
___
___
Ar
3d
4s
NOT AS STABLE AS IT CAN BE!!!
Greatly stabilized energetically due to less
electron repulsions
Cu
Ar
4s2
3d9
3d10
4s1
?
Ar
?
?
?
?
?
?
?
?
___
___
___
___
___
?
?
___
3d
4s
NOT AS STABLE AS IT CAN BE!!!
Greatly stabilized energetically due to less
electron repulsions
4
4d4
5s1
4d5

• Mo

Kr
5s2
?
?
?
?
?
?
?
___
___
___
___
___
___
Kr
4d
5s
NOT AS STABLE AS IT CAN BE!!!
Greatly stabilized energetically due to less
electron repulsions
Ag
Kr
5s2
4d9
4d10
5s1
?
Kr
?
?
?
?
?
?
?
?
___
___
___
___
___
?
?
___
4d
5s
NOT AS STABLE AS IT CAN BE!!!
Greatly stabilized energetically due to less
electron repulsions
5

• Dmitri Mendeleev (1869, Russian)
• Organized elements by increasing atomic mass.
• Elements with similar properties were grouped
  together.
• There were some discrepancies.

6
(No Transcript)
7

• Henry Mosely (1913, British)
• Organized elements by increasing atomic number.
• Resolved discrepancies in Mendeleevs arrangement.

8

• Representative Elements
• Transition Metals
• Inner Transition Metals

Elements in the same group have a similar
electron configuration therefore have similar
chemical properties
Group IA alkali metals (s1) Group IIA -
alkaline earth metals (s2) Group VIA chalcogens
(s2p4) Group VIIA halogens (s2p5) Group VIIIA
noble gases (s2p6)
VIIA
IA
IIA
VIA
VIIIA
9
Homework 1. Mendeleev arranged his periodic
table in order of increasing atomic ________. 2,
Moseley arranged his periodic table in order of
increasing atomic _________. 3a. What was wrong
with the order that Mendeleev used to arrange
elements?b. To what did Mendeleev attribute the
problem with his table?c. What did Mendeleev
predict through the use of his table? 4a. What
is the periodic law? b. Whose periodic table
arrangement agrees with the periodic law? 5a.
Write the electron configuration for Gold. b.
Draw the orbit filled diagram (up and down
arrows) for the electron configuration for
Gold. 6a. Write the electron configuration for
Samarium (Z 62). b. Show the orbit filled
diagram for this electron configuration.
10

• Periodic Trends and Reaction Tendencies
• Radii of Atoms (see fig. 12-14 pg. 150-152)
  size of atom (can be measured by ½ the distance
  between the nuclei of a diatomic molecule or
  subtract a known atomic radius from the distance
  between 2 nuclei for a molecule that is not
  diatomic)
• Group Trend Atoms get larger __________ a
  group due to the fact that electrons are being
  added to higher energy levels which are
  __________ from the nucleus.
• Periodic Trend Atoms get ___________ across a
  period (left to right), due to more protons being
  added to the nucleus as electrons are being added
  to the same or lower (for d and f electrons)
  energy levels. Since the electrons are no
  further from the nucleus and the nucleus has more
  overall charge, each electron experiences
  (feels) a ____________nuclear charge (this will
  cause all the electrons to be pulled closer to
  the nucleus).
• Sample problem E pg. 152 Practice 1-3 pg. 152
• Ionization Energy (pg. 153-156) A energy ?
  A e-
• Energy required to remove an electron from a
  gaseous atom. The first ionization energy is the
  energy required to remove the most loosely held
  _________ from a gaseous atom. Ionization energy
  is measured in kJ/mol (1 mole 6.02.1023
  particles).

down
further
smaller
greater
electron
Periodic Trend See figure 15 pg. 153 and fig. 16
pg. 154). In general, ionization energy tends to
increase to the right across a period due to a
greater nuclear charge holding on to electrons
that are added to the same energy level (this
same reason made the atoms of elements
__________ in size to the right across a
period). Sample exceptions to the trend Be to B
due to Be having a _____ s sublevel (while B has
a partially full p sublevel). There is also a
dip from N to O due to N having a half full p
sublevel while O has no special arrangement in
its orbital. Notice the large drop in ionization
energy from He to Li. The drop in ionization
energy from He to Li is due to 1. The valence
electron removed from Li is in a higher energy
level _________ from the nucleus. 2. The inner
shell electrons in Li shield the positive charge
from the valence shell electrons. 3. He having a
completely ______ s sublevel.
decrease
full
further
full
11

• Group Trend Ionization energy tends to decrease
  down a group. This is due to the valence
  electrons being located in higher energy levels
  __________ from the nucleus and the greater
  ____________ of the valence electron from the
  inner electrons. (See the noble gases to observe
  this trend)
• Shielding effect exhibited by inner shell
  electrons Draw the complete electron
  configuration of F and Br (do not use the noble
  gas configuration). Then draw a representation
  of the atom to observe the shielding exhibited by
  the inner shell electrons.

further
shielding
2s2
2p5
1s2
4p5
4s2
F
1s2
Br
2s2
2p6
3s2
3p6
3d10
7e-
7e-
2e-
18e-
8e-
2e-
We can conclude that metals have lower first
ionization energy values than non-metals
12

• Multiple Ionization Energy Values (see table 3
  pg. 155)
• After an atom loses 1 valence electron, it
  becomes more __________ to remove subsequent
  valence electrons from the atom (this is due to
  the loss of each electron increasing the relative
  numbers of protons to electrons in the atom).
• When an atom loses all of its valence electrons,
  the energy required to remove the subsequent
  electron from an inner shell (with full s and p
  sublevels) makes a large________.
• Notice the very large jump between the 1st and
  2nd ionization energy for ___ due to it only
  having
• 1 valence electron or between the 2nd and 3rd
  ionization energy for Be that has __ valence
  electrons.
• Example See table 3 pg. 155 for the large jumps
  ionization energy values for Al.

difficult
jump
Li
2
Elements will tend to gain or lose electrons so
as to obtain 8 valence electrons thereby making
the charges of elements in ionic compounds to be
1 for group IA 2 for group IIA 3 for group
IIIA -3 for group VA -2 for group VIA -1 for
group VIIA (1 for Ag 2 for Zn and Cd)
Al 1s22s22p63s23p1
Al1 1s22s22p63s2
Al2 1s22s22p63s1
Al3 1s22s22p6
e-
3e-
8e-
8e-
578 kJ/mol
2e-
8e-
1817 kJ/mol
2e-
2745 kJ/mol
?
2e-
2e-
?
?
8e-
2e-
Large jump in energy!!
7e-
2e-
11,578 kJ/mol
Al4 1s22s22p5
13
Homework 1. pg. 167 32 35 2. Give 2
factors to account for the large jump in
ionization energy for the 3rd electron for Be
(see table 3 pg. 155). 3. Given the following
electron configurations Q 3s23p5 R 3s1 T
2s22p4 a. Which are in the same period? b. Give
the name of each element and state to what group
it belongs. c. Which element would you predict to
have the smallest ionization energy? d. Which
element would you predict to have the largest
atomic radius? e. What charge of ion is each most
likely to obtain when making a compound?
14
Atomic Radius
Review

• Atomic Radius

• Increases to the LEFT and DOWN

15
Atomic Radius

• Why larger going down?
• Higher energy levels have larger orbitals
• Shielding - core e- block the attraction between
  the nucleus and the valence e-
• Why smaller to the right?
• Increased nuclear charge without additional
  shielding pulls e- in tighter

16

• First Ionization Energy

Tends to increase UP
Tends to increase up a group due to electrons
being closer to the nucleus and less shielded
Tends to increase RIGHT

• Tends to increase to the right due to more
  nuclear
• charge as electrons added to same energy
  level
• no further from the nucleus

Dip in IE half-filled sublevel (p3)
Dip in IE filled sublevel (s2)
17

• First Ionization Energy

• s2

• p3

He
Ne
Ar
Li
Na
K
18

• Successive Ionization Energies

• IE increases for each electron is removed

• Large jump in I.E. occurs after all valence
• electrons are removed

Al (1s22s22p63s23p1) 1st I.E. 578 kJ/mol
Al1 (1s22s22p63s2) Al1 (1s22s22p63s2)
2nd I.E. 1817 kJ/mol Al2
(1s22s22p63s1) Al2 (1s22s22p63s1 ) 3rd
I.E. 2745 kJ/mol Al3 (1s22s22p6)
Al3 (1s22s22p6) 4th I.E. 11,578 kJ/mol
Al4 (1s22s22p5)
19
Al (1s22s22p63s23p1) Al1
(1s22s22p63s2) Al1 (1s22s22p63s2)
Al2 (1s22s22p63s1)
Al2 (1s22s22p63s1 ) Al3
(1s22s22p6) Al3 (1s22s22p6)
Al4 (1s22s22p5)
20

• Periodic Trends Continued
• Electron Affinity (pg. 157-159) A e-? A-
  (energy may be a reactant or product)
• 1. The energy released or taken in when an
  ____________ is added to a gaseous atom.
• 2. A negative value of electron affinity
  indicates an attraction of an atom for an
  electron.
• 3. A ______________ electron affinity means that
  the electron is being repelled by the atom of the
  element (a positive value indicates an energy
  requirement for the atom to take the electron).
• 4. Metals tend to have more ______________
  electron affinity values than non-metals.
• Periodic Trend electron affinity values tend to
  become more negative to the right across of
  period (however there are many easily explained
  exceptions)
• Group IIA and VIIIA have positive values (groups
  with full sublevels in the valence shell) and the
  value of 0 for N with its half full p sublevel.
  Groups with half or completely full sublevels
  will not easily take in a new electron.
• Group Trend Down a group electron affinity
  values tend to become more positive (however this
  is not a well defined trend as an added electron
  to higher energy levels it is less attracted to
  the nucleus due to _____________ by inner
  electrons and the greater distance from the
  nucleus although the electrons have more room to
  move which makes their addition easier).
• Ionic Radii (see fig. 19 pg. 159)
• 1. When an atom loses an electron it becomes a
  positive ion called a __________. This positive
  ion will be ____________ than the parent atom.
• 2. When an atom gains an electron it becomes a
  negative ion called a __________. This negative
  ion will be ____________ than the parent atom.

electron
positive
positive
shielding
cation
smaller
anion
larger
21
Examples

• Which particle has the larger radius?

S or S-2
S-2
Al
Al or Al3
22
Which ion would you predict to be more stable
Fe2 or Fe3 ?
Fe Ar4s23d6
/
\
Fe2 formation Ar4s23d6
Fe2 Ar 3d6
/
/
\
Fe3 Ar 3d5
Fe3 formation Ar 4s23d6
Fe3 is more stable because of the ½ filled
3d5 sublevel being more stable than the 3d6
sublevel in Fe2
23
Electronegativity (pg. 161 fig. 20) The
tendency of an atom to attract a pair of
electrons to itself when bonded to another atom
is called electronegativitiy.
increase

• Periodic Trend Electronegativity values
  ___________ to the right across a period (except
  noble gases which dont readily form chemical
  bonds with other elements).
• Group Trend Electronegativity values
  ____________ down a group
• Sample problem G pg. 162
• Based on our observations of ionization energy,
  electron affinity and electronegativity we would
  conclude that __________________ would be the
  most reactive metal and _______________ would be
  the most reactive non-metal.

decrease
Fluorine (F)
Francium (Fr)
Increases UP and to the RIGHT
24

• Homework Day 1 pg. 162 Practice 1 pg.
  167 26 30 34 37 48 for atomic radius
• ionization energy and electronegativity
• 1. Hydrogen can be an atom (H) or ions (H-1) or
  (H1). Arrange these 3 in order of increasing
  size.
• Label one of them cation and one of them anion.
• Day 2
• Define the following
• ionization energy 2. shielding effect 3.
  electron affinity 4. electronegativity 5.
  atomic radius
• Short answer
• 1. Explain why carbon has such a large jump for
  the fifth ionization.
• 2. State four factors that affect ionization
  energy and briefly describe each
• 3. Write the following in order of increasing
  ionization energy.
• a. Be, Mg, Sr, Be2 b. Bi, Cs, Ba c. Na1,
  Al1, Na-1
• 4. Which of the following pairs is more
  electronegative?
• a. chlorine or fluorine b. carbon or nitrogen c.
  arsenic or calcium
• 5. Which of the following has the greatest atomic
  radius?
• a. sodium or lithium b. strontium or magnesium c.
  carbon or germanium
• 6. The first ionization energy tends to
  _____________ as atomic number increases in any
  horizontal row
• or period.
• 7. A column or group will show a decrease in
  first ionization energy as atomic number
  __________.

25

• Review Sheet on Periodic Trends
• What law states that the chemical properties of
  elements repeat in a periodic fashion based on
  their
• atomic number?
• 2. Who came up with the law in 1? 3. How did
  Moseley discover the atomic number?
• 4. What is the modern periodic table arranged
  by _______?
• Mendeleev arranged his periodic table by
  _________?
• Why did Te and I pose a problem for Mendeleev?
• 7. Fill in the following table.

26
8. Which group of elements has
a stable valence shell? The transition elements
are in the _____ block. 9.The
representative elements are in the __________
blocks. 10. The ________ are the
elements in the f block. 11. A
vertical column on the periodic table is called a
_______ or a ________. 12. A
horizontal row on the periodic table is called a
________. 13. Write the electron
configuration for Cu. 14.
__________ is the energy required to remove an
electron from a gaseous atom.
_____ across a period (to the right) and
______ down a group. 15.
___________ is the energy released when adding an
electron to a gaseous atom.
What 2 groups tend to have positive values?
Why? 16. ____ is the calculated
value determined by Linus Pauling for the
tendency of an atom to attract a pair
of electrons. _____ across a period (to
the right) and ______ down a group.
17. The atomic radii _____ left to right across
a period and ______ down a group.
18.Why is there a dip in ionization energy
between P and S? 19. Which
would you predict to be more stable Ru2 or
Ru3? 20. An anion has a
________ charge while a cation has a _________
charge. 21. When an atom loses
electrons it becomes a ___ ion and ___________ in
size. 22. When an atom gains
electrons it becomes a ___ ion and ___________ in
size. 23. Sublevels are most
stable when _______ or ______ full
24. _________ _________ decreases the attraction
for a valence electron for the nucleus
due to electrons existing between the
valence electron and the nucleus.
27
Practice Test
1. Which has a greater ionization energy? a. P
or S
Answer P
(dips from P toS due to half filled sublevel
stability of P - the greater the stability, the
higher the ionization energy)
b. Na or Mg
Answer Mg
(greater nuclear charge for Mg while adding the
electron to the same sublevel so no greater
shielding by inner electrons)
c. K or Na
Answer K
(electrons being added to energy levels that are
further from the nucleus so they feel less
nuclear charge and will be easier to remove, thus
a lower ionization energy)
d. Na or Na1
Answer Na1
(1 charge means an electron has been removed,
the ionization energy will increase for every
electron removed)
28
2. Which has a greater electronegativity? a. H or
He
Answer H
(even though electronegativity tends to increase
to the right, He has no value for IE due to full
1st energy level)
2. Which has a greater electronegativity? b. F or
Cl
Answer F
(electronegativity increases up a period)
2. Which has a greater electronegativity? c. Al
or Cl
Answer Cl
(Cl electronegativity increases to the right)
3. Non-metals tend to have _____________
ionization energy values than metals.
Answer larger (more positive)
4. Non-metals tend to have _____________ electron
affinity values than metals.
Answer more negative
29
5. Name the group IA IIA VIA VIIA and VIIIA
elements.
Answer alkali metals alkaline earth metals
chalcogens halogens
noble gases
6. What was wrong with Mendeleevs table?
Answer Mendeleev arranged elements in order of
increasing atomic mass, which would have some
elements arranged in the improper group by
reactivity.
Give an example off of the modern table that
demonstrates Mendeleevs problem.
Answer Ar and K Co and Ni Te and I
7. What is the name is given to a horizontal row
of the periodic table?
Answer period
8. What are vertical columns of the periodic
table called?
Answer Groups or Families
9. Which is larger? a. Ca or Ca2
b. F or F-1
Answer a. Ca
Answer F-1
30
10. What is the charge of a cation and an anion?
Answer cation (positive) anion (negative)
11a. What is the name of the s and p block (or
groups IA-VIIIA) elements?
Answer representative elements
b. What is the name of the d block elements?
Answer transition metals
c. What is the name of the f block elements?
Answer inner-transition metals
12. What is the electron configuration of Cr?
Answer Ar4s13d5
31
13. What are 4 factors that affect ionization
energy? Describe how each affects ionization
energy.
Answer energy level / atomic radius The higher
the energy level (the greater the atomic radius)
the further the electron will be from the
nucleus so the easier it will be to remove thus a
lower ionization energy.
nuclear charge The greater the nuclear charge
(from left to right across a period), the
greater the ionization energy in general (there
are exceptions)
shielding a greater shielding means there are
more electrons between the nucleus and outer
electrons, the greater the shielding the lower
the ionization energy will be.
sublevel stability full and half filled
sublevels are more stable than those with no
special electron configurations and thus
elements containing full or half fill sublevels
will have a larger ionization energy than
otherwise expected.
32
14. Why is there a large jump in ionization
energy when an element loses all of its valence
electrons?
Answer When an element loses all of its valence
electrons, the next available electron to be
removed is in a lower energy level (closer to
the nucleus) feeling more of the nuclear charge
and will require more energy to remove.
Further, the next electron is much less shielded
due to less electrons existing between it and
the nucleus also making it require more energy
to remove
Also, the next electron will be from a full p
sublevel which is very stable requiring more
energy to remove
(all of these factors contribute to the large
jump in ionization energy compared to the
smaller jumps experienced by merely losing an
electron).
15. What is the name of the energy change
associated with the addition of an electron to a
gaseous atom?
Answer Electron Affinity
33
16. What is the name of the tendency of an atom
to attract a shared pair of electrons?
Answer Electronegativity
17. What is the name of the energy required to
remove an electron from a gaseous atom?
Answer Ionization Energy
18. Why is the electron affinity values positive
for group VIIIA?
Answer Group VIIIA elements have full s and p
sublevels in there valence shell. The addition
of an electron would have to be to a higher
energy level, further from the nucleus
experiencing much less of the nuclear charge.
19. What is most reactive metal by periodic table
location?
Answer Francium (fluorine is the most reactive
non-metal by periodic table location)