Assignment 10/24
• Read Chapter 6, section 1. You have 10
minutes.
• Answer the questions below:
– 1. How is the modern periodic table arranged?
Who organized elements with this arrangement?
– 2. Who arranged the periodic table by increasing
atomic mass?
– 3. What is the name of Group 18 elements?
11/8 Activity
• 1. Write the electron configuration of your element.
• 2. Name the group (if applicable) that your element falls in.
List all properties of that group. Be detailed!
• 3. How many valence electrons does your element have in the
outer shell.
• 4. Draw the Lewis dot structure for your element.
• When finished…transfer all the information about the
elements and group to a piece of butcher paper. Show me
your work and hang it on the wall somewhere. We will do a
gallery walk of the groups at the end of the period today or
beginning tomorrow.
11/11 Today you will need…
• A clean sheet of paper, something to write
with, and the paper from the side table.
• We will do a gallery walk activity using the
information from Friday about the groups on
the periodic table.
• You will color code a periodic table when the
gallery walk is complete
Gallery Walk…
• When finished, answer the following questions on a clean
piece of paper and turn in:
– 1. Name Group 1, Group 2, Group 17, and Group 18 elements.
– 2. Identify the properties of each of the groups listed in number 1.
– 3. Determine the number of valence electrons in groups
1,2,13,14,15,16,17,18.
– 4. Draw the electron configuration for ONE element from each group
listed in number 1.
Chapter 5The
Periodic
Law
5.1-History of the Periodic Table
5.2-Electron Configuration & the Periodic Table
5.3-Electron Configuration & Periodic Properties
5.1-History of the Periodic Table
Pages 123-127
Mendeleev
• Dmitri Mendeleev (1869, Russian)
– Organized elements
by increasing
atomic mass.
– Elements with
similar properties
were grouped
together.
– There were some
discrepancies.
Mendeleev
• Dmitri Mendeleev (1869, Russian)
– Predicted properties of undiscovered elements.
Moseley
• Henry Moseley (1913, British)
– Organized elements by increasing atomic number.
– Resolved discrepancies in Mendeleev’s
arrangement.
– Periodic Law-the physical and chemical properties
of the elements are periodic functions of the
atomic numbers.
Organization of the Elements
• Periodic table-an arrangement of the
elements in order of their atomic numbers so
that elements with similar properties fall in
the same column, or group.
Additions to Mendeleev’s Periodic Table
• Noble gases
– Group 18
– Argon discovered in 1894
– Took so long to discover because very unreactive
• Lanthanides
– 14 elements with atomic numbers from 58-71
– Placed below the periodic table to conserve space
• Actinides
– 14 elements with atomic numbers 90-103
– Also placed below periodic table
Warm-Up 11/11
• Define Periodic Law.
• Where are the Lanthanides and Actinides
located on the periodic table.
• Describe, specifically what contributions
Mendeleev and Moseley made towards the
creation of the periodic table. (i.e dates,
arrangement, etc.)
5.2-Electron Configuration &
the Periodic Table
Pages 128-139
Periods & Blocks of the Periodic Table
• Length of period (row) determined by how
many electrons can occupy the sublevels
being filled.
– 1st period-1s sublevel being filled with 2 electrons
 2 elements, H & He
– 3rd period-3s & 3 p sublevels being filled with 2+6
electrons  8 elements
• Periodic table is divided into “blocks” based
on the filling of sublevels with electrons.
Blocks of the Periodic Table
Determining Period from Configuration
• An element’s period can be determined by
looking at its electron configuration
• The highest occupied energy level
corresponds to the element’s period
– As: [Ar]3d104s24p3
• 4 in 4p3 indicates that the highest energy level that
electrons occupy is the 4th. Therefore, As is located in
the 4th period of the periodic table.
Metallic Character
• Metals
• Nonmetals
• Metalloids
Areas of the Periodic Table
• Main Group Elements
• Transition Metals
• Inner Transition Metals
s-Block Elements: Groups 1 & 2
• Chemically reactive metals
• Include the alkali metals and the alkaline earth
metals
Alkali metals
•
•
•
•
•
Group 1 metals
ns1
Silvery appearance and very soft
Not found pure naturally because so reactive
Because of extreme reactivity with moisture,
usually stored under kerosene
• Video: Disposal of Surplus Sodium
• Video: Alkali Metals in Water
Alkaline-Earth metals
•
•
•
•
Group 2 metals
ns2
Harder, denser, & stronger than alkali metals
Also too reactive to be found free in nature
(but less reactive than Gp. 1)
• Video: Magnesium/silver nitrate mixture
reacting with water
d-Block Elements: Groups 3-12
• Metals with typical metallic properties
• Called “transition elements”
• Typically less reactive than Gps. 1&2, & some are extremely
unreactive
• d sublevels first appears at the 3rd energy level
• Fills after 4s
• Variations from expected in d-block, so elements in the same
group do not necessarily have the same outer e- configuration
p-Block Elements: Groups 13-18
• p and s-block elements together called “maingroup elements”
• Total number of electrons in highest energy
level=group # - 10
– Group 17 elements have 17-10=7 outer “valence”
electrons
• Properties of p-block elements vary greatly
since metals, nonmetals, and metalloids are
contained here
p-block Elements
• Halogens
– Group 17 nonmetals
– Most reactive nonmetals
• React with most metals to form salts
• Metalloids
– Fall on both sides of a “stair-step” line separating
metals and nonmetals
– Semi-conductors
f-Block Elements:
Lanthanides & Actinides
• Lanthanides
– Top row of f-block
– 14 elements
– Shiny metals similar in reactivity to the alkaline-earth
metals
• Actinides
–
–
–
–
Bottom row of f-block
14 elements
All radioactive
1st 4 elements found naturally on Earth; remainder
only lab-made elements
5.3-Electron Configuration &
Periodic Properties
Pages 140-154
Remember the Periodic Law
• When elements are arranged in order of
increasing atomic #, elements with similar
properties appear at regular intervals.
Atomic Radius (pm)
250
200
150
100
50
0
0
5
10
Atomic Number
15
20
Warm Up 11-12
Explain the joke:
Yo momma so fat if she was on
Mendeleev’s periodic table she
would be at the end.
Atomic Radius
½
the distance between the nuclei of
identical atoms that are bonded together
1
2
3
4
5
6
7
Atomic Radius
Atomic Radius (pm)
250
K
200
Na
Li
150
100
50
Ar
Ne
0
0
5
10
Atomic Number
15
20
Ionization Energy
• First Ionization Energy- the energy required to
remove one electron from a neutral atom
1
2
3
4
5
6
7
Ionization Energy
1st Ionization Energy (kJ)
• First Ionization Energy
He
2500
Ne
2000
Ar
1500
1000
500
Li
Na
K
0
0
5
10
Atomic Number
15
20
Ionization Energy
• Why opposite of atomic radius?
–In small atoms, e- are close to the
nucleus where the attraction is
stronger
• Why small jumps within each group?
–Stable e- configurations don’t want to
lose e-
Ionization Energy
• Successive Ionization Energies
 Large
jump in I.E. occurs when a CORE e- is
removed.
 Mg
Core e-
1st I.E.
736 kJ
2nd I.E.
1,445 kJ
3rd I.E.
7,730 kJ
Ionization Energy
• Successive Ionization Energies
 Large
jump in I.E. occurs when a CORE eis removed.
 Al
Core e-
1st I.E.
577 kJ
2nd I.E.
1,815 kJ
3rd I.E.
2,740 kJ
4th I.E.
11,600 kJ
Electron Affinity
• Energy change that occurs when an electron is
acquired by a neutral atom
• Tends to become less negative (less energy
released) DOWN and to the LEFT
1
2
3
4
5
6
7
Electronegativity
• A measure of the ability of an atom in a chemical
compound to attract electrons
• Most electronegative element is fluorine
– Given arbitrary value of 4; all others relative
1
2
3
4
5
6
7
A
Label the Groups:
B
E
F
G
C
D
Atomic Radius
Q: Why larger going down?
A: Higher energy levels have larger
orbitals
Q: Why smaller to the right?
A: Increased nuclear charge without
additional energy levels to shield the
electrons pulls them in tighter.
Ionization Energy
Q: Why is Ionization Energy
greatest in upper right corner?
A: In small atoms, e- are close
to the nucleus where the
attraction is stronger. More
protons = more positive
charge.
Ionic Radius: The distance
from the nucleus to the outer
orbital of an ion.
Atoms with an
Ions
electric charge
Cations
+ charge
Smaller
Lose
electrons
Anions
- charge
Bigger
Gain
Electrons
Ionic Radius
Q: Why are cations smaller than the
neutral atom?
A: When cations lose their electrons
they also lose an energy level.
Q: Why are anions larger than the
neutral atom?
A: As they gain electrons there is more
repulsion between them, which pushes
them farther from the nucleus.
Examples
• Which atom has the larger radius?
Be
or Ba
Ca
or Br
Examples
• Which atom has the larger radius?
Be
or Ba
Ba
Ca
or Br
Ca
Examples
• Which atom has the higher 1st I.E.?
N
or Bi
Ba
or Ne
Examples
• Which atom has the higher 1st I.E.?
N
or Bi
N
Ba
or Ne
Ne
Examples
• Which has the greater electonegativity?
K
or Li
Al
or Cl
Examples
• Which has the greater electonegativity?
K
or Li
Li
Al
or Cl
Cl
Examples
• Which particle has the larger radius?
S
or S2-
Al
or Al3+
Examples
• Which particle has the larger radius?
S
or S2-
S2-
Al
or Al3+
Al