Transcript Chapter 6 The Periodic Table and Periodic Law

UEQ
What else does the
Periodic Table tell us
about our atoms?
LEQ
What did the early
Periodic Tables look
like?
Chapter 6
The Periodic Table and Periodic
Law
A. Development of the Periodic Table (Read
details in the text)
1. Antonie Lavoisier: Compiled a list of 23
known elements.
2. John Newland: Arranged known
elements by mass resulted in a repeat of
properties every eight element. He called
this the Law of Octave.
3. Lothar Meyer and Dmitri Mendeleev: Like
Newland, arranged known elements in
increasing atomic mass. Mendeleev then
published and is given credit for the first
periodic table.
4. Henry Moseley: Arranged elements in
increasing atomic numbers (the # of
protons in the nucleus). The arrangement
still repeated in a regular pattern and is
called the periodic law.
LEQ
What are the different
parts of the Periodic
Table?
Periodic Table
• s, p, d, f
• s and p are representative elements
• # of electrons in valance
• Periods and Groups (Oxidation #,s)
• # of electrons in transitional elements
• Dot configuration
• Hybrid formation
• Metals (loss e-, cation)
• Non-metal (gain e-, anion)
The Modern Periodic Table
Metals (ion forms)
Non-metals (ion forms)
Periodic Box
Representative elements
Transitional elements
Alkali metals
Alkaline earth metals
Lanthanide series
Actinide series
Noble gases (inert gases)
Metalloids
Halogens
Periodic Trends
1. s, p, d, and f block
2. Valance electrons (fig 6.7)
a. electron configurations (shorthand)
b. ionic configuration
3. Dot configurations
4. Chemical and physical properties
based on repeatable trends.
Review Sampler
1. Who gave us the Law of Octave?
ans: John Newland
2. Who is the ‘Father of the Modern Periodic Table’?
ans: Henry Mosley
3. Who gave us the ‘Periodic Law’
ans: Henry Mosley
4. What does the Periodic Law state?
ans: properties of elements repeat in a regular pattern of
eight based on the atomic number of the element.
5. Who proposed a periodic table based on the atomic mass?
ans: Dmitri Mendeleev
Review Sampler
6. What part of the periodic table will you find the
representative elements?
ans: In the s and p block, Group A elements
7. Group 3A has how many valance electrons?
ans: 3, s2 p1
8. What is the oxidation number for Cl ?
ans: 1-
Review Sampler
9. What is the oxidation number for Mg and Sr ?
ans: both are 2+
10. Write the dot configuration for
Al
Mg
C
11. What is a cation? Give an example.
ans: + ion or metal. Any metal like Fe or K.
12. What is an anion? Give and example.
ans: - ion or non-metal. Any non-metal like Cl or O
Review Sampler
13. What Group # will we find the Halogens?
ans: Group 7A
14. Given an example of an alkali metal?
ans any element in Group 1A
15. Given an example of a metalloid ?
ans: Si, Ge, As, Sb, Te, Po, At
16. In which ‘series’ will you find Uranium?
ans: Actinium series
17. Why are the Noble gases also called inert?
ans: each has a filled valance shell with 8 electrons
18. Which has a 2+ oxidation number?
Na
Cr Ca N Cl
19. Write the oxidation number for the following:
O
K
Mg
Ni
Fe
Br
Ba
Li
S
B
P
Xe
LEQ
How are the different
trends in the Periodic
Table interpreted?
5. Atomic Radius: One-half the distance between
two adjacent nuclei. (fig 6.11 and 6-12)
a. Within the A Groups:
• Atomic radius decrease left to right
within a period.
Reason: the number of valance electrons
and the atomic number (the # of
positive charged particles) increase
left to right and the number inner
electrons (shielding effect) remains
the same.
Atomic radius increase from top to
bottom within a group.
Reason: the number of valance
electrons remains the same while
the number of inner shielding
electrons increases.
Atomic Radii
Ionization energy
Ionization energy: the amount of energy required to remove an
electron from a gaseous atom. Electrons removed are from
the outer most portion of the electron cloud.
• Trends: (fig 6-16 and 6-17)
• From left to right across within a single
period, the amount of ionization energy
increases (takes more and more energy to
remove an electron). How does this relate to
the atomic radius for the same atoms?
• Top to bottom within a group, the ionization
energy decreases (takes less energy to
remove an electron). How does this relate to
the atomic radius for the same atoms?
• Note: Table 6-5, page 192. What are the trend
in the values for the 1st to the 2nd to the 3rd
ionization energy relative to the position of the
atom on the periodic table?
Ionization Energy
Ionic Radius
A. The size of the atom as a result of the gain or loss of electron(s). Or
when an ion is formed. (fig 6-14 and 6-15)
•
Within the A Group:
• For positive ions, the ionic radius is
always smaller than the atomic radius
for the same atoms.
• For negative ions, the ionic radius is
always larger than the atomic radius
for the same atom.
Trend:
- From left to right within a single
period, the ionic radius decreases
according to the charge.
- From top to bottom within a single
group, the ionic radius increased
according to the charge.
Isoelectroic ions: same configuration
w/ different atoms
Electronegativity
Electronegativity: the ability of an atom to attract the electrons of
another atom.
The values range from less than one to 4.0. The larger the value, the
more likely the atom will attract electrons.
Fluorine is the most electronegative with a value of 4.0 and Francium
is the least electronegative with a value of 0.7.
Trends: The trends are more general
than other trends. (Fig 6-18)
• From left to right within a single
period, the electronegativity tends
to increase.
• Top to bottom within a single
group, the electronegativity tends
to decrease.
Calculate EN value: Used to define
type of bonds.
Electron Affinity
The amount of energy absorbed when an electron is added to an
atom to form an ion with a 1- charge.
* Elements with very negative electron affinities gain
electrons easily to form a negative ion (anion).
- non-metal >> metals
* A negative value for electron affinity will indicate that
energy is released.