Chapter 5
The Periodic Law
I: History of the Periodic Table
by 1860, sixty elements had been discovered
1860, a gathering of chemists in Germany
Cannizzaro
A: Mendeleev and Chemical Periodicity
use the new atomic masses and organize elements according to their properties
similarities in their chemical properties appeared at regular intervals
periodic - definition
Figure 5-1 page 123
elements with similar properties were grouped together
atomic masses as a guide but he let the properties determine the exact grouping of elements e.g. Tellurium (Te) and Iodine (I)
empty spaces for elements
Two questions remained
B: Moseley and the Periodic Law
Henry Moseley working with Ernest Rutherford
arranged in increasing order according to nuclear charge
atomic number
explained the tellurium/iodine question
Periodic Law - statement
C: The Modern Periodic Table
Periodic Table - definition
1. The Noble Gases
1894 - John Strutt and William Ramsay
1868 helium; 1895 - Ramsay
group 18
1898 - Ramsay
1900 - Friedrich Dorn
2. The Lanthanides
lanthanides - definition
similar in chemical and physical properties
part of period 6
3. The Actinides
part of period 7
at the bottom to save space
normally fit between groups 3 and 4
4. Periodicity
Figure 5-4 page 126
The differences in atomic numbers for the two groups follow the same pattern
key is looking for the patterns
Homework: 5.1
II: Electron Configuration and the Periodic Table
Group 18 - octet - stable
Configuration of valence electrons determines the atom's chemical properties
A: Period and Blocks on the Periodic Table
Vertical columns (group or family) have similar chemical properties
Length of each period is determined by the number of electrons that can occupy the sublevels being filled in that period - determined by the electron configuration of the elements
Table 5-1 page 128
1st period 2 elements filling s sublevel
2nd period: 8 elements filling s and p sublevel
3rd period: 8 elements filling s and p sublevel
4th period: 18 elements filling s and d and p sublevels
Figure 5-5 page 129
1. The s Block Elements: Groups 1 and 2
active metals - group 1 being more active than group 2
Group 1 = alkali metals
Group 1: each element has one electron in an s sublevel
Tend to lose that electron to form a +1 ion and achieve an octet
In the elemental state (uncombined) they have a silvery appearance and can be cut with a knife
Not found in nature in the elemental (uncombined) state. Instead they react vigorously with nonmetals and form compounds - the way they are found in nature
React strongly with water to form hydrogen gas and an aqueous solution of alkalis (solutions of a base)
Stored under kerosene to protect them from reacting with the moisture in the air
Moving down the column the melting points decrease
Group 2 = alkaline earth metals
each element has two electrons in an s sublevel
tend to lose two electrons to form a +2 ion and achieve an octet
harder, denser, and stronger than alkali metals
have higher melting points than alkali metals
less reactive than alkali metals
not found in nature in the elemental (uncombined) state
2. Hydrogen and Helium
one electron in the s sublevel
separated from Group 1 metals
Properties of hydrogen do not resemble those of any other group
Helium has two electrons in the s sublevel
part of group 18
ts filled s sublevel gives it special stability
Different than group 2 metals since they have an unfilled p sublevel of the energy level n
3. The d-block Elements: Groups 3-12
After the ns sublevel is filled, electrons go into the (n-1)d sublevel until the d sublevel is filled. After the (n-1)d sublevel is filled the np sublevel is filled.
d-block elements are metals
called transition elements or transition metals
less reactive than alkali metals and alkaline earth metals
some are unreactive and exist in nature in the elemental (uncombined) form e.g. palladium, platinum and gold
4. The p-Block Elements: Groups 13-18
electrons enter the p sublevel only after the s sublevel of the same energy level is filled
main-group elements
properties vary
right side
six metalloids
on the left, bottom
Group 17 - Halogens
most reactive nonmetals
forms salts
fluorine and chlorine
bromine
iodine
astatine
Metalloids or Semiconducting Elements
on red stepped line
mostly brittle solids
metals of p block
generally harder and denser than the s-block alkaline earth metals, but softer and less dense than the d-block metals
except for bismuth, found in nature in combined state
when purified as pure element, they are stable in the presence of air
The f-Block elements - Lanthanides and Actinides
between groups 3 and 4
fourteen f-block elements in each series - how many electrons can an f block hold?
similar in reactivity to the group 2
actinides are all radioactive
the first four of the actinides vs remaining actinides
Homework: Chapter 5: 5.2
III: Electron Configuration and Periodic Properties
periodic law and the electron configurations
A: Atomic Radii
atomic radius - definition
1. Periodic Trends
figure 5-13 page 141
figure 5-14 page 142
across second period
2. Group Trends
group 1
group 13
interplay between the distance of the added electrons from the nucleus and the increased number of protons in the nucleus
B: Ionization Energy
A + I.E. ---> A+ + e-
Ion - definition
Ionization - definition
compare ease with which atoms give up electrons
ionization energy - definition
figure 5-15 page 143
page 144
1. Period Trends
First element in series vs last element - for ionization energy
Low I.E. indicates
High I.E. Indicates
trend across period and why
ionization energy of metals vs nonmetals
2. Group Trends
trend for main group down family and why
3. Removing Electrons from Positive Ions
ions that result from applying first, second and third ionization energy
Size of first vs second vs third ionization energy and why
table 5-3 page 145
ionization energy of group 18 vs other groups
Difference between first and second I.E. of Li
Difference between second and third I.E. Of Be
Difference between I.E. Of beryllium and boron
Difference between I.E. Of nitrogen and that of oxygen
C: Electron Affinity
Electron Affinity - definition
A + e- ---> A- + electron affinity (energy)
Some atoms need to absorb energy to accept the electron
A + e- + electron affinity (energy) --->A-
The ions formed are unstable
figure 5-17 page 147
In tables not equations: energy will have a negative value for exothermic reactions; energy will have a positive value for endothermic reactions
1. Periodic Trends
Halogens gain electron most readily - size of electron affinity?
Why does giving off a large amount of energy indicate the atom wants the electron?
Across the p block of any series, adding an electron produces greater negative values.
Not the case between group 14 and group 15. Why?
2. Group Trends
not as regular as for I.E.
down the group what happens and why?
Many exceptions in the transition metals
3. Adding Electrons to Negative Ions
add an electron to a negative ion - repulsion
all second electron affinities are positive
D: Ionic Radii
Cation - definition
Radius of the ion vs radius of atom and why
Anion - definition
radius of anion vs radius of atom and why
Figure 5-19 page 149
1. Periodic Trends
ions that metals tend to form and nonmetals tend to form
Cationic radii decrease across a period - why
Beginning with group 15, anions tend to form
Anionic radii decrease across the period - why
2. Group Trends
Gradual increase in ionic radius down the group - why
E: Valence Electrons
Valence Electrons - definition
Usually found in incompletely filled main energy levels
Table 5-4 page 150
Elements in groups 13-18 have a number of valence electrons equal to the group number minus 10
F: Electronegativity
valence electrons hold atoms together in chemical compounds
Valence electrons are not always mid-way between the two atoms
This affects the chemical properties of the compound
Electronegativity - definition
fluorine
relative scale with 4.0 being highest
1. Periodic Trends
figure 5-20 page 151
increase going across the periods - generally
active metals are the least electronegative
active nonmetals are the most electronegative
Down a group the electronegativities tend to decrease down a group or remain about the same - why
Some noble gases do not form compounds and do not have electronegativities assigned - why
Highest electronegativities are in the upper right of the periodic table - why
Figure 5-21 page 152
G: Periodic Properties of the d and f Block Elements
d block elements vary less and with less regularity than the main group elements. e.g. flat curves in figure 5-14 and 5-16
For d block elements, electrons in both the Ns and (n-1)d sublevel are available to interact with the surroundings
Electrons in the incompletely filled d sublevels are responsible for many characteristic properties of the d block elements
1. Atomic Radii
generally decrease across a period
less than that for main group elements
in figure 5-14 the radii decrease then increase slightly across each of the four periods that contain d block elements
repulsion
2. Ionization Energy
ionization energy generally increase across the periods
ionization energy increases going down group
shielding of s electrons by d electrons
3. Ion Formation and Ionic Radii
Order of removing electrons.
Highest energy level then highest energy sublevel where s < p < d < f
e.g. Fe
Most d block elements form 2+ ions - why
Cations have smaller radii than atoms - why
4. Electronegativity
Values of electronegativity increase as radii decreases and vice versa
Summary of Periodic Trends
Homework: 5.3
end of notes
Definitions
Periodic means recurring at regular intervals.
Periodic Law state the physical and chemical properties of the elements are periodic functions of their atomic numbers.
The periodic table is an arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column or group.
The lanthanides are the fourteen elements with atomic numbers from 58 to 71.
The actinides are the fourteen elements with atomic numbers from 90 to 103.
Ionization energy is the energy required to remove on electron from a neutral atom of an element.
An ion is an atom or group of bonded atoms that has a positive or negative charge.
Any process that results in the formation of an ion is referred to as ionization.back
Atomic Radius is one-half the distance between the nuclei of identical atoms that are bonded together.
Electron affinity is the energy change that occurs when an electron is acquired by a neutral atom.
A cation is any positive ion.
An anion is any negative ion.
Valence electrons are those electrons available to be lost, gained, or shared in the formation of chemical compounds.
Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons.