Download Notes Chapter 5 “The Periodic Law” Section 1 Dmitri Mendeleev

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Notes Chapter 5 “The Periodic Law”
Section 1
Dmitri Mendeleev – father of the modern periodic
table. Listed elements by masses and properties.
Left a column for the noble gases. Publishes in 1869.
Henry Mosely working with Rutherford in 1911 –
helped come up with the “periodic law” – physical
and chemical properties are functions of the atomic
Today periodic table is arranged according to atomic
number and properties.
1894 Strutt and Ramsay discovered Argon the first
of the noble gases.
The Lanthanide – All end in 4f, numbers 58-71,
period 6.
Actinides – all end in 5f, number 90-103. Belongs in
period 7.
**to save space the lanthanides and actinides are
set off below the main table.
8, 8, 18,18, 32 –periodicity
Section 2
See table 1 and figure 5.
S block – groups 1 and 2 – group 1(alkali metals- end
in s1, not found in nature as free elements. Always
found in compounds. React strongly with water. Soft
silvery appearance can be cut with a knife, will have
a +1 charge). Group 2( alkaline earth metals, end in
s2, harder, denser, and stronger than group 1, not as
reactive as group 1, will have a +2 charge).
Hydrogen is not an alkali metal, it is located above
the group because of it electron configuration.
Helium possesses special chemical stability, and is
unreactive so it is above group 18.
D-block – transition elements all end in d. Metals
with typical metallic properties. They are good
conductors, electricity, and have luster. Less
reactive than group 1 and 2. They do not easily
form compounds. Will have a + charge.
P block – all end with P, mixed group. Metals form +
ions and nonmetals form – ions. Group 17(halogens)
most reactive group in the p block, have -1 charge.
Main- group elements- s and p block.
Section 3
Atomic Radii- half the distance between the
nuclei of atoms that are bonded together.
Atomic radii decreases from left to right and
across a period and increase down a group.
Ionization energy – energy required to remove
an electron. Ionization increase as electrons are
removes ex. 1st electron removed has the
lowest ionization energy. See figure 15
Electron Affinity – energy required to acquire
and electron. Most atoms release energy when
they acquire an electron. Some atoms have to
be forced to gain electrons, this will require
energy. See figure 17
Ionic Radii –
*cation – positive, usually metals, smaller
radius, because electrons have been removed.
*anion – negative, usually nonmetals, larger
radius, because electrons have been gained.
Valence electrons – outer most s and p; gained, lost,
or shared to form bonds. The P must follow the S.
Electronegativity – atoms ability to attract electrons
from another atom in a compound. See figure 20.
Tends to increase across a period.