Transcript Chapter 5 - Midway ISD

Chapter 5
The Periodic Law
Sect. 5-1: History of the
Periodic Table
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Stanislao Cannizzaro (1860) proposed
method for measuring atomic mass at First
International Congress of Chemists
Dmitri Mendeleev (1869) arranged elements
by atomic mass & similar chemical
properties; left blanks for undiscovered
elements
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Henry Moseley (1911) arranged periodic
table by atomic number
Periodic law – properties of elements are
periodic functions of their atomic #’s
Noble gas group, lanthanide and actinide
series added later
Sect. 5-2: Electron Configuration
and the Periodic Table
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S-block elements are highly reactive metals
because they easily give up their 1 or 2
valence electrons
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Group 1 – Alkali metals
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Silvery, can be cut with knife
Group 2 – Alkaline Earth metals
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Harder, denser, stronger, and slightly less reactive
than group 1
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Special cases:
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Hydrogen grouped with 1 because of electron
configuration, but doesn’t share their properties
Helium is grouped with 18 because it has similar
properties since its outside energy level is full,
even though it has the same electron
configuration as group 2
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D-block elements
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Total # electrons in d plus electrons in highest s
orbital = group #
Referred to as transition elements
Good conductors of heat/electricity
High luster
Not as reactive as s-block elements
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P-block elements
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Combined with s-block they are called main-group
elements
Contains metals, non-metals, and metalloids, thus
wide range of properties
Group 17 – halogens
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Most reactive nonmetals
Group 18 – noble gases
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nonreactive
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F-block elements
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Lanthanides
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Shiny, similar in reactivity to group 2
Actinides
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All radioactive
First 4 have been found naturally, all others are manmade
Sect. 5-3: Electron Configuration
and Periodic Properties
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Atomic Radius – one half the distance
between the nuclei of chemically bonded
identical atoms
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Decreases from left to right across a period due to
higher positive charge on right pulling electrons
closer
Increases going down a group because of adding
energy levels
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Ion – charged particle
Ionization energy (IE) – energy required to
remove an electron from a neutral atom in the
gas phase
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Increases as you move to the right because those
elements will less readily give up an electron
Decrease as you move down a group due to
electrons being further away from nucleus and
shielded by inside electrons
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2nd and 3rd ionization energies refer to
removing additional electrons from positively
charged ions
2nd and 3rd Ionization energies have a drastic
“jump” if the ion has the electron
configuration of a noble gas
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Electron Affinity – energy change when a
neutral atom gains an electron
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Reported as a negative # because of loss of
energy
Generally decreases as you move down a group
Generally decreases as you move left on a period
Exceptions for half-filled or filled sublevels
Adding additional electrons will always have a
positive value (requires energy)
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Ionic Radii
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Cation – positively charged ion (lost electron)
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Anion – negatively charged ion (gained electron)
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Will decrease radius because of loss of outer energy
level
Will increase radius because protons “pulling in” are
the same and with extra electrons they repel each
other and spread out
Cation & anion radius increases from the right to
the left across a period
Cation & anion radii increase down a group
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Valence Electrons – electrons in outermost
energy level (can be gained lost or shared)
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For s-block, # valence electrons is equal to group
number
For p-block, # valence electrons is equal to group
number minus 10
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Electronegativity – measure of ability of an
atom in a compound to attract electrons
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Generally decrease as you move to the left of a
period
Generally decrease or stay the same moving
down a group
Nitrogen, Oxygen, and Halogens are most
electronegative
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Trends for d- and f-blocks
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atomic radius trend is same a main group, but
with smaller changes
Ionization energy trend is same for period, but
increases going down a group
Ion formation – electrons are removed from the s
orbital 1st, then the d
Electronegativity trends are same