Transcript Periodic Table

Periodic Table –
Organizing the
Elements
Chapter 5.4
& Chapter 14
Dmitri Mendeleev
About 70 elements had
been found by the mid
1800’s
Mendeleev was the first to
organize them in a
systematic way
He listed the elements in
order of increasing
atomic mass
Arranged the elements
in columns so those with
similar properties were
side by side
He left blank spaces
where nothing fit
He predicted the
physical properties of the
missing elements
He was mostly correct
Henry Moseley
Moseley determined the
atomic # of the elements
and arranged the table
by atomic number
instead of atomic mass
The modern periodic table
is arranged by atomic
number
The periodic table has
atomic # increasing from
left to right & top to bottom
Periodic Law
The horizontal rows on
the periodic table are
called periods
Properties change as
you move across a
period
The properties repeat
when you move from one
period to the next
Periodic Law: there is a
periodic repetition of the
chemical & physical
properties of the elements
Groups
Each vertical column is
called a group or family
Elements in the same
group have similar
properties
Groups have a number and
a letter (pg 124)
The group with Li, Na, K etc
is called Group 1A
Group 1A elements are also
called the alkali metals
All group A elements are
called the representative
elements
They exhibit a wide range
of physical & chemical
properties
Elements on the left side
of the periodic table
(except for hydrogen) are
metals
Group 2A are the
Alkaline Earth Metals
Group B elements are the
transition & inner-transition
metals
Gold & silver are transition
metals
Uranium is an innertransition metal
The upper right hand
corner of the table has
the non-metals
Some are gases, some
are solids & some are
liquids at room
temperature
Bromine is a liquid,
Oxygen is a gas and sulfur
is a solid
Group 7A are called the
Halogens (F, Cl, Br, I)
Group 0 (8A) are the noble
gases (He, Ne, etc)
Metals have a shiny
appearance (luster) & are
good conductors of heat &
electricity, most are solids
Nonmetals do not have
luster & are poor
conductors
Elements bordering the
step-line are called
metalloids or semi-metals
Si & Ge are metalloids
They have properties inbetween metals &
nonmetals
The elements can also be
classified by their electron
configuration
Electrons play the most
important part in determining
the properties of elements
Write the electron
configurations for the Alkali
Metals
What similarities do you
see?
The Halogens?
The Noble Gases?
The noble gases have their
outermost s & p sublevels
filled completely
The Representative
Elements have their
outermost s & p sublevels
partially filled
The Transition Metals – their
outermost s & nearby d
sublevels contain electrons
The Inner Transition Metals
– their outermost s & nearby
f sublevels contain electrons
The Table can be broken up
into blocks - tell you the
outermost sublevels that are
filled
s block, p block, d block & f
block
Where are they?
Each period on the Table
corresponds to a principle
energy level being filled
# electrons can be determined
by counting left to right
d block is one less than the
period, f block 2 less
The electron configuration
can be determined for most
of the elements this way
“S” block
Groups 1 & 2
Electron Configuration ends in an
S Sub-level.
Highest energy level is equal to
the period number of the element.
–i.e. Calcium’s (in the 4th period)
electron configuration ends in 4s.
“P” block
Groups 13 thru 18
Electron Configuration ends in a P
Sub-level.
Highest energy level is equal to
the period number of the element.
–i.e. Silicon’s (in the 3rd period)
electron configuration ends in 3p.
“D” block
Groups 3 thru 12
Electron Configuration ends in a D
Sub-level.
Highest energy level one less than
the period number of the element.
–i.e. Silver’s (in the 5th period)
electron configuration ends in 4d.
“F” block
“Inner Transition Metals”
Electron Configuration ends in an
F Sub-level.
Highest energy level two less than
the period number of the element.
–i.e. Uranium’s (in the 7th period)
electron configuration ends in 5f.
Regardless of the “Block,” the number
of electrons in the highest sub-level is
equal to the element’s column number
within its block.
–Ex: Nitrogen is in the 3rd column of
the p block and its configuration
ends in p3.
–Ex: Iron is in the 6th column of the d
block and it ends in d6.
Examples:
Determine the last term in the
electron configurations of the
following elements:
–Chlorine:
3p5
–Potassium:
4s1
–Mercury:
5d10
Stable electron
configurations
1. Ending in a full p sublevel or 1s
– Noble gases are the most stable
2. Ending in a full sub-level other
than a p or the 1s.
– Magnesium (3s2) is more stable
than Sodium (3s1).
3.Ending in a half-filled multi-
orbital (p, d or f) sub-level.
– Nitrogen (2p3) is more stable
than Carbon (2p2) or Oxygen
(2p4).
– Stability due to un-paired
electrons in orbitals having same
“Spin.”
Periodic Trends
An element’s placement in the
periodic table determines
characteristics like the size of the
atom, its ability to attract electrons
and the stability of its electron
configuration.
Atomic Radius
Size of atoms of each element:
–How will the size of atoms change
as we proceed down a group?
i.e. Compare the sizes of Li and Na.
From Li to Na, we add an entire
energy level, therefore the size
increases.
How will the size of atoms change
as we proceed across a period?
–Compare C, N and O. Which is
largest?
Oxygen has the most electrons.
However, it also has the most protons.
The outermost electrons of Oxygen
are in the same sub-level as C and N.
Oxygen’s greater nuclear charge
attracts the electrons, causing the
atom to contract!
Oxygen is the smallest of the
three, Carbon is the largest.
Atomic Radius decreases as we
go across a period from left to
right and up a group.
Examples
Rank the following sets in order of
decreasing Radius.
–S, Cr, Se, Sr, Ne
Sr, Cr, Se, S, Ne
–Fe, N, Ba, Ag, Be
Ba, Ag, Fe, Be, N
Ionic Size vs. Atomic Size
When an atom becomes an ion, it will
either get smaller or larger
Metals lose electrons and will get
smaller (stronger pull from the
nucleus)
Nonmetals gain electrons and will get
bigger (more e- to repel one another)
Ionic Size vs. Atomic Size
Which is bigger?
–Na or Na+
–Cl or Cl–O or O-2
–Mg or Mg+2
Ionization Energy
Amount of energy required to
remove a valence electron from an
atom.
The more stable an element is, the
harder it will be (more energy is
required) to remove an electron.
Some elements become more
stable by losing an electron so they
lose electrons easily (less energy
needed).
How does ionization energy vary
within a group (compare Li and
Na)?
–The electron to be removed from Na
is further from the nucleus than
Lithium’s electron.
–Sodium’s electron is held more
loosely and therefore easier (less
energy) to remove.
How does ionization energy vary
across a period? (Compare
elements in 3rd period)
–Sodium attains a Noble Gas
configuration by losing an electron,
so little energy is required.
–Magnesium is somewhat stable due
to a full 3s sub-level, so more
energy is needed.
–Argon is a Noble Gas. Due to its
stability, it is very difficult (much
energy needed) to remove an
electron.
–Chlorine has no stability in its
configuration, so it is easier to
remove an electron.
Ionization energy increases
across a period and up a group.
Examples
Rank the following sets in order of
decreasing Ionization Energy.
–K, Zn, Cs, Ar, P
Ar, P, Zn, K, Cs
–C, He, Ag, Pt, Sn
He, C, Sn, Ag, Pt
nd
2
nd
2
Ionization Energy
The
ionization energy is the
amount of energy required to
remove the second electron on the
outside of an atom
Sometimes it is larger than the 1st
ionization NRG, sometimes, it is
smaller
nd
2
Ionization Energy
For elements like Na and the alkali
metals, the 2nd ionization NRG is
much higher than the 1st
WHY???
Na loses 1 e- and becomes like a
noble gas.
Losing the 2nd would be counterproductive and will not happen easily!
nd
2
Ionization Energy
For elements like Mg and the
other alkaline earth metals, the 2nd
ionization NRG is lower than the
1st
Losing 1 e- is relatively difficult
because of the s2 configuration
(somewhat stable) but losing the
next e- is super easy
nd
2
Ionization Energy
st
1
How do you think the
ionization
NRG and the 2nd ionization NRG
compare for the halogens?
The Noble gases?
WHY???
Electronegativity
Describes an element’s attraction
for an electron in a covalent bond.
Elements that need electrons to
complete an energy-level will have
a high electronegativity.
Elements that want to lose
electrons have low
electronegativities.
How does Electronegativity vary
within a group? (compare F and
Cl)
–Both elements need an electron to
complete a p sub-level.
–Fluorine’s p sub-level is closer to its
nucleus, so it has a greater
magnetic attraction for a free
electron.
–F has a higher electronegativity!
How does electronegativity vary
across a period? (period 2)
–Fluorine benefits the most by
gaining an electron, so it has the
highest electronegativity.
–Lithium, which wants to lose an
electron has very little attraction for
an additional electron.
–Carbon can gain electrons but
sometimes loses them as well, so
its electronegativity is between F
and Li.
–Noble Gases have no
Electronegativity!
Electronegativity increases across
a period and up a group (Noble
Gases omitted).
Examples
Rank the following sets in order of
decreasing Electronegativity:
–Cu, F, Mn, Sr, Si
F, Si, Cu, Mn, Sr
–Al, Ca, S, Cl, Fe
Cl, S, Al, Fe, Ca