Download Periodic Table – Organizing the Elements

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
Metals Vs Nonmetals
The table can be divided into 3 main
types of elements – metals,
nonmetals & metalloids
Metals reside on the left side of the
table
Nonmetals are on the upper right side
of the table
Metalloids are b/tw the metals &
nonmetals
Metals are typically shiny (luster) &
are good conductors of heat and
electricity
Metals are all solids, except Mercury
(Hg) which is a liquid
Nonmetals are usually dull (no luster)
& are poor conductors of heat and
electricity
Nonmetals can be solid (C & S), liquid
(Br) or gas (O & H)
Metalloids
Metalloids are their own category of
elements – there are only 7 metalloids!
–B, Si, Ge, As, Sb, Te, & At – they
border the stair-step line that most
periodic tables include – but not all.
Metalloids have properties in b/tw the
metals & nonmetals
–Si is used in computers b/c it is a
semi-conductor
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
called the alkali metals
Alkali metals are the MOST
active metals
Group 2A are the
Alkaline Earth Metals
Alkaline Earth Metals are
not as active as the alkali
metals, but react with
many substances
Group 7A are called the
Halogens (F, Cl, Br, I)
Halogens are the most active
nonmetals
Group 0 (8A) are the noble
gases (He, Ne, etc)
The noble gases are mostly
UNREACTIVE!
Groups 3A to 6A
Groups 3A to 6A do not have specific
names – they are usually named by
the top element of the group.
For instance, The Carbon family is
Group 4A.
These groups can contain metals,
metalloids and nonmetals!
All group A elements are
called the representative
elements
They exhibit a wide range
of physical & chemical
properties
Group B elements are the
transition & inner-transition
metals
Gold & silver are transition
metals
Uranium is an innertransition metal
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
Blocks
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?
“S” block
Groups 1 & 2
Alkali Metals and Alkaline Earth
Metals
Electron Configuration ends in an
S Sub-level.
“P” block
Groups 13 thru 18 (or 3A - 8A)
Includes halogens and noble
gases (except He)
Electron Configuration ends in a P
Sub-level.
“D” block
Groups 3 thru 12
Transition Metals
Au, Ag, Fe, Pt, etc
Electron Configuration ends in a D
Sub-level.
“F” block
“Inner Transition Metals”
Includes Uranium
Electron Configuration ends in an
F Sub-level.
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
increases going down 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
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 & decreases down
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
2nd Ionization Energy
After an element loses one electron, it
may lose another.
Sometimes it is easier to lose the second
electron than the first.
Sometimes the first electron is easier to
remove
This depends on the stability of the
electron configuration after it loses the
1st electron
2nd Ionization Energy
Magnesium will be more unstable
after losing 1 electron, so the 2nd
electron will be lost more easily
Sodium becomes very stable after
losing just 1 electron, so it will be
more difficult to lose the second
electron
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 decreases down 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
Atomic Radius vs. Ionic Radius
When atoms become ions, their
size will change:
– Metals will lose electrons and
become smaller than the neutral
atom (Na+ is smaller than Na)
– Nonmetals will gain electrons and
become larger than the neutral atom
(F- is larger than F)