The Periodic Table
Beyond protons, neutrons, and
electrons
It wasn’t always like this…
Early PT Folks
 Johann Dobereiner

Triads- groups of 3 with similarities/ trends


Cl, Br, I – the properties of Br were intermediate
to those of Cl and I
Limited to some groups, not effective with others
 JAR Newlands (1864) Law of Octaves


Every eight elements the pattern repeats itself,
similar to a musical scale repeating every 8 notes
Not generally well received; people thought him a
fool
The Modern Periodic Table
 The original PT was arranged by mass
 By Dmitri Mendeleev and J Lothar Meyer in
1869


Mendeleev predicted the existence of unknown
elements (which turned out to be Ge, Sc, and
Ga), and predicted their properties from the
patterns he saw
Mendeleev corrected the assumed atomic
masses for elements (In, Be, U)
 These are reasons why he is credited with the first
periodic table and is dubbed “The Father of the
Modern Periodic Table” over Meyer
Ekasilicon
Changes….
 Henry Mosley changed the table to be
organized by atomic number (Z) instead; it
then more closely followed trends/ patterns
e- configuration and the PT
 PT also shows trends in electron
configuration

Groups are based upon electron configuration






Alkali metals are #s1 (# is period)
Alkaline earth metals are #s2 (# is period)
Halogens #p5 (# is period)
Noble gases #p6 (# is period)
Transition metals d block (# is period -1)
Inner transition metals are f block (# is period -2)
Blocks and l
*
*
* orbital shape
The blocks you already know correspond to the orbital of the last
(outermost) e- , or valence e-s occupied
Patterns (Periods) and the PT
 We see patterns for many things, including
 Atomic number *(not a periodic pattern, but a pattern)
 Electron configuration
 Atomic radii
 Ionization energy
 Electron affinity
 Electronegativity
 Activity
 Density
The Periodic Law
 Mendeleev says "The properties of the
elements are a periodic function of their
atomic masses"
 We now say: “When atoms are arranged by
increasing atomic number, the physical and
chemical properties show a (repeating)
pattern”
Periodic…
 Summed up: Properties of elements are
periodic functions of their atomic
numbers.
 Hence, we call the table of elements the
PERIODIC table (go figure)
Octet Rule
 “Atoms gain, lose, or share electrons in order to create a
full outer shell”

This is typically going to be eight electrons
 H and He are exceptions; wanting to fill the 1s orbital
 H may want to go to no electrons, which is considered “full” even
though it is empty
 H gains an electron to become H- , (same e configuration as He
 The law can be used to predict several properties
Nuclear Charge
 Nuclear charge – the attraction felt for an electron
by the nucleus
 Electrons are both


attracted to the nucleus and
repelled by other electrons.
 The nuclear charge that an electron experiences
depends on both factors.
 This effects all periodic properties
Effective Nuclear Charge
The effective nuclear charge, Zeff, is found this way:
Zeff = Z − S,
Z = atomic number
S = inner core e-
© 2009, Prentice-Hall, Inc.
Atomic Radii
•Half the distance between adjacent nuclei
•½ (2R)= atomic radius
Atomic Radii
 INCREASES as you go down a group
 This is because n increases
 DECREASES as you go across a period
(Yes, this is counterintuitive)
 Due to the fact that you add e- as you add p+, so the
nucleus is more positively charged, (electrons have the
same negative charge)
 Results in each electron being more attracted to the
(increasingly) more positive nucleus, and being pulled
in closer
 Sort of like making a magnet more powerful- it will
decrease the distance where it will pull objects
towards it
Ionic Radii
 Cations (+)

Smaller than the neutral atom
 The electrons have less repulsion, and pull in
closer to the nucleus
 Anions (-)
 Larger than the neutral atom

More electrons = more repulsion = larger
electron cloud
Ionization Energy
(Heretofore called IE)
 The amount of energy needed to remove an
electron from an atom
 More specifically, an isolated atom of the
element in the gas phase
 Measure in kJ/ mol
Al Al + + eAl + Al +2 + e(g)
(g)
(g)
(g)
I = 580 kJ/mol
I = 1815 kJ/mol
1
2
Why IE?
 Electrons want to hang around the atom (due to the
protons in the nucleus pulling on them), so it takes
energy to remove electrons
 In general


The smaller that atom, the more energy it takes to remove
an electron
 Because the electron is closer to the nucleus than in a
larger atom
The fewer electrons that atom possess, the harder it is to
remove an electron
 Because it will hang on to them tighter as they are closer
to the + charged nucleus;
 There is less repulsion between electrons too
IE, continued
 1st IE: energy needed to remove the first
electron from an element
 2nd IE: energy needed to remove the second
electron from an element
1st IE
Successive IE
 There are also 3rd, 4th, 5th , and so on IEs (which are
successive IEs), until you can’t pull any more off
 It takes more energy to remove successive electrons
than to remove the first

Because…there are then more protons than electrons,
and the stronger positive charge will then act on the
remaining electrons to hold them to the atom
 (Remember that the charge on the nucleus increases
while the charge on each electron remains the same,
causing more pull by the nucleus on each individual
electron)
Things to keep in mind…
 Remember (from coming up with the
abbreviated electron configurations) that:


Inner core electrons are those electrons from
previous Noble Gas
Valence electrons are the electrons that are
on the exterior of an atom

These are the electrons that are responsible for
the behavior (properties) of the element
Successive IEs
 Are higher than the first
 Due to the fact that there is going to be more protons
than electrons at that point, resulting in a stronger
attraction on the remaining electrons than there was in
the first place

Basically increasingly larger jumps as each electron is
removed

One jump is usually much larger than the others,
because once the inner core configuration is reached,
electrons are removed from the inner core, taking a lot
more energy
Successive IEs
I1
I2
I3
I4
I5
I6
I7
Na
495
4560
Mg
735
1445
7730
Al
580
1815
2740
11600
Si
780
1575
3220
4350
16100
P
1060
1890
2905
4950
6270
21200
Si
1005
2260
3375
4565
6950
8490
27000
Cl
1255
2295
3850
5160
6560
9360
11000
Ar
1527
2665
3945
5770
7230
8780
12000
IE and the PT
Electron Affinity (EA)
 The energy change associated with the addition of an
electron to a gaseous atom
 Negative values mean that energy is released when
adding an e
more negative means more E released when adding
an electron
 Wants an electron more than something with a more
positive value
 Positive values mean that energy needs to be added
to add an e
More positive means more E needed to add the
electron
 Does not want an added electron; takes E to do it
The trend for EA is?
 EA becomes more positive moving down
the PT
 EA becomes more negative from left to
right

Farther from the nucleus


There are several exceptions to this
The smaller the atom, the more e--erepulsion when adding electrons
EA trends
Electronegativity (Eneg)
 The ability of an atom to attract electrons in a
bond


Some atoms share electrons easily, others are
electron hogs
The ability to share is rated (usually) from 0 to
4


Elements with 0 Eneg share easily
Elements with a high (close to 4) Eneg don’t
share e- well
Electronegativity Trends
 If it normally goes +, it has a low Eneg
 If it normally goes -, it is has a high Eneg


The smaller it is, the higher the Eneg
The larger it is, the lower the Eneg
 Noble gases, which normally take no charge,
we say have no Eneg values
Metallic character
 Metallic character is acting like a metal (conductive,
shiny, malleable, etc.))
 All elements possess from very low to very high
metallic character.
 The scale is from Fr to F.
 Fr has the most metallic character and F has the
least.
 In groups, metallic character increases with atomic
number because each successive element gets
closest to Fr.
 In periods, metallic character decreases when
atomic number increases because each successive
element gets closest to F.
Reactivity
 The nature (metal, non-metal, semi-metal)
makes a difference in how an element’s
chemical reactivity
 The trends are characterized by their nature
Metals reactivity trend
 In groups, reactivity of metals increases with
atomic number because the ionization energy
decreases.
 In periods, reactivity of metals decreases
when atomic number increases because the
ionization energy increases.
Nonmetals reactivity trend
 In groups, reactivity of non-metals decreases when
atomic number increases


because the electronegativity decreases
Relate to size- it increases.
 In periods, reactivity of non-metals increases with
atomic number

because the electronegativity increases.
 Relate to size- radii decreases
 Remember, the radii would have an effect on this
Density: in general
 Density of solids is greatest
 Measured in g/cm3
 Highest in center of table (d- block)
 Density of gases
 Measured in g/L at Standard Temp &Pressure (STP,
which is 1atm and 0°C)
 Increases as you go down a group
 Increases as you go across the table, then Decreases
 Density of liquids
 Measured in g/mL
 Density of Hg is greater than that of Br2
Density
To sum it up…
Summary chart again