Download atomic theory and the periodic table

Tatiana Roizer
Chapter 3
ATOMIC THEORY AND THE PERIODIC TABLE
Subatomic
particle
Electron (e)
Discoveries
Thomson, 1887
Properties
Present in all atoms
Extermely light (1/1836 mass of H atom)
Posses negative charge, assigned -1
Proton (p)
Thomson and Goldstein,
Present in all atoms
1907
About the same mass as H atom
Has positive charge equal in magnitude but
oppisite in sign to electron, assigned +1
Neutron (n)
Chadwick, 1932
About the same mass as a proton
Has no Charge (is electrically nuetral)
http://www.sci.tamucc.edu/pals/morvant/genchem/atomic/page5.htm
Dr. Mark Morvant
The nucleus was found to be composed of two kinds of particles


Some of these particles are called protons
o charge = +1
o mass is about the same as a hydrogen atom
The other particle is called a neutron
o has no charge
o has a mass slightly more than a proton
For an atom to be neutral,
# of Protons = # of Electrons
Atomic Number (Z) - The number of protons in the nucleus of
an atom. All atoms of particular element have the same atomic
number, which is indicated by a subscript to the left of the
element symbol
Mass Number - the number of protons plus neutrons in the
nucleus of an atom.
Isotopes - Different forms of an element having the same number
of protons but different numbers of neutrons (and therefore
different atomic weights).
Isotopes are identified by their Mass Number
Mass Number = Protons + Neutrons
Atomic Weight (Mass) - The mass of a particular atom relative to
the mass of an atom or carbon-12 (12C), which is arbitrarily
assigned a mass of exactly 12.
Average Atomic Weight - Average weight of an element based on the naturally
occurring isotopes and the relative abundance of these isotopes on Earth.
A unit of mass equal to the mass of a single atom of the most common isotope of carbon,
, divided by 12,
The atomic mass unit, also called the dalton after the chemist John Dalton, is a small
unit of mass used to express atomic masses and molecular masses. It is defined to be 1/12
of the mass of one atom of Carbon-12. The abbreviations "u", "amu" and "Da" are used
for this unit; often, atomic masses are written without any unit and then the amu is
implied.
The value is
1 amu ≈ 1.6605387 × 10-27 kilograms.
The unit is convenient because one hydrogen atom weighs approximately 1 amu, and
more generally an atom or molecule that contains n protons and neutrons will have a
mass approximately equal to n amu. This is only a rough approximation however, since it
doesn't account for the mass contained in the binding energy of the nucleus.
Another reason the unit is used is that it is much easier to compare masses of atoms and
molecules (determine relative masses) than to measure their absolute masses. Finding the
mass of a given molecule in amus is thus easier than to express 1 amu in terms of
kilograms.
Avogadro's number NA and the mole are defined so that one mole of a substance with
atomic or molecular mass 1 amu will weigh precisely 1 gram. As an equation:
1 amu = 1 gram/mole
or equivalently
1 gram = NA amu
For example, the molecular mass of water is 18.01508 amu, and this means that one mole
of water weighs 18.01508 grams, or conversely that 1 gram of water contains
NA/18.01508 ≈ 3.3428 × 1022 molecules.\n
Alkali Metals - Group IA
Alkaline Earth Metals - Group IIA
Chalcogens - Group VIA
Halogens - Group VIIA
Noble Gases - VIIIA


Columns are called Groups or Families
o
Elements with similar chemical and physical properties are in the same
column
Rows are called Periods
o
Each period shows the pattern of properties repeated in the next period
Main Group (Representative Group) - Groups IA - VIIIA
Transition Metals - Groups IB - VIIIB
Rare Earth Elements - Lanthanides (Ce - Lu) and Actinides (Th - Lr)
Metals


about 75% of all the elements
lustrous, malleable, ductile, conduct heat and electricity
Nonmetals

dull, brittle, insulators
Metalloids


also know as semi-metals
some properties of both metals & nonmetals
Law of Mendeleev:
Properties of the elements recur in regular cycles
(periodically) when the elements are arranged in
order of increasing atomic weight.
Periodic Law:
The properties of the elements are a periodic function of
atomic numbers.
ATOMIC ORBITALS
What is an atomic orbital?
Orbital /áwrbit'l/ noun. (Phys) Space in an atom occupied by an electron. A
subdivision of the available space within an atom for an electron to orbit the
nucleus. an atom has many orbitals, each of which has a fixed size and shape and
can hold up to two electrons. (Encarta)
When the a planet moves around the sun, you can plot a definite path for it which
is called an orbit. A simple view of the atom looks similar and you may have
pictured the electrons as orbiting around the nucleus. The truth is different, and
electrons in fact inhabit regions of space known as orbitals.
Orbits and orbitals sound similar, but they have quite different
meanings. It is essential that you understand the difference
between them.
To plot a path for something you need to know exactly where the object is
and be able to work out exactly where it's going to be an instant later. You
can't do this for electrons.
The Heisenberg Uncertainty Principle (not required at A'level) says loosely - that you can't know with certainty both where an electron is and
where it's going next. That makes it impossible to plot an orbit for an
electron around a nucleus. Is this a big problem? No. If something is
impossible, you have to accept it and find a way around it.
Each orbital has a name. The orbital occupied by the hydrogen
electron is called a 1s orbital. The "1" represents the fact that
the orbital is in the energy level closest to the nucleus. The "s"
tells you about the shape of the orbital. s orbitals are spherically
symmetric around the nucleus - in each case, like a hollow ball
made of rather chunky material with the nucleus at its centre.
The orbital on the left is a 2s orbital. This is similar to a 1s orbital
except that the region where there is the greatest chance of finding
the electron is further from the nucleus - this is an orbital at the
second energy level.
If you look carefully, you will notice that there is another region of
slightly higher electron density (where the dots are thicker) nearer
the nucleus. ("Electron density" is another way of talking about how
likely you are to find an electron at a particular place.)
2s (and 3s, 4s, etc) electrons spend some of their time closer to the
nucleus than you might expect. The effect of this is to slightly reduce
the energy of electrons in s orbitals. The nearer the nucleus the electrons get, the
lower their energy.
3s, 4s (etc) orbitals get progressively further from the nucleus.
p orbitals
Not all electrons inhabit s orbitals (in fact, very few electrons live in s orbitals). At
the first energy level, the only orbital available to electrons is the 1s orbital, but at
the second level, as well as a 2s orbital, there are also orbitals called 2p
orbitals.
A p orbital is rather like 2 identical balloons tied together at
the nucleus. The diagram on the right is a cross-section
through that 3-dimensional region of space. Once again, the
orbital shows where there is a 95% chance of finding a
particular electron.
Unlike an s orbital, a p orbital points in a particular direction - the one drawn
points up and down the page.
At any one energy level it is possible to have three absolutely
equivalent p orbitals pointing mutually at right angles to each
other. These are arbitrarily given the symbols px, py and pz.
This is simply for convenience - what you might think of as the
x, y or z direction changes constantly as the atom tumbles in space.
The p orbitals at the second energy level are called 2px, 2py and 2pz. There are
similar orbitals at subsequent levels - 3px, 3py, 3pz, 4px, 4py, 4pz and so on.
All levels except for the first level have p orbitals. At the higher levels the lobes
get more elongated, with the most likely place to find the electron more distant
from the nucleus.
d and f orbitals
In addition to s and p orbitals, there are two other sets of orbitals which become
available for electrons to inhabit at higher energy levels. At the third level, there is
a set of five d orbitals (with complicated shapes and names) as well as the 3s
and 3p orbitals (3px, 3py, 3pz). At the third level there are a total of nine orbitals
altogether.
At the fourth level, as well the 4s and 4p and 4d orbitals there are an additional
seven f orbitals - 16 orbitals in all. s, p, d and f orbitals are then available at all
higher energy levels as well.
For A'level purposes, you have to be aware that there are sets of five d orbitals at
levels from the third level upwards, but you will not be expected to draw them or
name them. Apart from a passing reference, you won't come across f orbitals at
all.
Fitting electrons into orbitals
You can think of an atom as a very bizarre house (like an inverted pyramid!) with the nucleus living on the ground floor, and then various rooms (orbitals) on
the higher floors occupied by the electrons. On the first floor there is only 1 room
(the 1s orbital); on the second floor there are 4 rooms (the 2s, 2px, 2py and 2pz
orbitals); on the third floor there are 9 rooms (one 3s orbital, three 3p orbitals and
five 3d orbitals); and so on. But the rooms aren't very big . . . Each orbital can
only hold 2 electrons.
A convenient way of showing the orbitals that the electrons live in is to draw
"electrons-in-boxes".
"Electrons-in-boxes"
Orbitals can be represented as boxes with the electrons in them shown as
arrows. Often an up-arrow and a down-arrow are used to show that the electrons
are in some way different A 1s orbital holding 2 electrons would be drawn as
shown on the right, but it can be written even more quickly as 1s 2. This is read as
"one s two" - not as "one s squared".
You mustn't confuse the two numbers in this notation:
The order of filling orbitals
Electrons fill low energy orbitals (closer to the nucleus) before they fill higher
energy ones. Where there is a choice between orbitals of equal energy, they fill
the orbitals singly as far as possible.
The diagram (not to scale) summarises the energies of the orbitals up to the 4p
Notice that the s orbital always has a slightly lower energy than the p orbitals at
the same energy level, so the s orbital always fills with electrons before the
corresponding p orbitals.
The real oddity is the position of the 3d orbitals. They are at a slightly higher level
than the 4s - and so it is the 4s orbital, which will fill first, followed by all the 3d
orbitals and then the 4p orbitals. Similar confusion occurs at higher levels, with
so much overlap between the energy levels that the 4f orbitals don't fill until after
the 6s, for example.
For A'level purposes you simply have to remember that the 4s orbital fills before
the 3d orbitals. The same thing happens at the next level as well - the 5s orbital
fills before the 4d orbitals. All the other complications are beyond A'level.
Writing electronic configurations


The two electrons in He represent the complete filling of the first electronic shell.
Thus, the electrons in He are in a very stable configuration
For Boron (5 electrons) the 5th electron must be placed in a 2p orbital because the
2s orbital is filled. Because the 2p orbitals are equal energy, it doesn't matter
which 2p orbital is filled
What do we do now with the next element, Carbon (6 electrons)? Do we pair it with the
single 2p electron (but with opposite spin)? Or, do we place it in another 2p orbital?
The second 2p electron in Carbon is placed in another 2p orbital, but with the same spin
as the first 2p electron:
Hund's rule: for degenerate orbitals, the lowest energy is attained when the number of
electrons with the same spin is maximized
Electrons repel each other, by occupying different orbitals the electrons remain as far as
possible from one another


A carbon atom in its lowest energy (ground state) has two unpaired electrons
Ne has filled up the n=2 shell, and has a stable electronic configuration
Electronic configurations can also be written in a short hand which references the last
completed orbital shell (i.e. all orbitals with the same principle quantum number 'n' have
been filled)


The electronic configuration of Na can be written as [Ne]3s1
The electronic configuration of Li can be written as [He]2s1
The electrons in the stable (Noble gas) configuration are termed the core electrons
The electrons in the outer shell (beyond the stable core) are called the valence electrons
Something curious
The noble gas Argon (18 electrons) marks the end of the row started by Sodium
Will the next element (K with 19 electrons) put the next electron one of the 3d orbitals?

Chemically, we know Potassium is a lot like Lithium and Sodium

What these elements (the alkali metals) have in common is an unpaired valence
electron in an s orbital


If Potassium has an unpaired electron in an s orbital it would mean that it is in the
4s orbital
Thus, the 4s orbital would appear to be of lower energy than the 3d orbital(s)
* http://wine1.sb.fsu.edu/chm1045/notes/Struct/EConfig/Struct08.htm
1996 Michael Blaber
Here is the summary of what I covered above:
To predict a ground state electronic configuration:



Aufbau principle - Lowest energy orbitals fill first
Pauli exclusion principle - No 2 electrons can have the same set of quantum
numbers (maximum of 2 electrons per orbital)
Hund's rule - When filling degenerate orbitals preserve the maximum multiplicity
(maximum number of unpaired electrons)
These rules often give the correct electron configuration for an atom or ground state ion.
A guide to the order of orbital energies:
Order of increasing energy:
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f
The following site is helpful in further understanding the orbital energy levels:
http://lectureonline.cl.msu.edu/~mmp/period/electron.htm
Work Cited
1. American Institute of Physics. www.aip.org/history/ curie/periodic.htm, 2004.
2. W. Bauer, Michigan State Univ. Online.
http://lectureonline.cl.msu.edu/~mmp/period/electron.htm, 1999.
3. Bentor, Yinon. Chemical Elements. Online.
http://www.chemicalelements.com/show/electronconfig.html, 2003
4. Blaber, Michael. Florida State Univ. Online
http://wine1.sb.fsu.edu/chm1045/notes/Struct/EConfig/Struct08.htm, 1996.
5. Holum, John. Fundamentals of General, organic, and Biological Chemistry. 6th Ed.
New York: John Wiley & Sons, 1998.
1998 - 2004, Inc.
6. Johnson, Charles www.the-periodic-table.com/.../
7. Morvant, Mark. http://www.sci.tamucc.edu/pals/morvant/genchem/atomic/index.htm,
1991
8. Ophardt, Charles. Virtual Chembook. Elmhurst College. Online
http://www.elmhurst.edu/~chm/vchembook/index.html, 2003
9. Weisstein, Eric. Wolfram Research
http://scienceworld.wolfram.com/physics/AtomicMassUnit.html, 2004.