Tatiana Roizer



Tatiana Roizer

Chapter 3

ATOMIC THEORY AND THE PERIODIC TABLE

|Subatomic particle |Discoveries |Properties |

|Electron (e) |Thomson, 1887 |Present in all atoms |

| | |Extermely light (1/1836 mass of H atom) |

| | |Posses negative charge, assigned -1 |

|Proton (p) |Thomson and Goldstein, 1907 |Present in all atoms |

| | |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) |



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

 

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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).

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Isotopes are identified by their Mass Number

Mass Number = Protons + Neutrons

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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, [pic], divided by 12, |

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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

[pic]

 Alkali Metals - Group IA

Alkaline Earth Metals - Group IIA

Chalcogens - Group VIA

Halogens - Group VIIA

Noble Gases - VIIIA

[pic]

• 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 1s2. This is read as "one s two" - not as "one s squared".

You mustn't confuse the two numbers in this notation:

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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 [pic]

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

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• 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?

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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

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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

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• 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)

* 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:

[pic]

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:



 Work Cited

1. American Institute of Physics. history/ curie/periodic.htm, 2004.

2. W. Bauer, Michigan State Univ. Online.

, 1999.

3. Bentor, Yinon. Chemical Elements. Online. , 2003

4. Blaber, Michael. Florida State Univ. Online , 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 .../

7. Morvant, Mark. , 1991

8. Ophardt, Charles. Virtual Chembook. Elmhurst College. Online , 2003

9. Weisstein, Eric. Wolfram Research , 2004.

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