Atoms

Chemistry I

© Ray Lovegrove
© Rajesh Rathod



CHAPTER 3:

Ideas about Atoms

 

In chapter one we looked briefly at the ideas of the Ancient Greeks about atoms. You will remember that the main idea involved tiny particles of matter that could not be broken down. This idea was, like many Ancient Greek ideas, held up over the centuries that followed; Shakespeare used the word 'atomies' in  Romeo and Juliet (Queen Mab speech). It was an idea that needed some

refinement however, and in 1809, the English Quaker chemist JOHN DALTON came up with the ATOMIC THEORY. This theory has been at the heart of modern chemistry for almost two hundred years.

Dalton's Atomic Theory was developed to marry old ideas with the (then) modern discoveries of chemistry; in particular, Dalton was convinced that quantitative chemistry could throw new light on how elements combined with each other.  The Atomic Theory is usually broken down into four statements:
  1. Atoms can neither be created nor destroyed.

  2. Atoms of the same element are the same as each other in terms of  mass, colour and size.

  3. Atoms of different elements are different from each other in terms of mass, colour and size.

  4. Atoms join together in small whole numbers to form molecules.

Now, if you know a little chemistry, you will already see that that Dalton was wrong about a few things, but his Atomic Theory is basically sound. At the end of the chapter you should be able to list the ways in which current knowledge seems to contradict Dalton on each of the four statements listed above.

One big problem occurs when chemists and physicists talk about atoms - they are so small that we can not see them. All our ideas about atoms involve us in developing SCIENTIFIC MODELS. Models help us picture what we cannot see, and build up an understanding of how things behave. Dalton's model of the atom pictured it pretty much as a solid, hard sphere, and because he believed them to be coloured (despite being colour-blind himself), this is often referred to as the BILLIARD BALL MODEL. Indeed, if you ask people to describe what they think of as an atom, most people will describe something approaching this model. Dalton also gave symbols to the elements to help him explain how they joined together. Some of the symbols can be seen in figure 1.

ATOMIC MASS
The mass of atoms is a difficult concept - if something is so small that you cannot see it, how can you find its mass?   Of course, to use grams is not helpful as they are not small enough, so chemists use a very helpful measurement called ATOMIC MASS UNITS, or amu for short.  We need a standard against which to measure atoms - the IUPAC (see chapter 1), have decided that the standard unit is carbon (it is the element present in more compounds than any other). One carbon atom has a mass of exactly 12 amu.  This is used to compare other atoms;  for instance, a magnesium atom has twice the mass of a carbon atom, so has a mass of  24 amu. One carbon atom has twelve times the mass of a single hydrogen atom, therefore hydrogen has a mass of 1 atomic mass unit. This is the basic idea of atomic mass, but we will need to refine it later on in this chapter and further investigate the concept in later sections.

PROTONS, NEUTRONS and ELECTRONS
Dalton's billiard ball  model of the atom was the accepted idea for almost a hundred years. With increasing interest in electricity and the continued development of the vacuum pump however, new and more sophisticated models needed to be constructed. For one reason, it was becoming obvious that atoms were composed of even smaller units called SUB-ATOMIC PARTICLES. The main evidence for this comes from experiments involving gases at very low pressures that are subjected to high voltage charges. When this takes place, positive and negative parts of the atom can be detected.  Quite a lot was discoved about these charges at the start of the 20th century; for instance, the negatively charged particles were called ELECTRONS (chemists are quite happy to think of electrons as very small particles, but much of their behaviour is wave-like) and it was found that they were very small but with a quite large charge. The positive particles were called PROTONS and although they also had a fairly large charge, were found to be much larger than electrons.  This meant that a new model was needed - for a short time it was thought that these protons and electrons were evenly distributed throughout each atom like the fruit in a cake. This was called the 'plum-pudding model'.

The mass of an atom is greater than the total mass of protons and electrons found within it. How can this be?  Eventually, a third particle, called the NEUTRON, was found.  It had a mass very much the same as a proton and no charge at all (the lack of charge made it difficult to detect). It was also established that protons and neutrons were to be found in the centre of the atom in a dense area called the NUCLEUS.  It was discovered that electrons were found moving very quickly, in an area well away from the nucleus, called the SHELL. This model of atoms was like little solar systems, with the electrons orbiting the nucleus rather as planets orbit the sun. As it happens, this model is not very well supported by evidence, but it has captured peoples' imagination, so seems to be still around.

As chemists we are not too bothered by the fact that we can not see atoms or sub-atomic particles - we are more interested in how they behave. To understand this behaviour we need two models to help us.

MODEL 1.   Sub-atomic particles table.    We cannot measure the mass of sub-atomic particles in any conventional unit - they are just too small.  So, we use a.m.u. (as mentioned above). It is also difficult to measure the charge on these particles, so we compare them to the charge on one proton (which has a single positive charge, illustrated by +).  This happens to be exactly equal and opposite to the charge on an electron (a single negative charge, illustrated by -). The follwing table is useful to compare the three particles in the atom.

Sub atomic particle
Mass in a.m.u.
Relative charge
Position in the atom
PROTON      1p+
1
+
Nucleus
NEUTRON        1n0  
1
0
Nucleus
ELECTRON           e–
1/1840 (almost 1/2000)
–
Shell


Note that the neutron must be given a zero.  If we were to represent its charge as a dash. this would look like a negative sign which we already use for electrons.

MODEL 2.   Electrons in shells model.  While the table contains a lot of information, it is not a working model that can be used, neither does it lend itself to helping us understand how atoms join together.


Electrons in KLMN Shells[P+N Nucleus]




We will be using this model very often in the coming chapters and it is a very good method for showing such things as bonding and the forming of ions. Note that the atom shows protons as P and neutrons as N in the nucleus.  The nucleus is often not shown at all; the electrons are shown as green dots (crosses also can be used to show electrons) in KLMN shells. These shells have sub-shells or more accurately termed orbitals; s, p, d, f, etc. Atomic Structure and Electronic Configuration 3ii 3vi 4

ISOTOPES
Dalton could not have dreamt of the existence of subatomic particles and he would have been suprised to discover that all atoms of the same element are exactly the same. Some atoms of carbon for instance, have a mass of 12 a.m.u. and some have a mass of 14 a.m.u. All carbon atoms have the same number of protons and electrons but the numbers of neutrons differ.  Atoms of the same element but with different atomic masses are called ISOTOPES.  Most elements have some naturally occuring isotopes including carbon, hydrogen and oxygen - which means that you have a few isoptopes in your body as you are reading this.

If you take a good look at the  Periodic table you will find that the elements all have their mass (weight) recorded in a.m.u.  The figure used is the RELATIVE ATOMIC MASS.  A definition of this figure is - the weighted mean of naturally occuring atoms of that element. It is very important to refer to individually named isotopes as having atomic mass but a named element as having a RELATIVE atomic mass. It is also important to remember that weighted mean is not the same as the average.  If we look at the two naturally occuring isotopes of chlorine, they are 35Cl and 37Cl, having masses of 35 a.m.u. and 37 a.m.u. respectively. The average of 35 and 37 is 36, but the relative atomic mass (Ar) of chlorine is in fact 35.5.  Not only the mass of the isopes but their abundance must be taken into account when calculating their relative atomic mass.  75% of all chlorine atoms are 35Cl, and 25% are 37Cl.  Thus the following calculation shows how the relative atomic mass is calculated.


Ar  Cl  =

(35 x 75) + (37 x 25)
         –––––––––––––––––– = 35.5
     100
        


If you check the periodic table above you will find a few elements have relative atomic masses that reflect some common isotopes; other elements have few or rare isotopes so the Ais very close to the atomic mass.

MORE ABOUT PROTONS, NEUTRONS AND ELECTRONS
Using the periodic table again, we can see what makes elements of different atoms unique - it is not the mass (note that cobalt and nickel both have the same Ar, rather it is the number of protons that they have.  Atoms are arranged in the periodic table according to the number of protons present. The number of protons in an atom is called the ATOMIC NUMBER. Note also that the number of protons is always equal to the number of electrons - this is because the overall charge on any atom is zero, thus the proton and electron charges cancel each other out.  You can calculate the number using these rules;

The number of PROTONS in an atom is equal to the ATOMIC NUMBER
The number of ELECTRONS in an atom is equal to the ATOMIC NUMBER
The number of NEUTRONS in an atom is equal to the MASS NUMBER minus the ATOMIC NUMBER

Using these rules, the following table can be constructed; the pattern for protons and electrons is obvious, but neutrons show only a vague increase in numbers as the proton number increases.

ELEMENT
PROTONS
ELECTRONS
NEUTRONS
H
1
1
0
He
2
2
2
Li
3
3
4
Be
4
4
5
B
5
5
6
C
6
6
6
N
7
7
7
O
8
8
8
F
9
9
10
Ne
10
10
10
Na
11
11
12
Mg
12
12
12
Al
13
13
14
Si
14
14
14
P
15
15
16
S
16
16
16

ATOMIC STRUCTURE AND THE ELECTRON 3i 3ii 3iii 3iv 3v
In 1913 Niels Bohr proposed a model of hydrogen atom, which retained earlier nuclear model of Rutherford and Thomson but made further progress towards the behaviour of the electrons. A dramatic explanation for Rydberg spectral expression resulted.


Louis de Broglie in 1923 suggested that particles like electrons could be associated with wave properties. De Broglie recognised that integers introduces by Bohr (to explain behaviour of electron in the hydrogen atom), enter naturally in problems dealing with the waves.

In 1926 Schrodinger developed new approach called wave, or quantum mechanics.


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