IB syllabus > atomic theory > 2.1 

These notes were written for the old IB syllabus (2009). The new IB syllabus for first examinations 2016 can be accessed by clicking the link below.

IB syllabus for first examinations 2016

2.1 - The Atom


2.1.1: State the position of protons, neutrons and electrons in the atom. TOK: None of these particles can be (or ever will be) directly observed. Which ways of knowing do we use to interpret indirect evidence gained through the use of technology? Do we believe or know of their existence.


Rutherford's experiments

Rutherford demonstrated this with his famous scattering experiments at the beginning of the 20th century.

Electrons and orbitals

The electrons are sometimes said to 'spin' around the nucleus like planets orbiting the sun. Although this is easy to envisage it presents an inaccurate picture.

The electrons behave as areas of negative charge that are spread out within the region of space that they occupy. This region of space is called the orbital and only two electrons are allowed to occupy any one orbital. We say that each electron within an orbital has a 'spin' state. This is just the word used by science to differentiate the two electrons and any 'spin', is and would be, impossible to demonstrate. The two electrons have slightly different character that's all.


2.1.2: State the relative mass and relative charge of protons, electrons and neutrons.
The accepted values are: Relative Mass (Charge), proton 1 (+1), neutron 1 (0), electron 1/1840 (-1)


The actual masses of these sub-atomic particles are very small and to the nearest whole number measured relative to the mass of a carbon 12 isotope being equal to 12 units they are:

Although these values suggest that the protons and neutrons are identical they do, in fact, have very slightly different masses, which is only of concern to us when considering changes in the nuclear structure in radio-chemistry. Compared to the mass of the protons and neutrons the electrons have neglegible mass and can be forgotten about when performing calculations involving mass.

Note: A full description of the electron is rather complicated as they behave both as particles, in some circumstances, and waves in others. This has lead quantum physicists to calling them (and some other particles) wavicles - things that possess the characteristics of both particles and waves.

Each orbital has a certain energy, the electrons fill the orbitals starting with the lowest energy orbital and working up.

The orbitals are grouped into energy shells 1, 2, 3, 4 etc


2.1.3: Define the terms mass number (A), atomic number (Z) and isotope of an element.


AZE Notation

The nuclear equation in the previous section shows how the particles within an atom or isotope can be represented. This is called the AZE notation. The number at the top left represents the mass number and the bottom left the atomic number.

Example

56 Fe -->
26

Represents an iron atom with a mass of 56 units and 26 protons. In some respects there is no need for the symbol as the fact that it has 26 protons DEFINES the atom as iron.

Using these values it is easy to find the number of subatomic particles.

The mass number = protons + neutrons = 56

The atomic number = number of protons = 26

Therefore the number of neutrons = 56 - 26 = 30

Atomic number

The atomic number is the number of protons in the atom. This is given the symbol Z.

As only the protons and neutrons have appreciable mass, the total mass of an atom is the sum of its protons and neutrons. This is called the mass number. It is given the symbol A.

Electrons

In a neutral atom the number of protons must be balanced out by the number of electrons and so the electron number must also be Z. It is important to distinguish between the neutral atom and ions when dealing with the number of electrons.


2.1.4: Deduce the symbol for an isotope given its mass number and atomic number. Use the AZE notation.


Isotopes

It seems likely that the neutrons role in an atomic nucleus is to glue the protons together so that the electrostatic repulsion of like charges doesn't make the nucleus fall apart. Certainly it is possible to have atoms of the same element with different numbers of neutrons - these are called isotopes.

Hydrogen, for example has three isotopes

Isotope number of protons number of neutrons mass
Hydrogen 1 0 1
Heavy hydrogen (deuterium) 1 1 2
Super-heavy hydrogen (tritium) 1 2 3

Tritium, however, is unstable and the nucleus decays (breaks apart) after a time and one of its neutrons changes into a proton and a beta particle (fast moving electron ) is emitted. This is an example of radioactivity - in this case, beta radiation.

This can be expressed by a nuclear equation:

3 T --> 3 He + 0 ß
1 2 -1

Notice how the top line is balanced (3 = 3 + 0) and so is the bottom line (1 = 2 - 1). The top line represents the total mass of each particle and the bottom line shows the number of positive charges (protons)


2.1.5: Calculate the number of protons, electrons and neutrons in atoms and ions from the mass number, atomic number and charge.


Calculating the sub-atomic particles

As stated above it is important to be careful when dealing with the numbers of fundamental (sub-atomic) particles in ions.

Species atomic number mass number protons neutrons electrons
11B atom 5 11 5 6 5
Na+ ion 11 23 11 12 10
35Cl- ion 17 35 17 18 18
37Cl atom 17 37 17 20 17
34S2- ion 16 34 16 18 18

Notice that the negative ions have extra electrons added and the positive ions have electrons taken away.


2.1.6: Compare the properties of the isotopes of an element.


Properties of isotopes

On the whole the properties of isotopes are almost identical to one another. There may be slightly different physical properties such as density or rate of diffusion. Some isotopes are radioactive and emit radiation (alpha, beta and gamma) - these are called radioisotopes.

Isotopes may also react with different rates. An example of this is the 'uptake' of deuterium by hydrogenated compounds.

Separation of isotopes

The separation of isotopes is carried out using the different rates of diffusion. For example uranium has two main isotopes uranium 235 adn uranium 238 (i.e. with masses 235 and 238)

The 235 isotope is needed for nuclear power (and explosives) but usually only occurs at very low percentages and must be seaprated from the other isotope.

This is carried out by reacting the uranium with fluorine making uranium hexafluoride UF6 (this is a volatile gas) . The uranium hexafluoride is then passed through many centrifuges where the heavier UF6 made from uranium 238 does not pass through as quickly as the UF6 made from uranium 235.

Eventually after passing through many centrifuges the final product is pure UF6, containing only uranium 235. It can then be chemically reduced back to pure uranium or, more often, to uranium oxide pellets for its final destination in the nuclear power industry.

Relative atomic mass

For elements which have more than one isotope (the majority), the measured relative atomic mass of a sample will be the average mass of all of the isotopes in the sample. Usually the relative proportions of a naturally occurring element do not vary throughout the world and so it is possible to calculation the average mass of one atom of an element from isotope and abundance data. These average values are the relative atomic masses recorded in the periodic tables and data books.

Example:

Chlorine has two common isotopes - chlorine 35 and chlorine 37 with relative percentage abundances of 77.5% and 22.5% respectively.

This means that in any naturally occurring sample of chlorine for every 100 atoms there are 77-78 atoms with a mass of 35 units and 22-23 atoms with a mass of 37 units.

To find the average mass of one atom we must add up the masses of all 100 atoms and then divide by 100.

Mass of 100 chlorine atoms = (77.5 x 35) + 22.5 x 37) = 3545

therefore average mass of one chlorine atom = 3545/100 = 35.45

Clearly, it is not possible to have an atom with a mass of 35.45 units, but this represents the average relative mass of a chlorine atom.

 

Example:

Bromine has two common isotopes - bromine 79 and bromine 81, with relative percentage abundances of 50% and 50% respectively.

This means that in any naturally occurring sample of bromine for every 100 atoms there are 50 atoms with a mass of 79 units and 50 atoms with a mass of 81 units.

To find the average mass of one atom we must add up the masses of all 100 atoms and then divide by 100.

Mass of 100 bromine atoms = (50 x 79) + (50 x 81) = 8000

therefore average mass of one bromine atom = 8000/100 = 80.00

Although, it is possible to have a bromine atom with a mass of 80 units, there are not any in a natural sample, this represents the average relative mass of a bromine atom.


2.1.7: Discuss the uses of radioisotopes


Uses of radioisotopes

Isotopes have a huge variety of uses from medicinal to industrial, from nuclear energy to fire alarms.

14C is used in in radiocarbon dating.

In this process use is made of the fact that living organisms take up carbon throughout their lives. The percentage of the isotope carbon 14 is fairly constant in our atmosphere as it is produced in the upper atmosphere by cosmic bombardment of naturally occurring carbon dioxide.

This means that the percentage of carbon 14 contained by all living organisms is also constant. However, when a living organism dies it stops taking up carbon 14. The isotope decays naturally with a half life of about 5,600 years. So a simple procedure involving counting the radioemissions due tocarbon 14 from a sample of material that was once alive, can be used to estimate its date.

Cobalt 60 is used in hospitals as a beta emission source in the treatment of cancer

Beta rays are fast moving electrons. They can be focussed onto cancerous tissue to destroy it using a cobalt 60 source. This form of treatment is known as radiotherapy.

131I and 125I are used as medical tracers

In several conditions the body can be scanned for problems using iodine, which is easily taken up by the body and transported through the lymphatic system. The isotopes 131I and 125I are easy to detect and short lived in the body.


Resources

Line spectra

 


http://www.chem.uoa.gr/applets/AppletMS/Appl_Ms2.html