ATOMIC STRUCTURE
WHAT ARE ISOTOPES? ARE THEY IMPORTANT?
Isotopes are atoms of the same element with different numbers of neutrons and therefore different masses (different nucleon/mass numbers).
o This gives each isotope of a particular element a different mass or nucleon number, but, being the same element they have the same atomic or proton number.
o They are also chemically identical, because they have the same number of electrons, hence the same electron structure.
o Study the diagrams of the isotopes of carbon further down the page.
o Relative Isotopic Mass is dealt with on a separate calculation page
The phrase 'heavier' or 'lighter' isotope means 'bigger' or 'smaller' mass number for a particular element.
There are small physical differences between the isotopes e.g. the heavier isotope has a greater density or boiling point, the lighter the isotope the faster it diffuses.
However, because they have the same number of protons (proton/atomic number) isotopes of a particular element have the same electronic structure and identical chemistry.
Examples of isotopes are illustrated and described below.
Caution Note: Do NOT assume the word isotope means the atom it is radioactive, this depends on the stability of the nucleus i.e. unstable atoms (radioactive) might be referred to as radioisotopes.
Many isotopes are extremely stable in the nuclear sense and NOT radioactive i.e. most of the atoms that make up you and the world around you!
hydrogen–1, hydrogen–2, and hydrogen–3 are the three isotopes of hydrogen with mass numbers of 1, 2 and 3, with 0, 1 and 2 neutrons respectively. All have 1 proton, since all are hydrogen! Hydrogen–1
is the most common, there is a trace of hydrogen–2 (sometimes called deuterium) naturally but hydrogen–3 (sometime called tritium) is very unstable and is used in atomic bombs – nuclear fusion weapons.
o They are sometimes denoted more simply as 1H, 2H and 3H since the chemical symbol H means hydrogen and therefore must have only one proton.
and or 3He and 4He, are the two isotopes of helium with mass numbers of 3 and 4, with 1 and 2 neutrons respectively but both have 2 protons. Helium–3 is formed in the Sun by the initial nuclear fusion process.
Helium–4 is also formed in the Sun and as a product of radioactive alpha decay of an unstable nucleus.
o An alpha particle is a helium nucleus (mass 4, charge +2) and if it picks up two electrons it becomes a stable atoms of the gas helium. For more details see Radioactivity Revision Notes Part 4
and or 23Na and 24Na, are the two isotopes of sodium with mass numbers of 23 and 24, with 12 and 13 neutrons respectively but both have 11 protons in the nucleus and 11 surrounding electrons. Sodium–23 is quite stable e.g. in common salt (NaCl, sodium chloride) but sodium–24 is a radio–isotope and is a gamma emitter used in medicine as a radioactive tracer e.g. to examine organs and the blood system.
and are the two nuclear symbols for the two most common and stable isotopes of the element chlorine. They both have 17 protons in the nucleus and 35–17 = 18 and 37–17 = 20 neutrons respectively (and both have 17 surrounding electrons).
and are the two nuclide symbols for the two most common and stable isotopes of the element bromine. They both have 35 protons in the nucleus and 79–35 = 44 neutrons and 81–35 = 46 neutrons respectively.
By coincidence, there are almost exactly 50% of each isotope present in naturally occurring bromine.
The three known isotopes of carbon
o
o o
isotope nuclide symbol protons neutrons electrons % abundance
carbon–12 126C 6 6 6 98.9%, stable
carbon–13 136C 6 7 6 1.1%, stable
carbon–14 146C 6 8 6 trace, unstable
radioactive
o The table of information on the three isotopes of carbon is illustrated by the diagrams above it.
o Now is an appropriate point to introduce the concept and definition of relative atomic mass (Ar), which is required for very accurate quantitative chemistry calculations.
o The relative atomic mass of an element is the average mass of all the isotopes present compared to 1/12th of the mass of a carbon–12 atom (12C = 12.00000 amu i.e. the standard).
When you average the masses of the isotopes of carbon, taking into account their relative
abundance (%), you arrive at a relative atomic mass of carbon of 12.011, Ar(C) = 12.011, though at this academic level 12.0 is accurate enough!
See also chemical calculations on how to calculate relative atomic mass
I've put this calculation on its own page because there is plenty on atomic structure already on this page!
Anything on this page relevant to the calculation of RAM is repeated on the page.
Knowledge of isotopes is important in modern science.
o Radioactive isotopes are used in medicine to trace aspects of body chemistry due to their radioactive emissions, and in chemical synthesis as tracers to follow how a reaction sequence occurs.
o Radioactive isotopes are used in radiotherapy to kill malignant cancer cells.
For lots more details see the RADIOACTIVITY NOTES
DO NOT CONFUSE ISOTOPES and ALLOTROPES – see Appendix 3.
3. The Electronic Structure of Atoms – rules to be learned
WHAT DO WE MEAN BY Electron configuration, electronic structure of atoms – arrangement in shells or energy levels?
The electrons are arranged in energy levels or shells around the nucleus and with 'orbits' on average increasing in distance from the nucleus.
o The lowest energy levels are always filled first, you can think of the lower the shell, the nearer the nucleus, and numbered 1, 2, 3 etc. as the shell gets further from the nucleus.
Each electron in an atom is in a particular energy level (or shell) and the electrons must occupy the lowest available energy level (or shell) available nearest the nucleus.
When the level is full, the next electron goes into the next highest level (shell) available.
There are rules to learn about the maximum number of electrons allowed in each shell and you have to be able to work out the arrangements for the first 20 elements (for GCSE students, upto at least 36 for Advanced level students).
o The 1st shell can contain a maximum of 2 electrons (electrons 1–2) o The 2nd shell can contain a maximum of 8 electrons (electrons 3–10) o The 3rd shell also has a maximum of 8 electrons (electrons 11–18)
o The 19th and 20th electrons go into the 4th shell, (required limit of GCSE knowledge).
o Remember the total electrons to be arranged equals the atomic/proton number for a neutral atom.
If you know the atomic (proton) number, you know it equals the number of electrons in a neutral atom, you then apply the rules to work out the electron arrangement (configuration).
For elements 1 to 20 the electron arrangements/configurations are written out in the following manner:
o Note that each number represents the number of electrons in a particular shell, dots or commas are used to separate the numbers of electrons in each shell. They are written out in order of increasing average distance from the positive nucleus which holds these negative electrons in their energy levels (shells).
o The electron configurations are summarised below with reference to the periods of the periodic table and in order of increasing atomic number.
For more see the Periodic Table and Electron Structure notes below.
o Period 1 – elements 1 to 2 (2 elements)
the electron arrangement is written out simply as 1 or 2 o Period 2 – elements 3 to 10 (8 elements)
electron arrangements of 2.1 to 2.8 (since 1st shell is full with 2 electrons i.e. the first number) o Period 3 – elements 11 to 18 (8 elements)
denoted by 2.8.1 to 2.8.8 (1st,2nd full shells with 2,8 electrons) o Period 4 – first two elements 19 to 20
written out as 2.8.8.1 and 2.8.8.2 (1st,2nd,3rd full shells with 2,8,8 electrons)
Reminder – this is as far as GCSE students need to know, after that things get more complicated, BUT only for advanced level students!
For example, after element 18, the 3rd shell can hold a maximum of 18 electrons!
o The above is summarised in the diagram below
o o The electron shell arrangements are quoted in numbers e.g. 2.4 for C (carbon) but you need to be able to draw electron diagrams showing the electronic structure of the atom.
Some examples are given below and GCSE/IGCSE/O level students need to be able to work and draw the electronic structures of the first 20 elements.
You should notice that the number of shells used equals the period number of the element in the periodic table.
They can be all worked by the 'shell filling' rules described above.
o For the rest of Period 4 and other Periods you need a more advanced electron configuration system upto at least Z=36 using s, p, d and f orbital notation BUT for advanced level chemistry students only!
Examples: diagram, symbol or name of element (Atomic Number = number of protons and the number of electrons in a neutral atom), shorthand electron arrangement and a diagram to help you follow the numbers.
Filling 1st shell, electron level 1 2 elements only, Period 1 of the Periodic Table
Filling 2nd shell, electron level 2 to to 3 of the 8 elements of Period 2
Filling 3rd shell, electron level 3 to 3 of the 8 elements of Period 3
The first 2 elements of the 4th shell to Kr [2.8.18.8], start of Period 4
Only the first 2 of the 18 elements of Period 4 are shown above, the rule for 3rd shell changes from element 21 Sc onwards (studied at Advanced level, so GCSE students don't worry!)
A few more 'snappy' examples – given atomic number, work out electron configuration (abbreviated to e.c.)
Z = 3 e.c. 2.1 or Z = 7 e.c. = 2.5 or Z = 14 e.c. = 2.8.4 or Z = 19 e.c. = 2.8.8.1 etc. upto Z = 20
4. Which electron arrangements are stable and which are not?
Both atoms and ions are considered
WHY ARE SOME ELECTRON ARRANGEMENTS ARE MORE STABLE THAN OTHERS?
WHICH ELECTRON ARRANGEMENTS ARE THE MOST STABLE AND WHICH ELECTRON ARRANGEMENTS THE LEAST STABLE?
HOW DO ELECTRON ARRANGEMENTS RELATE TO THE REACTIVITY OF CHEMICAL ELEMENTS?
When an atom has its outer level full to the maximum number of electrons allowed, the atom is particularly stable electronically and very unreactive.
o This is the situation with the Noble Gases: He is [2], neon is [2,8] and argon is [2,8,8] etc.
o There atoms are the most reluctant to lose, share or gain electrons in any sort of chemical interaction because they are so electronically stable.
o For all elements most of their chemistry is about what outer electrons do or don't!
o [2], [2,8] and [2,8,8] etc. are known as the 'stable Noble Gas arrangements', and the atoms of other elements try to attain this sort of electron structure when reacting to become more stable.
o More details on Electron configuration notes for Advanced Level Chemistry Students
The most reactive metals have just one outer electron.
o These are the Group 1 Alkali Metals, lithium [2,1], sodium [2,8,1], potassium [2,8,8,1]
o With one outer shell electron, they have one more electron than a stable Noble Gas electron structure.
o So, they readily lose the outer electron when they chemically react to try to form (if possible) one of the stable Noble Gas electron arrangements – which is why atoms react in the first place!
o When Group 1 Alkali Metal atoms lose an electron they form a positive ion because the positive proton number doesn't change, but with one negative electron lost, there is a surplus of one + charge e.g.
sodium atom ==> sodium ion
Na ==> Na+
is [2.8.1] ==> [2.8] electronically
in fundamental particles [11p + 11e] ==> [11p + 10e]
IONS are atoms or group of atoms which carry an overall electrical charge i.e. not electrically neutral.
The most reactive non–metals are just one electron short of a full outer shell.
o These are the Group 7 Halogens, namely fluorine [2,7], chlorine [2,8,7] etc.
o These atoms are one electron short of a stable full outer shell and seek an 8th outer electron to become electronically stable – yet again, this is why atoms react!
o They readily gain an outer electron, when they chemically react, to form one of the stable Noble Gas electron arrangements either by sharing electrons (in a covalent bond) or by electron transfer forming a singly charged negative ion (ionic bonding) e.g.
chlorine atom ==> chloride ion
Cl ==> Cl–
is [2.8.7] ==> [2.8.8] electronically
in fundamental particles [17p + 17e] ==> [17p + 18e]
the positive proton number of Cl doesn't change but the chloride ion carries one extra negative electron to give the surplus charge of a single – on the ion.
EXTRA NOTE ON 'ATOMIC' NOTATION – representation of isotopes of ions
Nuclide notation and ions (interpretation required for advanced level students only)
o sodium–24 isotope ion, 11 protons, 13 neutrons, 10 electrons (one electron lost to form a positive ion)
o sodium–23 isotope ion, 11, protons, 12 neutrons, 10 electrons (one electron lost to form a positive ion)
o isotope sulfur–32 in the form of the sulfide ion, 16 protons, 16 neutrons, 18 electrons (two electrons gained to form the double charged negative ion)
For more on electron structure and chemical changes and compound formation see ...
o GCSE/IGCSE/AS notes on CHEMICAL BONDING
and for more on metal and non–metal reactivity see
o GCSE/IGCSE notes REACTIVITY SERIES of METALS o GCSE/IGCSE notes Group 1 ALKALI METALS o GCSE/IGCSE notes GROUP 7 HALOGENS
5. The Periodic Table and Electronic Structure – more patterns!
Selected Elements of the Periodic Table are shown below with atomic number and chemical symbol.