- THE FIRST MILLENIUM
1923 – Toronto, Canada
Early research had shown that there was almost certainly a link between the pancreas and diabetes, but at the time it was not understood what it was.
We now know a hormone from the pancreas controls the flow of sugar into the blood stream. Diabetics lack this function and are gradually killed by uncontrolled glucose input into the body’s systems.
Banting believed that the islets of Langerhans might be the most likely site for the production of this hormone and began a series of tests using laboratory animals.
After successfully treating dogs – showing signs of diabetes after the pancreas had been removed – with a solution prepared from an extract from the islets of Langerhans, Banting’s team (Best, MacLeod and Collip) purified their extract and named it insulin.
Human trials successfully took place in 1923 and dying patients were restored to health. The same year, industrial production of insulin from pigs’ pancreas began.
In the Second World War Banting undertook dangerous research into poisonous gas and was killed in an air crash while flying from Canada to the United Kingdom.</p
1912 – England
X-rays scattered from a crystal will show constructive interference provided their wavelength ( λ ) fits the equation
2d sin θ = n λ
where d is the spacing between atoms of the crystal, θ the angle through which the rays have scattered and n is any whole number
This is the cornerstone of the science of X-ray crystallography.
1913 – Denmark
‘Electrons in atoms are restricted to certain orbits but they can move from one orbit to another’
Bohr’s was the first quantum model for the internal structure of the atom.
Bohr worked with RUTHERFORD in Manchester and improved upon Rutherford’s model, which said that electrons were free to orbit the nucleus at random.
Classical physics insisted that electrons moving around the nucleus would eventually expire and collapse into the nucleus as they radiated energy. Bohr resolved the issue surrounding Rutherford’s atomic structure by applying the concept of quantum physics set out by MAX PLANCK in 1900.
He suggested that the electrons would have to exist in one of a number of specific orbits, each being defined by specific levels of energy. From the perspective of quantum theory, electrons only existed in these fixed orbits where they did not radiate energy. The electrons could move to higher-level orbits if energy was added, or fall to lower ones if they gave out energy. The innermost orbit contains up to two electrons. The next may contain up to eight electrons. If an inner orbit is not full, an electron from an outer orbit can jump into it. Energy is released as light (a photon) when this happens. The energy that is released is a fixed amount, a quantum.
Quanta of radiation would only ever be emitted as an atom made the transition between states and released energy. Electrons could not exist in between these definite steps. This quantised theory of the electrons’ orbits had the benefits of explaining why atoms always emitted or absorbed specific frequencies of electromagnetic radiation and of providing an understanding of why atoms are stable.
Bohr calculated the amount of radiation emitted during these transitions using Planck’s constant. It fitted physical observations and made sense of the spectral lines of a hydrogen atom, observed when the electromagnetic radiation (caused by the vibrations of electrons) of the element was passed through a prism. The prism breaks it up into spectral lines, which show the intensities and frequencies of the radiation – and therefore the energy emissions and absorptions of the electrons.
Each of the elements has an atomic number, starting with hydrogen, with an atomic number of one. The atomic number corresponds to the number of protons in the element’s atoms. Bohr had already shown that electrons inhabit fixed orbits around the nucleus of the atom.
Atoms strive to have a full outer shell (allowed orbit), which gives a stable structure. They may share, give away or receive extra electrons to achieve stability. The way that atoms will form bonds with others, and the ease with which they will do it, is determined by the configuration of electrons.
As elements are ordered in the periodic table by atomic number, it can be seen that their position in the table can be used to predict how they will react.
In addition to showing that electrons are restricted to orbits, Bohr’s model also suggested that
Bohr called the jump to another orbit a quantum leap.
Although it contained elements of quantum theory, the Bohr model had its flaws. It ignored the wave character of the electron. Work by WERNER KARL HEISENBERG later tackled these weaknesses.
Bohr’s theory of complementarity states that electrons may be both a wave and a particle, but that we can only experience them as one or the other at any given time. He showed that contradictory characteristics of an electron could be proved in separate experiments and none of the results can be accepted singly – we need to hold all the possibilities in mind at once. This requires a slight adjustment to the original model of atomic structure, we can no longer say that an electron occupies a particular orbit, but can only give the probability that it is there.
In 1939 he developed a theory of nuclear fission with Jon Archibald Wheeler (b.1911) and realised that the 235uranium isotope would be more susceptible to fission than the more commonly used 238uranium.
The element bohrium is named after him.
1914 – Manchester, England
‘Moseley’s law – the principle outlining the link between the X-ray frequency of an element and its atomic number’
Working with ERNEST RUTHERFORD’s team in Manchester trying to better understand radiation, particularly of radium, Moseley became interested in X-rays and learning new techniques to measure their frequencies.
A technique had been devised using crystals to diffract the emitted radiation, which had a wavelength specific to the element being experimented upon.
In 1913, Moseley recorded the frequencies of the X-ray spectra of over thirty metallic elements and deduced that the frequencies of the radiation emitted were related to the squares of certain incremental whole numbers. These integers were indicative of the atomic number of the element, and its position in the periodic table. This number was the same as the positive charge of the nucleus of the atom (and by implication also the number of electrons with corresponding negative charge).
By uniting the charge in the nucleus with an atomic number, a vital link had been found between the physical atomic make up of an element and its chemical properties, as indicated by where it sits in the periodic table.
This meant that the properties of an element could now be considered in terms of atomic number rather than atomic weight, as had previously been the case – certain inconsistencies in the MENDELEEV version of the periodic table could be ironed out. In addition, the atomic numbers and weights of several missing elements could be predicted and other properties deduced from their expected position in the table.
‘An acid is a molecule or ion capable of donating a proton ( a hydrogen nucleus, H+ ) in a chemical reaction, while a base is a molecule or ion capable of accepting one’
The Bronsted-Lowry concept extends Arrhenius’ concept as it includes reactions that take place in the absence of water.
The strongest known acid is an 80 percent solution of antimony pentaflouride in hydrofluoric acid.
The strongest known base is caesium hydroxide.
<< top of page