New OCR Gateway specification from September
2016 Higher tier: grades 9 to 4:
In this and subsequent posts I’m simply going to explain and illustrate each learning objective as they come up in the topics in the new GCSE specification.
You can really get ahead of your class if you follow this blog and all the posts that will appear here about the new GCSEs over the coming months.
This rejigging of the specification is just that: there is nothing really new here since “atomic structure” and the particle model has been with us for the past half century at least.
That written in italics is for the higher tier paper only.
C1.2a describe how and why the atomic model has changed over time.
The first mention of a model of the atom is found in Greek literature.
The Greek philosopher Democritus believed (he had no experimental evidence at the time) that all matter was made up of small (though he had no idea how small) indivisible particles (neither did he have any idea that the atom could be split and contain sub–atomic particles).
He was the person who coined the word “atom” for these particles.
You can read more about the ancient Greeks and atoms here.
Atomic Models:
Atomic models are our human attempts to picture the invisible.
All pictures of atoms have limitations because of the ways in which the picture is obtained.
The models I describe below are those that are required for the GCSE Gateway course.
We work our way through them historically showing how each is improved upon by the next.
We start with the description of the atom from the British naturalist John Dalton and finish with the model developed in 1913 by the Danish chemist Niels Bohr.
Dalton:
John Dalton (1766-1844) first put forward his theory of the atom in the early 1800’s.
Dalton believed (note that word “believed” you see Dalton had very little experimental evidence to back up his theory at the time so he was effectively shooting in the dark which for a supposed scientist is acting very irrationally!!) nevertheless Dalton believed atoms to exist.
He believed that all stuff was made up of them.
He believed them to be indivisible and indestructible.
He believed some substances (he called them elements) were composed of just one type of atom and these atoms were all identical because they had the same mass and the same properties.
He believed compounds formed when two or more different kinds of atoms combined in some way.
He believed a chemical reaction must involve a rearrangement of the atoms of the substances taking part.
And the thing is John Dalton was really on the right lines.
It was only much later that atoms were shown to be divisible and could be destroyed and that atoms of the same element could have different masses that we now call “isotopes”.
JJ Thomson (1856-1940):
JJ Thomson has his place in the history of the development of atomic theory because he was the first person to suggest that there were particles smaller than the atom and that the atom was therefore divisible.
In experiments performed in 1897, he observed that electrically charged plates deflected cathode rays in a vacuum tube.
His brilliant intuition was to interpret the results to suggest that cathode “rays” were actually composed of particles smaller than the atom.
He also calculated the mass/charge ratio of these particles.
He also worked out an estimate of the size of the charge of these particles.
We now know the particles to be electrons and to carry a negative charge of 1.6×10–19 Coulombs.
He suggested that the atom consisted of a “blob” of positive charge in which the negative electrons were stuck much like sultanas in a Christmas pudding!!!
In fact, his model became known as the “Plum Pudding Model of the atom”.
You can find out more about JJ Thomson here.
Ernest Rutherford, Hans Geiger and Ernest Marsden:
Rutherford’s place in the pantheon of early physicists investigating the structure of the atom is secure.
It was Rutherford who suggested to Geiger and Marsden an experiment to test the JJ Thomson “Plum Pudding theory” of the atom.
Beginning I September 1907 Rutherford began investigating the scattering of alpha particles.
Rutherford said at the time that, “I was brought up to look at the atom as a hard fellow, red or grey in colour according to taste.”
But by 1907 he had become convinced that the atom was largely empty space and not a hard fellow at all.
A German physicist in 1903, Philipp Lenard, bombarding atoms with electrons (cathode rays) had shown as much.
As Lenard put it “the space occupied by a cubic meter of solid platinum was as empty as the space of stars beyond the earth”!!!
That was not all, passing alpha particles through mica Rutherford obtained result to suggest that the alphas were being scattered i.e. the atoms in the mica were deflecting the alpha particles.
What was going on?
Working in Manchester with Hans Gieger, the two men came up with a way of measuring the alpha scattering that allowed them to see the effects of the scattered alpha particles.
Gieger set to work with his 18 year old assistant Ernest Marsden shooting alpha particles at foils of different metals.
As expected from the Thomson Plum Pudding Model, they observed gentle deflection of the stream of alpha particles passing through the thin metal foils.
But what they did not expect was to see stray particles not passing through.
Rutherford suggested to the two young experimenters that they repeat the experiment but this time look for alpha particles that had reflected from the metal foils.
When they did so they found to their surprise that the gold foil reflected alpha particles back from the foil. Not all the particles were passing straight through.
Diagram Rutherford experiment
AB is the radiation source a glass the full of radioactive gas
RR is the gold foil
P is a lead sheet to prevent alpha particles hitting the scintillation screen directly.
S is the scintillation screen to detect reflected alpha particles
M is the microscope used to observe flashes on the screen from the alpha particles.
To them all, this was a momentous result for how could it be that a soft blob of positive charge could reflect small, solid alpha particles?
As Rutherford famously wrote at the time:
“It was quite the most incredible event that has ever happened to me in my life….It was almost as incredible as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you… I realised that this scattering backwards must be the result of a single collision (where) the greatest part of the mass of the atom was concentrated in a minute nucleus.”
Rutherford had blown apart the Plum Pudding Model of the atom.
His new model had a very small nucleus at its centre with electrons “orbiting” around it.
Rutherford
model of the atom
Rutherford went public with his historic findings on March 7th 1911 at a meeting of the Manchester Literary and Philosophical Society!
But there was a problem with the Rutherford model of the atom.
How did the electrons work?
If two or more electrons rotated around the nucleus they would interact with each other, start oscillating and quickly tear the atom apart!
There’s more about the Rutherford atomic model here.
Niels Bohr:
Niels Bohr was a brilliant Danish scientist. He arrived at Rutherford’s Manchester University laboratory in late 1911 looking to improve his understanding of the latest developments in physics.
He took up with the problem of the Rutherford atomic model flying apart.
And by late 1913 he had come up with a novel solution.
He proposed that there must be what he called “stationary states” in the atom.
These were orbits that the electrons could occupy without the atom becoming unstable, without radiating light and without spiralling into the nucleus.
Bohr’s model prevailed because experiments in spectroscopy confirmed his theoretical predictions.
When atoms of metals in their salts are heated they emit light in what is called a flame test
Flame test colours.
The light is seen not as a rainbow like light from the sun but as a series of coloured lines in the spectrum.
Bohr’s model was able to explain how the lines of colour in the spectra were formed.
Electrons can move up to higher orbits and then, letting out the energy gained as light, drop back to lower orbits.
It is Bohr’s model of the atom that we accept as our model of the atom at GCSE level.
There’s more about the Bohr atomic model here
You can also find a very detailed timeline of the development of the model of the atom here
C1.2b describe the atom as a positively charged nucleus surrounded by negatively charged electrons, with the nuclear radius much smaller than that of the atom and with most of the mass in the nucleus.
Bohr’s model of the atom is the model we use today at GCSE.
The mass of the atom is concentrated in the very small nucleus in the centre of the atom.
The nucleus is so small it about 100000 times smaller than the atom itself.
The electrons orbit the nucleus in shells or energy levels. (Bohr called them stationary states.)
Particles in the nucleus are called protons and neutrons.
The number of protons is equal to the number of electrons.
And most of the atom is empty space literally there is nothing, no air, no stuff, no eather between the protons and electrons.
Get your head round that!!!
So when your kind teacher tells you “you are empty space” there is more than a grain of truth in their statement!!!
C1.2c recall the typical size (order of magnitude) of atoms and small molecules.
The concept that typical atomic radii and bond length are of the order of 10–10m.
As you probably now realise atoms and molecules are incredibly small.
They are 10–10m in diameter.
That’s 10 billion of them lined up on a meter stick!!
In an atom if a grain of sand is the nucleus then put the nucleus at the centre of a football stadium and the electrons are in the stands!!.
C1.2d recall relative charges and approximate relative masses of protons, neutrons and electrons.
C1.2e calculate numbers of protons, neutrons and electrons in atoms and ions, given atomic number and mass number of isotopes.
The definitions of an ion, atomic number, mass number and an isotope, also the standard notation to represent these.
Atomic number (Z): This is the number of protons in the atomic nucleus
Mass Number (A): This is the number of protons and neutrons in the atomic nucleus
Atoms are particles that contain equal numbers of protons and electrons.
Ions
An atom with more protons than electrons is positively charged and is called an Ion i.e. a cation (pronounced kat–ion)
An atom with more electrons than protons is negatively charged (it is also an Ion) and is called an anion (pronounced an–ion)
Isotopes are atoms of the same element but they contain different numbers of neutrons in their nuclei.
Isotopes are shown like this:
So you’ll be shown the symbol for an isotope and asked to work out the numbers of protons, neutrons and electrons in its atom.
So here is an isotope symbol:
It has 84 protons and neutrons (the mass number)
It has 36 protons (the atomic number) and therefore it has 36 electrons
Therefore it also has 84 – 36 neutrons = 48 neutrons
In this and subsequent posts I’m simply going to explain and illustrate each learning objective as they come up in the topics in the new GCSE specification.
You can really get ahead of your class if you follow this blog and all the posts that will appear here about the new GCSEs over the coming months.
This rejigging of the specification is just that: there is nothing really new here since “atomic structure” and the particle model has been with us for the past half century at least.
That written in italics is for the higher tier paper only.
C1.2a describe how and why the atomic model has changed over time.
The first mention of a model of the atom is found in Greek literature.
The Greek philosopher Democritus believed (he had no experimental evidence at the time) that all matter was made up of small (though he had no idea how small) indivisible particles (neither did he have any idea that the atom could be split and contain sub–atomic particles).
He was the person who coined the word “atom” for these particles.
You can read more about the ancient Greeks and atoms here.
Atomic Models:
Atomic models are our human attempts to picture the invisible.
All pictures of atoms have limitations because of the ways in which the picture is obtained.
The models I describe below are those that are required for the GCSE Gateway course.
We work our way through them historically showing how each is improved upon by the next.
We start with the description of the atom from the British naturalist John Dalton and finish with the model developed in 1913 by the Danish chemist Niels Bohr.
Dalton:
John Dalton (1766-1844) first put forward his theory of the atom in the early 1800’s.
Dalton believed (note that word “believed” you see Dalton had very little experimental evidence to back up his theory at the time so he was effectively shooting in the dark which for a supposed scientist is acting very irrationally!!) nevertheless Dalton believed atoms to exist.
He believed that all stuff was made up of them.
He believed them to be indivisible and indestructible.
He believed some substances (he called them elements) were composed of just one type of atom and these atoms were all identical because they had the same mass and the same properties.
He believed compounds formed when two or more different kinds of atoms combined in some way.
He believed a chemical reaction must involve a rearrangement of the atoms of the substances taking part.
And the thing is John Dalton was really on the right lines.
It was only much later that atoms were shown to be divisible and could be destroyed and that atoms of the same element could have different masses that we now call “isotopes”.
JJ Thomson (1856-1940):
JJ Thomson has his place in the history of the development of atomic theory because he was the first person to suggest that there were particles smaller than the atom and that the atom was therefore divisible.
In experiments performed in 1897, he observed that electrically charged plates deflected cathode rays in a vacuum tube.
His brilliant intuition was to interpret the results to suggest that cathode “rays” were actually composed of particles smaller than the atom.
He also calculated the mass/charge ratio of these particles.
He also worked out an estimate of the size of the charge of these particles.
We now know the particles to be electrons and to carry a negative charge of 1.6×10–19 Coulombs.
He suggested that the atom consisted of a “blob” of positive charge in which the negative electrons were stuck much like sultanas in a Christmas pudding!!!
In fact, his model became known as the “Plum Pudding Model of the atom”.
You can find out more about JJ Thomson here.
Ernest Rutherford, Hans Geiger and Ernest Marsden:
Rutherford’s place in the pantheon of early physicists investigating the structure of the atom is secure.
It was Rutherford who suggested to Geiger and Marsden an experiment to test the JJ Thomson “Plum Pudding theory” of the atom.
Beginning I September 1907 Rutherford began investigating the scattering of alpha particles.
Rutherford said at the time that, “I was brought up to look at the atom as a hard fellow, red or grey in colour according to taste.”
But by 1907 he had become convinced that the atom was largely empty space and not a hard fellow at all.
A German physicist in 1903, Philipp Lenard, bombarding atoms with electrons (cathode rays) had shown as much.
As Lenard put it “the space occupied by a cubic meter of solid platinum was as empty as the space of stars beyond the earth”!!!
That was not all, passing alpha particles through mica Rutherford obtained result to suggest that the alphas were being scattered i.e. the atoms in the mica were deflecting the alpha particles.
What was going on?
Working in Manchester with Hans Gieger, the two men came up with a way of measuring the alpha scattering that allowed them to see the effects of the scattered alpha particles.
Gieger set to work with his 18 year old assistant Ernest Marsden shooting alpha particles at foils of different metals.
As expected from the Thomson Plum Pudding Model, they observed gentle deflection of the stream of alpha particles passing through the thin metal foils.
But what they did not expect was to see stray particles not passing through.
Rutherford suggested to the two young experimenters that they repeat the experiment but this time look for alpha particles that had reflected from the metal foils.
When they did so they found to their surprise that the gold foil reflected alpha particles back from the foil. Not all the particles were passing straight through.
AB is the radiation source a glass the full of radioactive gas
RR is the gold foil
P is a lead sheet to prevent alpha particles hitting the scintillation screen directly.
S is the scintillation screen to detect reflected alpha particles
M is the microscope used to observe flashes on the screen from the alpha particles.
To them all, this was a momentous result for how could it be that a soft blob of positive charge could reflect small, solid alpha particles?
As Rutherford famously wrote at the time:
“It was quite the most incredible event that has ever happened to me in my life….It was almost as incredible as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you… I realised that this scattering backwards must be the result of a single collision (where) the greatest part of the mass of the atom was concentrated in a minute nucleus.”
Rutherford had blown apart the Plum Pudding Model of the atom.
His new model had a very small nucleus at its centre with electrons “orbiting” around it.
Rutherford went public with his historic findings on March 7th 1911 at a meeting of the Manchester Literary and Philosophical Society!
But there was a problem with the Rutherford model of the atom.
How did the electrons work?
If two or more electrons rotated around the nucleus they would interact with each other, start oscillating and quickly tear the atom apart!
There’s more about the Rutherford atomic model here.
Niels Bohr:
Niels Bohr was a brilliant Danish scientist. He arrived at Rutherford’s Manchester University laboratory in late 1911 looking to improve his understanding of the latest developments in physics.
He took up with the problem of the Rutherford atomic model flying apart.
And by late 1913 he had come up with a novel solution.
He proposed that there must be what he called “stationary states” in the atom.
These were orbits that the electrons could occupy without the atom becoming unstable, without radiating light and without spiralling into the nucleus.
Bohr’s model prevailed because experiments in spectroscopy confirmed his theoretical predictions.
When atoms of metals in their salts are heated they emit light in what is called a flame test
Flame test colours.
The light is seen not as a rainbow like light from the sun but as a series of coloured lines in the spectrum.
Bohr’s model was able to explain how the lines of colour in the spectra were formed.
Electrons can move up to higher orbits and then, letting out the energy gained as light, drop back to lower orbits.
It is Bohr’s model of the atom that we accept as our model of the atom at GCSE level.
There’s more about the Bohr atomic model here
You can also find a very detailed timeline of the development of the model of the atom here
C1.2b describe the atom as a positively charged nucleus surrounded by negatively charged electrons, with the nuclear radius much smaller than that of the atom and with most of the mass in the nucleus.
Bohr’s model of the atom is the model we use today at GCSE.
The mass of the atom is concentrated in the very small nucleus in the centre of the atom.
The nucleus is so small it about 100000 times smaller than the atom itself.
The electrons orbit the nucleus in shells or energy levels. (Bohr called them stationary states.)
Particles in the nucleus are called protons and neutrons.
The number of protons is equal to the number of electrons.
And most of the atom is empty space literally there is nothing, no air, no stuff, no eather between the protons and electrons.
Get your head round that!!!
So when your kind teacher tells you “you are empty space” there is more than a grain of truth in their statement!!!
C1.2c recall the typical size (order of magnitude) of atoms and small molecules.
The concept that typical atomic radii and bond length are of the order of 10–10m.
As you probably now realise atoms and molecules are incredibly small.
They are 10–10m in diameter.
That’s 10 billion of them lined up on a meter stick!!
In an atom if a grain of sand is the nucleus then put the nucleus at the centre of a football stadium and the electrons are in the stands!!.
C1.2d recall relative charges and approximate relative masses of protons, neutrons and electrons.
| Proton |
Electron |
Neutron |
|
Relative Mass
|
1 amu |
1/1837
amu |
1 amu |
| Relative
charge |
+1 |
–1 |
0 |
| Position
in atom |
nucleus |
shells |
nucleus |
C1.2e calculate numbers of protons, neutrons and electrons in atoms and ions, given atomic number and mass number of isotopes.
The definitions of an ion, atomic number, mass number and an isotope, also the standard notation to represent these.
Atomic number (Z): This is the number of protons in the atomic nucleus
Mass Number (A): This is the number of protons and neutrons in the atomic nucleus
Atoms are particles that contain equal numbers of protons and electrons.
Ions
An atom with more protons than electrons is positively charged and is called an Ion i.e. a cation (pronounced kat–ion)
An atom with more electrons than protons is negatively charged (it is also an Ion) and is called an anion (pronounced an–ion)
Isotopes are atoms of the same element but they contain different numbers of neutrons in their nuclei.
Isotopes are shown like this:
So you’ll be shown the symbol for an isotope and asked to work out the numbers of protons, neutrons and electrons in its atom.
So here is an isotope symbol:
It has 84 protons and neutrons (the mass number)
It has 36 protons (the atomic number) and therefore it has 36 electrons
Therefore it also has 84 – 36 neutrons = 48 neutrons
Here
are a few more for you to work out:
