Unit 3 - Investigative and Pratical Skills in AS Chemistry
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Summary of AQA AS-Level Chemistry
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Unit 3 - Investigative and Pratical Skills in AS Chemistry
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Unit 3 - Investigative and Pratical Skills in AS Chemistry
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The Atom
Atoms are the stuff all elements and compounds are made up of. they’re made of three types of
particle- protons, neutrons and electrons. Electrons have a 1- charge and they whizz around the
nucleus in orbitals. Most of the mass of the atom is concentrated in the nucleus. The diameter
of the nucleus is rather tiny compared to the whole atom. The nucleus is where you find the
protons and neutrons. The mass and charge of these subatomic particles is really small, so
relative mass and relative charge are used instead:
Particle Relative Mass Relative Charge
Proton 1 +1
Neutron 1 0
Electron 1/2000 -1
You can figure out the number of protons, neutrons and electrons from the nuclear symbol.
A- Mass (nucleon) Number. This tells you the total number of protons and neutrons
in the nucleus.
Z- Atomic (Proton) Number. This is the number of protons in the nucleus- This
identifies the element. All atoms of the same element have the same number of
protons
X- Element Symbol
For neutral atoms, which have no overall charge, the number of electrons is the same as the
number of protons. The number of neutrons is just mass number minus atomic number.
Ions have different numbers of protons and electrons. Negative ions have more electrons
than protons. E.g. Br- The negative charge means that there is 1 more electron than there are
protons. Br has 35 protons, so Br- has 36 electrons. Positive ions have fewer electrons than
protons. E.g. Mg2+ The 2+ charge means that there’s 2 fewer electrons than there are protons.
Mg has 12 protons, so Mg2+ must have 10 electrons.
Isotopes are atoms of the same element with different numbers of neutrons but the same
number of protons. It is the number and arrangement of electrons that decides the chemical
properties of an element. Isotopes have the same configuration of electrons, so they’ve got th
same chemical properties. Isotopes of an element do have slightly different physical properties
though, such as different densities, rates of diffusion, etc. This is because physical properties
tend to depend more on the mass of the atom.
,Atomic Models
The accepted model of the atom has changed throughout history. The model of the atom that
is currently accepted fits all the observations and evidence we have so far, so we assume its
true until someone shows that it is incomplete or wrong. In the past, completely different models
were accepted, because they fitted the evidence at the time.
Some ancient Greeks thought that all matter was made from indivisible particles. At the start of
the 19th century John Dalton described atoms as solid spheres, and said that different types of
sphere made up the different elements. But as scientists did more experiments, our currently
accepted models began to emerge, with modifications or refinements being made to take
account of new evidence.
Experimental evidence showed that atoms weren’t solid spheres. In 1897 JJ Thompson did
a whole series of experiments and concluded that atoms weren’t solid and indivisible. His
measurements of charge and mass showed that an atom must contain smaller, negatively
charged particles. He called these particles ‘corpuscles’- we call them electrons. The ‘solid
sphere’ idea of atomic structure had to be changed. The new model was known as the plum
pudding model- a positively charged sphere with negative electrons embedded in it.
Rutherford showed that the plum pudding model was wrong however. In 1909, Ernest
Rutherford and his students Hans Geiger and Ernest Marsden conducted the famous Gold Foil
experiment. They fired alpha particles, which are positively charged, at an extremely thin sheets
og gold. From the plum pudding model, they were expecting most of the alpha particles to be
deflected very slightly by the positive pudding that made up most of an atom. In fact, most of the
alpha particles passed straight through the gold atoms, and a very small number were deflected
backwards. This showed that the plum pudding model couldn’t be right. So Rutherford can up
with a model that could explain this new evidence- the nuclear model of the atom:
A few alpha particles Most of the alpha
particles
are deflected very pass straight through the
strongly by the empty space
nucleus
There is a tiny, positively charged nucleus at the center of the atom, where most of the atom’s
mass is concentrated. The nucleus is surrounded by a ‘cloud’ of negative electrons and most of
the atom is empty space.
Rutherford’s model was modified several times. It seemed pretty convincing but the scientists
of the day, continued with their experiments, wanting to be sure of the truth. Henry Moseley
discovered the charge of the nucleus increase from one element to another in units of one.
,This led Rutherford to investigate the nucleus further. He finally discovered that it contained
positively charged particles that he called protons. The charges of the nuclei of different atoms
could then be explained- the atoms of different elements have a different number of protons
in their nucleus. there was still one problem with the model though- the nuclei of atoms were
heavier than tney would be if they contained just protons. Rutherford predicted that there
were other particles int h nucleus, that had mass but no charge- and the neutron as eventually
discovered by James Chadwick.
The Bohr model was a further improvement. Scientists realised that electrons in a ‘cloud’ around
the nucleus of an atom would spiral down into the nucleus and cause the atom to collapse. Niels
Bohr proposed a new model of the atom with four basic principles:
1) Electrons can only exist in fixed orbits, or shells, and not anywhere in between
2) Each shell has a fixed energy
3) When an electron moves between shells electromagnetic radiation is emitted or absorbed
4) Because the energy of shells is fixed, the radiation will have a fixed frequency
The frequencies of radiation emitted and absorbed by atoms were already known from
experiments. The Bohr model fitted these observations. It also explained why some elements
(the noble gases) are inert. He said that the shells of an atom can only hold fixed numbers of
electrons, and that an element’s reactivity is due to its electrons. When an atom has full shells of
electrons it is stable and does not react.
There is more than one model of atomic structure in use today. We now know that the Bohr
model is not perfect- but it’s still widely used to describe atoms because it’s simple and explains
many observations from experiments, like bonding and ionisation energy trends. The most
accurate model we have today involves complicated quantum mechanics. Basically, you can
never know where an electron is or in which direction it’s going in at any moment, but you can
say how likely it is to be at any particular point in the atom.
This model might be more accurate but it’s a lot harder to get your head round and visualise.
It does explain some observations that can’t be accounted for by the Bohr model though. So
scientists use which ever model is most relevant to what they’re investigating.
, Relative Mass
Relative masses are masses of atoms compared to carbon-12. The actual mass of an atom is
very, very tiny. So, the mass of one atom is compared to the mass of a different atom. This is its
relative mass.
The Relative Atomic Mass, Ar, is the average mass of an atom of an element compared to 1/
12 of the mass of an atom of carbon-12
The Relative Isotopic Mass is the mass of an atom if an isotope of an element compared to 1/
12 of the mass of an atom of carbon-12
The Relative Molecular Mass, Mr, is the average mass of a molecule or formula unit compared
to 1/12 of the mass of an atom of carbon-12
Ar can be worked out from isotopic abundances. Different isotopes of an element occur in
different isotopic abundances. For example, 76% of chlorine atoms found on earth have
a relative isotopic mass of 35, while 24% have a relative isotopic mass of 37. The relative
isotopic mass of chlorine is the average mass of all chlorine atoms. If you’ve got the isotopic
abundances as percentages, the easiest way to calculate relative atomic mass is to:
Step 1: Multiply each relative isotopic mass by its % relative isotopic abundance and add up the
results- (76x35) + (24x37) = 2660 + 888 = 3548
Step 2: Divide by 100- 3548/100 = 35.5
So the relative atomic mass of chlorine is 35.5.
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