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A level Chemistry Revision

Useful as a recap of A level material. Mainly quantum.
Module

General and Organic Chemistry (AOC106DI)

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Academic year: 2018/2019
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Atomic structure and the Periodic table

1 Atomic structure

Structure of the atom and isotopes (1) + Relative masses (2)

Understanding

How can you find the mass of something as small as an atom?

Why are atomic masses unitless?

Finding the mass of an atom must be possible, since a quick glance at any chemistry book

will show very accurate atomic masses, Hydrogen at 1, Oxygen at 15, and so

on. Look closely and you'll see no unit of measurement on those atomic masses.

This is because they are relative masses where the mass of one atom is compared to the

other.

Although we can't find the mass of any individual atom, we can easily find the mass of a

bunch. If we could find the mass of the same sized bunch of atoms of two different

elements, say Hydrogen and Oxygen, then we could compare their masses. Provided that

we have the same number of atoms in each bunch, this is easy to do.

Fortunately we can do this if we believe in Avogadro's hypothesis that "equal volumes of

gases at the same temperature and pressure contain equal numbers of molecules."

Suppose that we take two equal sized flasks of gas, one filled with hydrogen at room

temperature and atmospheric pressure, and the other filled with oxygen at the same

conditions.

Hydrogen is the lightest substance so it makes sense to assign it a mass of 1. You can

see that an oxygen atom is 16 times as heavy. It has a mass of 16 on a relative scale.

If hydrogen is the lightest of all substances, then why not give it a mass of exactly 1 on our

relative mass scale?

The main reason is historical:

 John Dalton, the father of atomic theory, initially proposed using hydrogen as the basis. Issues of measurability and repeatability quickly cropped up. So did mistakes. Dalton, for example, thought water was HO rather than H 2 O.  Those issues were one of the reasons chemists switched to an Oxygen-based standard. The other reason was that Oxygen combines with a lot of things to form oxides, giving greater ease of chemical analysis.  Also, with Oxygen the atomic masses of most of the elements become approximately whole numbers but with Hydrogen as standard the atomic masses of most of the elements are fractional.

 Physicists' investigations at the atomic level caused them to develop their own standard in the 20th century, based on 16 O rather than the natural mix of O 16 O, 17 O, and 18 O used by chemists.

 As the composition of naturally occurring oxygen is not constant, the chemists’ scale became very confusing.  Physicists had a solution: Switch to their isotopically pure 16 O standard.  This would have represented an unacceptably large change in chemistry's natural oxygen-based standard. It would have required textbooks, reference books, and perhaps most importantly, the recipes used at refineries and other chemical factories to have been rewritten. The commercial costs would have been immense.

 The carbon-12 based standard presented a nice compromise. By chance, defining 1amu as 1/16th of the mass of a atom of a natural mix of O 16 O, 17 O, and 18 O is very close to defining it as 1/12 the mass of a atom of 12 C

So we could have set hydrogen to be exactly 1, but then we'd have had to really revise the atomic weight table back in 1961. If hydrogen was assigned a mass of 1 exactly, then oxygen would have become 15, quite a difference from the mass chemists were using.

If a hydrogen atom has only one proton, and carbon-12 has 6 protons and 6 neutrons to make up its mass of twelve, why isn't the mass of hydrogen 1/12 of that of carbon-12?

Mass of a

hydrogen

atom

Mass of sub-atomic particles

Mass of a carbon- atom

Relative atomic mass = the weighted mean of the relative isotopic masses of an element

An amu = 1/12 of the mass of an carbon-12 atom

Avogadro’s constant = the number of atoms in 12 grams of carbon-

We have to use relative isotopic and atomic mass, instead of mass number, due to mass defect (where binding energy creates a difference between the mass number of a nucleus and its true measured mass)

Mass defect is why the mass of a proton or neutron by itself is more than (and not equal to) 1 amu

Mass spectrometry (2)

Understanding

How a mass spectrometer works

The basic principle

If something is moving and you subject it to a sideways force, instead of moving in a

straight line, it will move in a curve - deflected out of its original path by the sideways

force.

The amount of deflection you will get for a given sideways force depends on the mass

of the object. If you knew the speed of the ball and the size of the force, you could

calculate the mass of the ball if you knew how much it deflected. The less the

deflection, the heavier the ball.

You can apply exactly the same principle to atomic sized particles.

An outline of what happens in a mass spectrometer

Atoms and molecules can be deflected by magnetic fields - provided the atom or

molecule is first turned into an ion. Electrically charged particles are affected by a

magnetic field although electrically neutral ones aren't.

Stage 1: Ionisation

The atom or molecule is ionised by knocking one or more electrons off to give a

positive ion. This is true even for things, which you would normally expect to form

negative ions (chlorine, for example) or never form ions at all (argon, for example).

Most mass spectrometers work with positive ions.

Stage 2: Acceleration

The ions are accelerated so that they all have the same kinetic energy.

Stage 3: Deflection

The ions are then deflected by a magnetic field.

Stage 4: Detection

The beam of ions passing through the machine is detected electrically.

A full diagram of a mass spectrometer

Understanding what's going on

The need for a vacuum

It's important that the ions produced in the ionisation chamber have a free run through

the machine without hitting air molecules.

The number of isotopes The 5 peaks in the mass spectrum shows that there are 5 isotopes of zirconium - with relative isotopic masses of 90, 91, 92, 94 and 96 on the 12 C scale.

The abundance of the isotopes The relative abundances are given as percentages. You can find these relative abundances by measuring the lines on the stick diagram.

The mass spectrum of diatomic elements

The mass spectrum of chlorine

Chlorine has two isotopes, 35 Cl and 37 Cl, in the approximate ratio of 3 atoms of 35 Cl to 1 atom of 37 Cl. You might suppose that the mass spectrum would look like this:

You would be wrong!

The problem is that chlorine consists of molecules, not individual atoms.

When chlorine is passed into the ionisation chamber, an electron is knocked off the molecule to give a molecular ion, Cl 2 +.

These ions won't be particularly stable, and some will fall apart to give a chlorine atom and a Cl+ ion. The term for this is fragmentation.

The Cl+ ions will pass through the machine and will give lines at 35 and 37 and you would get exactly the pattern in the last diagram.

The problem is that you will also record lines for the unfragmented Cl 2 + ions.

Think about the possible combinations of chlorine-35 and chlorine-37 atoms in a Cl 2 + ion. Both atoms could be 35 Cl, both atoms could be 37 Cl, or you could have one of each sort. That would give you total masses of the Cl 2 + ion of:

35 + 35 = 70
35 + 37 = 72
37 + 37 = 74

That means that you would get a set of lines in the m/z = 70 region looking like this:

What you can't do is make any predictions about the heights of the lines at 35/37 compared with those at 70/72/74. That depends on what proportion of the molecular ions break up into fragments.

The overall mass spectrum looks like this:

An orbital is a sort of 3D map of the places that the electron is likely to be found. In the hydrogen case, the electron can be found anywhere within a spherical space surrounding the nucleus. What is the electron doing in the orbital? We don't know, we can't know, and so we just ignore the problem!

4 values, called quantum numbers, describe the characteristics of a particular quantum state (of an electron).

Summary of the Quantum Numbers

The principle quantum number n This describes the average distance of the orbital from the nucleus, the size of the orbital and the energy of the electron in it. It’s similar to Bohr’s energy levels.

The quantum number l (angular momentum)

The shape of the orbital (or the angular momentum) is limited by the principle quantum number (the energy level).

E. if the n value is 3, this means that the 3rd electron shell has 3 sub-shells - 0,1 and 2.

The magnetic number m 1

This is read as "one s two" - not as "one s squared".

Filling orbitals

Pauli Exclusion Principle The principle states that, in an atom or molecule, no two electrons can have the same four electronic quantum numbers. As an orbital can contain a maximum of only two electrons, the two electrons must have opposing spins. This means if one is assigned an up-spin (+1/2), the other must be down-spin (-1/2).

Hund's rule Hund's rule states that:

 Where there is a choice between orbitals of equal energy (i. orbitals in the same subshell), electrons occupy singly as far as possible.  All of the electrons in singly occupied orbitals have the same spin.

Electrons are negatively charged and, as a result, they repel each other. Electrons tend to minimize repulsion by occupying their own orbitals, rather than sharing an orbital with another electron.

Furthermore, quantum-mechanical calculations have shown that the electrons in singly occupied orbitals with parallel spins are less effectively screened or shielded from the nucleus.

The order of filling orbitals - the Aufbau Principle Aufbau is a German word meaning building up or construction. Electrons fill low energy orbitals before they fill higher energy ones.

The diagram below summarises the electron configurations that you may need to know when you are using the Aufbau Principle in A-level.

The aufbau diagram above gives the overall configuration correctly in all but about 20 cases.

It is only when one questions the order of filling that this approach gives the wrong answer.

When predicting the way that the electrons fill in Scandium, we might suppose that the final three electrons would all enter into some 3d orbitals to give 1s 2 , 2s 2 , 2p 6 , 3s 2 , 3p 6 , 3d 3.

It is natural to question why one or two electrons are usually pushed into a higher energy orbital. The answer is because 3d orbitals are more compact than 4s, and as a result any electrons entering 3d orbitals will experience greater mutual repulsion.

NOTE: The slightly unsettling feature is that although the relevant s orbital can relieve such additional electron-electron repulsion, different atoms do not always make full use of this form of sheltering (e. Chromium and Copper) because the situation is more complicated than just described. One thing to consider is that nuclear charge increases as we move through the atoms in the periodic table, and there is a complicated set of interactions between the electrons and the nucleus as well as between the electrons themselves. This is what ultimately produces an electronic configuration and, contrary to what some educators may wish for, there is no simple qualitative rule of thumb that can cope with this complicated situation.

Looking at the table above, you can see the electronic configuration of the transition metals (from Scandium to Copper). The definition of a transition metal = an element that has at least 1 stable ion with a partly filled d orbital. Zinc, therefore, is not a transition metal.

Ionisation energies (4)

Things to remember:

Defining ionisation energies

The ionisation energy of an element = the energy required to remove an electron from each atom in one mole of that gaseous atom.

First ionisation energy = the energy required to remove the most loosely held electrons (those in the outermost energy levels).

This can also be shown in symbol form: −¿ +¿+ e ¿ X ( g ) ⟶X (¿ g )

Second ionisation energy: −¿ 2 +¿+ e ¿ +¿ ⟶X (¿ g ) X (¿ g )

Things to note

The state symbol - (g) - is essential. When you are talking about ionisation energies, everything must be present in the gas state.

Ionisation energies are measured in kJ mol-1 (kilojoules per mole).

All elements have a first ionisation energy - even atoms that don't form positive ions in experiments. The reason that helium doesn't normally form a positive ion is because of the huge amount of energy that would be needed to remove one of its electrons.

Patterns of first ionisation energies in the Periodic Table

The first 20 elements

First ionisation energy shows periodicity. That means thatit varies in a repetitive way as you move through the Periodic Table. For example, look at the pattern from Li to Ne, and then compare it with the identical pattern from Na to Ar.

Factors affecting the size of ionisation energy

Ionisation energy is a measure of the energy needed to pull a particular electron away from the attraction of the nucleus.

Group 3

Group 2

Electron-electron repulsion raises the energy of the involved electrons above the amount they would have if there were no repulsion.

This repulsion exists between:

  • two electrons in the same orbital (usually the weakest)
  • electrons in different orbitals, within a given quantum shell or subshell
  • electrons in adjacent quantum shells (usually the strongest)

The orbital of the electron.

The lower the quantum shell number (n), the higher the probability of finding the electron close to the nucleus (or as the textbook states: the closer the electron to the nucleus).

This is because the 1s electrons shield the 2s and 3s electrons - and the 2s electrons shield the 3s electrons.

Understanding:

The orbital shape (angular momentum) also affects the nuclear charge felt by an electron.

The distance of maximum probability for a 2 p electron is slightly less than that for a 2 s electron. However, in contrast to 2 p curve, there is a small additional maxima in the 2 s curve, which lies at or around the maxima for a 1 s orbital. This indicates that the electron in 2 s orbital spends some of its time near the nucleus. In other words, the 2 s electron penetrates into the inner 1 s shell and therefore, is held more tightly than the 2 p electron.

That is the reason why 2 s electron is more stable and has lower energy than a 2 p electron.

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A level Chemistry Revision

Module: General and Organic Chemistry (AOC106DI)

26 Documents
Students shared 26 documents in this course

University: Coventry University

Was this document helpful?
Atomic structure and the Periodic table
1.1 Atomic structure
Structure of the atom and isotopes (1) + Relative masses (2)
Understanding
How can you find the mass of something as small as an atom?
Why are atomic masses unitless?
Finding the mass of an atom must be possible, since a quick glance at any chemistry book
will show very accurate atomic masses, Hydrogen at 1.00794, Oxygen at 15.9994, and so
on. Look closely and you'll see no unit of measurement on those atomic masses.
This is because they are relative masses where the mass of one atom is compared to the
other.
Although we can't find the mass of any individual atom, we can easily find the mass of a
bunch. If we could find the mass of the same sized bunch of atoms of two different
elements, say Hydrogen and Oxygen, then we could compare their masses. Provided that
we have the same number of atoms in each bunch, this is easy to do.
Fortunately we can do this if we believe in Avogadro's hypothesis that "equal volumes of
gases at the same temperature and pressure contain equal numbers of molecules."
Suppose that we take two equal sized flasks of gas, one filled with hydrogen at room
temperature and atmospheric pressure, and the other filled with oxygen at the same
conditions.
Hydrogen is the lightest substance so it makes sense to assign it a mass of 1.0. You can
see that an oxygen atom is 16 times as heavy. It has a mass of 16 on a relative scale.
If hydrogen is the lightest of all substances, then why not give it a mass of exactly 1 on our
relative mass scale?
The main reason is historical:

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