Atomic Structure Notes
Atomic Structure Notes
Atoms are tiny particles which make up matter. They contain a nucleus which is surrounded by electrons. The nucleus itself contains smaller particles called protons and neutrons.
The nucleus, and protons and neutrons, are much more massive than electrons.
Protons and electrons have a very small electrical charge. Protons are positively charged; electrons are negatively charged.
Atoms of a certain element all have the same number of protons and electrons, e.g. all atoms of lithium have 3 protons and 3 electrons.
Atoms can gain or lose electrons so that they have unequal numbers of protons and electrons. These new particles are called ions.
They are indicated by a charge.
e.g 27Al3+ this ion has lost 3 electrons.
18O2- this ion has gained 2 electrons.
The number of neutrons can vary. Isotopes are atoms of the same element with different numbers of neutrons and different mass numbers.
e.g. chlorine, carbon, hydrogen
Isotopes of an element have identical chemical properties, but they can have different physical properties. Isotopes of the same element with more neutrons have:
- higher density
- higher melting and boiling points
- slower rate of diffusion
Many isotopes are radioactive; these are called radioisotopes (radioactive isotopes) The nuclei of radioisotopes breaks down spontaneously emitting radiation – either α, β or γ. Their uses include nuclear power generation, sterilisation of medical equipment, finding cracks and stresses in metal, and the preservation of food. Some radioisotopes of certain elements are particularly useful.
A mass spectrometer is a machine that can:
- measure the relative masses of different isotopes
- measure the relative abundances of the different isotopes in a sample of an element.
It separates positively charged ions according to their mass. A sample that is going to be tested in the mass spectrometer has to be in the gaseous state. If the sample is a liquid or a solid, a heater is used to vaporise some of it. There are five main stages in the operation of the mass spectrometer: vapourisation, ionisation, acceleration, deflection and detection.
A sample is heated at low pressure to turn it into a gas. It is then injected into the instrument.
An electron gun produces a stream of high-energy electrons from a heated metal filament. These high-energy electrons bombard the vaporised sample. When a high-energy electron hits an atom in the sample it knocks an electron out of the atom, forming an ion with a single positive charge:
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X(g) + e– → X+(g) + 2e–
An electric field accelerates the positive ions into the rest of mass spectrometer. Negatively charged plates attract the positive ions from the ionised sample, and the ions are focussed through slits in the plates into a narrow beam.
The beam of ions is deflected around a bend by a magnetic field. The lighter the ion, the more it will be deflected. Electromagnets are used so that the strength of the magnetic field can be varied.
When an ion hits the detector, a tiny electrical current is released. This is amplified and the signal is fed to a computer that analyses the signal. A graph of relative abundance against mass is produced, called a mass spectrum. The more of a particular ion that reaches the detector, the bigger the signal and the higher its bar in the mass spectrum.
When the mass spectrometer is working, its magnetic field strength is gradually increased until all the different ions produced by a sample have been detected. Different isotopes of an element have different masses, and so different masses in the mass spectrometer. If all the peaks in the mass spectrum come from ions with a single positive charge:
- the mass/charge (m/z) ratio is also the mass number of each isotope
- the height of a peak is the proportional to the relative abundance of the isotope.
This information can be used to work out the relative atomic mass, Ar, of the element in the sample.
Relative atomic masses
The relative atomic mass of an element is the weighted mean mass of all the isotopes of that element relative to the mass of 1/12 of an atom of carbon-12
Σ(mass of isotope x % abundance)
The mean atomic mass of an element = ___________________________________________
e.g. Chlorine has two isotopes.
75% of chlorine atoms are 35Cl and 25% of chlorine atoms are 37Cl.
(75 x 35) + (25 x 37)
The mean atomic mass of chlorine = ______________________________ = 35.45
e.g.2 magnesium has three isotopes. 78.6% are 24Mg; 10.11% are 25Mg and 11.29 are 26Mg. Calculate the relative atomic mass to 4 significant figures.
(78.6 x 24) + (10.11 x 25) + (11.29 x 26)
The mean atomic mass of Mg = ______________________________________________________
Definition of ionisation energy
Ionisation energy is the energy required to remove one mole of electrons from one mole of atoms in the gaseous state.
e.g the first ionisation energy of sodium is the energy to complete the following reaction for one mole of sodium atoms.
Na(g) → Na+(g) + e-
Subsequent ionisation energies describe the removal of one mole of electrons from positive ions, e.g:
second ionisation energy Na+(g) → Na2+(g) + e-
third ionisation energy Na2+(g) → Na3+(g) + e-
fourth ionisation energy Na3+(g) → Na4+(g) + e-
Factors influencing ionisation energy
a. Distance of electron from nucleus
The closer an electron is to the nucleus, the greater the force of attraction to the nucleus, and the more difficult it is to remove. Electrons in the first principal quantum shell are closer to the nucleus, so are harder to remove, and the energy needed to remove them is greatest. Ionisation energy decreases in each subsequent principal quantum shell because they are further from the nucleus.
Electrons are all negatively charged and repel each other. The more electrons there are, the greater the repulsion and the easier they are to remove. As electrons are successively removed, there is less shielding: the ionisation energy increases because the effective attraction to the positive nucleus is greater.
c. Size of positive charge
Removing an electron from a positive ion is more difficult than removing an electron from a neutral atom. The greater the positive charge on an ion, the harder it is to remove an electron, so the successive ionisation energies increase.
Prediction of electron structure from ionisation energies
The subsequent ionisation energies of an element can be used to confirm the electron configuration of the principal quantum shells.
e.g. for sodium
A graph of ionisation energy against log10ΔHi will show a jump when an electrons start to be removed from a lower principal quantum shell.
Ocean waves travel on the surface of the water. You can see them and you can feel them. As you swim through the water, you can even make your own waves. The wind creates waves in the flag. Both the waves in the flag and the ocean waves are waves that you can see.
- There are other kinds of waves. We cannot see these waves, but we experience them every day.
- These waves are called electromagnetic waves.
- Electromagnetic waves have a wide range of wavelengths ranging from low energy radio waves to high energy γ-radiation. Visible light occupies a very narrow part of the spectrum.
Radio waves, television waves, and microwaves are all types of electromagnetic waves. They differ from each other in wavelength. Wavelength is the distance between one wave crest to the next.
Electromagnetic waves can travel through space and, depending on the wavelength, also through matter.
The velocity of travel (c), is related to its wavelength (λ) and its frequency (f). Velocity is measured in m s-1, wavelength in m and frequency in s-1, so it is easy to remember the relationship between them:
c = λ x f
Electromagnetic radiation is a form of energy. The smaller the wavelength and thus the higher the frequency the more energy the wave possesses.
Visible light waves are the only electromagnetic waves we can see. We see these waves as the colours of the rainbow. Each colour has a different wavelength. Red has the longest wavelength and violet has the shortest wavelength. When all the waves are seen together, they make white light.
White light is made of all colours of the spectrum. When white light pass through a prism, the white light is broken apart into the colours of the visible light spectrum. All the colours can be obtained and that is continuous spectrum of colours. Water vapour in the atmosphere can also break apart wavelengths creating a rainbow.
The atomic emission spectrum of an is the set of of the waves emitted by of that element.
If an element is heated it emits electromagnetic radiation of certain wavelengths. This is because elctrons are excited into higher energy levels. When they return to a lower energy level, electromagnetic radiation is emitted. The emission spectrum produced is a series of discrete lines. When electrons return to level 1, ultraviolet radiation is emitted.
When they drop to level 2, visible radiation is emitted, and a series of lines can be seen, rersulting in a line spectrum.
Each atom's atomic emission spectrum is unique and can be used to determine if that element is part of an unknown compound.
Light consists of Electromagnetic radiation of different wavelengths. Therefore, when the elements or their compounds are heated either on a flame they emit energy in form of light. Analysis, of this light, with the help of spectroscope gives us a discontinuous spectrum.
A spectroscope is a instrument which is used for separating the components of light, which have different wavelengths. The spectrum appears in a series of line called line spectrum. This line spectrum is also called the Atomic Spectrum because it originates in the element.
Each element has a different atomic spectrum. The production of line spectra by the atoms of an element, indicates that an atom can radiate only certain amount of energy. This leads to the conclusion that electrons cannot have any amount of energy but only a certain amount of energy.
The visible emission spectrum of hydrogen
The visible emission spectrum of sodium
The frequencies of light that an atom can emit are dependent on states the electrons can be in. When energy is supplied to an atom electrons are excited.
When an electron is excited, electron moves to a higher energy level/orbital. When the electron falls back to its ground level they emit energy. This energy corresponds to a particular wavelength and shows up as a line in the spectrum.
When electrons return to the first level, the series of lines occurs in the ultraviolet region as this involves the largest energy change.
The visible region spectrum is formed by electrons dropping back to the second energy level and the first series in the infrared is due to electrons falling to the n=3 level.
There can be many types of transitions from higher to lower energy levels:
The highest energy levels converge at the ionisation energy of the hydrogen atom. An electron with this energy or more is free to leave the atom.
Energy, matter and motion
A molecule has energy associated with:
- translation – the whole molecule moving around
- rotation – the whole molecule rotating
- vibration – the bonds vibrating
- electrons – excitation, dissociation and ionisation
Microwave energy in the range 1 x 10–22 to 1 x 10–20 J causes changes in the level of rotational energy. For example, microwave ovens emit energy at a frequency of 2.45 x 109 Hz, which corresponds to the energy needed to rotate the water molecules in food.
Infrared energy in the range 1 x 10–20 to 1 x 10–19 J causes changes in the level of vibrational energy of bonds.
Visible and ultraviolet energy
Visible light and ultraviolet energy in the range 1 x 10–19 to 1 x 10–16 J causes changes in the level of electronic energy.
At low absorbed energies, electrons may be excited to a higher energy level, then fall back to a lower energy level.
At higher absorbed energies, bonding electrons separate. Covalent bonds break by homolytic fission and free radicals are formed.
For example: Cl2 → ∙Cl + ∙Cl
At very high absorbed energies, electrons can leave a molecule causing it to be ionised.
For example: Cl2 → Cl + e–
First Ionisation Energies of Successive Elements
The first ionisation energy generally increases across a period because the nuclear charge increases, so the electrons are held more tightly, and more energy is needed to remove them. Nuclear charge also increases across a period, which pulls the electrons closer, decreasing the atomic radius. This also causes first ionisation energy to increase, because the attraction between the positive nucleus and negative electrons is greater when they are closer.
There is a decrease between group 2 and group 3 elements. The electrons in the p sub-shell are at a higher energy level than those in the s sub-shell so are more easily removed.
There is a decrease in first ionisation energy between groups 5 and 6. In group 6, there are two electrons in the same orbital of a p sub-shell. It is easier to remove one of them because the other electron in the orbital repels it.
First ionisation energy decreases down a group, because there are successively more principal energy levels (shells) so electrons are further away from the nucleus, more shielded, and therefore less energy is needed to remove them