Atomic Structure and the Periodic Table. revision notes for o level chemistry

Atomic Structure and the Periodic Table.

Cathode rays; they are produced when a gas is enclosed in glass tube at very low pressure and high voltage is applied across the tube. They move from the cathode towards the anode.

NB: they are not necessarily attracted by the anode because when the anode is placed at a different position, these rays just pass the anode.

However when an electric field is applied across their path, the rays get deflected towards the positive plate which indicates that they are negatively charged.

Properties of cathode rays.

• When they strike the opposite end of the cathode ray tube, the tube fluoresces with green light.
• An object placed in the path of the rays gives a sharply defined shadow at the extreme end of the tube. This shows that the rays travel in a straight line.
• When blades of a paddle wheel mounted on an axle are placed in the path of the rays, the wheel rotates in a direction away from the cathode. This indicates the rays have mass and hence momentum.
• The rays are deflected by a magnetic field and electric field in the direction which shows that the rays are negatively charged.

Protons.

When different gases are used in the cathode ray tube and a perforated cathode was employed, positively charged particles are observed to move in a direction opposite to that of electrons.

When hydrogen gas was used, it resulted in the production of positively charged protons H⁺ which were found to be identical in mass and charge to the particles in the nucleus of an atom of an element. These positive ions are produced by collision of the electrons from the cathode with the gaseous atoms or molecules of a gas in the tube.

Neutrons.

These were later discovered around 1932 when beryllium was bombarded with alpha-particles.

Equations.

They were found to have no charge, neutral, but had mass of the magnitude of the proton. This mass is also found within the nucleus of an atom.

Natural Radioactivity.

Radioactivity is the spontaneous disintegration of certain atomic nuclei with the formation of new elements accompanied by emission of alpha-particles (helium nuclei), beta-particles are electrons or gamma rays (short-wave electromagnetic waves/radiations).

Natural radioactivity is a result of spontaneous disintegration of naturally occurring radial isotopes. Eg. Polonium, radon, actinium, uranium, thorium, protactinium.

Nature of Radioactive Emissions.

There are three types of radiations;

• Alpha-particles; they are helium nuclei. They are positively charged. They have mass number 4 and atomic number 2. They have very low penetrating power.
• Beta-particles; they are electrons so are denoted by (symbol).i.e. single negative charge and zero mas. They have a greater penetrative power than alpha-particles.
• Gamma rays; just electromagnetic waves like the x-rays having short wave length but very high frequency. They therefore have no mass and are not charged. They have the highest penetrating power. When passed through a magnetic or electric field.

The particles are deflected as follows;

Diagram.

The gamma rays have no charge hence not affected.

Beta particles are deflected more than alpha-particles.

Unstable isotopes.

• Alpha emitters

Are all atoms of atomic number greater than 83 and mass number greater than 209. They are all unstable. This is because they have too much mass that the nuclear forces cannot hold the nuclear forces together. Alpha-emission which involves loss of two protons and two neutrons from the atomic nucleus reduces the mass.eg.

Equation.

The exception is beryllium, it undergoes alpha-particle decay.

• Beta-decay.

For isotopes which have more neutrons than the stable isotopes of the same element. The heaviest isotopes of an element are therefore more likely to emit a beta-particle. Here mass number remains unchanged while atomic number increases by one. In the process of beta-decay, a neutron from the nucleus disintegrates to give an electron and a proton.

Equations

• Gamma radiation.

The product nucleus after alpha or beta decay often possesses too much energy which energy can be dissipated by emitting a gamma ray.

Equations.

Examples.

Stability of Radioactive Isotopes.

The stability of an atom appears to depend on the number of protons and neutrons in the nucleus of an atom.

Therefore stability of the nucleus is the ability of the nuclear forces to hold the nucleus together. Information concerning stability is illustrated by plotting the number of neutrons against the number of protons present in the nucleus of all known nucleoids. The rest is shown below.

When this is done, a continuous line can be drawn through those nucleoids that are stable. The position of the neutron-proton ratio for unstable nucleoids lies in the shaded area on either side of the line.

Graph.

• From the diagram, the following information can be deduced;
• The lighter stable nucleoids mainly have equal numbers of protons and neutrons.
• The heavier stable nucleoids have more neutrons than protons.
• Many stable nucleoids have even numbers of protons and neutrons.
• Unstable nucleoids disintegrate in such a way as to produce isotopes nearer to the stability line.eg a nucleoid with a neutron-proton ratio above the stability line decays so as to gain an increase in atomic number.ie. By beta-emission which changes a neutron to a proton thereby decreasing neutron proton ratio.
• A nucleoid with neutron-proton ratio below the stability line disintegrates in such a way as to decrease its atomic number. In heavy nucleoids, this can be achieved by alpha-emission.

Rate of Decay.

It’s expressed as disintegration per second.ie. Number of atoms that decay per second.

NB: not all unstable isotopes in sample decay at once. Decay is a matter of chance or probability. We can predict when an atom will decay, but we cannot predict with a high degree of certainty the time required for a fraction of a very large number of atoms to decay.

Decay is a first order process, expressed as (equations)

Half-life.

It’s the time taken for a radioactive sample to decay to half its initial value. It is constant for any particular nuclei.

Equations.

Examples.

Nuclear Fission and Fusion.

Nuclear Fission.

It’s the spontaneous splitting of a radioactive nucleus with or without excitement to give smaller nuclei and some other particles.

Energy is also released in most of the nuclear fission.

Equation.

The produced nuclei undergo fission by colliding with the produced neutrons. The produced neutrons may also collide with the unfusioned uranium 235. In most of these fissions, energy is produced hence the total mass of products will be slightly less than that of the starting material.

There’s overall loss in mass and it’s called mass defect. This mass is given by Einstein’s equation; ².

Nuclear Fusion.

It’s the combination of two light nuclei to produce a large one with the release of very high amount of energy. The reaction needs energy (got from nuclear fission) but energy produced is far much greater.

Equation.

A hydrogen bomb works on the basis of nuclear fusion. There is fissionable material to produce nuclear energy which brings about nuclear fusion.

NB: nuclear reactions are chemical changes involving the disintegration of both the nucleus and electrons of an atom. They occur when the nuclei of atoms are bombarded with various kinds of high speed particles.

Uses of radioactive isotopes.

• Eradication of cancer; cobalt 27 is used and phosphorus is used in treatment of leukaemia.
• Sterilization of medical equipment; x-rays may be used.
• Carbon dating; carbon 14 is an invaluable tool in the study of ages of old specimen/objects.

Uses of carbon 14.

• Study of many organic reactions and their mechanisms.
• Study of the path taken by carbon dioxide during photosynthesis.
• Radio carbon dating.
• Tracer technique.
• Following reaction mechanisms.

Hazards of radiations.

• All radiations produce biological damage due to their ionising properties. The extent of the damage depends on nature of radiation, intensity and duration.
• Some are more harmful than others due to their intensity.
• All people are exposed to cosmic and background radiations. The damages involved include; burns (destruction of body pigment and loss of body hair) and genetic damage.ie. Subsequent change of the chromosomes.
• Destruction of lymphoid tissues or glands and bone marrow.

Electronic Structure of Atoms (Emission Spectrum of Hydrogen).

When electricity was passed through a discharge tube containing hydrogen gas at low pressure, many of the molecules break up into atoms. These atoms emit visible and invisible radiations some of which are in infra-red and some are in the ultra violet part of the spectrum. If the radiations are analysed, a line spectrum is observed consisting of three prominent lines.

The significance of the line spectrum of hydrogen.

This line spectrum gives the fact that within the atoms of elements, there exists different energy levels in which electrons are distributed around the nucleus. These levels become closer and closer as the distance from the nucleus increases.

Explain why a spectrum of an element is compared to the fingerprint of a human being.

This is because each element has its own clear charge. As a result it emits different energy quanta. In different electron transition hence distinct spectro lines similar to the distinct finger prints of a human being.

Sub-energy Levels.

As already noted, examination of line spectra using spectrometer of high dissolving power many are found to consist of more than one line, very close to each other. This indicates that within a given energy level, there exist sub-energy levels. Here electrons differ slightly in energy. The number of electrons in each main level and how they are sub-divided at different sub-energy levels can be deduced from the measurement of ionisation energy.

Theoretically, for a given quantum number (n), there are n sub-energy levels. The various sub-energy levels are indicated by the letters s, p, d and f.

Ionisation Energy.

Ionisation energy is the minimum amount of energy required to remove one mole of electrons from outer most energy level of one mole of gaseous positively charged ions to form one mole of positively charged ions.

First ionisation energy is the minimum amount of energy required to remove completely a loosely bound electron from the outermost energy level of a gaseous atom to form a unipositively charged gaseous ion.

Equation.

The value of this energy depends on the energy level and sub energy level occupied by the electron in its sub-energy state. I.e. it depends on;

• Nuclear charge.
• Distance from the nucleus of the electron to be removed.
• Screening effect.

The outermost electrons are called the valence are called the valence shell electrons. They are responsible for the chemical nature of an element for an element to react, it must lose or gain electrons or share its valence electrons with other atoms in order obtain stable electronic configuration similar to that of noble gases. The elements require sublimation energy/atomisation energy and then ionisation energy before it takes part in chemical reaction.

Sublimation energy is the energy required to change one mole of atoms of an element into one mole of free gaseous (molecules) atoms of the same element. It’s the energy involved in the change; equations.

NB: when an electron is removed from a gaseous atom, the number of protons in the nucleus remains the same hence the nuclear charge remains the same, but the screening effect decreases and this results into effective nuclear charge. The remaining electrons are therefore strongly attracted towards the nucleus and to remove the next electron requires more energy than the previous one. Generally, the successive ionisation energies of an element increase. However there is always a sharp increase (a big jump) when an electron from a new energy level closer to the nucleus is being removed.

Example.

Uses of ionisation energy.

• Enables us to determine the group in the table to which the element belongs.
• It can be used to predict whether an element is a metal or non-metal.
• Enables us to determine the period to which the element belongs.

Periodic Table.

Elements in the modern form of the periodic table are arranged in order of their atomic numbers which for most parts run parallel to the RAM. This was due to the observation that the properties of elements depend on the number and arrangement of electrons.

The periodic law states that the properties of the elements are a periodic function of their atomic numbers.

• Atomic/covalent radius.

Since the electron cloud of an atom has no definite limit, the size of an atom cannot be defined simply and easily. However, the radius of an atom is often defined as the distance of closest approach to another identical atom in any bonding situation.

This means that the radius of an atom is determined by the effective radius of the outer electrons. For a solid element, the atomic radius is the mean distance between two nuclei of any identical atoms which are assumed to be spherically shaped and in contact at equilibrium.

Variation.

• Atomic radius decreases along the period from left to right.
• Within a group as we move from the top to the bottom, the atomic radius increases.

Explanation.

Atomic radius depends on mainly two factors;

• Nuclear charge (attraction of the positively charged nucleus for the electrons.
• Screening/shielding effect.ie. Repulsion between the electrons from the inner energy levels and the valence electron.

From left to right across the period, atomic number increases therefore nuclear charge increases and this causes great attraction for the valence electrons which are pulled closer hence the reduction in size. The increase in nuclear charge outweighs screening effect because the new electrons added go to the same energy level.

The increase in atomic radius from the top to the bottom of the group is due to the increase in the shielding/screening effect because of addition of a new full electron level, when moving from one element to another. The screening effect outweighs the increased nuclear charge.

NB: in period IV (the first transition series), the normal trend in which the atomic radius is decreasing from left to right across the period is interrupted by the transition elements. This is true for other periods also. The increased nuclear charge is almost balanced by the greater screening effect produced by adding an extra electron to the penultimate shell (energy level before valence shell).

Transition elements have almost the same atomic radius because the effect produced by nuclear charge is almost balanced by the greatest screening effect produced by an electron to the penultimate shell.

• Trend 2; Ionisation Energy.

This is the minimum amount of energy required to remove one mole of loosely bound electrons from one mole of gaseous atoms to form one mole of unipositively charged gaseous ions.

Ionisation energy depends on;

• Nuclear charge; the higher the nuclear charge, the higher the ionisation energy.
• Screening/shielding effect; the greater the screening effect, the lower the ionisation energy.ie. The electron that is effectively screened from the nuclear attraction requires less energy to be removed and vice versa.
• Distance from the nucleus of the electron to be removed; the nearer the electron to be removed from the nucleus, the greater it is attracted towards the nucleus and hence the greater the ionisation energy and vice versa.
• Variation of ionisation energy; across the period from left to right, there is a general increase in the first ionisation energy. This is due to the fact that nuclear charge is increasing across the period from one element to the next. This increase causes a decrease in the atomic radius and thus a decrease in the distance of the outer electron from the nucleus. The screening effect remains almost the same from one element to the next across a period since the electrons are added successively to the same energy level. Such electrons screen each other very little from the increasing nuclear charge.

NB: the increase in the ionisation energy across a period is not uniform.

Examples.

Consider a graph of first ionisation energy against atomic number for periods 2 & 3 below.

Graphs.

NB: the stability of orbitals/sub-energy levels containing electrons is such that a completely full sub-energy level is more stable than a half full sub-energy level. Also half full and full sub-energy levels are more stable than less than half full and more than half full but less than full sub-energy levels.

The reversal in period 2.

It occurs from Beryllium to Boron then from Nitrogen to Oxygen.

Electronic configurations.

In beryllium, the first electron is being removed from a full 2s sub-energy level which is stable and moreover it is nearer the nucleus thus experiencing a much greater pull. Removal of this electron requires more energy than for Boron where the first electron is being removed from a less half full 2p sub-energy level which is unstable. The 2p sub-energy level is also slightly further away from the nucleus thus it experiences a less nuclear pull compared to the 2s.

In nitrogen, the first electron is being removed from a half full 2p sub-energy level which is stable more energy is therefore required to remove the first electron in nitrogen than for oxygen where the first electron is being removed from more than half full but less than full 2p sub-energy level which is unstable.

Example.

The table below shows the first ionisation energy of the elements from sodium to argon.

Element	Na	Mg	Al	Si	p	S	Cl	Ar
Atomic radius	11	12	13	14	15	16	17	18
First ionisation energy (kj/mol	495	738	577	787	1060	1000	1255	1520

• What is meant by the first ionisation energy?
• Write an equation for the process involving first ionisation energy.
• Plot a graph of first ionisation energy against atomic number for the elements given in the table above.
• Explain the shape of the graph in terms of the electronic structure of these elements.

Answers.

• Refer to notes.
• Equation.
• Graph.
• The shape of the graph in terms of the electronic structure of the elements;

Electronic configurations.

There is a reversal in period 3 which occurs from magnesium to aluminium then from phosphorus to sulphur.

Electronic configuration.

In magnesium, the first electron is being removed from a full 3s sub-energy level which is stable and nearer the nucleus, thus experiencing a much greater pull. Removal of this electron requires more energy than for aluminium where the first electron is being removed from less than half full 3p sub-energy level which is unstable and slightly further away from the nucleus thus it experiences a less nuclear pull compared to the 2s.

In phosphorus, the first electron is being removed from a half full 3p sub-energy level which is stable and more energy is required to remove the first electron in phosphorus as it is nearer the nucleus, thus experiencing a much greater pull. In sulphur, the first electron is being removed from more than full 3p sub-energy level which is unstable and slightly further away from the nucleus thus experiences a less nuclear pull.

The graph below shows a plot of successive ionisation energies of sodium against successive electron being removed from the atom.

Graph.

What information can be obtained from the graph about the sodium atom?

• The atom belongs to group I of the periodic table.
• The atom belongs to period 3 of the periodic table.

Explain the trend in the ionisation energy of sodium.

The ionisation energies of sodium increase as the number of electrons reduce on the atom. When an electron is removed from a gaseous atom, the number of protons in the nucleus remains the same but the screening effect reduces therefore resulting into effective nuclear charge. The remaining electrons are therefore strongly attracted towards the nucleus and to remove the next electron, it requires more energy than the previous one. Generally, the ionisation energies of sodium increase, however there is a big jump in the ionisation energy when an electron from a new energy level closer to the nucleus is being removed.

Ionisation within a Group.

There is a decrease in ionisation energy in moving from top to the bottom of each group. This is because down a group, a full energy level of electrons is introduced in moving from element to the next. The screening effect therefore outweighs the increase in nuclear charge leading to increase in atomic radius and therefore increase in the distance of the electron to be removed from the nucleus.

NB: ionisation energy is important in determining properties of an element.eg. When the first ionisation energy is greater than 800kj/mol suggests that the elements are non-metals and first ionisation energy is less than 800kj/mol suggests that the elements are metals.

Metallic character increases down the group because the ionisation energy decreases down the group and decreases across a period because the ionisation energy increases.

The result of these two trends is that the non-metals numbering to about 2o are confined mostly at the top right hand corner of the periodic table while the metals are confined at the bottom left hand corner of the periodic table.

Electronegativity and Electropositivity.

It is the ability or tendency of an atom of an element when covalently bonded to another atom to attract the shared electrons towards itself.

In a covalent molecule, that contains two atoms with different electronegativities, the more electronegative atom acquires a partial negative charge while the other attains partial positive charge.

Variation.

In the periodic table, there is a general increase in electronegativity in moving from left to right across a period and a decrease in electronegativity (increase in electropositivity) down the group.

Electronegativity depends on;

• The nuclear charge.
• The atomic radius.

Increase in nuclear charge and the simultaneous decrease in the atomic radius result in the increase in electronegativity across a period.

Down the group, electronegativity decreases due to increase in atomic radius which is a result of increased screening effect which outweighs the nuclear charge. The outermost electrons are therefore weakly attracted.

NB; the higher the electronegativity of an element the more difficult it is to remove an electron from an atom and to remove an electron from its ions.

Effect of differences in electronegativity.

Electronegativity difference between any two elements A & B in a compound AB will affect the type of bond formed when A reacts with B.

• When the electronegativity of A is similar to that of B, the bond formed between the two is covalent and the compound is non-polar.eg. Cl₂, N₂, Br₂, O₂, F₂.
• When there’s a slight difference in electronegativity between A and B, the type of bonding is covalent but the molecule will be polar.
• When the electronegativity difference is very big, ionic bonding occurs.eg. Sodium and chlorine form an ionic bond.ie. Bonding is via electron transfer where the most electropositive element (metal) loses electrons and these are taken up by most electronegative element (non-metal).

Electron Affinity.

It is the energy change that occurs when one mole of gaseous atoms gains one mole of electrons to form one mole of negatively charged gaseous ions. It is the energy released for the energy attachment reaction below;

Equation.

It depends on the following factors;

• Nuclear charge.
• Atomic radius.
• Variation of trend.

It increases from left to right across the period, due to increased nuclear charge and decreased atomic radius. It decreases down the group because of increased atomic radius and as a result of screening effect, outweighs nuclear charge.

A value for electron affinity which is really the first molar electron affinity is negative.ie. There is release of energy.

Once the atom has gained one electron, the negative charge on the anion formed opposes the addition of a second electron so that the second electron affinity has a positive value.

NB: the more negative the electron affinity, the more stable is the anion formed, thus from the values of electron affinity shown below, chloride ion is more stable than the bromide ion with respect to their corresponding atoms and the bromide is more stable that the iodide.

Equations.

NB: in general, electron affinities become more exothermic as a period is crossed from left to right because the incoming electron is attracted more strongly by the progressively smaller atoms with unincreasing positive charge in their nucleus.