Message-ID: <070313Z08071994@anon.penet.fi> Newsgroups: alt.drugs From: firstname.lastname@example.org (Dalamar) Date: Fri, 8 Jul 1994 06:59:04 UTC Subject: CHEMISTRY: Atomic Structure The Structure of the Atom _________________________ To obtain a model for the atom we must first examine the three basic types of 'building-blocks' from which atoms are constructed. These 'building-blocks' are known as the proton, neutron and the electron. You will sometimes see these referred to as 'subatomic particles'. Each of these particles has different properties and plays a different role in an atom. Protons are positively charged, each carrying a charge of +1. Neutrons, as the name might suggest, are electrically neutral particles of about the same mass as a proton. Electrons are negatively charged, each carries a charge of -1, exactly opposite and equal to that on a proton. However, electrons are tiny when compared to the proton or neutron - electrons have around 1/1836 the mass of a proton. This information is presented in the table below. Mass is measured in atomic mass units, where 1 amu is equivalent to the mass of a proton or neutron. Particle Charge Mass Symbol _________________________________________________ Proton +1 1 amu p Neutron 0 1 amu n Electron -1 1/1836 amu e _________________________________________________ It has been determined that in an atom the protons and neutrons bind together to form a nucleus around which the electrons orbit. It is easy to see why this model of the atom has been likened to a minature solar system. The nucleus of the atom is the 'sun' and the electrons are the small orbiting 'planets'. The number of protons in the nucleus of an atom is known as the _atomic number_. The atomic number of an atom tells us which element it is from. For example an atomic number of 3 tells us we are looking at a lithium atom and an atomic number of 9 tells us we are looking at a fluorine atom. Atoms, when taken as a whole, are electrically neutral. This means that the number of protons in the nucleus must be matched by an equal number of orbiting electrons. Any excess or deficiency in the number of electrons orbiting the nucleus, compared to the number of protons in the nucleus, gives an overall charge imbalance. This imbalance will be -1 extra for each surplus electron supplied above the number of protons. If we add two electrons to a neutral atom it will acquire a net charge of -2. If electrons are stripped away from a neutral atom we are left with an excess in the number of protons over the number of electrons. As each proton carries a +1 charge, each electron deficiency gives a +1 extra charge on the atom. If we take three electrons away from a neutral atom it acquires a net charge of +3. These charged atoms are known as _ions_. Positive ions are known as _cations_ and negative ions are known as _anions_. Before moving on a few examples will help to illustrate these ideas. If possible find a copy of the periodic table of the elements. The elements in the table are listed in order of increasing atomic number from left to right. The horizontal rows are also known as _periods_. Each element in a period has one more proton in its nucleus than the element to its immediate left. When the far right end of a period is reached the addition of the next proton moves us back to the left and one row down. The vertical columns of the table are known as _groups_. Elements which make up groups are found to have very similair properties to each other and this is not just mere coincidence, it has its reasons rooted in something we shall go on to consider - the way in which an atoms electrons are positioned around its nucleus. Find fluorine in the periodic table, symbol F. In the box which details this element will be its symbol, atomic number and _mass number_. The _mass number_ is the total number of protons plus neutrons in the nucleus, or the total number of _nucleons_, a term which collectively refers to both protons and neutrons. Sometimes these two numbers appear as superscript and subscript to the left of the elements symbol. The superscript is the _mass number_, the total number of protons plus neutrons. The subscript is the _atomic number_, the total number of protons alone. For fluorine these values are 9 and 19. Now we have all the information we need to formulate a picture of a fluorine atom. In the nucleus, as indicated by the atomic number, are 9 protons. The mass number 19 tells us that the total number of nucleons is 19, so the number of neutrons must be (19 - 9) = 10 neutrons. Atoms are electrically neutral, therefore to balance the +9 charge which the 9 protons introduce, there must be 9 orbiting electrons giving a cancelling charge of -9. The electrons are held in orbit by the electrostatic attractive force they feel from the positively charged nucleus. Remember that charges of the opposite sign _attract_ one another, whilst charges of the _same_ sign repel. If we now add an electron to the fluorine atom the total number of electrons becomes 10, one more than the number of protons in the F nucleus. This extra electron brings with it a -1 charge which has no cancelling +1 proton in the nucleus. The fluorine 'atom' now carries a net negative charge of -1. We no longer have a fluorine 'atom', but a _fluoride ion_, in this case an _anion_ because it is negatively charged. A diagram will illustrate these points further. Mass number = 19 FFFFFFFF F FFFF F F Atomic number = 9 F x x x x In these diagrams the F x x x represents the nucleus x F x x x F x with its 9 protons and x x x 10 neutrons. Each x x x x x represents an orbiting electron. A fluorine atom A fluoride ion Electrically neutral Net charge of -1 Isotopes ________ The protons and neutrons (nucleons) of an atom are held tightly bound together by a force known as the _strong nuclear force_. This force is extremely strong and is required to overcome the repulsive forces that the protons exert on one another due to their close proximity. Remember that the closer you try and bring charges of opposite sign together, the greater is the replusive force they exert on each other - much like trying to put the north pole ends of two magnets together. To alleviate some of this repulsion is the function of the neutrons. The neutrons act by 'diluting' the concentration of positive charge in the nucleus by forcing the protons to be on average further apart. As the atomic number rises, so do the repulsive forces present in the nucleus, with the result that more neutrons are needed to 'dilute' the charge concentration. Some elements display varying numbers of neutrons in the nuclei of their atoms. For example, an atom of hydrogen has one proton in its nucleus and no neutrons. But what if we introduce a neutron to the nucleus ? Remember, it is the number of protons which determines which element we have, not the number of neutrons. So what is this new atom we have created which has one proton, one neutron and one orbiting electron ? The new atom is known as an _isotope_ of hydrogen. Isotopes are elements with identical numbers of protons but differing numbers of neutrons in their nuclei. In the case of hydrogen the isotope with the 1 neutron is known as _deuterium_. There also exists a hydrogen atom with 1 proton and 2 neutrons, known as _tritium_. However, in the case of hydrogen, the fraction of deuterium atoms in any given sample is miniscule compared with the number of 'normal' hydrogen atoms. We say that the natural abundance of deuterium is small compared with the natural abundance of hydrogen. If you look at the mass numbers for the elements you will see that alot of them are _not_ whole number values. This is due to the presence of isotopes. The number indicated as the mass number is an average of the isotopic masses weighted for natural abundance. For example, chlorine exists as a mixture of Cl-35 and Cl-37. When these mass numbers are averaged, taking into account the percentage of each isotope present in a sample, the mass number comes out as 35.45. Because they are the same element, isotopes are identical in terms of chemical reactivity, hence we never notice that chlorine is a mixture of 2 isotopes. Electron Energy Levels ______________________ So far you have seen that the atom consists of the proton, the neutron and the electron. The protons and neutrons together form the nucleus of the atom, around which orbit the electrons. The number of electrons must exactly match the number of protons in order for overall electrical neutrality to be achieved. The function of the neutrons is to stabilise the nucleus by diluting the repulsive forces of the protons and that elements whose atoms can have differing numbers of neutrons are known as isotopes. When we come to examine the arrangement of the electrons around the nucleus a distinct pattern emerges. It is found that the electrons occupy 'shells' which are of well defined energy and distance from the nucleus. Electrons occupying different shells are of different energies and distances from the nucleus. The number of electrons a shell can hold is fixed and this number cannot be exceeded. The first shell filled is the K shell, which can hold a maximum of two electrons. The K shell is also the closest to the nucleus, which means that electrons in it will be the most tightly held. When the K shell has been filled by 2 electrons the next shell to fill is the L shell. The L shell is capable of holding _eight_ electrons before it becomes full. The electrons in the L shell are further away from the nucleus than those in the K shell, so are not held so tightly by the attractive force from the nucleus. To build up a picture of the occupancy of these shells in an atom whose atomic number we know we use the following rules. 1. The shells are filled in order from lowest energy (closest to nucleus) to higher energy (further from nucleus). 2. The current shell _must_ be completely filled before moving on to fill the next one of higher energy. When this is done the atom is said to be in its _ground state_, the atom is at a minimum of energy, all electrons occupy the lowest energy levels available. The number of electrons in each shell can be indicated by listing the shells in order of increasing energy, together with the number of electrons in that shell. Hydrogen has one proton in its nucleus, so it must also have only one electron. This single electron must occupy the shell of lowest energy - the one nearest the nucleus - and this is the K shell. This may be written as K1, indicating the lone occupancy of the K shell. The next element, helium, has an atomic number of 2 indicating 2 protons in its nucleus. This is matched by 2 orbiting electrons. Following our rules we must place _both_ of these electrons in the K shell, which is then full. The electronic configuration of helium is therefore K2. With the third element, lithium, we begin the filling of the L shell which is capable of holding 8 electrons. The start of the new shell can be noticed in the periodic table, where we jump from helium on the far right, to lithium on the far left. If you count all the elements in the Li row, including Li, you will see that there are 8, the same as the number of electrons the L shell may hold before becoming full. The electronic configuration of Li, atomic number three, is therefore K2 L1. The L shell will continue to fill as we traverse the row, until we reach the element with the configuration K2 L8 (neon). Neon, like helium, has a _full_ outer shell of electrons. It is the electrons in the outermost shell of an atom which is responsible for the elements chemical reactivity. The next shell to fill is the M shell which is capable of holding 18 electrons before becoming full. The element after neon, sodium, with atomic number 11, therefore has the electronic configuration K2 L8 M1. Sodium, like lithium, has only one electron in its outermost shell. Also, both sodium and lithium are, like the rest of the group, soft metals with similar reactivity. If you were to sit down and work out the electronic configurations of all the group I metals (Li, Na, K etc) you would see that they all have one electron in the outermost shell of their neutral atoms. It is this similarity in electronic structure which causes the similarity in properties in the group I metals and for other groups in the periodic table as well. If you were to work out the electronic configurations for the atoms of the noble gases (He, Ne, Ar etc), you would see that they all have their outermost shells completely full. The noble gases are also extremely unreactive. This can be attributed to the full outer shell of electrons, which provides stability and unreactivity. This idea of a full outer shell of electrons providing stability can be used as a powerful rationalising tool when discussing bonding between atoms, where an atom will strive to acquire a full outer shell, either by the gaining of electrons, loss of electrons or the sharing of electrons. I shall cover bonding theory in another file, but first we need to look at a few more of the properties of atoms which will aid us in predicting reactivity. * * * * * * * C * * * Ne * * * * * * Carbon K2 L4 Neon K2 L8 Ionisation Energy _________________ The first ionisation energy of an atom is the amount of energy required to remove one electron from the outermost shell to an infinite distance. This may be represented by the equation : E ========> E(+) + e(-) Note that the total charge on either side of any equation is always equal, in this particular case both sides are neutral (the positive charge on the cation balances the negative charge of the electron). The second ionisation energy is the energy required to remove a second electron from the now unipositive ion. This process may be represented by the equation : E(+) ========> E(2+) + e(-) Again, the charges on each side of the equation balance, in this case there is a plus one charge on each side (the -1 charge on the single electron cancels one of the two positive charges on E(2+) leaving a net +1. Removing electrons from an atom requires us to do work, that is we must supply sufficient energy in order to overcome the attractive force between nucleus and electron. As we have already seen, the electrons occupy shells which are of varying distance from the nucleus. Consequently electrons in different shells experience different attractive forces from the nucleus and they will therefore differ in the amount of energy needed to remove them. Remember, the closer the electrons are to the nucleus, the harder it will be to remove them. If we examine the first ionisation energy as a function of atomic number a regular pattern emerges. 1. Across a period there is a steady _increase_ in first ionisation energy, which peaks at each noble gas. 2. Down a group the first ionisation energy markedly _decreases_ from element to element. The increase in I.E. across a period is due to the increasing nuclear charge exerting a greater force on the orbiting electrons. Across a period the electrons are being fed into the same shell, so they are all no further away from the nucleus. However, the nuclear charge is _increasing_ and this naturally has the effect of binding those electrons more tightly. This then leads to the increase in I.E. which is observed in crossing a period. Based on the argument of increasing nuclear charge you may have expected the I.E. to increase down a group too, as each group member has more protons in its nucleus than the one above it. This, you would reason, would cause an increase in the attractive forces those outer electrons are going to feel and hence a rise in I.E. However, we are forgetting that for each successive group member the outermost electrons are in shells which are progressively further from the nucleus. This increase in electron to nucleus distance produces a drop in the attractive force which outweighs the increase in atomic number. The result is a decrease in I.E. on descending any group. From the above discussion it should now be clear that the elements with the highest ionisation energies are those to the top and right of the periodic table (eg O, F, Ne, Cl). These elements have ionisation energies in excess of 15 eV. The elements with the lowest I.E.s are those to the left and bottom of the periodic table (eg Cs, Fr,). These elements have I.E.s around or below below 5 eV. Knowing the exact figures isn't important as long as you have an idea of the trends. Knowledge of an elements I.E. can allow us to predict, for example, whether that element will be an oxidising or reducing agent. As an example of how I.E.s differ down a group here are the first and second I.E.s of the group I metals. Metal First I.E. Second I.E. Li 520 7296 The measurements here are in kilo-joules per mole. The mole Na 496 4563 is a unit of measurement of substance. K 419 3069 Rb 403 2650 Cs 375 2420 For sodium the first I.E. is 496 kJ/mol, this represents the amount of energy required to remove the single M electron to leave the Na+ cation (K2 L8). The amount of energy required to remove the second electron is huge compared with the first - 4563 kJ/mol. There are two reasons for this. 1. The second electron is being removed from a _full_ orbital shell which contains electrons closer to the nucleus than the original single M electron already removed ie the process is K2 L8 ====> K2 L7 in which we are breaking into a _full shell_ in which the electrons are closer to the nucleus. 2. The second electron is being removed from an already positively charged cation, with the result that we need to do more work in order to overcome this extra attractive force. Na ========> Na(+) + e(-) requires _less_ energy than : Na(+) ========> Na(2+) + e(-) Size of Atoms and Ions ______________________ Across a period in the periodic table, electrons are being fed into the same shell, so you may have expected no change in atomic size as we cross the period. However, in traversing the period we introduce more and more positive nuclear charge, with the result that the electrons being fed into the current shell feel the pull of the nucleus more strongly, thus there is a contraction in atomic size. Down groups there is an _increase_ in atomic size as, going from one element to the next in the group, the outermost electrons are in shells progressively further from the nucleus. Anions (negative ions) are always larger than their parent atoms. The reason being that the addition of an electron to the atom will cause an increase in the replusive field that the orbiting electrons mutually feel. This increase causes the electrons to spread out more in space thus increasing the size of the ion in comparison to the size of the atom. Cations (positive ions) are always smaller than their parent atoms. A loss of one or more electrons causes a reduction in the repulsive forces between the electrons and thus an overall contraction in radius. Also, the electrons which are lost may totally empty the outer shell, which will naturally lead to a reduction in radius as the next inner shell is closer to the nucleus. A sodium atom, for example, has the electronic configuration K2 L8 M1. Loss of a single electron gives a sodium ion, Na+, which has the stable noble gas electronic configuration of neon, K2 L8. The loss of the single electron from the M shell gives a natural reduction to the radius of the cation vs atom, as the outermost electrons are now in the L shell and not the M shell. In addition, the ratio of positive charges on the nucleus to the number of orbital electrons is increased. Thus the effective nuclear charge is increased and the electrons are pulled in. The greater the charge on the cation, the smaller it becomes. For Sodium: Atomic Radius Na (K2 L8 M1) = 1.57 Angstroms 1 angstrom = Ionic Radius Na+ (K2 L8) = 0.98 Angstroms 0.0000000001 metres Electronegativity _________________ The electronegativity of an atom is a measure of its ability to attract electrons to itself when the atom is bonded to others as part of a compound. An atoms ability to attract electrons to itself depends greatly upon its size. Generally, the smaller the atom, the greater is its electronegativity ie the better is its ability to attract electrons. We have already seen that across a period there is a decrease in atomic size which corresponds to increasing nuclear charge. Down groups in the table there is a marked increase in size as the outermost electrons are in orbital shells progressively further from the nucleus. These trends indicate that across periods there is an _increase_ in electronegativity and down groups there is a _decrease_ in electronegativity. Therefore the most electronegative elements are to be found at the top right of the periodic table (N, O, F, Cl) and the least electronegative are at the bottom left (Rb, Cs). The electronegativities of the elements can be placed on a scale of 0-4, with fluorine, the most electronegative element, assigned the value of 4. The following partial periodic table lists some electronegativity values. __________________________________________________________________________ H (2.1) __________________________________________________________________________ Li (1.0) Be (1.5) B (2.0) C (2.5) N (3.0) O (3.5) F (4.0) __________________________________________________________________________ Na (0.9) Cl (3.0) __________________________________________________________________________ K (0.8) Br (2.8) __________________________________________________________________________ Rb (0.8) I (2.5) __________________________________________________________________________ Cs (0.7) __________________________________________________________________________ This particular scale is known as the Pauling scale after its inventor. Atoms whose electronegativity falls below the 2.1 mark compete poorly for electrons, in fact these elements are sometimes referred to as electropositive because they have very little pulling power. They also happen to be the elements with low ionisation energy. The lower the value of EN below 2.1 the more electropositive the element will be, so that Cs, with an EN value of around 0.7 is very electropositive indeed (has very low ionisation energy and competes poorly for electrons when it is part of a compound). Electronegativity is a useful concept for chemists. For example, the difference in electronegativity between two bonding atoms can be used to predict whether that bond will be predominantly _ionic_ or _covalent_. If you do not understand what is meant by these two terms then don't worry - I shall cover them in the next file : Bonding and Structure. Dalamar. ------------------------------------------------------------------------- To find out more about the anon service, send mail to email@example.com. Due to the double-blind, any mail replies to this message will be anonymized, and an anonymous id will be allocated automatically. You have been warned. 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