453 lines
25 KiB
Plaintext
453 lines
25 KiB
Plaintext
Message-ID: <070313Z08071994@anon.penet.fi>
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Newsgroups: alt.drugs
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From: an58264@anon.penet.fi (Dalamar)
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Date: Fri, 8 Jul 1994 06:59:04 UTC
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Subject: CHEMISTRY: Atomic Structure
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The Structure of the Atom
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_________________________
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To obtain a model for the atom we must first examine the three basic types of
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'building-blocks' from which atoms are constructed. These 'building-blocks' are
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known as the proton, neutron and the electron. You will sometimes see these
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referred to as 'subatomic particles'. Each of these particles has different
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properties and plays a different role in an atom. Protons are positively
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charged, each carrying a charge of +1. Neutrons, as the name might suggest, are
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electrically neutral particles of about the same mass as a proton. Electrons are
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negatively charged, each carries a charge of -1, exactly opposite and equal to
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that on a proton. However, electrons are tiny when compared to the proton or
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neutron - electrons have around 1/1836 the mass of a proton. This information is
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presented in the table below. Mass is measured in atomic mass units, where 1
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amu is equivalent to the mass of a proton or neutron.
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Particle Charge Mass Symbol
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_________________________________________________
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Proton +1 1 amu p
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Neutron 0 1 amu n
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Electron -1 1/1836 amu e
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_________________________________________________
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It has been determined that in an atom the protons and neutrons bind together
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to form a nucleus around which the electrons orbit. It is easy to see why this
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model of the atom has been likened to a minature solar system. The nucleus of
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the atom is the 'sun' and the electrons are the small orbiting 'planets'.
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The number of protons in the nucleus of an atom is known as the _atomic number_.
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The atomic number of an atom tells us which element it is from. For example an
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atomic number of 3 tells us we are looking at a lithium atom and an atomic
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number of 9 tells us we are looking at a fluorine atom. Atoms, when taken as a
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whole, are electrically neutral. This means that the number of protons in the
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nucleus must be matched by an equal number of orbiting electrons. Any excess or
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deficiency in the number of electrons orbiting the nucleus, compared to the
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number of protons in the nucleus, gives an overall charge imbalance. This
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imbalance will be -1 extra for each surplus electron supplied above the number
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of protons. If we add two electrons to a neutral atom it will acquire a net
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charge of -2. If electrons are stripped away from a neutral atom we are left
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with an excess in the number of protons over the number of electrons. As each
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proton carries a +1 charge, each electron deficiency gives a +1 extra charge on
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the atom. If we take three electrons away from a neutral atom it
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acquires a net charge of +3. These charged atoms are known as _ions_.
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Positive ions are known as _cations_ and negative ions are known as _anions_.
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Before moving on a few examples will help to illustrate these ideas.
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If possible find a copy of the periodic table of the elements. The elements in
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the table are listed in order of increasing atomic number from left to right.
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The horizontal rows are also known as _periods_. Each element in a period has
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one more proton in its nucleus than the element to its immediate left. When
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the far right end of a period is reached the addition of the next proton moves
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us back to the left and one row down. The vertical columns of the table are
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known as _groups_. Elements which make up groups are found to have very
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similair properties to each other and this is not just mere coincidence, it
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has its reasons rooted in something we shall go on to consider - the way in
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which an atoms electrons are positioned around its nucleus.
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Find fluorine in the periodic table, symbol F. In the box which details this
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element will be its symbol, atomic number and _mass number_. The _mass number_
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is the total number of protons plus neutrons in the nucleus, or the total number
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of _nucleons_, a term which collectively refers to both protons and neutrons.
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Sometimes these two numbers appear as superscript and subscript to the left of
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the elements symbol. The superscript is the _mass number_, the total number
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of protons plus neutrons. The subscript is the _atomic number_, the total number
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of protons alone. For fluorine these values are 9 and 19. Now we have all the
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information we need to formulate a picture of a fluorine atom. In the nucleus,
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as indicated by the atomic number, are 9 protons. The mass number 19 tells us
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that the total number of nucleons is 19, so the number of neutrons must be
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(19 - 9) = 10 neutrons. Atoms are electrically neutral, therefore to balance
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the +9 charge which the 9 protons introduce, there must be 9 orbiting electrons
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giving a cancelling charge of -9. The electrons are held in orbit by the
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electrostatic attractive force they feel from the positively charged nucleus.
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Remember that charges of the opposite sign _attract_ one another, whilst
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charges of the _same_ sign repel. If we now add an electron to the fluorine
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atom the total number of electrons becomes 10, one more than the number of
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protons in the F nucleus. This extra electron brings with it a -1 charge which
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has no cancelling +1 proton in the nucleus. The fluorine 'atom' now carries a
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net negative charge of -1. We no longer have a fluorine 'atom', but a
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_fluoride ion_, in this case an _anion_ because it is negatively charged.
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A diagram will illustrate these points further.
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Mass number = 19 FFFFFFFF
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F
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FFFF
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F
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F
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Atomic number = 9 F
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x x x x In these diagrams the F
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x x x represents the nucleus
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x F x x x F x with its 9 protons and
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x x x 10 neutrons. Each x
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x x x x represents an orbiting
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electron.
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A fluorine atom A fluoride ion
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Electrically neutral Net charge of -1
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Isotopes
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________
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The protons and neutrons (nucleons) of an atom are held tightly bound together
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by a force known as the _strong nuclear force_. This force is extremely strong
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and is required to overcome the repulsive forces that the protons exert on one
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another due to their close proximity. Remember that the closer you try and bring
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charges of opposite sign together, the greater is the replusive force they exert
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on each other - much like trying to put the north pole ends of two magnets
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together. To alleviate some of this repulsion is the function of the neutrons.
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The neutrons act by 'diluting' the concentration of positive charge in the
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nucleus by forcing the protons to be on average further apart. As the atomic
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number rises, so do the repulsive forces present in the nucleus, with the result
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that more neutrons are needed to 'dilute' the charge concentration.
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Some elements display varying numbers of neutrons in the nuclei of their atoms.
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For example, an atom of hydrogen has one proton in its nucleus and no neutrons.
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But what if we introduce a neutron to the nucleus ? Remember, it is the number
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of protons which determines which element we have, not the number of neutrons.
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So what is this new atom we have created which has one proton, one neutron and
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one orbiting electron ? The new atom is known as an _isotope_ of hydrogen.
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Isotopes are elements with identical numbers of protons but differing numbers
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of neutrons in their nuclei. In the case of hydrogen the isotope with the
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1 neutron is known as _deuterium_. There also exists a hydrogen atom with
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1 proton and 2 neutrons, known as _tritium_. However, in the case of hydrogen,
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the fraction of deuterium atoms in any given sample is miniscule compared with
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the number of 'normal' hydrogen atoms. We say that the natural abundance of
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deuterium is small compared with the natural abundance of hydrogen.
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If you look at the mass numbers for the elements you will see that alot of them
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are _not_ whole number values. This is due to the presence of isotopes. The
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number indicated as the mass number is an average of the isotopic masses
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weighted for natural abundance. For example, chlorine exists as a mixture of
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Cl-35 and Cl-37. When these mass numbers are averaged, taking into account the
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percentage of each isotope present in a sample, the mass number comes out as
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35.45. Because they are the same element, isotopes are identical in terms of
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chemical reactivity, hence we never notice that chlorine is a mixture of 2
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isotopes.
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Electron Energy Levels
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______________________
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So far you have seen that the atom consists of the proton, the neutron and
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the electron. The protons and neutrons together form the nucleus of the atom,
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around which orbit the electrons. The number of electrons must exactly match
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the number of protons in order for overall electrical neutrality to be achieved.
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The function of the neutrons is to stabilise the nucleus by diluting the
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repulsive forces of the protons and that elements whose atoms can have differing
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numbers of neutrons are known as isotopes.
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When we come to examine the arrangement of the electrons around the nucleus a
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distinct pattern emerges. It is found that the electrons occupy 'shells' which
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are of well defined energy and distance from the nucleus. Electrons occupying
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different shells are of different energies and distances from the nucleus.
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The number of electrons a shell can hold is fixed and this number cannot be
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exceeded. The first shell filled is the K shell, which can hold a maximum of
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two electrons. The K shell is also the closest to the nucleus, which means
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that electrons in it will be the most tightly held. When the K shell has been
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filled by 2 electrons the next shell to fill is the L shell. The L shell is
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capable of holding _eight_ electrons before it becomes full. The electrons in
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the L shell are further away from the nucleus than those in the K shell, so are
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not held so tightly by the attractive force from the nucleus. To build up a
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picture of the occupancy of these shells in an atom whose atomic number we
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know we use the following rules.
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1. The shells are filled in order from lowest energy (closest to nucleus) to
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higher energy (further from nucleus).
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2. The current shell _must_ be completely filled before moving on to fill the
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next one of higher energy.
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When this is done the atom is said to be in its _ground state_, the atom is
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at a minimum of energy, all electrons occupy the lowest energy levels available.
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The number of electrons in each shell can be indicated by listing the shells in
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order of increasing energy, together with the number of electrons in that shell.
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Hydrogen has one proton in its nucleus, so it must also have only one electron.
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This single electron must occupy the shell of lowest energy - the one nearest
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the nucleus - and this is the K shell. This may be written as K1, indicating
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the lone occupancy of the K shell. The next element, helium, has an atomic
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number of 2 indicating 2 protons in its nucleus. This is matched by 2 orbiting
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electrons. Following our rules we must place _both_ of these electrons in the
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K shell, which is then full. The electronic configuration of helium is therefore
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K2. With the third element, lithium, we begin the filling of the L shell which
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is capable of holding 8 electrons. The start of the new shell can be noticed in
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the periodic table, where we jump from helium on the far right, to lithium on
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the far left. If you count all the elements in the Li row, including Li, you
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will see that there are 8, the same as the number of electrons the L shell may
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hold before becoming full. The electronic configuration of Li, atomic number
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three, is therefore K2 L1. The L shell will continue to fill as we traverse the
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row, until we reach the element with the configuration K2 L8 (neon). Neon, like
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helium, has a _full_ outer shell of electrons. It is the electrons in the
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outermost shell of an atom which is responsible for the elements chemical
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reactivity. The next shell to fill is the M shell which is capable of holding
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18 electrons before becoming full. The element after neon, sodium, with atomic
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number 11, therefore has the electronic configuration K2 L8 M1. Sodium, like
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lithium, has only one electron in its outermost shell. Also, both sodium and
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lithium are, like the rest of the group, soft metals with similar reactivity.
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If you were to sit down and work out the electronic configurations of all the
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group I metals (Li, Na, K etc) you would see that they all have one electron
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in the outermost shell of their neutral atoms. It is this similarity in
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electronic structure which causes the similarity in properties in the group I
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metals and for other groups in the periodic table as well.
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If you were to work out the electronic configurations for the atoms of the noble
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gases (He, Ne, Ar etc), you would see that they all have their outermost shells
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completely full. The noble gases are also extremely unreactive. This can be
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attributed to the full outer shell of electrons, which provides stability and
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unreactivity. This idea of a full outer shell of electrons providing stability
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can be used as a powerful rationalising tool when discussing bonding between
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atoms, where an atom will strive to acquire a full outer shell, either by the
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gaining of electrons, loss of electrons or the sharing of electrons. I shall
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cover bonding theory in another file, but first we need to look at a few more
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of the properties of atoms which will aid us in predicting reactivity.
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* * *
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* *
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* * C * * * Ne *
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* *
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* * *
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Carbon K2 L4 Neon K2 L8
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Ionisation Energy
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_________________
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The first ionisation energy of an atom is the amount of energy required to
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remove one electron from the outermost shell to an infinite distance.
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This may be represented by the equation :
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E ========> E(+) + e(-)
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Note that the total charge on either side of any equation is always equal, in
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this particular case both sides are neutral (the positive charge on the cation
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balances the negative charge of the electron).
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The second ionisation energy is the energy required to remove a second electron
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from the now unipositive ion. This process may be represented by the equation :
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E(+) ========> E(2+) + e(-)
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Again, the charges on each side of the equation balance, in this case there is
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a plus one charge on each side (the -1 charge on the single electron cancels one
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of the two positive charges on E(2+) leaving a net +1.
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Removing electrons from an atom requires us to do work, that is we must supply
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sufficient energy in order to overcome the attractive force between nucleus and
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electron. As we have already seen, the electrons occupy shells which are of
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varying distance from the nucleus. Consequently electrons in different shells
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experience different attractive forces from the nucleus and they will therefore
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differ in the amount of energy needed to remove them. Remember, the closer the
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electrons are to the nucleus, the harder it will be to remove them.
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If we examine the first ionisation energy as a function of atomic number a
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regular pattern emerges.
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1. Across a period there is a steady _increase_ in first ionisation energy,
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which peaks at each noble gas.
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2. Down a group the first ionisation energy markedly _decreases_ from element to
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element.
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The increase in I.E. across a period is due to the increasing nuclear charge
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exerting a greater force on the orbiting electrons. Across a period the
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electrons are being fed into the same shell, so they are all no further away
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from the nucleus. However, the nuclear charge is _increasing_ and this naturally
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has the effect of binding those electrons more tightly. This then leads to the
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increase in I.E. which is observed in crossing a period.
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Based on the argument of increasing nuclear charge you may have expected the
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I.E. to increase down a group too, as each group member has more protons in its
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nucleus than the one above it. This, you would reason, would cause an increase
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in the attractive forces those outer electrons are going to feel and hence a
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rise in I.E. However, we are forgetting that for each successive group member
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the outermost electrons are in shells which are progressively further from the
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nucleus. This increase in electron to nucleus distance produces a drop in the
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attractive force which outweighs the increase in atomic number. The result is
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a decrease in I.E. on descending any group.
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From the above discussion it should now be clear that the elements with the
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highest ionisation energies are those to the top and right of the periodic
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table (eg O, F, Ne, Cl). These elements have ionisation energies in excess of
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15 eV. The elements with the lowest I.E.s are those to the left and bottom of
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the periodic table (eg Cs, Fr,). These elements have I.E.s around or below
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below 5 eV. Knowing the exact figures isn't important as long as you have an
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idea of the trends. Knowledge of an elements I.E. can allow us to predict, for
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example, whether that element will be an oxidising or reducing agent. As an
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example of how I.E.s differ down a group here are the first and second I.E.s of
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the group I metals.
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Metal First I.E. Second I.E.
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Li 520 7296 The measurements here are in
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kilo-joules per mole. The mole
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Na 496 4563 is a unit of measurement of
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substance.
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K 419 3069
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Rb 403 2650
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Cs 375 2420
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For sodium the first I.E. is 496 kJ/mol, this represents the amount of energy
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required to remove the single M electron to leave the Na+ cation (K2 L8).
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The amount of energy required to remove the second electron is huge compared
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with the first - 4563 kJ/mol. There are two reasons for this.
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1. The second electron is being removed from a _full_ orbital shell which
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contains electrons closer to the nucleus than the original single M
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electron already removed ie the process is K2 L8 ====> K2 L7 in which
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we are breaking into a _full shell_ in which the electrons are closer
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to the nucleus.
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2. The second electron is being removed from an already positively charged
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cation, with the result that we need to do more work in order to overcome
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this extra attractive force.
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Na ========> Na(+) + e(-) requires _less_ energy than :
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Na(+) ========> Na(2+) + e(-)
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Size of Atoms and Ions
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______________________
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Across a period in the periodic table, electrons are being fed into the
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same shell, so you may have expected no change in atomic size as we cross
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the period. However, in traversing the period we introduce more and more
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positive nuclear charge, with the result that the electrons being fed into the
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current shell feel the pull of the nucleus more strongly, thus there is a
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contraction in atomic size. Down groups there is an _increase_ in atomic size
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as, going from one element to the next in the group, the outermost electrons
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are in shells progressively further from the nucleus.
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Anions (negative ions) are always larger than their parent atoms. The reason
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being that the addition of an electron to the atom will cause an increase in the
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replusive field that the orbiting electrons mutually feel. This increase causes
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the electrons to spread out more in space thus increasing the size of the ion in
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comparison to the size of the atom.
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Cations (positive ions) are always smaller than their parent atoms. A loss of
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one or more electrons causes a reduction in the repulsive forces between the
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electrons and thus an overall contraction in radius. Also, the electrons which
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are lost may totally empty the outer shell, which will naturally lead to a
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reduction in radius as the next inner shell is closer to the nucleus. A sodium
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atom, for example, has the electronic configuration K2 L8 M1. Loss of a single
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electron gives a sodium ion, Na+, which has the stable noble gas electronic
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configuration of neon, K2 L8. The loss of the single electron from the M shell
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gives a natural reduction to the radius of the cation vs atom, as the outermost
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electrons are now in the L shell and not the M shell. In addition, the ratio of
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positive charges on the nucleus to the number of orbital electrons is increased.
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Thus the effective nuclear charge is increased and the electrons are pulled in.
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The greater the charge on the cation, the smaller it becomes.
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For Sodium:
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Atomic Radius Na (K2 L8 M1) = 1.57 Angstroms 1 angstrom =
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Ionic Radius Na+ (K2 L8) = 0.98 Angstroms 0.0000000001 metres
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Electronegativity
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_________________
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The electronegativity of an atom is a measure of its ability to attract
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electrons to itself when the atom is bonded to others as part of a compound.
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An atoms ability to attract electrons to itself depends greatly upon its size.
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Generally, the smaller the atom, the greater is its electronegativity ie the
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better is its ability to attract electrons. We have already seen that across
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a period there is a decrease in atomic size which corresponds to increasing
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nuclear charge. Down groups in the table there is a marked increase in size
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as the outermost electrons are in orbital shells progressively further from the
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nucleus. These trends indicate that across periods there is an _increase_ in
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electronegativity and down groups there is a _decrease_ in electronegativity.
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Therefore the most electronegative elements are to be found at the top right
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of the periodic table (N, O, F, Cl) and the least electronegative are at the
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bottom left (Rb, Cs).
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The electronegativities of the elements can be placed on a scale of 0-4, with
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fluorine, the most electronegative element, assigned the value of 4. The
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following partial periodic table lists some electronegativity values.
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__________________________________________________________________________
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H (2.1)
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__________________________________________________________________________
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Li (1.0) Be (1.5) B (2.0) C (2.5) N (3.0) O (3.5) F (4.0)
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__________________________________________________________________________
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Na (0.9) Cl (3.0)
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__________________________________________________________________________
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K (0.8) Br (2.8)
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__________________________________________________________________________
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Rb (0.8) I (2.5)
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__________________________________________________________________________
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Cs (0.7)
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__________________________________________________________________________
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This particular scale is known as the Pauling scale after its inventor.
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Atoms whose electronegativity falls below the 2.1 mark compete poorly for
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electrons, in fact these elements are sometimes referred to as electropositive
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because they have very little pulling power. They also happen to be the elements
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with low ionisation energy. The lower the value of EN below 2.1 the more
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electropositive the element will be, so that Cs, with an EN value of around
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0.7 is very electropositive indeed (has very low ionisation energy and competes
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poorly for electrons when it is part of a compound).
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Electronegativity is a useful concept for chemists. For example, the difference
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in electronegativity between two bonding atoms can be used to predict whether
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that bond will be predominantly _ionic_ or _covalent_. If you do not understand
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what is meant by these two terms then don't worry - I shall cover them in the
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next file : Bonding and Structure.
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Dalamar.
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