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Atomic Theory . Unit 7. The Atom. An atom is the smallest part of any element that can take part in a chemical reaction. An atom is the smallest particle of an element which still retains the properties of that element. Atoms are extremely small. Structure of the Atom. - PowerPoint PPT Presentation

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  • Atomic Theory Unit 7

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    The AtomAn atom is the smallest part of any element that can take part in a chemical reaction.An atom is the smallest particle of an element which still retains the properties of that element.Atoms are extremely small.

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    Structure of the Atom

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    Elements and CompoundsAn element is a substance that cannot be split into a simpler substance by chemical means.A compound is a substance which is made up of two or more different elements combined together chemically.Most elements from the periodic table are found in nature as compounds very few are found in elemental form.

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    MoleculesA molecule is a group of atoms joined together. It is the smallest part of an element or compound that can exist independently.There are two types of molecules:Molecules of elements:(All the same atoms in the molecule)e.g. Cl2, S8.Molecules of compounds:(Different atoms in the molecule)e.g. H2O, HCl, H2SO4.

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    Diatomic ElementsElements that end in ine and gen exist as diatomic molecules in their natural form.

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    Atoms, Molecules, Elements & CompoundsAtomsMoleculesCompoundsElementsofofof(Only 1)(2 or more)(all the same)(different)

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    Summary

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    The Periodic TableThe Periodic table is the chemists compass to a vast array of compounds.In the modern periodic table the elements are listed in order of increasing atomic number, and arranged in order of the numbers of electrons in their outer main energy level.These are referred to as the group number and the period number. The periodic table is divided into two main sections metals and non metals. A memory aid to where this division occurs is :Being Silly Assures Teachers Attention.

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    The Periodic TableP.T.EUnit 7, *

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    The Named GroupsSome of the Named GroupsP.T.EUnit 7, *

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    Group NumberThe group number of an element is equal to the numbers of electrons in the outer main energy level (shell) of an atom of that element.The named groups are:Group one Alkali metal.Group two Alkaline earth metals.Group seven Halogens. Group eight Noble gases.

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    Period The period number of an element is the number of main energy levels (shells) in an atom.

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    Solids, Liquids and GasesSolids, Liquids and GasesSolidsLiquidsGasesUnit 7, *

    HHeLiBeBCNOFNeNaMgAlSiPSClArKCaScTiVCrMnFeCoNiCuZnGaGeAsSeBrKrRbSrYZrNbMoTcRuRhPdAgCdInSnSbTeIXeCsBaLaHfTaWReOsIrPtAuHgTlPbBiPoAtRnFrRaAcRfDbSgBhHsMt

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    IonsIons are atoms or groups of atoms which have lost or gained electrons and hence have a charge.A positive ion (cation) is formed when an atom(s) loses electrons, they are denoted by the following notation Xn+.A negative ion (anion) is formed when an atom(s) gains electrons, they are denoted by the following notation Xn-.Simple ions (ions of 1 atom only) can be worked out from their valency (and from the periodic table).The valency of an element is equal to the number of electrons which an atom must lose or gain in order to attain noble gas configuration.

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    Simple Ions Simple IonsP.T.EUnit 7, *

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    Transition Metal Ions

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    Polyatomic Ions

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    Making Formulas for CompoundsGet the correct ions from:Periodic Table of Elements.Transition metals table.Polyatomic ions table.Put the positive ion first.If the charges are different cross over the magnitude (number part) of the charges.If you require more than one of a polyatomic ion make sure to put it in brackets before adding the subscript.There should be no charges on a compound.

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    P.T.EUnit 7, *

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    P.T.EUnit 7, *

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    P.T.EUnit 7, *

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    N.B. If you require more than one of a polyatomic ion you must put it in brackets.P.T.EUnit 7, *

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    N.B. When we require only one of an item there is no need to fill in a one, it is assumedP.T.EUnit 7, *

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    P.T.EUnit 7, *Try Some!

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    Naming Compounds Summary

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    Atomic NumberThe modern periodic table has the elements arranged in order of increasing atomic number. (Smaller of the two)The atomic number of an element in the Periodic Table tells us the number of protons in the nucleus. i.e. sodium is the eleventh element in the periodic table and therefore it has 11 protons in its nucleus.

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    Atomic Mass Number The atomic mass number of an element is the sum of the numbers of protons and neutrons in the nucleus of an atom.For example, sodium has 11 protons and 12 neutrons in the nucleus. Therefore its mass number is 23 or we say that the mass of the atom is 23 atomic mass units.A special unit called the atomic mass unit is used as the masses of the atoms are so small.

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    Nuclear Formula

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    Nuclear Formula It is called a nuclear formula as it only gives information about the nucleus of the atom or ion, you have to deduce the information for electrons. Remember:All atoms are neutral.Positive ions have lost electrons.Negative ions have gained electrons.

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    Unit 7, *

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    IsotopesIn 1919 an English chemist called Francis William Aston built an instrument called a mass spectrometer to measure the masses of atoms. He started work with a sample of neon gas and discovered something very unusual. He found that that neon gas consisted of two varieties of neon atom.One type of neon atom had a mass number of 20 and the other type had a mass number of 22. He concluded that neon gas consisted of atoms of neon that differed in the number of neutrons in the nucleus.

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    Isotopes of Chlorine & Carbon

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    Isotopes of Lithium

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    Isotopes of Hydrogen

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    % of Each Isotope PresentFor his work on the discovery of the isotopes Aston received the 1922 Nobel Prize in chemistry.Not only did Aston detect the presence of isotopes, but he also determined the percentages of each of the isotopes present. For example in his study of chlorine gas he found that there were approximately three times as many chlorine 35 atoms as there were chlorine 37 atoms. He was then able to work out the average mass of an atom of chlorine.The periodic table contains this figure for the average mass of a atom of the element concerned. This is why the values given in the periodic table are all decimals. His method of calculation is shown in the next example.

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    Unit 7, *

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    Relative Atomic MassIf you locate chlorine in the periodic table you will find that 35.45 is the number given under the symbol for chlorine.This average mass of an atom is measured relative to the mass of the carbon 12 isotope. For this reason it is called the relative atomic mass. The symbol for relative atomic mass is Ar.Since the relative atomic mass is the ratio of two masses, it has no units. Therefore, we say that the average mass of an atom of chlorine is 35.5 a.m.u but its relative atomic mass is 35.5.

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    Unit 7, *

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    Unit 7, *The response of the ion detector (intensity of lines on photographic plate) is converted to a scale of relative numbers of atoms. (i.e. % of each isotope present)

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    Mass spectrum for mercury.The percent natural abundances of the mercury isotopes are:

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    Average Mass of Mercury

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    Calculating Molecular Mass

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    Unit 7, *

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    The Arrangement of ElectronsUp to this we said that electrons orbit the nucleus in shells (Bohrs theory). Each shell contains electrons with a certain amount of energy and electrons normally occupy the lowest available energy level.The ground state of an atom refers to its state when all of its electrons are in their lowest available energy levels.

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    No. of Electrons in each ShellIn this situation the first shell n = 1 could hold 2n2 = 2 electrons.In this situation the second shell n = 2 could hold 2n2 = 8 electrons.In this situation the third shell n = 3 could hold 2n2 = 18 electrons.In this situation the fourth shell n = 4 could hold 2n2 = 32 electrons.This works well to explain some topics in chemistry, but a more detailed arrangement is sometimes necessary.

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    Bohr Model of the AtomIn a Bohr atom, the electron is a particle that travels in specific, fixed orbits, but never in the space between orbits. This arrangement expresses energy quantization, and accounts for the atomic emission spectra.

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    A Ladder for Distance, Bohr Model for EnergyBohr orbits are like steps in a ladder. It is possible to be on one step or another, but it is impossible to be between steps.Unlike the ladder Bohr orbitals do not have equal spacing between the orbits.

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    Energy Absorption and EmissionWhen a hydrogen atom absorbs energy, an electron is excited to a higher energy level.The electron is then in an unstable and temporary level.The electron falls back to the lower energy level and emits a photon of light (or some other form of radiation).

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    Evidence for the Existence of Energy LevelsWhen atoms are excited (i.e. given energy) by heating them or subjecting them to electrical discharge, they usually emit light or some other form of radiation.A spectrum of light from a bulb would produce a continuous spectrum from red through to the various colours to violet.The spectrum from excited atoms are not continuous but consist of a number of distinct lines each corresponding to a definite frequency of light. (Each colour of light has a different frequency)

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    Spectrum of Ordinary White Light

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    The Emission Spectrum of Helium

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    Spectrum of Hydrogen

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    Line Spectrum of Selected Elements

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    Visible Lines in the Hydrogen SpectrumEach Bohr orbit has a distinct, fixed energy. When an electron relaxes from a high energy orbit to a low energy orbit, light (or some other form of radiation) is released.The 486 nm line corresponds to an electron falling from the n = 4 to the n = 2 orbit.The 657 nm line corresponds to an electron relaxing from the n = 3 to n = 2 orbit.

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    Model of Hydrogen & Emission SpectrumA portion of the hydrogen atom model is shown with the nucleus at the centre of the atom and with the electron in one of a set of discrete orbits.When the atom is excited, the electron jumps to a higher orbit (black arrows).Transitions in which the electron falls to the second level are accompanied by the emission of light.

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    Orbital Energy Levels for Hydrogen

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    Energy Levels and Spectral lines for hydrogen

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    Evidence for the existence of energy levelsContinuous Spectrum (bulb)Emission Spectrum (Hydrogen)The fact that the spectrum of hydrogen consists of distinct lines indicates that only certain energy emissions are possible.

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    This was accounted for as follows by BohrIn an atom the electrons revolve around the nucleus only in certain allowed orbits or shells.While in a particular shell, an electron has a definite amount of energy.Electrons normally occupy the lowest available energy level or ground state.When energy is given to an atom one or more electrons are promoted to a higher energy level or excited state, such a state is unstable and temporary.

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    Absorption & EmissionWhen the electron falls back to the lower energy level the energy difference between the two levels is emitted as a unit of light (a photon) or some other form of radiation, such as infrared or ultra violet.

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    Equation for Energy DifferenceThe emitted radiation has a definite frequency corresponding to the energy difference between the two levels.E2 E1= hfE2 energy of the higher level.E1 energy of the lower level.h Planck's constant.f frequency of the emitted radiation.

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    Setup to Examine Hydrogen Spectrum

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    Apparatus for Observing SpectraSpectrometer

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    Energy Levels & SublevelsBohrs theory explains the various lines in the hydrogen spectrum very well, but hydrogen is relatively simple, with only one electron, for the other elements a lot of modification was required.A close study showed the main energy level (principle level) had to be split into one or more sublevels.The number of sublevels is equal to the number of the main energy level.

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    Energy Levels Energy levels are split into sublevels as follows:n = 1 has 1 sublevel, called 1s.n = 2 has 2 sublevels called 2s and 2p.n = 3 has 3 sublevels called 3s, 3p and 3d.n = 4 has 4 sublevels called 4s, 4p, 4d and 4f.A s sublevel holds 2 electrons.A p sublevel holds 6 electrons.A d sublevel holds 10 electrons.A f sublevel holds 14 electrons.An energy level is the fixed amount of energy an electron has due to its position in an atom.

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    Shells are Organised into Subshells

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    Arrangement of Energy Levels

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    Electronic ConfigurationsArrangement of electrons in an atom or ion.Memory Aid:

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    Examples:Li1s2, 2s1F1s2, 2s2, 2p5Cl1s2, 2s2, 2p6, 3s2, 3p5Ar1s2, 2s2, 2p6, 3s2, 3p6Two exceptions are:Cr1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 3d5 Instead of 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d4Cu1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 3d10 Instead of 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d9 (Full and half-full sublevels have extra stability)

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    Valence ElectronsNote: that the valence electrons are in the highest main energy level. The valence electrons are the first ones to be removed.

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    Electronic Configurations of ions:A positive ion has lost electronsA negative ion has gained electronsIons lose or gain electrons to become more stable, so when the electronic configuration is written out - the part in square brackets should be a noble gas.

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    ExamplesMg2+[1s2, 2s2, 2p6]2+

    1s2, 2s2, 2p6 = NeCl-[1s2, 2s2, 2p6, 3s2, 3p6]-

    K+[1s2, 2s2, 2p6, 3s2, 3p6]+

    1s2, 2s2, 2p6, 3s2, 3p6 = Ar

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    Filling of Outer Levels

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    Blocks of Sublevels

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    d block From Sc to Zn (This is just the first period of them)The d block are so called as the change/build up in electron structure takes place in the d orbitals.The d block contains the transition elements.

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    Transition MetalsForm coloured compoundsCr3+ Green; Cu2+ BlueHave variable valenciesFe2+, Iron (II) ion; Fe3+ Iron (III) ion.Can act as catalystsA catalyst speeds up or slows down a reaction and does not itself take part in the reaction.Exceptions to these are Sc and Zn as they do not form coloured compounds or exhibit variable valencies.Sometimes transition metals are said to go from Ti to Cu ( excluding Sc and Zn)

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    Atomic orbitalsRegions in space where a electron is most likely to be found.Each orbital can only hold a maximum of two electrons.This they can only do if they have opposite spin (Hunds rule).Electrons fill orbitals of equal energy singly before filling them in pairs (Alfbau principle).

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    s - orbitalss orbitals are spherical in shape.The 2s orbital is similar in shape to the 1s orbital, but larger in size.

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    p - orbitalsDumb-bell in shape3 p orbitals to make up a p sublevel.pxpypz

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    Electronic Configurations with Orbitals of Equal EnergyRemember Hunds Rule and Aufbau Principle.The following is the electron distribution of the first ten elements in the periodic tableA represents an orbital.The arrows represent the spin direction.

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    1s2s2px2py2pZUnit 7, *

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    Quantum Numbers:Set of numbers used to describe each electron uniquely in an atom.There are four in total:Principle Quantum Number.Secondary Quantum Number.Magnetic Quantum Number.Spin Quantum Number.

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    Principal Quantum Number:The number of the main energy level that an electron is in.Values are 1, 2, 3 and 4 corresponding to the main energy levels n =1, n = 2, n = 3 and n = 4 respectively.

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    Secondary Quantum NumberThe number of the energy sublevel (that an electron is in).Values 0, 1, 2 and 3 corresponding to the sublevels s, p, d and f, respectively.

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    Magnetic Quantum NumberThe number of the orbital (that an electron is in).An s sublevel, has one orbital, assigned a magnetic quantum number of 0.A p sublevel has 3 orbitals, assigned magnetic quantum numbers of either 1, 0, +1, corresponding to px, py or pz respectively.

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    Spin Quantum NumberGives the direction of the spin of an electron in an orbital.The first electron in the orbital is assigned a value of + and the second -.

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    Summary Illustration

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    Quantum Number Summary Table

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    Assigning Quantum Numbers

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    Assigning Quantum Numbers

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    Ionisation Energy The first ionisation energy of an element is the minimum energy required to remove the most loosely bound electron from an isolated atom of that element in its gaseous state.

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    Graph of the first Twenty Ionisation Energies

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    Explaining the GraphThe maximum values are for the noble gases.Reason: Their atoms are very stable because of their electronic configuration [full outer (sub) level], so it is difficult to remove an electron.The minimum values are for the group one metals (alkali metals).Reason: Their atoms have only one electron in their outer level, so it is easily removed (as when this is lost it will have noble gas configuration.) This is why group one are so reactive.In general, ionisation energies increase in moving across a period from the alkali metal to the next noble gas.Reason: 1. Increase in nuclear charge. (greater pull for electrons)2. Decrease in atomic radius.Ionisation energies gradually decrease in moving down a group.Reason: 1. Increase in atomic radius. 2. Screening effect.(This is where the inner shell or shells of electrons help to shield the outer electrons from the positive charge in the nucleus.

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    Exceptions to Rule to 3 Across a PeriodThere are two exceptions to this generalisation:(a)Group two elements (e.g. Be, Mg) have abnormally high values. This is because the most loosely bound electron comes from a full s orbital. (e.g. 1s2, 2s2, 2p6, 3s2 in Mg) which is a relatively stable state. When the next element in each case (B, Al) is be ionised, the electron being removed is the single electron in the p orbital (e.g. 1s2, 2s2, 2p6, 3s2, 3p1 in Al).(b)Group five elements also show abnormally high values (e.g. N and P). The reason here is that the electrons being removed are from exactly half filled p orbitals, (e.g. 1s2, 2s2, 2p6, 3s2, 3p3 in P) and the half filled orbitals are the next most stable state after that of completely filled orbitals.

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    Ionisation Energy Trends - SummaryIncrease going across a period.Increase in nuclear charge.Decrease in atomic radius.Decrease going down a group.Increase in atomic radius.Screening effect.Exceptions, Group 2.Full (outer) sublevel.Exceptions, Group 5.Half full (outer) sublevel.

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    Ionisation Energy TrendsP.T.EUnit 7, *Trends in Ionisation Energies

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    Atomic RadiiP.T.EUnit 7, *Atomic Radii

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    Atomic Radii Versus Ionic Radii

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    Unit 7, *

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    Unit 7, *

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    Higher Ionisation Energy Levels for the Third Period

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