Periodicity

PeriodicityContentsThe Periodic TablePhysical PropertiesChemical PropertiesThe Periodic TableObjectives:Describe the arrangement of elements in the periodic table in order of increasing atomic numberDistinguish between the electron arrangement of elements and their position in the periodic table up to Z = 20Apply the relationship between the number of electrons in the highest occupied energy level for an element and its position in the periodic table.ActivityWatch video from Science Bank 3 DVD Disc 2 on;The Periodic TableElectronic StructurePhysical PropertiesObjectives:Describe basic trends in electron arrangement, structure and electrical conductivity in the periodic table.Describe and explain the trends in melting points for the alkali metals (Li Cs) and the halogens (F I)Describe and explain the trends in melting points for elements across period three.Periodicity and propertiesPeriodicity is explained by the electron structures of the elements:Group 1, 2 and 3 are metals, with up to three electrons in their outer shells. They have giant metallic structures. They give away their outer electrons (up to three) to form ionic compounds.Group 4, with four electrons in their outer shells, are semi-metals. They have giant structures. They only form covalent compounds.Group 5 to 7 are non-metals with five, six or seven electrons in their outer shells. They either accept electrons (up to three) to form ionic compounds, or share their outer electrons to form covalent compounds.Group 0 are the noble gases atoms that have full outer shells and are unreactive.Metallic StructureIn metallic bonding each metal atom loses its outer (valence) electrons. Forming cations surrounded by a sea of delocalised electrons. The greater the number of valence electrons in the metal, the greater the strength of the metallic bond.++++++++------------2+2+2+2+2+2+2+2+------------Giant Covalent StructureGroup 4 elements do not form metallic bonds, instead they tend to share electrons to obtain a full outer shell of electrons, this type of bonding is called Covalent bonding.In Giant Covalent structures the strong covalent bonds occur in all directions making this type of structure particularly strong.Diamond is a classic example of a giant covalent structure.

Simple covalent structures.Elements in Groups 5 to 7 form simple covalent molecules. This means that the atoms in each molecule are bonded covalently, these are strong bonds formed from the sharing of electrons.The molecules are held together by weak forces of attraction. These forces are called van der Waals forces and are produced by the electrons in the molecules.The bigger the atoms in the molecule or the more atoms in the molecule, the greater the number of electrons and thus the greater the strength of the van der Waals forces and the higher the melting and boiling point of the substance.Van der Waals Van der Waals forces are caused by the electrons in a molecule.As molecules move the electrons get dragged behind them, as they are much lighter than the heavy nucleus. This means that a temporary imbalance of electrons exists forming a temporary polarity (this means that for a brief moment in time the molecule will have a slightly positive end and a slightly negative end). This temporary polarity (dipole) causes a brief electrostatic attraction between one molecule and another. This is a Van der Waals force.Activity;Analyse the graph from the worksheet Trends in melting points and boiling points of Period 3 Elements then describe and explain this trend using your knowledge of bonding.

Metallic bonding, cations in a seaof delocalised electrons. As the number of outer electrons (valence electrons) increase the strength of the metallic bonding increases.Giant covalent bonding, there are strong covalent bonds in each direction.Simple covalent molecules. There are weak intramolecular forces called ;van der Waals between the simple molecules. The strength of the van der waals depends on the size of the molecule. The bigger the molcule the greater the van der Waals. Atomic only there are only very weak forces of attraction between the particles

Notice the difference in the trends as we go down each group, can you explain it?Comparing Tm and Tb trends down groups 1 and 7Explaining Tm and Tb trends down groups 1 and 7In Group 1 the melting points __________ down the group. This is because, the atoms get ________ down the group and so the forces of attraction between them become ________.In Group 7 the melting points show the opposite trend, and _________ down the group.This is because the solid crystals of the halogens contain non-polar diatomic molecules, which are only attracted to each other by weak forces, the strength of these forces increase with the mass of the molecule (chapter 4 - bonding)decreasebiggerstrongerincreaseActivityReread through your notes including page 47- 48 of the Chemistry Course Companion and answer the question from the worksheet; Melting point trends in Groups 1 & 7 & Period 3Find the answers by clicking here.Trends in physical propertiesElectrical conductivityThe pattern is good conductivity on the ________ and poor conductivity on the _________.This trend is due to bonding. M_________ bonding on the ______ and c__________ bonding on the _______.Melting point and boiling pointThere is a clear break in the middle of the table between elements with high melting points (on the ______) and those with low melting points (on the ______).These trends are due to giant ionic lattice structures on the _____ and simple molecular structures on the _______

leftpooretallicleftovalentrightleftrightleftrightAtomic RadiusObjectives:Describe and explain the trends in atomic radii for the alkali metals and the halogensDescribe and explain the trends in atomic radii for elements across period three.Measuring atomic radiiThe atomic radius should be defined as the distance from the centre of the nucleus to the outermost electron. This is not possible, why?What can be measured, by a technique known as X-ray diffraction, is half the distance between two nuclei of bonded atoms. If they are bonded covalently it is known as the covalent radius of an atom, whereas if metallic bonding is involved it is the metallic radius.

P45 ib bookSTUDENT NOTESAtomic propertiesSome key properties of atoms, such as size and ionisation energy are periodic.Activity: (in pairs)Look up the atomic radii of the elements on page 7 of the book of data.Think of a suitable scale and draw scale drawings of the atoms down Group 1 (dont include Hydrogen or Francium) and along Period 3, using compasses.Label each atom with its name, symbol and actual radius. Use a key to show the scaling you have used.Describe the trend that you can see; down a group and along a period. Can you explain it?Trends in atomic radiiAtoms get larger as we go down any group.

And get smaller as you go along the period from left to right.

Trends in atomic radiiAtomic radius is a periodic property , it decreases across each period and where there is a jump we start the next period.

Down a groupAs we go down group 1 in the periodic table, each element has one extra complete shell of electrons than the one before.So, for example, the outer electron in potassium is in shell 4, while that in sodium is in shell 3.Both outer electrons experience the same shielded nuclear charge of 1+.Shielded nuclear charge refers to the charge from the resultant nuclear charge on the electrons in the outer shell.Activity:NaMgAlSiPSClElectron arrangementAtomic radius/ nmNuclear chargeEffective nuclear chargeUsing your data booklet and your notes fill in the information in the table below. Can you explain the trend in atomic radius along period 3?Activity:NaMgAlSiPSClElectron arrangement2,8,12,8,22,8,32,8,42,8,52,8,62,8,7Atomic radius/ nmNuclear chargeEffective nuclear chargeUsing your data booklet and your notes fill in the information in the table below. Can you explain the trend in atomic radius along period 3?The Shielding effectThe inner shells of electrons shield the outer electrons from the pull of the nucleus.This doesnt only affect the atomic radii but also affects reactivity as you will see in the next topic.Across a periodAs we move from sodium to chlorine we are adding protons to the nucleus and electrons to the outer shell the third shell.The charge on the nucleus increases from 11+ to 17+ (or +1 to +7 after allowing for shielding of the inner shells, which remains the same).This increased charge pulls the electrons in closer to the nucleus. So the size of the atom decreases as we go across a period.Ionic PropertiesObjectives:Describe and explain the trends in ionic radii for the alkali metals and the halogensDescribe and explain the trends in ionic radii for elements across period three.Ionic propertiesActivity: Look up the ionic radii of the elements on page 7 of the book of data.Think of a suitable scale and draw scale drawings of the atoms down Group 1 (dont include Hydrogen or Francium) and along Period 3, using compasses.Label each atom with its name, symbol and actual radius. Use a key to show the scaling you have used.Describe the trend that you can see; down a group and along a period. Can you explain it?Compare these to the atomic radii. Can you explain the differences?Trends in ionic propertiesPositive ions (cations) vs atomsWhen an atom of a Group 1 element loses an electron, the ion that is formed has a much smaller radius almost half the value. There are two reasons for this.One fewer electrons than there are protons, so the nucleus attracts the remaining electrons much more strongly.One fewer energy level, because the outer shell has effectively been removed, and the remaining electrons have the noble gas electron arrangement of the preceding element

Trends in ionic radii down a groupThe size of the ions _______ down group 1 as the outer energy level becomes progressively _________ from the nucleus.

How to compare ionic radii along a periodNot all atoms form unipositive ions so we cannot make a direct comparison of ionic radii, but we can compare ions with the same number of electrons these are called iso-electronic ions.Sodium ions, Na+, magnesium ions, Mg2+, and aluminium ions, Al3+, all contain 10 electrons and have the electron configuration of neon (2,8), with the second shell completely full, these are iso-electronic ionsTrend in ionic radii along the metals in Period 3As we move along the metals in period 3 the iso-electronic ions become much smaller. Can you think why?Sodium has __ protons in the nucleus, magnesium has __ and aluminium __. The __ protons in the aluminium nucleus will attract the eight electrons in its outer shell ____ than the __ protons in the nucleus of the sodium ion.

Trends in ionic propertiesNegative ions (anions)When the atoms of elements in group 7 gain one electron to form a negative ion, there will be one more electron in the outer shell and hence more electron-electron repulsion.The number of protons in the nucleus is unchanged, each of the electrons will be attracted less strongly, and the radius of the ion increases to almost twice the radius of the atom.

Trends in ionic radii down a groupThe size of the ions _______ down group 7 as the outer energy level becomes progressively _________ from the nucleus.

Trend in ionic radii along the non-metals in Period 3As we move from metals to non-metals in Period 3 the iso-electronic ions become much _______, but from right to left they get __________. Can you think why? Compare the electron arrangements and protons of silicon and phosphorus.

Trend in ionic radii along the metals in Period 3Silicon has ___ protons and ____ electrons, phosphorus has ___protons and ___ electrons. Silicon has the electron arrangement _____ and phosphorus has the arrangement ______. There will be a greater pull on the outer electrons in __________ which makes the ion much _________.Phosphorus has __ protons in the nucleus, sulphur has __ and chlorine __. The __ protons in the chlorine nucleus will attract the eight electrons in its outer shell ____ than the __ protons in the nucleus of the phosphorus ion.

Ionisation EnergiesObjectives:Define the terms first ionisation energy.Describe and explain the trends in first ionisation energies for the alkali metals and the halogensDescribe and explain the trends in first ionisation energies for elements across period three.Ionisation EnergiesA graph of the first ionisation energies of the elements against atomic number illustrates periodicity very clearly.Can you state the definition for the first ionisation energy of an element, can you write an equation to show this?The first ionisation energy of an element is the energy required to remove one mole of electrons from one mole of atoms of the element in the gaseous state.M(g) M+(g) + e-It is measured in Kilojoules per mole kJ mol-1Trends in first ionisation energiesActivity:Look at your graph of first ionisation energies.Describe (state) the trend in first ionisation energies down groups 1 and 7 and across period 3.Using your knowledge of nuclear charge, shells, pull of nucleus, electron arrangement try and explain the trend in first ionisation energies down groups 1 and 7 and across period 3.

Explanation of trend down a groupThe elements in group 1 have the _______ values in each period and the noble gases have the ________values.As we go down group 1 the values _______, because the outer electron is ________ from the nucleus and is therefore already in a higher energy level, so ____ energy is required to remove it.This is the same for all groups.

highestlowestincreasedecreasefurthernearermorelesshardereasier

Explanation of trend across Period 3.Across a period, as each energy level is successively filled with electrons an equal number of protons are also being added to the nucleus.As each electron is added the level (shell) is attracted _____ to the nucleus and therefore it becomes successively _______to remove an electron, so that ionisation energies generally ________ across a period.highestlowestincreasedecreasefurthernearermorelesshardereasier

What does first ionisation energy imply?Apart from Hydrogen most elements have more than one electron, but will the energy required to remove successive electrons be the same?The second ionisation energy of an element is defined as the energy required to remove one mole of electrons from one mole of unipositive ions in the gaseous state.M+(g) M2+(g) + e-What would be the definition of the third ionisation energy of an element?

To remove the first electron is relatively easyTo remove the next 8 electrons gets progressively harderTo remove the last 2 electrons is incredibly hardElectronegativityObjectives:Define the terms electronegativityDescribe and explain the trends in electronegativity for the alkali metals and the halogensDescribe and explain the trends in electronegativity for elements across period three.Compare the relative electronegativity values of two or more elements based on their positions in the periodic table.What is electronegativity?The forces that hold atoms together are all about the attraction between positive and negative chargesIn ionic bonding we have complete transfer of electrons from one atom to another.Even in covalent bonds, the electrons shared by the atoms will not be evenly spread if one of the atoms is better at attracting electrons than the other.The ability of an atom to attract a bonding pair of electrons to itself is known as electronegativity.What is the trend?Electronegativity is a relative value not an absolute value, and so there are different scales of electronegativity in use. The values used by the IB are attributed to the North American chemist Linus Pauling (1901 1994).

The Pauling scale is used to measure electronegativity. It runs from 0 to 4. The greater the number, the more electronegative the atom.

What is the most electronegative element on the periodic table?What is the least electronegative element on the periodic table?Elements in groups 1 and 2 are often called electropositiveElements in groups 5, 6 and 7 are called electronegativeExplaining electronegativityTwo factors affect electronegativity:The size of the atom. The smaller the atom, the larger the electronegativity (remember your trend in atomic radius)!The shielded nuclear charge. The greater the shielded nuclear charge, the larger the electronegativity (remember how to calculate shielded nuclear charge, proton number minus number of inner electrons, e.g. shielded nuclear charge of sodium = 11 10 = +1)Chemical propertiesObjectives:Discuss the similarities and differences in the chemical properties of elements in the same group.Discuss the changes in nature, from ionic to covalent and from basic to acidic, of the oxides across period three.Activity Watch the video on Patterns of Reactivity of Group I from Science Bank Disc 1.

63STUDENT NOTESGroup 1The group 1 elements are all:ShinySilveryReact vigorously with waterStored under oil to stop them reacting with air.Soft (cut easily with a knife)Once cut they tarnish easilyThe reactivity increases down the group.Can you explain the reactivity of the alkali metals down the group?

Reaction with waterAll the alkali metals react ______________ with water to form hydrogen and ions of the metal hydroxide in water:2M(s) + 2H2O(l) 2M+(aq) + 2OH-(aq) + H2(g)Where M represents Li, Na, K, Rb or CsAlkali metals are considered good reducing agents as they readily donate an electron.vigorouslyLithium reacts the most ________, and retains its shape as it reacts.The heat produced when sodium reacts is enough to _____ it into a ball so that it darts about on the surface of the water. Occasionally the hydrogen released may ______ and produce a _______flame characteristic of sodium ions.Potassium is much ______ reactive, and the hydrogen evolved usually burns steadily with a _______ flame because of the presence of ___________ ions.Rubidium is even more reactive, and the reaction with caesium may be so violent that the glass vessel _______.

ignitepotassiumlesslilacyellowexplodesslowlymeltredmoreReactivity of Group IAll the Group I elements have just one electron in their outer shell, and when they react they lose this outer electron to form the unipositive ion.The reactivity increases down the group as the outer electron becomes successively easier to remove, as the attraction from the nucleus decreases. ActivityWatch the demonstration of sodium reacting with chlorine and then sodium reacting with bromine to form the sodium halides.Notice the appearance of the reactants and the products.Reaction with HalogensThe alkali metals also react readily with the halogens, with the reactivity of the metals increasing down the group. If a piece of heated sodium is lowered into a gas jar of chlorine it will burst into flames, and white fumes will be seen, caused by the formation of the ionic salt sodium chloride.Write the reaction between sodium and chlorine.Similarly, if bromine vapour or iodine vapour is passed over heated sodium, a vigorous reaction occurs and the two elements combine to produce sodium bromide and sodium iodide respectively.Write the equations for these reactions.The halogensThe four halogens (F2), chlorine (Cl2), bromine (Br2) and Iodine (I2) all exist as diatomic molecules. When they react, the single bond between the atoms in the halogen molecule is broken and the two atoms then each gain one electron to form a halide ion.The halogens are strong oxidising agents, they accept electrons readily, Fluorine is the most oxidizing and reactivity decreases down the group.The halogen gases are very poisonous and this was first utilized in the first world war.Oxidising effect of the Halogens.In aqueous solution; chlorine is a clear, colourless solution, bromine is pale yellow to orange, (depending on the concentration of the bromine), similarly, iodine can vary from pale yellow to brown. Consequently, it can be quite difficult to distinguish between dilute, aqueous solutions of bromine and iodine simply on their colour.As the halogens are non-polar they dissolve readily in hydrocarbon solvent. This forms the basis of a useful test to distinguish between the two halogens as the colours in hydrocarbon solvent are very different. Iodine forms a purple solution and bromine forms an orange solution.Water (control)Sodium chlorideSodium bromide

Sodium iodide

Chlorine water (Cl2)Bromine water (Br2)Iodine water (I2)Copy this table into your books, then fill in as you watch the video on the following slide.ActivityLook at the following video and try to decide whether a chemical reaction has taken place, write in your table which halogen is present in the hydrocarbon solvent.We use hydrocarbon solvent to distinguish between the halogensChlorine is clear, colourlessBromine is orange brownIodine is pink

In this experiment we see some examples of displacement reactions. The experiment instructions ask you to use water as a control. The purpose of the water is to show the effect of dilution. If the brown iodine solution is added to a colourless solution, with which it does not react, it will become paler in colour. This is because the same amount of coloured iodine is now spread out over a larger volume. Students often mistake this for a chemical reaction, saying that as the solution is paler there must be less iodine there. In this experiment we see some examples of displacement reactions. The experiment instructions ask you to use water as a control. The purpose of the water is to show the effect of dilution. If the brown iodine solution is added to a colourless solution, with which it does not react, it will become paler in colour. This is because the same amount of coloured iodine is now spread out over a larger volume. Students often mistake this for a chemical reaction, saying that as the solution is paler there must be less iodine there.

STUDENT NOTESWater (control)Sodium chlorideSodium bromide

Sodium iodide

Chlorine water (Cl2)Bromine water (Br2)Iodine water (I2)Cl2Cl2Br2I2Br2Br2Br2I2I2I2I2I2Write a word and balanced symbol equation for each of the reactions in the table.Chemical Properties of elements in the same periodThere is a very noticeable change in the properties of the elements across the period.Metals can be distinguished from non-metals by their chemical properties. The oxides of metals tend to consist of the metal ions and oxide ions, and so are said to be ionic. Write down the formulas of sodium oxide, magnesium oxide and aluminium oxide.When they are in the liquid state they will conduct electricity, and are decomposed to their elements in the process.Silicon dioxide, SiO2, has a giant covalent structure and so has a very high melting and boiling point, but it does not conduct electricity when molten.The oxides of the non-metals sulphur, phosphorus and chlorine are all simple covalent, and have relatively low melting and boiling points.Argon does not form an oxide.The oxides of the period three elements also show a pattern in their acidity/basicity copy out the table on the next slide and fill in whilst watching the video. ElementNaMgAlSiPSClArState at rtpNature of bondingFormula of oxidepHCopy out this table and fill in as you watch the video clip ILPAC VIDEO 4 and during discussion after.ActivityILPAC VIDEO 4 VC4 (properties of the chlorides and oxides of the period 3 elements).77STUDENT NOTESElementNaMgAlSiPSClArState at rtpsssssggNature of bondingFormula of oxidepHIonicGiant CovalentSimple covalentAtomicSiO2Al2O3MgOP4O10SO2SO3Cl2OCl2O7Na2O10.5?9.16.31.41.8?

What do you predict to be the pH of Na2O and Cl2O7?pH 12pH 2.5Reaction of metal oxides with waterMetals on the left of the periodic table typically have basic oxides. For example;Magnesium oxide reacts with water to give magnesium hydroxide, which is sparingly soluble in in water and produces a somewhat alkaline solution of pH about 10:MgO(s) + H2O(l) Mg(OH)2(s) Mg2+(aq) + OH-(aq)As we move to the right along the period we reach aluminiumAluminium oxide is virtually insoluble in water. But, aluminium oxide will react with both acids and alkalis, it is an amphoteric oxide.Non-metals on the right of the table typically have acidic oxides. For example;Sulphur dioxide reacts with water to give a strongly acidic solution of sulphuric(IV) oxide acid (sulphurous acid). This dissociates producing H+ ions, which cause acidity of the solution:SO2(g) + H2O(l) H2SO4(aq) H+(aq) + HSO3-(aq)Reaction of non-metal oxides with waterIn summaryThe overall pattern is that; Metal oxides, on the left of the periodic table, form alkaline solutions in water.Non-metal oxides, on the right of the periodic table, form acidic solutions in water Semi-metals, in the middle of the period show amphoteric behaviourTrends across period three HL83STUDENT NOTESObjectives:Explain the physical states (under standard conditions) and electrical conductivity ( in the molten state) of the oxides of the elements in period 3 in terms of their bonding and structure.Properties of Period 3 oxidesPropertyFormulaNa2O

MgOAl2O3SiO2P4O10 P4O6SO3 SO2Cl2O7 Cl2OTm oC12752852202716102417-92Tb oC-3600298022301754580State at r.t.pElectrical conductivity when moltengoodgoodgoodVery poornonenonenoneStructureionicGiant covalentSimple covalentNature of oxidebasicamphotericacidicPhysical propertiesMelting point;Sodium oxide, magnesium oxide and aluminium oxide are all ionic, this accounts for their high melting and boiling points.Silicon dioxide has a diamond like macromolecular structure with a very high melting point and boiling point.The oxides of sulphur, phosphorus and chlorine are covalent because the relatively small difference in electronegativity between the elements and oxygen.Physical propertiesElectrical conductivity when molten;Sodium oxide, magnesium oxide and aluminium oxide are good conductors of electricity when molten but do not conduct when solid. This is because they are ionic, so when molten, the ions are free to move.Silicon dioxide is a very poor conductor of elelctricity.The covalent nature of the non-metal oxides mean they do not conduct electricity when molten as there are no free ions or electrons to move.Chemical propertiesActivity:Read through page 53 IB course companion.Under suitable sub-headings write the chemical reactions for; the metal oxides with water. the amphoteric nature of aluminium oxide.the non-metal oxides with water.Objectives:Explain the physical states (under standard conditions) and electrical conductivity ( in the molten state) of the chlorides of the elements in period 3 in terms of their bonding and structure.Describe the reactions of chlorine and the chlorides referred to above with waterPhysical Properties of the metal chloridesThe physical properties of the chlorides are related to their structuresSodium chloride Na+ Cl- and magnesium chloride Mg2+ (Cl-)2, are ionic so both;Conduct electricity when molten Have high melting points and boiling pointsAluminium chloride has a more covalent character, because of this;It is a poor conductor of electricity when moltenIt has a relatively low melting point, in fact it sublimesSpecial characteristics of Aluminium ChlorideThere is evidence that in the solid state, it exists as aluminium and chloride ions (in an ionic lattice) but as it melts there is a dramatic change in the bonding, and the covalent dimer Al2Cl6 is formedMolten aluminium chloride is a poor conductor of electricity.In the gaseous state there is an equilibrium between aluminium chloride monomer and aluminium dimer2AlCl3(g) Al2Cl6(g)Insert diagram of dimer91STUDENT NOTESPhysical Properties of the other period 3 chloridesThe remaining chlorides of period 3 all have simple molecular structures.

These molecules are held together by weak forces of attraction called van der Waals (chapter 4), which result in low melting points and boiling points.

Insert pictures of simple molecular structures of non-metallic chlorides92STUDENT NOTESChemical properties of the Period 3 metal chloridesWhen sodium chloride dissolves into water it gives a pH of 7, it is the only chloride to be neutral all the others are acidic.Magnesium chloride gives a slight acidic solution, when it dissolves in water the small, more densely charged magnesium ion, attracts the water molecules and causes some of them to dissociate to form hydrogen ions, this makes the solution acidic.What do you expect to happen when aluminium is dissolved in water?Aluminium chlorideThe aluminium ion is even smaller and even more densely charged, so when the anhydrous aluminium chloride is added to water, a very exothermic reaction takes place, and hydrochloric acid is formed:AlCl3(s) + 3H2O(l) Al2O3(s) + 6HCl(aq)Strictly speaking the aluminium ion becomes hexahydrated to form; hexaaquaaluminium (III) ion, [Al(H2O)6]3+The water molecules are strongly attracted to the high charge density on the small aluminum ion nd three of them successively dissociate to give hydrated aluminium hydroxide and hydrochloric acid:[Al(H2O)6]3+Al(H2O)3(OH)3(s) + 3H+(aq)Chemical properties of non-metallic period 3 chloridesAll the other chlorides react vigorously with water to produce acidic solutions of hydrochloric acid, together with the fumes of hydrogen chloride.Write the equations for these reactions, find them on page 54 IB Course Companion.Chlorine itself reacts with water to some extent to form an acidic solution.Write this equation.

PropertyNaClMgCl2AlCl3 Al2Cl6SiCl4PCl3 PCl5S2Cl2Cl2Tm oC801714178 sublimes-70-112-80-101Tb oC14131412-5876136-35StatesolidliquidLiquid (solid)liquidgasElectrical conductivity when moltengoodpoornoneStructureionicSimple covalentNature of solutionneutralWeakly acidicacidicObjectives:List the characteristic properties of transition elements.Explain why Sc and Zn are not considered to be transition elements.Quick questionsHow many electrons can exist in an orbital?How many d orbitals does the 3d level have?What is the maximum number of electrons that can be held if all the d orbitals of the 3d level are filled?What is Hunds rule? Put these levels in order of distance from the nucleus starting with the level closest to the nucleus: 4s, 1s, 3d, 2p,Write the electronic configuration of manganese using the s,p,d,f notation.d-block elementsThe d-block elements are typical metals:They are good conductors of heat and electricity.They are hard, strong and shiny, and have high melting points and boiling points, (one notable exception is mercury which is a liquid at room temperature).These physical properties, together with low chemical reactivity, make transition metals extremely useful, e.g. iron for vehicle bodies and to reinforce concrete, copper for water pipes and titanium for engine parts that need to resist high temperatures.Definition of a transition elementTransition elements are defined as d-block elements with an incomplete d level of electrons in one or more of their oxidation states.Write the electron configuration for the elements scandium to zinc.Zn, which can form Zn2+ ions, is not a transition element, can you think why?Although the 4s level fills up before the 3d, it is also the first to lose electrons, so neither Zn or its ion Zn2+ have incomplete d orbitals.Write out the electronic configuration for the Zn 2+ ionIs Scandium a transition element?Scandium is also considered a non-transition element, its common oxidation state is +3, and the tripositive ion has the electronic configuration of argon.What is the electronic configuration of argon?The electron configuration of scandium metal is [Ar]4s23d1, so in a sense the metal is a transition element but its compounds are not.

Other electronic configurationsAs the atomic number increases, the 3d sub-level fills up regularly, except for chromium and copper.These have the configurations [Ar] 4s1 3d5 and [Ar] 4s1 3d10 respectively, because it is more energetically favourable to half-fill and completely fill the 3d sub-level rather than spin-pair two of the electrons in the 4s orbital.Characteristic properties of the transition elementsTransition elements show characteristic properties, although they are not unique to transition elements. These properties include:Variable oxidation statesColoured compoundsComplex ion formationGood catalystsObjectives:List the characteristic properties of transition elements.Know that the transition elements can exist in different oxidation states.Explain why some complexes of d-block elements are coloured.Variable oxidation statesGroup 1 metals lose their outer electron to form only 1+ ions and Group 2 lose their outer electrons to form 2+ ions in their compounds.A typical transition metal can use its 3d as well as its 4s electrons in bonding and this means that it can have a greater variety of oxidation states in different compounds. This is because the 4s and 3d levels are very close in energy.The 4s level electrons are removed before the 3d level.Oxidation numbers shown by the elements of the first d-series in their compoundsScTiVCrMnFeCoNiCuZn+I+I+I+I+I+I+I+I+II+II+II+II+II+II+II+II+II+III+III+III+III+III+III+III+III+III+IV+IV+IV+IV+IV+IV+IV+V+V+V+V+V+VI+VI+VI+VIIThe most common oxidation states are shown in red, though they are not all stable.Coloured compoundsMany transition elements have coloured compounds .This is a result of spaces in the d-orbitals.These orbital's are not of exactly the same energy except in isolated gaseous atoms, so electrons can move from one orbital to another of higher energy.To do this they must absorb electromagnetic energy of a frequency in the visible region, e.g. if a substance absorbs green light, it lets through red and blue and thus appears purple.The wavelength of light that is absorbed depends on several factors, all of them due to the amount of energy required to promote a d electron from the lower split level to the higher split level.E.g., a compound containing the ion [Ar]3d9 has 9 electronsWhen white light passes through, a wavelength in the visible region (e.g green) is absorbed and an electron is promoted to the higher level.The rest of the light is reflected and the complimentary colour seen (red and blue are reflected which makes purple).higher energylower energyThere are 5 d - orbitalsIn some compounds these are not of equal energyActivityFollow the instructions of the experiment to investigate the different oxidation states and coloured compounds of Vanadium compounds.vanadate(V), VO2+(aq), is yellow vanadate(IV), VO2+(aq), is bluevanadium(III), V3+(aq), is greenvanadium(II), V2+(aq), is mauveReducing Vanadium ionsCollect all the apparatus, including safety spectacles.Measure 5cm3 acidified ammonium vanadate (V) NH4VO3 (in H2SO4) and pour into a test tube.Add a small piece of zinc metal and warm gently. Observe the different colours.When the solution becomes green add a spatulaful of zinc powder to speed up the reaction, and observe the final mauve colour of V2+

Oxidising Vanadium ionsCarefully filter the remaining mauve solution into a clean test tube.Add acidified potassium manganate (VII) dropwise. Shake after each addition and note the colour changesZinc metal is used as the reducing agent. What does reduction mean?Zinc is powerful enough to reduce the vanadate(V) all the way to vanadium(II). The reaction is sufficiently slow that you can see the colours of each of the intermediate oxidation statesAcidified Potassium manganate (VII) is used as the oxidising agent.What does oxidation mean?Potassium manganate (VII) can oxidise Vandium (II) to vanadate (V). Draw arrows to show the direction of oxidation and reduction.oxidationreduction

Objectives:Define the term ligand.Describe and explain the formation of complexes of d-block elements.Define coordination number and be able to use the coordination number to predict the shape of the complexComplex ion formationBecause of their small size, transition metal ions attract species that are rich in electrons.Such species are called ligands.Ligands are neutral molecules or negative ions that contain a non-bonding pair of electrons.These electron pairs can form coordinate covalent bonds with the metal ion to form a complex.Transition metal ions are good Lewis acids because they can accept a pair of electrons.Ligands are good Lewis bases because they can donate a pair of electrons.LigandsWater is a common ligand, and most (but not all) transition metal ions exist as hexahydrated complex ions in aqueous solutions, for example, hexaaquairon (III) ion [Fe(H20)6]3+Other examples of ligands include:Ammonia (NH3)Chloride ion (Cl-)Cyanide ion (CN-)Ligands can replace each other depending on the concentration of the solution and how much energy is given to the solution.Coordination numberThe number of lone pairs bonded to the metal ion is known as the coordination number.Compounds have a different shape depending on the coordination number,Coordination number of 6 are octahedralCoordination number of 4 are tetrahedral or square planar Coordination number of 2 are usually linear

Objectives:Recall what a catalyst is.State the difference between heterogeneous and homogeneous catalystsState examples of the catalytic action of transition elements and their compoundsCatalytic PropertiesCatalysts are substances that increase the rate of a chemical reaction without themselves being chemically changed at the end of the reaction.Essentially they work by providing an alternative pathway for the reaction-one with a lower activation energy.They do this by helping the two reacting species to come into closer contact with each other.Catalysts can be; heterogeneous and homogeneous.Heterogeneous CatalystsHeterogeneous catalysts, is where the catalyst is in a different phase from the reactants and products, they may do this by adsorbing reacting molecules onto the surface of the metal.Hydrogenation using catalystsAn example of a heterogeneous catalyst is in the use of nickel or palladium in hydrogenation.Compounds containing carbon-carbon double bonds (C = C) are said to be unsaturated, when hydrogen is added across the double bond they become saturated.The catalyst works by adsorbing hydrogen and the unsaturated compound (e,g, ethene) onto its surface and aligning them so that they are in the correct orientation to react.

Transition metals are particularly good at adsorbing small molecules, so the metals themselves make good heterogeneous catalysts.Other examples include:Iron in the Haber process, where ammonia is manufactured from nitrogen and hydrogen, (write the balanced symbol equation for this reaction)

Rhodium, platinum and palladium in catalytic converters in cars. These catalysts convert carbon monoxide, oxides of nitrogen and unburnt hydrocarbons into less polluting gases carbon dioxide, nitrogen and water.

Homogeneous catalystsHomogeneous catalysts are in the same phase as the reactants and products.In homogeneous catalysis often the two reacting species bond chemically together, and then leave.During the process the oxidation state of the central element in the catalyst will increase and then decrease. One of the reasons why transition metal compounds are such good catalysts is that they have variable oxidation states: that is that they can be relatively easily oxidized and reduced.Examples of compounds of transition metals that are important catalysts are:The use of manganese (IV) oxide to decompose hydrogen peroxide:2H2O2(aq) 2H2O(l) + O2(g)Vanadium (V) oxide in the conversion of sulphur dioxide into sulphur trioxide during the manufacture of sulphuric acid in the contact process:2SO2 (g) + O2 (g) 2SO3 (g)Iron and cobalt in biological catalysts, e.g. cobalt in vitamin B12. Vitamin B12 is essential for the production of red blood cells and the for the correct functioning of the central nervous system.MnO4 (s)V2O5 (s)

The structure of vitamin B12 showing the central cobalt atomR represents different groups (e.g. a methyl group, -CH3) to which the cobalt atom can bond in different forms of the vitamin.Vitamin B12 is found in foods of animal origin (fish, meat, liver, eggs and milk)A lack of vitamin B12 causes pernicious anaemia.Chart1371115692213639362740168326283175533927181722398487

Tm/KTb/KAtomic numberMelting and boiling points/KTrends in the melting and boiling points of the period 3 elements.

Sheet1Ar1112131415161718Tm/K371922936168331739217284Tb/K115613632740262855371823987

Sheet1

Tm/KTb/KAtomic numberMelting and boiling points/KTrends in the melting and boiling points of the period 3 elements.

Sheet2

Sheet3

Chart1454160037111563371047312961302952300950

Tm / KTb / KAtomic numberTm and Tb / KTrends in Tm and Tb down Group 1

Sheet1Ar1112131415161718Tm/K371922936168331739217284Tb/K115613632740262855371823987Ar31119375587Tm / K454371337312302300Tb / K160011561047961952950

Sheet1

Tm/KTb/KAtomic numberMelting and boiling points/KTrends in the melting and boiling points of the period 3 elements.

Sheet2

Tm / KTb / KAtomic numberTm and Tb / KTrends in Tm and Tb down Group 1

Sheet3

Chart25485172239266332387458575610

Tm / KTb / KAtomic numberTm and Tb / KTrends in Tm and Tb for Group 7

Sheet1Ar1112131415161718Tm/K371922936168331739217284Tb/K115613632740262855371823987Ar31119375587Tm / K454371337312302300Tb / K160011561047961952950Ar917355385Tm / K54172266387575Tb / K85239332458610

Sheet10000000000000000

Tm/KTb/KAtomic numberMelting and boiling points/KTrends in the melting and boiling points of the period 3 elements.

Sheet2000000000000

Tm / KTb / KAtomic numberTm and Tb / KTrends in Tm and Tb down Group 1

Sheet30000000000

Tm / KTb / KAtomic numberTm and Tb / KTrends in Tm and Tb for Group 7

Chart11310237051990079910901400131016802080494736577786106010001260152041859063266164865371676275773674590857776296694111401350

First Ionisation Energy / kJ molAtomic NumberIonisation Energy kJ/molA graph to show the First Ionisation Energies of the first 36 elements.

Sheet1ArIE1131022370351949005799610907140081310916801020801149412736135771478615106016100017126018152019418205902163222661236482465325716267622775728736297453090831577327623396634941351140361350

Sheet1

First Ionisation Energy / kJ molAtomic NumberIonisation Energy kJ/molA graph to show the First Ionisation Energies of the first 36 elements.

Sheet2

Sheet3

Chart14964563691395441335216611201152549128934141367159079

IONISATION ENERGYNumber of ElectronsEm kJ/molA Graph to show the Successive Ionisation Energies of Sodium

Sheet1Table showing the Em1 for the 1st 20 elements.ATOMIC NUMBER1st IONISATION ENERGY113122237235204900580161086714028131491681102081114961273813578147891510121610001712511815211941920590Succesive Ionisation Energies for SodiumNO.OF ELECTRONSIONISATION ENERGY14962456336913495445133526166117201158254919289341014136711159079

Sheet100000000000

IONISATION ENERGYNumber of ElectronsEm kJ/molA Graph to show the Successive Ionisation Energies of Sodium

Sheet2

Sheet3

Chart41110.59.16.31.41.817

pHAtomic number of the elements across period 3pHA graph to show the pH of period 3 oxides.

Sheet1Chlorides across Period 3Ar11121314151617pH5.95.73.20.30.70.3Oxides across Period 3pH10.59.16.31.41.8

Sheet1

Atomic number of period 3 elementspHA graph to show the trend in pH for the period 3 chlorides

Sheet2

pHAtomic number of the elements across period 3pHA graph to show the pH of period 3 oxides.

Sheet3

Chart15.95.73.20.30.70.3

Atomic number of period 3 elementspHA graph to show the trend in pH for the period 3 chlorides

Sheet1Chlorides across Period 3Ar11121314151617pH5.95.73.20.30.70.3Oxides across Period 3pH10.59.16.31.41.8

Sheet1

Atomic number of period 3 elementspHA graph to show the trend in pH for the period 3 chlorides

Sheet2

Sheet3