Presentation on theme: "Section 2 Periodicity Bonding in the Elements 1-20 (a)"— Presentation transcript:
1Section 2PeriodicityBonding in the Elements 1-20 (a)
2L.I. To learn about Bonding in the Elements 1-20 S.C. By the end of this lesson you should be able todescribe the metallic bondexplain what is meant by the term monatomicexplain what London dispersion forces are and howthey ariseexplain what happens to the strength of LDF as theatom size increasesexplain the difference between covalent network andcovalent molecular in terms of bpt and mptgive examples of metallic, covalent molecular, covalentnetwork and monatomic elements
3Periodic Pattern Johan Wolfgang Dobereiner – triads the atomic mass of the central element was approximately the mean of the other two.
4What does the periodic table and the sound of music have in common? John Newlands – octaves based on atomic mass (musical notes). oEvery eighth element showed similarities
5fun Period I II III IV V VI VII 1 H 2 Li Be B C N O F 3 Na K Mg Ca Al The modern Periodic Table is based on the work ofDimtri Mendeleev in 1869He arranged the elements based on: atomic mass, similar propertiesHe left gaps and made predictions for missing elementsPeriodIIIIIIIVVVIVII1H2LiBeBCNOF3NaKMgCaAl*SiTiPSCrClMnFe Co Ni4CuRbZnSrYZrAsNbSeMoBrRu Rh Pd5AgCdInSnSbTefunactivity in workbook
6Bonding in metalsactivity in workbookMetallic bonding is the electrostatic attraction betweenthe positively charged ions and the delocalised electrons.
7Bonding in metalsThe outer electrons are delocalised and free to move throughout the lattice, making metals good conductors of electricity.The greater the number of electrons in the outer shell the stronger the metallic bond.So the melting point of Al>Mg>Na
8Bonding in Monatomic elements He++Noble gases have full outer electron shellsThey do not need to combine with other atoms.The noble gases occur as single atoms, they aresaid to be monatomic.Since they can be liquefied and solidified there must be some weak attraction between the atoms.
9Bonding in monatomic elements ++The electrons in an atom “wobble” and become unevenlydistributed causing one side of the atom becomes slightlynegative while the other side becomes slightly positive.These slight charges are given the symbol δ ‘delta’.δ-δ+A temporary dipole is therefore formed.A dipole can induce other atoms to form dipoles, resulting in weak attractions between particles.δ-δ+London dispersion forcesactivity in workbook
10Bonding in monatomic elements London dispersion forces are a type of van der Waal force. They are very weak attractive forces.
11Group 8 element Electron arrangement Boiling Point oC Helium Argon Comparing Boiling pointsThe melting and boiling point of a substance gives an indication of the strength of the forces of attraction holding atoms or molecules together.activity in workbookGroup 8 elementElectron arrangementBoiling Point oCHeliumArgonKryptonXenon
12Noble gases b.p.’s20406080100120140160180166121HeliumNeonb.p / K87ArgonKryptonXeon274activity in workbookB.p.’s increase as the size of the atom increasesThis happens because the London dispersion forces increases with increasing size of atoms.
13Covalent Molecular Elements Many non-metals exist as discrete covalent molecules held together by covalent bonds.Discrete molecules have a definite number of atoms bonded together.9+Fluorine atomFluorine molecule F2diatomicactivity in workbook
14Examples of discrete molecules: Clactivity in workbook - table
15weak London dispersion forces strong covalentbondsMelting point low – why?activity in workbook - paragraph
16Halogen Boiling point (oC) Comparing Boiling pointsactivity in workbookHalogenBoiling point (oC)
17Halogens b.p.’s50100150200250300350400450500457332Fluorineb.p./ K238ChlorineBromineIodine85activity in workbookAs the size of the halogen molecule increases the boilingpoint increases. The bigger the molecule the strongerthe London dispersion forces between the halogen molecules.
18Fullerenes, molecules of carbon Fullerenes exists as large covalent molecules with a definite formula.Fullerenes were discovered in 1985 by Buckminster Fuller. Fullerenes are spherical in shape and usually contain sixty or seventy carbons.C60 is known as Buckminster fullereneactivity in workbook
19Covalent Network Elements Carbon - diamondDiamond has acovalent networkstructureEach of the outer electrons in a carbon atom canform a covalent bond with another carbon atom.So every C bonds to 4 others.m.p.’s C > 3642oCIt is high because many covalent bonds have to be broken.
20Carbon - Graphite Carbon bonded to only 3 other Carbons The spare (4th) electron is delocalised and so free to move. Graphite is a conductor of electricity.Van der Waals forces between thelayers allows layers to slide overeach other.Graphite can be used as a lubricant
21Properties of graphite and diamond PropertyDiamondGraphiteAppearanceColourless transparent solidConductionNoFeelSmooth, not slipperyHardnessVery hard
22Other Network Structures In the first 20 elements, only Boron, Carbon and Silicon have covalent network structures.m.p.’s B 2300oC, C > 3642oC and Si 1410oCactivity in workbook
23BONDING IN ELEMENTS - A SUMMARY activity in workbook
24Bonding patterns of the 1st 20 elements CovalentMolecularSiCBMetalliclatticeLiLiBeBeBNNCOOFFNeMonatomicNaNaMgMgAlPPSiSSClClArCovalentNetworkKKCaCaC , in the form of fullerenes, is covalent molecular
25animations/bonding_structure.aspThis interactive animation provides a visual representation of the bonding and structure of the first twenty elements in the periodic table, taking into account both the intra- and inter-molecular forces involved.
26Questions on elements – bonding and structure Explain why the covalent network elements have high melting and boiling points.Explain why the discrete molecular and monatomic elements have low melting and boiling points.Does diamond conduct electricity? Explain.Does graphite conduct electricity? Explain.How does the hardness of diamond compare with graphite? Explain.Give a use for both diamond and graphite.Complete the following table:
27Questions on elements – bonding and structure 7. Complete the following table:Type of bonding and structurePropertiesMetallic solids……………. of electricityCovalent network solids……….. …. melting pointsexception ……………….Covalent molecular solids………….. melting points…………… of electricityCovalent molecular (diatomic) gasesand monatomic gases…………… boiling points
28Patterns in the Periodic Table (b) Section 2PeriodicityPatterns in the Periodic Table (b)
29L.I. To learn about covalent radius S.C. By the end of this lesson you should be able todescribe the term covalent radiusexplain the changes in covalent radius down a groupexplain the changes in covalent radius across a periodexplain why there is no stated covalent radius for thenoble gases
30Covalent RadiusThe size of an atom is indicated by its covalent radius. (Page 7 of data booklet).There is no definite edge to an atom.266pmHowever, bond lengths can be worked out.The covalent radius of an element is half the distance between the nuclei of 2 of its bondedatomsFrom above the covalentradius would be 133 pm.Covalent radius – picometres (pm) 1pm =1 X 10 – 12 m
31Trends in covalent radius - Across a period activity in workbookNa 154 pm, Mg 145 pm, Al 130 pm, Si 117 pm, P 110 pm, S 102 pmGoing across a period the covalent radius (atomic size)decreases.Why? Going across a period the nuclear charge increases. The attraction between the outer electrons and the positive nucleus increases. Thus the outer electrons are more strongly attracted and so the atom size is smaller.
32Trends in covalent radius – Down a group activity in workbookLi 134 pmNa 154 pmK pmRb 216 pmGoing down a group the covalent radius (atomic size) increases.On moving down a group from one element to the next the number of electron shells increases.So the outer electrons are further from the nucleus and the atom size increases.
33Why is there no covalent radius value for the noble gases?
34L.I. To learn about ionisation energies S.C. By the end of this lesson you should be able todescribe the term 1st ionisation energywrite equations for the 1st ionisation energyexplain the trend in 1st ionisation energy down a groupexplain the trend in 1st ionisation energy across a perioddescribe the term 2nd ionisation energycarry out calculations involving ionisation energy
35Ionisation energies Na(g) Na+ (g) + e The first ionisation energy of an element is the energy required to remove one mole of electrons from one mole of atoms in the gaseous state.Units are kJmol-1.This is anendothermicprocess.(Page 11 of thedata booklet.)Na(g) Na+ (g) e
37Trends in 1st ionisation energy – Across a period activity in workbookGoing across a period the ionisation energy increases.Going across a period the nuclear charge increases.The attraction between the negative electrons and the positive nucleus increases. Thus the electrons are more tightly held and so more energy is needed to remove the outer electrons.
38Trends in 1st ionisation energy – Down a group activity in workbookGoing down a group the ionisation energy decreases.The explanation for this is(i) on moving down a group from one element to the next the number of electron shells increases and so the outer electron is further from the nucleus and less tightly held.(ii) the inner shells provide a screening effect which also decreases the attractive forces between the outer electrons and nucleus.
39The 2nd Ionisation Energy The second ionisation energy of an element is the energy required to remove the second mole of electrons.First IonisationSecond IonisationMg(g) Mg+(g) + e-Mg+(g) Mg2+ + e-ΔH = +738 kJ mol-1ΔH = kJ mol-1Third IonisationMg2+(g) Mg3+ + e-ΔH = kJ mol-1activity in workbook
40L.I. To learn about electronegativity S.C. By the end of this lesson you should be able todescribe the term electronegativityexplain the trend in electronegativity down a groupexplain the trend in electronegativity across a periodexplain why there are no quote values of electronegativityfor the noble gases
41ElectronegativityThe electronegativity is a measure of the attraction an atom involved in a bond has for the shared pair of electrons.Electronegativity values are based on the Pauling Scale, devised by Linus Pauling an American Chemist.Values on the Pauling Scale range from 0 to 4. A list of these values can be found in the data booklet on page 11.The higher the number on the Pauling scale is, the greater the attraction an atom has for the bonding electrons.
42Electronegativity – Across a Period Electronegativity values can be useful in predicting which type of bonding is most likely between two elements. (More about this later)Electronegativity – Across a Periodactivity in workbookOn crossing a period, electronegativity values increase. This is caused by an increase in nuclear charge as you move across a period from left to right.Electronegativity – Down a Groupactivity in workbookAs you go down a group, electronegativity values decrease.This is caused by the addition of another energy level of electrons as you go down a group which shields the bonded electrons from the nucleus; therefore they are not attracted as strongly.
43Electronegativity - The Monatomic Gases Why no values for group 8 elements?
44Section 3Structure and BondingBonding in Compounds
45L.I. To learn about bonding in compounds (a) S.C. By the end of this lesson you should be able todescribe the bonding and structure in ionic compoundsexplain the melting point of ionic compoundsdescribe the bonding and structure in covalent networkcompoundsexplain the melting point of covalent network compoundsdescribe the bonding and structure in covalent molecularexplain the melting point of covalent molecular compounds
46Three different types of compound - ionic, covalent molecular or covalent network. Ionic BondingNa Cl Na Cl-2)8)1 2)8)7 2)8 2)8)8In ionic compounds atoms achieve a full outer shell by either losing or gaining electrons and so form charged particles called ions.
47Complete for sodium, chlorine, bromine, oxygen, aluminium and ElementAtom electron arrangementIon electron arrangementIon symbolMg2)8)22)8Mg2+Complete for sodium, chlorine, bromine, oxygen, aluminium andnitrogen.Metal atoms always lose electrons to form positive ionse.g Na+Non-metal atoms always gain electrons to form negativeions e.g F-Glow: ionic bondingionic compounds
49Sodium chlorideLithium fluorideMagnesium oxideAluminium nitrideCalcium chlorideNow write ionic formula for the above.On show me boards – work out how these elements forman ionic compound
50The electrostatic attraction between positive and Ionic BondingNaCl3D lattice – regularrepeating patternof ionsThe attraction between positive and negative ions holds the compound together.The electrostatic attraction between positive andnegative ions is an ionic bond.+-ionic bond
52Ionic CompoundsNaCl as with all ionic compounds have many strong ionic bonds which are broken on melting thus the melting points are high (801 0C)Complete workbook
53shared pair of electrons COVALENT BONDINGIn covalent bonding the atoms share electrons.Covalent bonding is the electrostatic attraction between the shared electrons and the positive nuclei.shared pair of electronspositive nuclei
56Silicon dioxide and silicon carbide exist as a covalent network. All network structures have very high melting and boiling points.It is the strong covalent bonds that are broken on melting.
57Molecular Compounds Write formula for the following compounds: carbon monoxide, sulphur trioxide, carbon tetrafluoride, dinitrogen tetraoxide, phosphorus trifluoride
58Draw electron dot cross diagrams for the following molecules and structural formula CH4SCl2CO2
59Weak intermolecular forces COVALENT MOLECULARWeak intermolecular forcesStrong covalent bondsCovalent molecules tend to have low melting and boiling points as it is the weak intermolecular forces that are broken on melting.Complete workbook
60Plot a graph of melting points of the carbon tetrahalides against the covalent radius of the halogen in each molecules (see data book)CF4 = -184oCCCl4 = -23oCCBr4 = 90oCCI4 = 171oC
61m.p.’s of the carbon halides COVALENT MOLECULARm.p.’s of the carbon halides-18390-23171Temp/ oCincreasing size of moleculesCF4CCl4CI4CBr4What happens to the melting point as the size of themolecule increase?Why?As the molecule size increases the m.pt.s increase. This isbecause the strength of the London dispersion forcesincrease, so more energy is needed to separate molecules.Complete workbook
62L.I. To learn about polar covalent bonds (b) S.C. By the end of this lesson you should be able touse electronegativities to explain the difference betweenpure covalent and polar covalent bondsexplain the term permanent dipoleuse the data book to assign δ+ and δ+ partial chargeson atoms
63POLAR COVALENT BONDSTwo types of covalent bond can be formed:Pure covalent (or non-polar covalent)Polar covalent.
64Covalent Bonding Picture a tug-of-war: If both teams pull with the same force the mid-point of the rope will not move.
65Pure Covalent BondHeHThis even sharing of the rope can be compared to a pure covalent bond, where the bonding pair of electrons are held at the mid-point between the nuclei of the bonding atoms.
66Covalent Bonding What if it was an uneven tug-of-war? The team on the right are far stronger, so will pull the rope harder and the mid-point of the rope will move to the right.
67Polar Covalent BondA polar covalent bond is a bond formed when the shared pair of electrons in a covalent bond are not shared equally.This is due to different elements having different electronegativities.
68I Polar Covalent Bond δ- δ+ H e.g. Hydrogen Iodide e If hydrogen iodide contained a pure covalent bond, the electrons would be shared equally as shown above.This makes iodine slightly negative and hydrogen slightly positive. This is known as a dipole.However, iodine has a higher electronegativity and pulls the bonding electrons towards itself (winning the tug-of-war)
69H - H PURE COVALENT (OR NON-POLAR COVALENT) A pure covalent bond is formed when the atoms involved in the bond have an equal share of the bonding electrons. They have the same electronegativity.2.22.2H - HComplete workbook
70POLAR COVALENTWhen atoms with different electronegativity values join together, a polar covalent bond is formed.The dipole produced is permanent.
71A polar covalent bond is a bond where the electrons are not shared equally, one atom in the bond has a greater attraction than the other for the bonded electrons.Complete workbook
72L.I. To learn about the bonding continuum (c) S.C. By the end of this lesson you should be able toexplain the relationship between differences inelectronegativities and type of bondinguse data from the properties of compounds to deducethe type of bonding and structure
73Electronegativity difference Actual difference in electronegativity BONDING CONTINUUMElectronegativity Difference and Bond Type:Electronegativity differenceBond typeExampleActual difference in electronegativitycovalent (non polar)H-H0.0covalent (polar)H-Cl0.9H2O0.7covalent (very polar)H-F220.127.116.11ionicNaCl2.1
74The greater the difference in electronegativity the greater the polarity between two bonding atoms and the more ionic in character.A bonding continuum can be used to help us understand the differences in bonding.Complete workbook
75Bonding Continuum “Covalent compounds are formed by non-metals only” IS NOT AN ABSOLUTE LAW!Some compounds break this rule….
76Tin(IV)iodide – covalent or ionic? Predict its melting point.Complete workbookMelting point of tin(IV)iodide is 143oC.Tin electronegativity of 1.8Iodine has electronegativity of 2.6Molecule contains polar covalent bonds, but the symmetry cancels out the dipoles, therfore only weak London’s forces so low melting an boiling point.
77L.I. To learn about intermolecular forces (d) S.C. By the end of this lesson you should be able toexplain the difference between intramolecular andintermolecular forcesname the three types of van der Waals forcesexplain how London dispersion forces arise
78INTERMOLECULAR FORCES Intramolecular bonds are bond between atoms within a molecular – covalent bond.Intermolecular bonds are bonds which occur between molecules.Intermolecular bonds are calledvan der Waals’ forces. They are namedafter the Dutch Chemist JohannesDiderik van der Waals.
79There are three types of van der Waals’ forces: London dispersion forcesDipole-dipole interactions (permanent dipoles)Hydrogen bonding
801. London Dispersion Forces Electrons ‘wobble’ and temporary dipole occur. These cause induced dipoles on other atoms.The attraction between atom resulting from the temporarydipoles are known as London dispersion forces.
81London Dispersion Forces London dispersion forces are very weak attractive forces.
82L.I. To learn about intermolecular forces S.C. By the end of this lesson you should be able toexplain how dipole-dipole interactions arisedescribe a test that can be used to determine if amolecule is polarexplain the connection between symmetry and polaritydescribe how most hydrocarbons are classified in termsof polarity
832. Dipole-Dipole Interactions Polar MoleculesAre all molecules with polar bonds polar?HO--+Water has a polar covalent bonding between O and H.Is water polar?
84see scholar animation on polarity test Complete activity – testing polarity, and complete the tablesee scholar animation on polarity test
85Asymmetrical molecules e.g. H2O are POLAR Symmetry and polarityAsymmetrical molecules e.g. H2O are POLARIn an asymmetrical molecule there is a permanent dipoleworkbook activity
86Symmetrical molecules e.g. CCl4 are NON-POLAR In a completely symmetrical molecule the polarities cancel each other out so there is no permanent dipole.workbook activity
88H – H H - H Dipole-Dipole Interactions Dipole-dipole interactions are intermolecular forces which occur between polar molecules.polar moleculeDipole – dipole interactionsH – H H - Hnon - polar moleculeLondon dispersion forces
89Dipole-dipole interactions are stronger than London dispersion forces. Polar molecules have higher melting and boiling points than non-polar molecules.b.p. 56 o Cb.p. -1 o Cpolar moleculenon - polar molecule
90L.I. To learn about intermolecular forces S.C. By the end of this lesson you should be able toexplain how H-bonds arisewhat is necessary in a molecule to allow H-bonds to arise
91Hydrogen BondingHydrogen bonding is a special type of dipole-dipole interaction involving; H-N, H-O or H-F bondsworkbook activityFor H-bonds to exist between moleculesthe molecules must have a strong polar covalent bond2. the polar covalent bond must be between a hydrogen atom and either nitrogen, fluorine and oxygen (NOF)
92Hydrogen bonds are stronger than normal dipole-dipole interactions and London dispersion forces. Molecules which contain hydrogen bonding have much higher melting and boiling points than those with dipole- dipole interactions or London dispersion forces.
93L.I. To learn about relating properties of compounds to intermolecular forces (e)S.C. By the end of this lesson you should be able toexplain the connection between size of molecule andstrength of London dispersion forcesdescribe the evidence that proves the existence ofpermanent dipole-dipole interactions.H-bondsexplain why ice is less dense that waterexplain how intermolecular forces affect bpts, mpts,viscosity and solubilityexplain the term “like dissolves like”
94RELATING PROPERTIES TO INTERMOLECULAR BONDING 1. Melting and boiling points give an indication of the amount of energy needed to overcome the van der Waal’s forces between molecules.London dispersion forces – between non polar moleculesworkbook activityAlkane nameAlkane formulaboiling point(oC)pentanehexaneheptaneoctane
95From the table we can see that as the molecular size increases (number of electron shells) the boiling point increase and so the strength of the London dispersion forces increase.
96Dipole-dipole interactions – between polar molecules b.p. 56 o Cb.p. -1 o Cpolar moleculenon - polar moleculeFormula mass 58Formula mass 58workbook activityFor polar molecules, the melting and boiling points are higher than those of non-polar molecules. More energy is needed to overcome the dipole-dipole interactions between polar molecules. Molecules with a similar mass are used to allow us to ignore the London Dispersion Forces.
97Hydrogen bondingFor polar molecules that contain H-N, H-O or H-F the melting points are related to the strong hydrogen bonds between molecules.workbook activityFor polar molecules which contain H-N, H-O or H-F bonds the melting point and boiling points are higher than those of other polar and non-polar molecules. More energy is required to overcome the strong hydrogen bond.
107The stronger the van der Waals’ forces between a liquid are, the more viscous it will be. Liquids containing hydrogen bonding will be more viscous than molecules containing dipole-dipole interactions or London dispersion forcesThe most viscous liquid is propane-1,2,3-triol.
108Why does the ice float in Mrs Brown’s gin and tonic? Why do icebergs float?Why do fish in ponds not die in winter when the water freezes?
109DENSITY OF WATER/ICElook at candle wax and iceThe density of ice is unusual. Normally solids sink in their liquids.On cooling, water contracts but at 4oC it expands. At freezing point an open structure exists as a result of H-bonds.So ice floats!!