Chapter 7

Chapter 8: Periodic PropertiesNerve cells must pump Na+ ions out of the cell and K+ ions into the cell to complete the transmission of a signal. Both are group 1A elements yet the cells can distinguish them based on their size!Size of atoms or ions is an example of a periodic property.

HistoryElements like gold (Au) can be found in their pure state in nature and were known to the ancient humans.Others are more rare and occur only in compounds, which until the 19th century were difficult to isolate.In 1800, only 31 elements had been identified. By 1865, the number was 63.www.ptable.comHistoryIn 1869, Dmitri Mendeleev proposed a classification scheme based on the fact that physical and chemical properties begin to repeat when the elements are arranged in increasing weight.As many elements were still not yet discovered, Mendeleev left blank spaces where he predicted that these elements would be eventually isolated.History

HistoryIn 1913, Henry Moseley found that when an element is bombarded with high energy electrons it will produce an X-ray with a unique frequency. He then arranged them in order of frequency and assigned them a whole number, which we now know as their atomic number.OrbitalsSubshells are described by the combination of the n and l quantum numbers.Energy increases for each subshell according to Aufbau filling order.1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, etc.

Electron ConfigurationsGround state = lowest possible energy state for all of the electrons.Each subshell in the Aufbau order fills completely before starting the next subshell. Remember, s = 1 orbital, p = 3 orbitals, d = 5 orbitals, and d = 7 orbitals and EACH orbital can hold 2 electrons per quantum number rules.Electron ConfigurationsUse a superscript after each subshell to indicate the number of electrons in that subshell. H: 1s1He: 1s2Li: 1s2 2s1Be 1s2 2s2B: 1s2 2s2 2p1Electron ConfigurationsImagine doing a configuration for Sr or W or Pb!!!Shorthand methodLEP #8.Exceptions.Subshells and the Periodic Table

Core and Valence ElectronsCore electrons are very close to the nucleus and are the completed shell(s) of electrons.Valence electrons are the outermost shell of electrons. These are loosely held and provide for all of the bonding interactions.For the main group elements, the number of valence electrons = the group number.Orbital DiagramsHunds Rule when electrons are placed into degenerate orbitals, they will fill them singularly with parallel spins.For outermost shells, use boxes, circles, or lines to represent each orbital.Use a for spin = +1/2 and use a for spin = -1/2.LEP #1Periodic PropertiesElements in the same group tend to behave similarly.Ex) Group 1A metals Li, Na, K, Rb, CsSome, though, have major differences.Ex) O and SElectronic properties of valence electrons are always the same.The chemical properties of the elements are largely determined by the number of valence electrons they contain.Effective Nuclear ChargeThe attraction of the electron to the nucleus depends on two factors.Factor one is the magnitude of the charge.Factor two is the distance between the charge.In a many electron atom, each electron is attracted to the nucleus and repelled by the other electrons.Effective Nuclear ChargeThe core electrons versus the valence electrons.Argon, 10 core and 8 valence electrons.

Effective Nuclear ChargeCore electrons shield the valence electrons from the nucleus.Zeff = Number of Protons Number of Core Electrons Zeff for Na = +1, for Cl = +7Trend is: ________________________Size of AtomsAtoms are sometimes treated as hard spheres in models.Defining atomic size is tricky, though, because atoms do not have defined boundary layers.Non-bonding radii.Bonding radii.

Size of AtomsWhat trends do you see?Any exceptions?H37Atomic Radiipm (picometers)He32Li152Be112B85C77N70O73F72Ne70Na186Mg160Al143Si118P110S103Cl99Ar98K227Ca197Ga135Ge122As120Se117Br114Kr112Rb248Sr215In166Sn140Sb141Te143I133Xe130Electron Configurations of IonsWhen a main group element gains or loses electrons, it does so with predictability.Na (and other group 1A metals) always lose one electron.Na: [Ne] 3s1 Na+: [Ne] + 1e-Cl (and other group 7A halogens) always gain one electron.Cl: [Ne] 3s2 3p5 + 1e- Cl-: [Ne] 3s2 3p6Electron Configurations of IonsFor transition metal ions, their configurations are NOT what is expected.When filling, the 4s fills BEFORE the 3d orbital.When removing electrons, though, they do not remove in reverse order!ALWAYS remove the s electrons first.Ex) Fe+2, Co+3Diamagnetic = no unpaired electrons, paramagnetic = unpaired electrons.Size of IonsWhat trends do you see?

Isoelectronic SeriesIons that contain the same number of electrons are said to be isoelectronic.Place the following in order of increasing size: Na+, F-, Mg+2, O-2, and Al+3.Ionization Energy (IE)IE is the minimum amount of energy required to remove an electron from a gaseous atom or ion.IE1 is the energy to remove the first electron from a neutral atom. Ex) Na(g) Na+(g) + 1e-IE2 is the energy to remove the next electron.Ionization EnergyWhat trends do you see?Any exceptions?H1312First Ionization Energies, kJ/moleHe2372Li520Be899B801C1086N1402O1314F1681Ne2081Na496Mg738Al578Si786P1012S1000Cl1251Ar1521K419Ca590Ga579Ge762As947Se941Br1140Kr1351Rb403Sr549In558Sn709Sb834Te869I1008Xe1170Cs376Ba503Tl589Pb716Bi703Po812Rn1037Successive IEsWhat do you see?Why is it easier to remove the second electron from Cl than the second electron from Na?elementIE1IE2IE3IE4IE5IE6IE7Na4964,562Mg7381,4517,733Al5781,8172,74511,577Si7861,5773,2324,35616,091P1,0121,9072,9144,9646,27421,267S1,0002,2523,3574,5567,0048,49627,107Cl1,2512,2983,8225,1596,5429,36211,018Electron Affinity (EA)EA is the amount of energy released (exothermic) when a gaseous atom gains an electron.Ex) Cl(g) + e- Cl-1(g) ; DH = -349 kJAn important note about sign and size of DH.Some atoms will not accept an electron without having to absorb energy these are said to have no measurable EA.Electron AffinityOrdered by groups.Which group is the best?Why?Can we make predictions for transition metals?

MetalsRecall from Chapter 2 that metals:Are usually solids at room temperature.Are shiny and have lusterAre excellent conductors of heat and electricity.Tend to lose electrons easily (form cations).Metals oxides are said to be basic.Metal oxide + waterMetal oxide + acidMetalsSome metal ions can be identified by a flame test.Seen below from left: Li, Na, and K.

FireworksNon-metalsRecall from Chapter 2 that non-metals:Have a variety of appearances some are gases, some are solids, and one is a liquid. Some are single atoms like metals, while others are diatomic or larger.That are solids are brittle and poor conductors except carbon!Will either gain electrons (ionic) or share electrons (molecular).Non-metal oxides are said to be acidic.Non-metal oxide + waterNon-metal oxide + baseMetallic Character

Groups 1A and 7A are highly reactive.Group 1A loses 1 electron quite easily.Thus, reactivity driven by I.E.Reactivity increases from top to bottom within this group. Why?Group 7A gains 1 electron quite easily.Thus, reactivity driven by E.A.Reactivity greatest for F2 and Cl2 because they are gases.Reactivity and Periodic TableThe reaction of the alkali metals with water are all the same.2 Na(s) + 2 H2O(l) 2 NaOH(aq) + H2(g) + heat. Reaction is highly exothermic.

The reaction of Al with the liquid Br2.2 Al(s) + 3 Br2(l) 2 AlBr3(s) + heat.Reaction is also highly exothermic.Alkali metals and HalogensAlkali metalsHalogensGroups 2A and 6A are reactive, but less than the groups 1A and 7A.Group 2A metals will still react with water, but much more slowly.Group 6A non-metals (O2 and S) are similar.Groups 2A and 6ACalcium metalNa Na+ + 1e- ; H = +496 kJ

Ca Ca+2 + 2e- ; H = +1,735 kJ

Similarly, if you compared energy to gain one electron for group 7A versus gaining two electrons for group 6A, then the latter is much harder.Why the differences?Fairly unreactive with a few exceptions like Phosphorus.Phosphorus comes in several forms including both white and red phosphorus.White is the most unstable.Groups 3A, 4A, and 5A

Red Phosphorus and BromineGroup 8A forms no compounds at all with the exception of Xenon with Fluorine and Oxygen.XeF2, XeF4, XeF6, XeO3, and XeO4 have been produced.This is due to the relatively low I.E. of Xe. Group 8AReactivity on the left-hand side of the periodic chart (metals!) is driven by the ease of losing electrons.

Reactivity on the right-hand side of the periodic chart (non-metals) is driven by the ease of gaining electrons.Summary