periodicity

Arrange elements to reflect the trends inchemical and physical properties.First attempt (Mendeleev and Meyer) arrangedthe elements in order of increasing atomic weight.Certain elements were missing from this scheme.Example: In 1871, Mendeleev noted that As properlybelonged underneath P and not Si, which left a missing elementunderneath Si. He predicted a number of properties for this element.In 1886 Ge was discovered. The properties of Ge matchMendeleev’s predictions well.Mendeleev is known as the “father of the periodic table”.

Development of thePeriodic Table

Effective Nuclear Charge

Nuclear Charge- given by the atomic number of an atom.Outer e- of an atom (determine many of an atomsphysical and chemical properties) do not experience the fullattraction of this nuclear charge b/c they are shielded from thenucleus and repelled by inner electrons.Effective nuclear chargeis the charge experienced by an electron on amany-electron atom and is less than the full nuclear charge.

Development of thePeriodic Table

Electrons are attracted to the nucleus, butrepelled by the electrons that screen it from the nuclearcharge.The nuclear charge experienced by an electrondepends on its distance from the nucleus and the numberof core electrons.

Effective NuclearCharge

Moving to a new period increases theshielding and decreases the nuclear pull.

As we move across a period; effective nuclearcharge increases with nuclear charge as there is no changein the number ofinner electrons.

Increase in nuclear charge is largely offset by increasein the number ofinner electrons; both increase by eight between successiveelements.Pull (ENC) experienced by outer electrons remainsapproximately the same down a group.

Consider a simple diatomicmolecule.The distance between the twonuclei is called the bond distance.If the two atoms which make upthe molecule are the same, thenhalf the bond distance is called theatomic radius of the atom.Atomic Radius - determined bythe attraction of the nucleus for itselectrons

Sizes of Atoms and Ions

Periodic Trends in Atomic RadiiAs the principal quantum number increases, thesize of the orbital increases.As we movedown a group, the atoms becomelarger,(radius increases) b/cmore energy levelsare occupied.As we moveacross a period, atoms become



Periodic Trends in Atomic RadiiAs the principle quantum number increases (i.e.,we move down a group), the distance of theoutermost electron from the nucleus becomes larger. Hence,the atomic radius increases and there is lessattraction between the nucleus and the outer e-.As we move across the periodic table, thenumber of core electrons remains constant. However, thenuclear charge increases. Therefore, there is an increasedattraction between the nucleus and the outermostelectrons. This attraction causes the atomic radius todecrease.


Trends in the Sizes of IonsIon size is the distance between ions in an ioniccompound.Ion size also depends on nuclear charge, number ofelectrons, and orbitals that contain the valence electrons.Cationsvacate the most spatially extended orbital and aresmaller than the parent ion. (it has less e-’s) Ex. When Na becomes Na+. Na loses 1 e-from it’s 3rd shell, now only having 2 which decreases the

Cationis alwayssmallerthan atom from which it is formed.Anionis alwayslargerthan atom from which it is formed.


Trends in the Sizes of IonsFor ions of the same charge, ion size increasesdown a group.All the members of anisoelectronic serieshave the same number of electrons.Asnuclear charge increasesin an isoelectronic series theions become smaller:

O2-

Ionization Energy

The first ionization energy,I1, is the amount of energy required to remove 1mol of electrons from the ground state of 1 mol of thegaseous atom.

Na(g)→Na+(

Prentice Hall © 2003 Chapter 7

Ionization Energy

There is a sharp increase in ionization energywhen a core electron is removed. (Valence e- are easierto remove)

Ionization Energy

Periodic Trends in Ionization EnergiesIonization energydecreases down a group.This means that the outermost electron is morereadily removed as we go down a group.As the atom gets bigger, it becomes easier toremove an electron from the most spatially extendedorbital.Ionization energy generallyincreases across a period.As we move across a period,Z

Electron Affinities

Electron affinity is the opposite of ionizationenergy.Electron affinity is the energy change when 1 molof a gaseous atom gains 1 mol of gaseous electronsto form a gaseous ion:

Cl(g

) + e-

→Cl

Electron Affinities

Metals, Nonmetals, andMetalloids


MetalsMetallic character refers to the properties ofmetals (shiny or lustrous, malleable and ductile, oxidesform basic ionic solids, and tend to form cations inaqueous solution).Metallic character increases down a group.Metallic character decreases across a period.Metals have low ionization energies.Most neutral metals are oxidized rather thanreduced.


MetalsWhen metals are oxidized they tend to formcharacteristics cations.All group 1A metals form M+ions.All group 2A metals form M2+ions.Most transition metals have variable charges.


Metals


MetalsMost metal oxides are basic:

Metal oxide + water→

metal hydroxideNa2O(s

) + H2O(l


NonmetalsMost nonmetal oxides are acidic:

nonmetal oxide + water→

acidP4O10(s

) + 6H2

Group Trends for theActive Metals

Group 1: The Alkali MetalsAlkali metals are all soft and silver in color. Veryreactive and usually stored in oil.


Group 1: The Alkali MetalsChemistry dominated by the loss of their singlevalence electron:

M→M+

+ e-(

M is Metal)


Group 1: The Alkali MetalsThe name alkali comesfrom the fact that theyproduce an OH when reactingwith water - this is known asan “alkaline” solution.(basic) more on this later.


Group 1: The Alkali MetalsAlkali metals produce different oxides whenreacting with O2:

4Li(s

) + O2(g)

Na line (589 nm):3p

→3stransition

Li line: 2p

→2stransition

K line: 4p

→4stransition


Group 1: The Alkali Metals


Group 1: The Alkali Metals


Group 2: The Alkaline Earth Metals

Do not memorize - Just forcomparison.

Group Trends forSelected Nonmetals

Group 6: The Oxygen Group

Do not memorize - Just forcomparison.


Group 7: The Halogens


Group 7: The HalogensThe chemistry of the halogens is dominated bygaining an electron to form an anion: (good oxidizingagents)

X2

+ 2e-

→2X-X



Reactivity decreases going down the group. (oxidizingability) So there are more e- between the outer shell andnucleusas you go down a group which decreases the attractionforoutere-.Because chorine is a stronger oxidizing agent thanbromine, it canremove the e- from bromide ions in soln to form chlorideionsand bromine.



Also- chlorine and bromine can oxidize ions to formiodine. remember activity series of halogens.Cl2(aq) + 2I-(aq)�2Cl-(aq) +I

We can test for halides in solution by addingsilver nitrate.Silver ions react with X-to form a silver halide precipitate.Ag+(aq) + X-(aq)�AgX (s)AgCl (s) = white

Melting Point- Ability to change from solid to liquid state.When a substance melts, the attractive forces holding theparticlestogether are overcome and the particles are free to movearound inthe liquid state.Two factors determine mp of a solid and the temp at whichit willmelt.

Alkali metals – mp decreases down the group – this isbecause as theatoms get larger the attractive forces decrease.Halogens – mp increases down the group – this is becausehalogensare held together by very weak forces, but these forcesincrease as the mass

of the atom increases.Across period 3 – there is a large pattern because of thebonding typechanges. Generally rise until it reaches a peak at Group 4(Si)

There is a very noticeable change in chemical propertiesacross periods.Ex. Period 3Na, Mg, and Al are metals – they are shiny, good conductorsofheat and electricity.Si is a semiconductor of electricity. It is a metalloid.P, S, Cl2, and Ar are non-metals and do not conduct electricity.

Metals can be distinguished from non-metals by chemicalproperties.Metal oxides, such as Na2O, MgO, and Al2O3are white solids at r.t.and have high mp and bp’s. In their liquid form they canconductelectricity and decompose into their elements in theprocess. These oxides are giant ionic structures.

Metal oxides are basic.(b/c when they react with water they form OH-)Na2O(s) + H2O(l)�2 NaOH (aq)MgO(s) + H2O(l)

The oxides of P, S, and Cl can also form the followingreactions:P4O6(s) + 6H2O(l)�4H3PO

Al2O3is insoluble in water and will not change pH. It will behaveas a base when it reacts with an acid and it behaves as anacid when it reacts with a base. This is calledamphoteric -(can act as acid or base)Al2O3

periodicity

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