Chem Notes Chapter_8

8/13/2019 Chem Notes Chapter_8

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Periodic Properties of the

Elements

Chapter 8


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The Periodic Table Based upon periodic law

Independently proposed in 1869 by L. Meyer andD. Mendeleev

Similar properties recur periodically when

elements are arranged according to increasingatomic number*

Meyer used graphical methods of properties andatomic mass to find periodicity

Needed several different plots to confirm that there wererecurrent themes

* From General Chemistry Petrucci et al. Prentice Hall, New Jersey, 2007, page342.


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The Periodic Table Mendeleev used table format

Left spaces for elements not yet found Predicted properties of proposed elements Noble gases had not been discovered Organization was based on the formation of the

oxide or halide ratio and the atomic mass Atomic masses were converted to atomic number

by H. G. J. Moseley using x-ray analysis

Other formats of the periodic table have beenprepared since Mendeleevs version


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Electron Configurations e-configurations are linked to periodic table groups

Valence e-

Are part of the outermost or valence shells Usually have the highest principle quantum number Are involved in reactions

Core e- Are completed energy levels, i.e. the e- from the previous

noble gas Are not involved in reactions

Have rules for the orbital assignment of the e-


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Non-H Elements For multi-electron elements, e- may be counted and

assigned to orbitals using spdf notation or orbitaldiagramming

Orbitals have the same approximate shape as forhydrogen

e-with the same are no longer degenerate Shielding or screening

inner e-act as a barrier to the p+pull of the nucleus e-that are further away from the p+are less tightly held (are

loosely held)

The amount of p+

charge that the shielded e-

feels isreduced and known as Zeffor the effective nuclear charge


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Non-H Elements Effective Nuclear Charge

Zeff= Z - S p+charge that the e-actually feels Is higher the closer the electron is to the nucleus The inner core/noble gas e-shield the

outer/valence e-

from the nuclear charge but donot shield each other Penetration

Is an increased p+pull/feeling of the nucleus on

the e-when the e-is near to the nucleus Caused by overlap of the orbitals


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Electron Configurations Rule 1 Aufbau Principle

The e-go into the ground state/lowest possible energy conformation

in the order of increasing e-energy Start with filling the orbitals of lower energy first (inner shells), then

the valence shells with e- The orbital filling order was determined by experiment

Rule 2 Pauli Exclusion Principle

No two e-

in an atom can have the same set of four quantumnumbers Results in 2e-per orbital and the 2e-have opposite spin Determined by the Stern-Gerlach experiment

Rule 3 Hunds Rule

There should be as many unpaired e-

as possible Orbitals are first filled with 1e-, and then doubled up due to e- e-repulsions


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Electron Notation and Orbital

Diagrams Two ways to write the e-configurations (e.g.

16O) Spdf notation

Expanded version 16O: 1s2, 2s2, 2px2, 2py1, 2pz1

Condensed version16

O: 1s2

, 2s2

, 2p4

Inner electron version 16O: [He] 2s2, 2p4

Orbital diagrams

16O: 1s 2s 2p


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Diagrams ransition metal spdf notation may be written

two ways: [Noble gas] 3dx, 4sy [Noble gas] 4sy, 3dx

Exceptions to the three rules are: common for larger transition metals, the

lanthanides and the actinides Figure 8.7 has e-configurations for the elements Cr: [Ar] 3d5, 4s1 Cu: [Ar] 3d10, 4s1


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Diagrams he periodic table may be used to identify the

orbital type that is being filled: s block

Group 1A alkali metals Group 2A alkaline earth metals

p block Groups 3A-8A Group 7A halogens Group 8A noble gases, have completely filled shells

d block Groups 1B-8B transition metals f block

Inner transition metals Lanthanides and actinides


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Periodic Properties and TrendsAtomic Size

Work with effective atomic radius distance at 90% of the e- charge densitymeasured in terms of the internuclear

distance (distance between 2 atoms) Measured originally in angstroms (),currently in pm

1 nm = 1000 pm = 1 x 10-9m = 10 Have several types of radii and trends


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Periodic Properties and Trends Bonding atomic radius or covalent radius

measured as distance between 2 covalently bound or crystalline

atoms (single bond) Non-bonding atomic radius or van der Waals radius measured as distance between 2 nuclei in a solid as determined

by the density of the solid, not bound atoms Ionic radius

measured as distance between 2 ionically bound atoms takinginto account the difference between anions and cations Metallic radius

measured as distance between 2 metal atoms in the crystallinemetallic solid

Atomic radius Is an average bonding atomic radius determined using differentbound elements

Is smaller than the van der Waals radii


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Periodic Properties and Trends Atomic Ra ii tren 1 The atomic radius increases as you go down a group Larger values of n give larger atoms

Atomic Radii trend 2 The atomic radius decreases as you go across a row (left to

right) e-are entering the same subshell (amount of shielding is the

same) but the nucleus is growing (more p+pull) Atomic Radii trend 3

Transition and inner transition radii change very little The e- are entering into inner subshells which have similar

energy to the surrounding filled subshells

Atomic Radii trend 4 Fully filled and filled orbital sets have decreased radiiwhen compared to those around them

Related to shielding, Zeffand e-e-repulsions


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Periodic Properties and Trends Ionization energy (IE)

Is the amount of energy required for a gaseousatom to lose a loosely held e- Gives a positive energy value

The easier an element is to ionize, the more

metallic the cation is The first e- is easy to remove, the second e- is

difficult since it is more tightly held due to a cationhaving been formed

Valence e- are relatively easily removed, whereasinner e-are very difficult as a set of filled orbitalsmust be opened/broken


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Periodic Properties and Trends Ionization Energy (cont.)

Trends Increases across rows (left to right, same n)

Related to Zeff Decreases when going down groups (larger n)

Related to increased size and shielding Decreases with increasing atomic radii

Exceptions Al is smaller than Mg removal of e-from the p versus s

subshell S is smaller than P filled orbitals require less energy

to form than do other arrangements


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Periodic Properties and Trends Electron Affinity (EA)

Amount of energy lost upon gaining an e-

in thegaseous state to form an anion Stronger likelihood of gaining an e-high e-affinity

low EA value large negative value (most negativenumber)

Need the new e- configuration to be more stablethan the original (either filled or completelyfilled orbitals)

First EA value commonly negative (spontaneous) Second EA value commonly positive (notspontaneousenergy is required to add the e-)


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General Properties Metals

Good conductors of heat and electricity Malleable and ductile solids Tend to lose e-to form cations (positively charged

ions), some metals form more than one cation Require energy to lose the e-(not spontaneous) May be

Main group lose from the s orbital to obtain thenearest noble gas e-configuration

Transition lose from the s subshell first, then the dsubshell to form either filled or completely filledsubshells (rarely form noble gas e-configuration)


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General Properties Non-metals

Groups 3A-8A Poor at conducting heat and electricity Brittle solids, liquids or gases

Gain e-

to make negatively charged ions (anions)with e-configurations of the nearest noble gas First e- gained spontaneously, second e- requires

some energy for addition (not spontaneous)


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General Properties Noble gases

Group 8A Inert Complete e-configurations Other elements gain/lose e- to have the same e-

configuration Metalloids

Tend to look like metals Have some metal properties and some non-metalproperties Tend to gain e-to become the nearest noble gas,

although some lose e-

Chem Notes Chapter_8

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