Periodicity and Atomic Structure

Israel Schek ʬʠʸʹʩ ʷʹ 1

http://www.tau.ac.il/chemistry/undergraduate/undergraduate-courses.html

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Israel Schek ° �Ú ¥�±²¢�¢�� ¥³ ³¡¢ª±�¢©��

Tel Aviv University, Israel

General and Inorganic Chemistry:Atom Structure and the Periodic Table

: ʤ ʩʮ ʩʫ3 ̋ ʩʸ ʥʦʧ ʮʤ ʤʫʸʲ ʮʤʥ ʭ ʥʨʠʤ ʤʰʡʮ


Periodicity – Mendeleev Table (a)

¾ Dmitri Ivanovich Mendeleev (1837-1907) in 1869 (along with the

German chemist Julius Lothar Mayer) came out with the idea:

The known elements may be arranged in a periodic rectangular chart

according to their atomic mass (not yet atomic number – an

unknown concept at their current time).

¾ This periodicity would reflect their chemical and physical properties.

¾ To us, who are “scientifically educated”, it looks obvious that

chemical and physical properties are of periodic nature.

¾ But one should appreciate the genius insight of the periodic

arrangement rather than in the “more natural” linear arrangement.


¾ One of the first great successes of the periodicity was the prediction

of accurate characteristic properties of then (1871) concurrent

unknown elements, according to empty entries in the table.

¾ One hole was immediately below Silicon, and was temporarily

termed Eka-Silicon, later known as Germanium.

¾ The present order is according to the atomic number (nuclear

charge), not the atomic mass, as shown in 1913 by Henry Moseley

(1887-1915 Gallipoli, Dardanelles with other 55-65,000 Allied

Soldiers).

He measured wavelengths of the X-ray

spectral lines of a number of elements.

Periodicity – Mendeleev Table (b)


Earlier Versions (a)


¾ Dmitri Ivanovich Mendeleev

Centenary of Periodic Table,

Soviet Union 1969

showing his Notes

Earlier Versions (b)


Periods – General Structure

Block structure of the electron orbitals


Groups (Columns) at the Periodic Table

¾ The present table is of two-dimensional rectangular shape (with

obvious stairs and additional rows).

¾ Rows are called periods and columns are called groups.

¾ Elements belonging to a common group have in general similar

properties, which change gradually from one period to the next

one.

¾ The main reason is that they share common valence electronic

structure. The number of valent electrons defines the group

number.


Groups (Columns) at the Periodic Table

¾ http://www.privatehand.com/flash/elements.html


Alkaline Group¾ The elements of the first column are of the alkaline group, or

alkali metals or Group I (³ �¢¥°¥� ³�¤³§):

¾ Lithium, Sodium - ¨±³©, Potassium - ¨�¥²�, Rubidium, Cesium, Francium

¾ Electron configuration [inert gas]ns1

¾ These metals (high electron and thermal conductivities) are:

– Soft

– Low melting points

– Silver color

– All react drastically with water and replace the hydrogen. ³� ±¢��!

– Hardly found in their free state, but in compounds

– Metallic properties grow stronger down the column


Alkaline Earth Group

¾ The elements of the second column are of the alkaline earth group, or alkali earth metals or Group II ( ³�¢¥°¥� ³�¤³§ ³�¢±� ±«):

¾ Beryllium, Magnesium, Calcium - ¨�¢ª, Strontium, Barium, Radium

¾ Electron configuration [inert gas]ns2

¾ These metals have quite similar properties to their left neighbors:

– But less aggressive

– Soft

– Low melting points

– Hardly found in their free state, but in compounds

– Metallic properties grow stronger down the column


Group III

¾ Group III (no specific name) is at the 13th column (beyond the

large gap).

¾ Boron, Aluminum - ¨±§ , Gallium, Indium, Thallium

¾ Electron configuration [inert gas]ns2np1

¾ Apart from boron, which is semi-metal, other group elements are

metals.


Group IV

¾ Group IV (no specific name) is at the 14th column (beyond the large gap) is):

¾ Carbon - ¨§ , Silicon - ¨±�¯, Germanium, Tin - ¥¢��, Lead –³±�«


¾ Carbon is nonmetal and comprises Life

¾ Silicon (2nd most abundant element in the Earth mantle after oxygen), germanium - semimetals, and tin and lead are metals.

¾ Carbon, silicon, and germanium are usually in oxidation state +4, whereas tin and lead are +2.


Group V

¾ Group V (no specific name) is at the 15th column (beyond the

large gap).

¾ Nitrogen - ¨°© , Phosphorus - ¨ ±�, Arsenic, Antimony, Bismuth


¾ Nitrogen and phosphorus are nonmetals, arsenic and antimony

are semimetals, and bismuth is a metal.

¾ Oxides of the first three elements are acidic, that of bismuth is

basic and the oxides of antimony are in between, i.e. amphoteric.

¾ Nitrogen makes 80% of our atmosphere.


Group VI

¾ Group VI (no specific name) is at the 16th column (beyond the

large gap).

¾ Oxygen - ¨¯§ , Sulfur - ³¢±��, Selenium, Tellurium, Polonium


¾ Oxygen, sulfur, and selenium are nonmetals, tellurium is semi-

metal, and polonium is a metal (though not strong).

¾ Oxygen is the most abundant element on Earth, and the third most

abundant in Universe (after hydrogen and helium).

And, we breathe it.


Halogens Group

¾ Halogens (“producers of salts”), Group VII are at the 17th column

(beyond the large gap).

¾ Fluorine, Chlorine, Bromine, Iodine, Astatine


¾ React readily with most metals

Hence, found in Nature in compounds.

¾ Fluorine and chlorine are gases under regular conditions,

bromine is a liquid, and iodine is solid.


Inert Gases Group

¾ Noble gases or inert gases are at the 18th column

(beyond the large gap):

¾ Helium, Neon, Argon, Krypton, Xenon, Radon

¾ Electron stable configuration [previous inert gas]ns2np6

¾ Since they are extremely inert (though not absolutely),they are found in Nature as monatomic gases.

¾ Argon (left over after nitrogen and oxygen are removed from dry air) was first discovered by spectroscopic analysis in 1894 by Lord Rayleigh (recall Blackbody Radiation) and William Ramsay (1852-1916). Both won Nobel Prize in 1904.

Lord Rayleigh

William Ramsay


Periodic Table Groups (a)

¾ The “inertia” of inert gases suggested the explanation for chemical

bonding.

¾ As Gilbert Lewis (1875-1946) observed, atoms tend to arrange their

electronic constitutions in bonds, so as to reach an octet -

the stable octal inert gas configuration.

¾ William Ramsay (1852-1916) (left) and Physiologist Ivan Petrovich Pavlov (1849-1936)

60th Anniversary of Nobel Prize, Sweden 1964

Gilbert Lewis


¾ This scope was later extended when quantum mechanical tools

were applied to chemical bonds and the octet was rationalized.

¾ Ideas like molecular orbitals (MO) were incorporated and the

region of the electron wave function was then extended to the

whole molecule, rather than the sole atom.

¾ Moreover, it was found that in some molecules the inert gas octet

rule is not kept – more than eight electrons are populated around

some atoms (e.g. in SF6, PCl5).

¾ Although inert, mostly the heavier ones are capable to take part in

chemical bonds, though not very stable (e.g. Ar-Ar, Xe-Xe and

clusters like Xen (n>3)).

Periodic Table Groups (b)


Transition Metals

¾ The columns 3rd to 12th, between Group II and Group III are of

the transition metals.

¾ Their name is derived from their intermediate properties between

the active alkali and alkaline earth metals of the first two groups

and the mildly active metals of Group III and Group IV.

¾ First period of transition metals:

Scandium, Titanium, Vanadium, Chromium, Manganese,Ferum - ¥�±�, Cobalt, Nickel, Copper - ³²� ©, Zinc - ®��

¾ Electron configuration is of a partially filled d-orbital

[inert gas]ndk(n+1)s2


Consequences of the Orbital Energies (a)

¾ As a result of the shielding effect the 4s-orbitals, which penetrate

deeper into the nuclear zone than do the 3d-orbitals, have lower

energy than the latter.

¾ Hence, potassium (K) and calcium (Ca) have their electrons

exterior to the argon configuration in 4s - K: [Ar]4s1 and

Ca: [Ar]4s2.

¾ The next 10 electrons occupy in their turn the 3d-orbital to form

the transition metals, starting with scandium Sc: [Ar]3d14s2 up to

zinc Zn: [Ar]3d104s2.


¾ The energies of 4s- and 3d-orbitals are very close, and due to e-e

interaction, the system prefers energetically to have half-occupied

(d5) of the complete-occupied (d10) orbitals.

¾ Thus, an electron would be located in 3d-orbital rather than in the

4s-orbital.

For example, the configuration of chromium is Cr: [Ar]3d54s1

instead of [Ar]3d44s2, which is also acceptable, but a bit more

energetic.

¾ The copper configuration is Cu: [Ar]3d104s1 instead of [Ar]3d94s2,

where the 4s-orbital is only half filled.

Consequences of the Orbital Energies (b)


¾ This energetic nearness of 4s and 3d makes the transition metals

have different valences.

¾ After filling up the 4p-orbital from gallium Ga: [Ar]3d104s24p1

to krypton Kr: [Ar]3d104s24p6, the first long period is complete.

¾ Then we have the fifth row (or fifth period) with its 18 elements.

¾ It starts with rubidium Rb: [Kr]5s1, through the transition metals

yttrium Y: [Kr]4d15s2 to cadmium Cd: [Kr]4d105s2, via indium

In: [Kr]4d105s25p1 up to xenon Xe: [Kr]4d105s25p6, which

completes the fifth row and is an inert element.

Consequences of the Orbital Energies (c)


¾ Change of orbital

energies

of 4s and 3d between

Ca and Sc.

¾ Beyond Ca, 3d-orbital

energy falls abruptly.

Orbital Energies


¾ The sixth row starts with cesium Cs: [Xe]6s1 and barium Br:

[Xe]6s2.

¾ Then starts the filling up of the orbital 5d with lanthanum La:

[Xe]5d16s2.

¾ Next to lanthanum starts filling up the next 14 elements in which the upper electrons are of the orbital 4f (Ɛ=3). They are called

lanthanides.

¾ They start with cerium Ce: [Xe]4f15d06s2, through europium Eu:

[Xe]4f75d06s2, to lutetium Lu: [Xe]4f145d16s2.

Consequences of the Orbital Energies (d)


¾ Then going back to filling up the 10 transition metals with 5d

from hafnium Hf: [Xe]4f145d26s2, to mercury Hg: [Xe]4f145d106s2.

¾ Following are the main elements of 6p, starting with thallium Tl:

[Xe]4f145d106s26p1, and ending with the next inert element radon

Rn: [Xe]4f145d106s26p6, completing the sixth row.

¾ Following is the seventh row starting with francium Fr: [Rn]7s1

and radium Ra: [Rn]7s2.

¾ Then comes actinium with its 6d-electron Ac: [Rn]6d17s2.

¾ Then come the actinides with their 5f-electrons, starting with

thorium Th: [Rn]5f16d07s2, through americium Am: [Rn]5f76d07s2,

to lawrencium Lw: [Rn]5f146d17s2.

Consequences of the Orbital Energies (e)


¾ Verification of the configurations may be executed by passing a

beam of atoms through an inhomogeneous magnetic field, as done

by Stern and Gerlach.

¾ Since a spin is a magnetic moment attached to the electron,

it feels the magnetic field and interacts with it.

¾ According to the amount of deviation of the atomic beam from the

initial direction, one can determine the overall spin state of the

system.

¾ One may naively claim that the total spin moment is a sum of the

individual electronic spin moments.

Consequences of the Orbital Energies (f)


¾ For example, the [Ar]3d54s1 configuration of chromium has 6

unpaired spins (altogether 61/2=3 units of magnetic moment).

¾ The other, [Ar]3d44s2 configuration of chromium, has 4 unpaired

spins (altogether 41/2=2 units of magnetic moment.

¾ A beam of atoms in the first configuration would be deflected more

intensively than the second one, which is detected on the collecting

screen.

¾ A system with unpaired electron is a paramagnet and a system with

no unpaired electron is a diamagnetic.

Consequences of the Orbital Energies (g)


Lanthanides

¾ There are two more families of metals:

The 14 lanthanides starting with Lanthanum, which are

characterized by partially filled 4f-orbitals: [inert gas]6s24fk

¾ The 14 actinides starting with actinium, which are characterized

by partially filled 5f-orbitals: [inert gas]7s25fk

¾ They are very similar in their properties and therefore hard to

separate.

¾ The actinides include the heavy radioactive elements among them

the trans- uranium (Z>92).


Periodic Properties


Periodic Properties - Atomic Radii (a)

¾ Periodicity of properties is pronounced mostly in sizes and

energies. Following are some typical examples.

¾ The notion of atomic radius of an element is not a sharp concept

due to the wave-like nature of the elementary particles.

¾ One may locate the major part of the existence of the "electronic

cloud" in a certain radius around the nucleus.

¾ There are orders of magnitudes, and definitely, the uranium atom

is larger than a hydrogen atom.

¾ The atomic radius is defined as half the distance between

neighboring equivalent atoms (e.g. in H2 molecule or Fe solid).

¾ They are obtained in spectroscopic or crystallographic

measurements.


Atomic Radii in picometer (taken from Atkins & Jones)

Periodic Properties - Atomic Radii (b)


¾ Atomic radii decrease with growing atomic number across the

row, due to growing of the effective nuclear charges.

¾ Radii are higher at alkaline metals, lower at the earth-alkaline,

then decrease more gradually towards the inert elements.

Periodic Properties - Atomic Radii (c)


Periodic Variation of Atomic Radii in picometer (from Atkins & Jones)

Periodic Properties - Atomic Radii (d)


Periodic Properties - Mendeleev

Dmitri Mendeleev Centenary of Periodic Table,

Soviet Union 1969


Periodic Properties - Ionic Radii (a)

¾ As a basis to determination of the ionic contributions to inter-

nuclear distances one takes the radius of the oxygen ion O-2 as

0.14nm (=1.4Å).

¾ Cations are smaller than their parent elements due to loss of their

valent electrons, leaving behind much smaller cores.

¾ Anions on the other hand are larger than their parent atoms due to

gain of electrons.


Ionic Radii in picometer (taken from Atkins & Jones)

Periodic Properties - Ionic Radii (b)


Periodic Properties - Ionization Potential (a)

¾ Ionization potential is the minimum energy needed to remove an

electron from the ground state of the element

¾ ǻhionization is also called the ionization enthalpy (energy to be

invested at constant pressure of the gaseous atom).

¾ Since the electron is attracted to the rest of the atom, even if it is a

far electron in highly lying orbitals, energy always is invested to

remove an electron from the neutral atom.

ionizationǻH;e(g)EE(g) �� o


Ionization Potentials in kJoule/mole (taken from Atkins & Jones)

Periodic Properties - Ionization Potential (b)


¾ The ionization potential increases across a row from the alkaline

metals low values towards the inert elements high values.

¾ Than it falls back abruptly down to the next low value at the next

alkaline element at the next row.

¾ When the peak values at the inert elements are connected there is

a smooth decrease in the line starting at the highest value of

Helium (2370. kJ/mole) towards that of Radon (1040. kJ/mole).

¾ It is much harder to ionize the helium atom rather than the larger

inert atoms.

Periodic Properties - Ionization Potential (c)


¾ Atoms become smaller across a period (see the atomic radii

graph).

¾ Due to increasing effective charges electrons are more strongly

attracted to the nuclei when going from the alkaline element

toward the inert element in a period.

¾ Going down a group the outermost electrons are farther away

from the nuclei and therefore less attracted and less bound.

Periodic Properties - Ionization Potential (d)


¾ Since on the average outermost d-orbitals are deeper in location

than their preceding s-orbitals, moving from one element to the

next one across a transition elements group does not show a steep

change in relative values.

¾ This point is more emphasized across the lanthanides and

actinides, where there is hardly a change in the ionization potential

inside the group.

Periodic Properties - Ionization Potential (e)


Periodic Variation of Ionization Potentials (from Atkins & Jones)

Periodic Properties - Ionization Potential (f)


Periodic Properties - Electron Affinity (a)

¾ Electron affinity (meaning attraction towards electrons) is the

energy released when an electron is added to an atom or ion

¾ ǻHgain is positively defined when there is a gain of energy when

the extra electron joins the neutral atom, and negatively when

energy must be invested to force an electron into the neutral atom.

¾ Contrary to ionization potential, the process of gaining an electron

may be exothermic as well as endothermic.

gainǻH(g);EeE(g) �� o�


¾ The upper right corner is electron-philic (love of electrons).

¾ It is mostly pronounced for the fluorine, and other halogens, but also the VI group of oxygen and sulfur. The electron is welcomed.

¾ Inert elements do not "feel" affinity towards an extra electron,mostly neon.

¾ Nitrogen and some earth-alkaline dislike the idea to be negative.

Periodic Properties - Electron Affinity (b)


Electron Affinity in kJoule/mole (taken from Atkins & Jones)

Periodic Properties - Electron Affinity (c)


¾ Electronegativity is the extent of the element attraction towards

the electron.

¾ Linus Carl Pauling (1901-1994) and Robert Sanderson Mulliken

(1896-1986) defined a numerical measure for electronegativity.

¾ Pauling: 1954 Nobel Prize laureate for his studies on the nature of

the chemical bond and molecular structure of proteins.

¾ Mulliken: 1966 Nobel Prize laureate for his studies on chemical

bonds and the electron structure of molecules by means of the

orbital method.

Periodic Properties - Electronegativity (a)


¾ The electronegativity considers the resistance to lose an electron

(expressed by the ionization potential) as well as the agreement to

have an extra one (expressed by the electron affinity).

¾ The simplest expression (due to Mulliken (1934) is the average of

both values

EA)/2(IPȤ �

Periodic Properties - Electronegativity (b)


¾ Very electronegative elements (halogens, small VI group-atoms)

at the upper right corner zone.

¾ Electropositive elements (heavy alkaline, heavy earth-alkaline)

at the lower left corner (“they dislike electrons”).

Atkins & Jones

Periodic Properties - Electronegativity (c)


¾ The scale runs as

¾ The electronegativity would serve defining the polarity of a

chemical bond.

¾ The electric dipole moment of a simple diatomic molecule AB is

approximately linearly related to the electronegativity difference.

¾ The larger this difference, the more polar the bond is.

4ȤȤ0.79Ȥ FCe dd

BA ȤȤȝ �|

Periodic Properties - Electronegativity (d)


The periodicity of the Pauling electronegativity

Linus Carl Pauling,

Upper Volta 1977

Periodic Properties - Electronegativity (e)


Finally, in 3-Dim representation:

Periodic Properties - Electronegativity (f)

Periodicity and Atomic Structure

Documents

columnisrael schek

group number

alkaline group

periodic table http

number of elements

electron orbitalsisrael

present table

common group