General and Inorganic Chemistry I. - Lecture 1nlcd.elte.hu/szalai/pdf/lecture-4-handout.pdf ·...

General and Inorganic Chemistry I.Lecture 1

Istvan Szalai

Eotvos University

Istvan Szalai (Eotvos University) Lecture 1 1 / 40

Outline

1 Periodic Table


Periodic Table

Electron Configurations

The electrons occupy the orbitals in the way that gives the lowestenergy for the atom.

Pauli Exclusion Principle: No two electrons in an atom may haveidentical sets of four quantum numbers.

Hund’s Rule: Electrons occupy all the orbitals of a given subshellsingly before pairing begins. These unpaired electrons have parallelspins.

Electrons are assigned to orbitals in order of increasing value of(n + l).


Periodic Table

Electron Configurations

Li 1s22s1

C 1s22s22p2

Fe 1s22s22p63s23p64s23d6

Cu 1s22s22p63s23p64s13d10

Ce 1s22s22p63s23p64s23d104p65s24d105p66s25d14f 1

Valence shell: the electrons in the outer shell, those that were not presentin the precding noble gas orbitals.


Periodic Table

Periodic Table


Periodic Table

The properties of the elements are periodic functions oftheir atomic number

2 Li + 2 H2O → 2 Li+ + 2 OH− + H2

2 Na + 2 H2O → 2 Na+ + 2 OH− + H2

2 K + 2 H2O → 2 K+ + 2 OH− + H2


Periodic Table

History of the periodic table

About 330 B.C Aristotle proposed that everything is made up of amixture of one or more of four ”roots”. The four elements wereearth, water, air and fire.

In 1661 Boyle defined an element as a substance that cannot bebroken down into a simpler substance by a chemical reaction.

Lavoisier published a list of elements in 1789, or substances thatcould not be broken down further, which included oxygen, nitrogen,hydrogen, phosphorus, mercury, zinc, and sulfur. Lavoisier’sdescriptions only classified elements as metals or non-metals.

In 1817, Johann Wolfgang Dobereiner began to formulate one of theearliest attempts to classify the elements. He found that someelements formed groups of three with related properties. He termedthese groups ”triads”. (Cl-Br-I, Ca-Sr-Ba, S-Se-Te. . . )


Periodic Table

History of the periodic table

Alexandre-Emile Beguyer de Chancourtois, a French geologist, wasthe first person to notice the periodicity of the elements — similarelements seem to occur at regular intervals when they are ordered bytheir atomic weights. He devised an early form of periodic table,which he called the telluric helix. With the elements arranged in aspiral on a cylinder by order of increasing atomic weight, deChancourtois saw that elements with similar properties lined upvertically.

John Newlands was an English chemist who in 1865 classified the 56elements that had been discovered at the time into 11 groups whichwere based on similar physical properties.


Periodic Table

Dmitri Mendeleev (1837-1907)

Dmitri Mendeleev, a Siberian-born Russian chemist, was the first scientistto make a periodic table much like the one we use today. His table waspublished in 1869. It stated:

The elements, if arranged according to their atomic weights, exhibitan apparent periodicity of properties.

Elements which are similar as regards to their chemical propertieshave atomic weights which are either of nearly the same value (e.g.,Pt, Ir, Os) or which increase regularly (e.g., K, Rb, Cs).


Periodic Table

Dmitri Mendeleev (1837-1907)

The arrangement of the elements, or of groups of elements in theorder of their atomic weights, corresponds to their so-called valencies,as well as, to some extent, to their distinctive chemical properties; asis apparent among other series in that of Li, Be, Ba, C, N, O, and Sn.

We must expect the discovery of many yet unknown elements–forexample, elements analogous to aluminium and silicon–whose atomicweight would be between 65 and 75.


Periodic Table

Germanium

Properties Mendeleev prediction Observed values

Atomic weight 72 72,6Density 5,5 g/cm3 5,35 g/cm3

Oxide EO2 GeO2

Chloride ECl4 GeCl4


Periodic Table

Refinements to the periodic table

In 1914 Henry Moseley found a relationship between an element’sX-ray wavelength and its atomic number (Z), and thereforeresequenced the table by nuclear charge rather than atomic weight.Thus Moseley placed argon (Z=18) before potassium (Z=19) basedon their X-ray wavelengths, despite the fact that argon has a greateratomic weight (39.9) than potassium (39.1). The new order agreeswith the chemical properties of these elements, since argon is a noblegas and potassium an alkali metal.

In 1945, Glenn T. Seaborg proposed a significant change toMendeleev’s table: the actinide series. Seaborg’s actinide concept ofheavy element electronic structure, predicting that the actinides forma transition series analogous to the rare earth series of lanthanideelements.


Periodic Table

Classification of Elements

Name Electron structureRepresentative Elements

Alkali Metals ns1

Alkaline Earth Metals ns2

Boron group ns2np1

Carbon group ns2np2

Nitrogen group ns2np3

Oxygen group ns2np4

Halogens ns2np1

Noble Gases ns2np6

Transition Metals ns2(n − 1)d1−10

Lanthanides and Acthinides ns2(n − 1)d1(n − 2)f 1−14


Periodic Table

Periodic Properties

atomic radii, ionic radii

ionization energy

electron affinity

electronegativity

periodic chemical properties


Periodic Table

Atomic radii

Covalent radius: the nominal radius of the atoms of an element whencovalently bound to other atoms, as deduced the separation betweenthe atomic nuclei in molecules. In principle, the distance between twoatoms that are bound to each other in a molecule (the length of thatcovalent bond) should equal the sum of their covalent radii.

Ionic radius: the nominal radius of the ions of an element in a specificionization state, deduced from the spacing of atomic nuclei incrystalline salts that include that ion. In principle, the spacingbetween two adjacent oppositely charged ions (the length of the ionicbond between them) should equal the sum of their ionic radii.


Periodic Table

Atomic radii

Metallic radius: the nominal radius of atoms of an element whenjoined to other atoms by metallic bonds.

van der Waals radius: in principle, half the minimum distancebetween the nuclei of two atoms of the element that are not bound tothe same molecule.


Periodic Table

Atomic radii


Periodic Table

Atomic radii


Periodic Table

Atomic radii

Effective nuclear charge:Zeff = Z − S where S is the shielding effect of the inner electrons.

l . . . ni−1 ni

0 1,0 0,85 0,31 1,0 0,85 0,352 1,0 1,0 0,353 1,0 1,0 0,35

Na 1s22s22p63s1 Zeff = 11− (2× 1 + 8× 0, 85) = 2, 2Al 1s22s22p63s23p1 Zeff = 13− (2× 1 + 8× 0, 85 + 2× 0, 3) = 3, 6Na 1.86× 10−10m (186 pm)Al 1.43× 10−10m (143 pm)


Periodic Table

Ionic radii

Isolectronic species (10e−)Na+ Mg2+ Al3+

radius (nm) 0,097 0,066 0,051


Periodic Table

Ionization Energy

First ionization energy: the minimum amount of energy required toremove the most loosely bound electron from an isolated gaseous atom.

Be(g) → Be+(g) + e− 899 kJ/molBe+(g) → Be2+(g) + e− 1757 kJ/mol IE1 < IE2

Be2+(g) → Be3+(g) + e− 14849 kJ/mol IE1 < IE2 � IE3


Periodic Table

Ionization Energy


Periodic Table

Ionization Energy


Periodic Table

Electron Affinity

The amount of energy absorbed when an electron is added to an isolatedgaseous atom.

Cl(g) + e− → Cl−(g) −349 kJ/molO(g) + e− → O−(g) −141 kJ/molO−(g) + e− → O2−(g) +744 kJ/mol

Generally, nonmetals have more positive EA than metals. Atoms whoseanions are more stable than neutral atoms have a greater EA.


Periodic Table

Electron Affinity


Periodic Table

Electron Affinity


Periodic Table

Electronegativity

The electronegativity of an element is a measure of the relative tendencyof an atom to attract electrons to itself when it is chemically combinedwith another atom.

Pauling proposed the concept of electronegativity in 1932 as anexplanation of the fact that the covalent bond between two differentatoms (A–B) is stronger than would be expected by taking the average ofthe strengths of the A–A and B–B bonds.


Periodic Table

Electronegativity

To calculate Pauling electronegativity for an element, it is necessary tohave data on the dissociation energies of at least two types of covalentbond formed by that element.

∆ = EAB −√

(EAA × EBB)ENA − ENB = 0.102

√∆

ENF = 4.0


Periodic Table

Electronegativity

Mulliken proposed that the arithmetic mean of the first ionization energyand the electron affinity should be a measure of the tendency of an atomto attract electrons.

EN =IE + EA

130


Periodic Table

Electronegativity

Allred and Rochow considered that electronegativity should be related tothe charge experienced by an electron on the ”surface” of an atom: thehigher the charge per unit area of atomic surface, the greater the tendencyof that atom to attract electrons.

EN = 0, 359Zeff

r2+ 0, 744


Periodic Table

Electronegativity


Periodic Table

Chemical properties: hydrides

Ionic hydride → covalent hydrides

LiH BeH2 B2H6 CH4 NH3 H2O HF

NaH MgH2 (AlH3)x SiH4 PH3 H2S HCl

KH CaH2 Ga2H6 GeH4 AsH3 H2Se HBr

LiH + H2O → Li+ + OH− + H2

HCl + H2O → H3O+ + Cl−


Periodic Table

Chemical properties: oxides

Metal oxides → Nonmetal oxides

Li2O BeO B2O2 CO2 N2O5 OF2

Na2O2 MgO Al2O3 SiO2 P4O10 SO3 Cl2O7

KO2 CaO Ga2O3 GeO2 As2O5 SeO3 Br2O7

CaO + H2O → Ca2+ + 2 OH−

SO3 + 3 H2O → 2 H3O+ + SO2−4


Periodic Table

Ionic Bonding

2 Na(s) + Cl2(g) → 2 NaCl(s)


Periodic Table

Ionic Bonding

2 Na(s) + Cl2(g) → 2 NaCl(s)

Na [Ne] ↑ → Na+ [Ne]

3sCl [Ne] ↑↓ ↑↓ ↑↓ ↑ → Cl− [Ne] ↑↓ ↑↓ ↑↓ ↑↓

3s 3p 3s 3p

EB =Q2

4πε0

α

r

Electrostatic interaction(Coulomb force)α(NaCl) = 1.7475


Periodic Table

Ionic Bonding


Periodic Table

Types of Ions

Noble gas configurations2: H−, Li+, Be2+

s2p6: pl. Na+, Ca2+, Sc3+, Cl−, O2− . . .

d10s2 configurationSn ([Kr]5s24d105p2) → Sn2+ ([Kr]5s24d10) + 2e−

Tl+,Pb2+, Bi3+, . . .


Periodic Table

Types of Ions

Ions of transition metals

The first transition series is the result of the 3d orbitals being filled afterthe 4s orbital. However, once the electrons are established in their orbitals,the energy order changes - and in all the chemistry of the transitionelements, the 4s orbital behaves as the outermost, highest energy orbital.The reversed order of the 3d and 4s orbitals only applies to building theatom up in the first place. In all other respects, the 4s electrons are alwaysthe electrons you need to think about first.


Periodic Table

Types of Ions

Ions of transition metals

d10 configurationZn ([Ar]4s23d10) → Zn2+ ([Ar]3d10) + 2e−

Cu+, Ag+, Cd2+, Tl3+, . . .

[Ar]3d1: Ti3+

[Ar]3d2: V3+

[Ar]3d3: Cr3+

[Ar]3d4: Mn3+

[Ar]3d5: Mn2+, Fe3+

[Ar]3d6: Fe2+, Co3+

[Ar]3d7: Co2+

[Ar]3d8: Ni2+

[Ar]3d9: Cu2+


Periodic Table

Types of Ions


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