CH105 Part II: Inorganic Chemistry The optimist sees the glass - - PowerPoint PPT Presentation

CH105 Part II: Inorganic Chemistry

The optimist sees the glass half full. The pessimist sees the glass half empty. The chemist see the glass completely full, half in the liquid state and half in the vapor state.

A proton and a neutron are walking down the street. The proton says, "Wait, I dropped an electron help me look for it.” The neutron says "Are you sure?" The proton replies …… What is the most important rule in chemistry? Never lick the spoon!

What did the scientist say when he found 2 isotopes of helium? HeHe Did you hear Oxygen went on a date with potassium? It went OK

IUPAC Nomenclature of elements With atomic number above 100

• Digit

Name Abbreviation

nil n

• 1

un u

• 2

bi b

• 3

tri t

• 4

quad q

• 5

pent p

• 6

hex h

• 7

sept s

• 8
• ct
• 9

enn e

114 Ununquadiu m Uuq 118 Ununoctium Uuo

Money has recently been discovered to be a not-yet-identified super heavy element.

The proposed name is: Un-obtainium.

Factors Affecting Atomic Orbital Energies

• The energies of atomic orbitals are affected by

– nuclear charge (Z) and – shielding by other electrons

• Higher nuclear charge increases nucleus-electron

interactions and lowers sublevel energy

• Shielding by other electrons reduces the full nuclear

charge to an effective nuclear charge (Zeff).

Zeff is the nuclear charge an electron actually experiences. True Love !!

• Orbital shape also affects sublevel energy.

Shielding

The energy order of orbitals for a given quantum number depends on shielding effects (σ), effective nuclear charge (Z*) & penetration of orbitals

Z* = Z - σ

(inner electrons !!!)

How to determine or estimate the Z*?

P . S. There may be other ways of calculating these as given in the literature. Please stick to this procedure as far as this course is concerned.

• 1. All e-’s in higher principal shell contribute 0 to σ
• 2. Each e- in the same principal shell contribute 0.35 to σ

{If the electron resides in s or p orbital}

• 3. Electrons in (n-1) shell: each contribute 0.85 to σ
• 4. Electrons in deeper shell: each contribute 1.00 to σ

Calculate the Z* for the 2p electron: Fluorine (Z = 9) 1s2 2s2 2p5 Z* = Z – σ

P . S.: There may be other ways of calculating these as given in the literature. Please stick to this procedure as far as this course is concerned.

Screening constant for one of the outer electron (2p): 6 six (2s2 2p4 two 2s e- and four 2p e-) = 6 X 0.35 = 2.10 2 (1s2 two) 1s e- = 2 X 0.85 = 1.70 σ = 1.70 + 2.10 = 3.80 and Z* = 9 - 3.80 = 5.20

• 1. All e-’s in higher principal shell contribute 0 to σ
• 2. Each e- in the same principal shell contribute 0.35 to σ
• 3. All inner shells in (n-1) and lower contributes 1.00

How to determine or estimate the Z*? {If the electron resides in d or f orbital}

• P. S.

There may be other ways of calculating these as given in the literature. Please stick to this procedure as far as this course is concerned.

Z* increases very slowly down a group for the “valence electron”. Example of Valence configuration as ‘ns1’

n Z Z* H 1 1 1.0 Li 2 3 1.3 Na 3 11 2.5 K 4 19 2.2 Rb 5 37 2.2 Cs 6 55 2.2

H He Li Be Be C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr

Z* is effective nuclear charge A neutron walks into a bar. He asks the bartender, "How much for a beer?" The bartender offers him a warm smile and says, "For you, no charge".

Z* increases rapidly along a period

For example, take period two

Li Be B C N O F Ne 3 4 5 6 7 8 9 10 1.3 1.9 2.4 3.1 3.8 4.5 5.1 5.8 2s1 2s2 2p1 2p2 2p3 2p4 2p5 2p6

Z* is effective nuclear charge

Penetration of Orbitals

Orbital shape causes electrons in some orbitals to “penetrate” close to the

• nucleus. Penetration increases nuclear attraction and decreases shielding

Radial probability

Penetration of Orbitals

The penetration potential of an orbital varies as: ns > np > nd > nf The energy of the orbitals for a given n varies as: ns < np < nd < nf The penetration of 2s electron through the inner core is greater than that of a 2p electron because the latter vanishes at the nucleus. Therefore, the 2s electrons are less shielded than the 2p electrons.

penetration of 2s e- is greater than 2p

penetration increases nuclear attraction and decreases shielding

The electrons present in f are much less influenced by the nucleus as compared to d, those present in d much less influenced as compared to p, than s, etc.

Influence of nucleus on electrons

Two electrons (2e-) present in the same d-orbital repel each other more strongly than do two electrons in the same s-orbital.

It is essential to consider all contributions to the energy

• f

a configuration, and just not

• ne-electron
• rbital

energies.

(Hostel Room Mates)

Order of filling of orbitals

Penetration and shielding have enabled atomic orbitals to be arranged in rough order of increasing energy.

How do you fill electrons ?

Depicting orbital occupancy for the first 10 elements.

Two electrons (2e-) present in the same d-orbital repel each other more strongly than do two electrons in the same s-orbital. Therefore, occupation of orbitals of higher energy can result in a reduction in the repulsion between electrons (for eg., 4s), otherwise the repulsion will be more if the lower-energy 3d orbitals were occupied. It is essential to consider all contributions to the energy of a configuration, and just not one-electron orbital energies. How do you fill electrons? Justification of 4s first over 3d

Experimental data show that d-block elements are

• f the form 3dn4s2, with 4s orbital fully occupied.

Sc (at. No. 21) is [Ar]3d14s2 This order is followed in most cases

• but not always! (some exceptions)

Z = 24 Cr [Ar] 3d54s1; not [Ar] 3d44s2 Z = 29 Cu [Ar] 3d104s1; not [Ar] 3d94s2 Two atomic configurations do not follow the sequence of filling of orbitals

As atomic number increases, energy of 3d orbitals decreases relative to both 4s and 4p At z = 29, energy of 3d becomes much lower than 4s Hence order of filling 3d < 4s < 4p

Filling & removal in Transition elements

• Transition series: filling order: 4s, 3d
• removal order (cation formation): 4s, 3d (not 3d, 4s)

e.g. Ti [Ar] 3d2 4s2

• Ti2+ [Ar] 3d2

(not [Ar] 4s2) Why?

Ti2+ [Ar] 3d2 4s0 (not [Ar]3d04s2) Why?

• When 2 electrons are removed, regardless of where they

come from, all atomic orbitals contract (Z* increases because of net ionic charge and reduced shielding)

• Contraction has a small effect on 4s orbital which owes its

low energy to its deep penetration

• Contraction in d orbital causes a considerable decrease in

energy – this decrease is evidently enough to lower the energy of 3d well below 4s in the ion that results from this.

“A lion runs the fastest when he is hungry.” “In life go straight and turn right.”

r decreases r increases

Metallic Radius

Metallic radii of 5d- block elements are expected to be larger than that of the 4d-elements, but found that these are not larger. Of course these are larger than 3d- block elements.

Lanthanide Contraction

f-orbitals have poor shielding properties; low penetrating power. So Zeff (Z*) increases (more significantly) from left to right (for 5d) across the period leading to more compact atoms.

Ionisation Energy (IE)

The minimum energy needed to remove an electron from a gas phase atom

Depends on:

(a) Size, IE decreases as the size of the atom increases (b) Nuclear Charge (NC), IE increases with increase in NC (c) The type of electron Shielding effect

Reasons:

(1) Average distance of 2s electron is greater than that of 1s (2) Penetration effect (3) Electronic configuration

1st IE: H = 1312 KJ mol-1 Li = 520 KJ mol -1

On moving across a period

1. the atomic size decreases 2. nuclear charge increases Thus IE increases along a period

Ionisation Energy (IE)

I would like to apologize for not adding more jokes... but I only update them.... periodically!

Electron affinity (EA)

The amount of energy associated with the gain of electrons

The greater the energy released in the process of taking up the extra electron, greater is the EA

The EA of an atom measures the tightness with which it binds an additional electron to itself.

Electron affinity (EA)

On moving across a period: As the size decreases, the force of attraction by the nucleus increases. Consequently, the atom has a greater tendency to attract added electron, i.e., EA increases Generally the EA’s of metals are low while those of non- metals are high Halogens have high EA. This is due to their strong tendency to change their configuration to ns2np6 On moving down a group, the atomic size increases and therefore, the effective nuclear attraction decreases and thus electron affinity decreases

The process can be exo or endothermic

Electronegativity

measure of the tendency of an element to attract electrons to itself (from its neighbour)

On moving down the group

• Z increases but Z* almost remains constant
• number of shells (n) increases
• atomic radius increases
• force of attraction between added electron and

nucleus decreases

Therefore EN decreases down the group

Electronegativity On moving across a period left to right

• Z and Z* increases
• number of shells remains constant atomic

radius decreases

• force of attraction between added electron

and nucleus increases

Hence EN increases along a period

Trends in three atomic properties.

Hardness and Softness

[Chemical but not mechanical] An important concept of compounds formed

Rich people's dines with richer ones !! High IE, smaller size, low polarizability -- makes Harder Low IE, larger size, high polarizability -- makes softer Chemical Hardness or Softness of an atom can be correlated with ionization energy (IE), electron affinity (EA), size and

• polarizability. If the IE > EA, the EA can be ignored.

The lighter atoms of a group are chemically harder The heavier atoms of a group are chemically softer Happiness is state of mind !!

SCN- can bind through either S or N depending upon the HSAB nature of the metal ion. For e.g., Si or Pt

Trends are exhibited, By keeping the metal same and changing the anion/ligand By keeping the anion/ligand same and changing the metal S will prefer Pt due to Soft … Soft type interactions, since ‘S’ is soft Lewis base & ‘Pt’ is soft Lewis Acid N will prefer Si due to Hard … Hard type interactions, since ‘N’ is hard Lewis base, & ‘Si’ is hard Lewis Acid.

Oxidation States

[In atomic state they are all zero].

Alkali atoms show +1 & alkaline earth shows +2 More electronegative atoms tend to form anions and lesser electronegative atoms tend to form cations when combined with others

Tendency of an atom to form ions with different oxidation states (negative or positive) would depend on solvation or hydration

• r ligation and lattice formation energies of the corresponding
• ions. Compare this with the IE.

Why did the noble gas cry?

Because all his friends Ar- gon.

Bonding (Interaction) types: Covalent, Non-covalent, Ionic

Non-covalent interactions are WEAK Interactions between (atoms, molecules, compounds) Atoms  Molecules  Supramolecules  Materials/Solids

Hydrogen bonding interactions Ion –molecular interactions Vander Waal’s interactions

What do you call a tooth in a glass of water? A one molar solution.

The name's Bond. Ionic Bond. Taken, not shared

Van der Waal’s Interactions

Three types of Van der Waal’s interactions:

a) Dipole – Dipole Interactions b) Dipole – Induced Dipole Interactions c) Induced Dipole – Induced Dipole Transient Dipole – Transient Dipole (London Dipersion Forces)

What do dipoles say in passing?

"Have you got a moment?”

Coordination numbers

Number of neighbours in interaction with the central ion

• - Can be primary (closely interacting and/or

bonding)

• - Can be secondary (distant than the primary but

interacting – mostly no bonding)

• - All this affects the reactivity, conductivity,

electronic and magnetic properties

Coordination Geometry

The way the nearest neighbours are arranged in space, a variety of geometries emerge: (Main group, Transition and Lanthanides) Linear (2) Trigonal (3) Tetrahedral (4) Square planar (4) Trigonal bipyramidal (5) Square pyramid (5) Octahedral (6) Pentagonal bipyramid (7) Singly capped octahedron (7) Doubly capped octahedron (8) Capped pentagonal bipyramid (9) Decahedron (10) Dodecahedron (12)

Linear (2) Trigonal plane (3) Square planar (4) Tetrahedral (4) Number Geometry Polyhedron

If I could rearrange the periodic table, I'd put Uranium and Iodine together.

Square pyramid (5) Trigonal bipyramid (5)

Number Geometry Polyhedron

Coordination No. 5

Did you hear oxygen and magnesium got together? OMg!

Triagonal prism Octahedral (6) Coordination No. 6

Singly capped

• ctahedron (7)

Pentagonal bipyramidal (7)

Number Geometry Polyhedron

Coordination No. 7

Doubly capped

• ctahedral (8)

Heptagonal dipyramid Tricapped triagonal prism

• Coordn. No. 8 (only for information, not for the exam)
• Coordn. No. 9 (only for information, not for the exam)

Pentagonal Prism Bicapped square Prism Octadecahedron Hendecahedron

• Coordn. No. 10 (only for information, not for the exam)
• Coordn. No. 11 (only for information, not for the exam)

Icosahedrons Cuboctahedrons Hexagonal prism Hexagonal antiprism

• Coordn. No. 12 (only for information, not for the exam)