ch4252 inorganic chemistry 1b spring 2013 semester...

Nanotechnology Group

1

Dr. Kevin M. Ryan Department of Chemical and Environmental Sciences

University of Limerick

CH4252 Inorganic Chemistry 1B Spring 2013 Semester

Introduction


2 Time Slots 2013

(2 Lecture slots per week and 1 lab) •  Lectures 2-3 on Monday, P1033

4Lecture 11-12 on Tuesday, FG042 Labs beginning Wk 2

2A 9-12 Wednesday B3053 2B 11-2 Thursday B3053

Tutorials TBA


3 Lecture Notes

•  Lecture notes will be available from the my research group website www.ifnano.com and are password protected

•  Electronic Lecture Notes will only cover 70% of course content –remainder during lecture, on board or on slides marked with red square in lecture which are not in download

•  Notes will be available for download week before lectures beginning week 2


4 Assessment

•  Final examination – 70% Format of examination will be given in Tutorial 1

•  Coursework – 30% –  1 Test 10% (half way through the course) – Lab reports 20%


5 Recommended texts

•  You will need to refer to a General Chemistry book (Any) Get your own !!

•  Recommended Oxford Chemistry Primers *(15) Chemical Bonding Mark J Winter

•  Concise Inorganic Chemistry 4th/5th. Edition, J.D.

Lee chs. 1-7 should be used for reference. •  Shriver and Atkins Inorganic Chemistry 4th

Edition


6 Aims

q  To provide a modern understanding of chemical bonding in molecules.

q  To describe the relation between bonding, structure and properties in solids

q  To understand the factors determining the structure of ionic compounds.

q  To describe the structure of simple compounds in terms of close-packing.

q  To calculate the lattice energies of ionic compounds. q  To describe the bonding in transition metal

complexes.


7 Syllabus:

A. Revision of Lewis structures and VSEPR theory. Covalent bonding: simple molecular orbital treatment for diatomic

molecules;valence bond theory. hybridisation, resonance and electron delocalisation.

Comparison of valence bond and molecular orbital approaches. Polarity in bonds. B. Bonding in transition metal complexes: crystal field and ligand field theory. C. Relationship between structure, bonding and properties in solids: metallic,

ionic, molecular and covalent solids. Structure of metals and close-packing. Structures of simple compounds in

terms of close-packing. Ionic crystals: Factors affecting crystal structure; ionic radii, radius ratio and its

importance; Madelung constants and estimation of lattice energies; the Born-Haber cycle.

Influence of bonding on the physical properties of materials is emphasised throughout the module.


8 Basic chemical ideas

•  We will be revising some basic ideas underpinning this course and these are covered in General Chemistry 1 notes and in any General Chemistry textbook.

•  It is ESSENTIAL to master these ideas in order to understand the chemistry covered in this course, both theory and practical.

•  Tutorial exercises should be done as requested and written up in a Problems Book. The questions will be discussed subsequently in a tutorial.


9 Basic ideas #1

You should be able to: □ Recognise and use the names and symbols of the s and

p block elements and the first row of the d block. □ Know the names and formulae for common laboratory

acids and alkalis, and the names and formulae of common ions.

□ Write the formulae and names for simple ionic and covalent molecular compounds.

□ Identify and balance simple non-redox and redox equations.

□ Work out oxidation numbers for elements in compounds or ions.


10 Basic ideas #1(contd.)

□ Work out the GMM of a pure compound from its formula and calculate the number of moles and number of particles in a given mass of substance.

□ Work out the number of moles of solute in a solution or the concentration of solution knowing the amount dissolved in a given volume.

□ Work out the number of moles and number of molecules in a given volume of gas at a given pressure and temperature and vice versa (ideal gas law).

□ Work out volumetric calculations involving two solutions or a solution and a solid using a first principles approach.

□ Be able to work out how to dilute a solution down to a specified concentration.

Tick off these items as you think you have mastered them and can use

the ideas. □


11

Lecture 1-3 CH4252

Dr Kevin M. Ryan


12


13 What is an atom ?

•  Rutherfords α particle Experiment

Expected Result uniform distribution

Experiment Actual Result


14 The Bohr Model

•  Bohr was trying to account for the emission spectra of Hydrogen

•  Bohr suggested that negatively charged electrons revolve around a nucleus which is positively charged at certain fixed distances called orbits

•  However 2 problems with this model 1: electron will spiral into nucleus due to electromagetic interactions and energy loss and all values of r are allowed which is not the case in the spectrum

•  His solution is that only certain orbits are allowed (quantised)

•  This implies that the energy is also quantised


15 Bohr Model (Contd)

•  When atoms are heated they emit light – colours are characteristic but the emitted light is not continuous i.e not all frequencies are emitted

•  Lyman n’ à 1 •  Balmer n’ à 2 •  Paschen n’ à 3

• Bohr model e.g. orbits (Leaving cert) breaks down in certain cases e.g atomic spectrum in magnetic field is different to normal therefore a new model is needed to fully explain properties


16 Hydrogen


17 The nature of atomic orbitals

The modern picture of electrons in atoms

Orbit (Bohr) Orbital (known position Energy known, and energy/velocity) position uncertain - only probability known Electron in an atom described by Schrodinger wave equation → wave equation, ψ Probability of finding an electron is ψ2

Solutions ψ1 , ψ2 ... corresponding to E1, E2 ...


18 Particle or wave from Orbits to Orbitals

•  de Broglie λ = h/mv •  Suggests that particle and wave phenomena are different attributes

of all matter •  Heisenberg Uncertainty Principle : It is not possible to

simultaneously measure the position and momentum of an electron •  Use probabilities: The probability of finding an electron at a point

equates to the electron density at that point •  The orbitals we draw reflect boundary conditions beyond which we

expect the electron density to be zero


19 Quantum Numbers

•  Each wavefunction is obtained by solving the Shrodinger equation is labelled by a set of three integers called quantum numbers

n, l and ml

n = principal quantum number l = orbital angular momentum quantum number ml = magnetic quantum number

The principal quantum number defines a series of shells, The orbitals belonging to each shell are classified into subshells distinguished by l, For a particular value of n, l can have values l=0, 1, …n-1


20 Quantum Numbers

Value of l = 0, 1, 2, 3 Orbital designation s p d f Subshell with quantum number l consists of 2l +1 individual orbitals, distinguished by the magnetic quantum number ml which has 2l +1 integer values from +l to –l. This quantum number specifies the component of orbital angular momentum around an arbritary axis passing through the nucleus. e.g, s subshell 1 orbital ml = 0 p subshell, three orbitals ml = +1, 0, -1

d subshell, five orbitals ml = +2, +1, 0, -1, -2


21 Shells, Subshells and Orbitals

Classification of orbitals into subshells (same value of l) and shells (same value of n)


22 3d Orbital view

s orbital

p orbitals


23 d-orbitals


24 f- orbitals


25 Drawing

•  s Orbital

•  p Orbitals

•  d Orbitals


26 Consider the first 20 elements

•  Know the First 20 Elements


27 Janet Form Periodic Table


28 Basic Rules

•  Pauli Exclusion Principle: An orbital can accommodate only two electrons not any more and if two do occupy the orbital their spins must be paired

•  Occupancy is either 0, 1 or 2 •  Hund’s rule: When more than one orbital

has the same energy, electrons occupy separate orbitals and do so with parallel spins


29 Building up Principle

•  The ground state electron configuration of many electron atoms are determined experimentally by spectroscopy

•  Variances 4s is lower in energy than 3d up to Sc


30 E.g. Phosphorous

•  Phosphorous Atomic Number 15 1s2 2s2 2p6 3s2 3p3

↑↓

↑↓

↑↓ ↑↓

↑↓

↑

↑↓

↑ ↑

1s

2s

2p

3s

3p

4s

3d

Hund’s rule


31

•  Neon

•  Phosphorous Atomic Number 15 1s2 2s2 2p6 3s2 3p3

[Ne] 3s2 3p3 or [Ne] 3s2 3px1 3py1 3pz1

↑↓

↑↓

↑↓ ↑↓

↑↓

↑

↑↓

↑ ↑

1s

2s

2p

3s

3p

4s

3d

E.g. Phosphorous


32 Lewis Structures of Atoms

•  Writing elements in shorthand form e.g. [Ne] 3s2 3p3 tells us the number of valence electrons as the 1s2 2s2 2p6 portion denoted [Ne] is a closed shell with its full complement of electrons

•  The valence electrons of phosphorous can be represented as a lewis structure

P


33 Atomic Structure

•  Core Electrons: Innermost electrons corresponding to the electron count of the previous noble gas

•  Valence Electrons: All the electrons that are not core electrons


34 Lewis Structures

Fill in the Lewis structures of the first 18 elements

H

Li Be B C N O F Ne

He

Na Mg Al Si P S Cl Ar


35


36 Revision of Trends

•  Atomic Radii –increase down a group and within s and p blocks decrease from left to right across a period (Lanthanide contraction will be dealt with later!)

•  Effective nuclear charge, the energy of an electron in an orbital and the orbital size are dependent on the effective nuclear charge Zeff , the nuclear charge experienced by an electron is reduced by shielding from other electrons.

•  The first ionisation energy, I1 is the energy required to remove the least tightly bound electron

•  Electronegativity of an element is the power of an atom of the element to attract electrons when it is part of a compound; there is a general increase across a period and a general decrease down a group


37 Effective Nuclear Charge


38

Atomic Radius


39


40


41


42 Tutorial Questions 1.1

1.  Complete the following

table

2.  Give the ground state electron configuration of (a) C, (b) Ca, (c) Bi, (d) Ca2+

3.  Identify the elements that have the ground state electron configurations (a) [Ne] 3s23p4, (b) [Kr]5s2, (c) [Ar] 4s2 3d3 , (d) [Kr] 5s24d105p1

4. Account for the decrease in first ionisation energy between phosphorous and sulphur

n l ml Orbital designation

No. of Orbitals

2 2p

3 2

4s

4 +3,+2…,-3


43 Tutoral Questions 1.1

•  5. Suggest a reason for why the increase in Zeff for a 2p electron is smaller between N and O than between C and N


44

CH4252 2 Lewis Molecular

Structures


45 Covalent and Ionic Bonding

•  2 types of non-metallic chemical bonding covalent and ionic

Ionic bonding occurs when a compound such as common salt NaCl adopts a lattice structure consisting of a regular array of +ve charged ions (cations) and negatively charged ions –ve (anions)

Lewis suggested that a covalent bond is where two atoms are held together by shared pairs of electrons


46 IONIC versus COVALENT BONDING

← (XA-XB) 3.0_____________________________________0 Ionic Polar Covalent bonding covalent bonding bonding 3-D network 3-D network Molecular solids solids substances Metal+ Non-metal + Non-metal non-metal N.B. XA and XB are the electronegativities of the combining atoms.


47

Chemical Bonding

Compound, AxBy

/ \ XA-XB > 1.7 XA-XB < 1.7

IONIC COVALENT BONDING BONDING

↓ ↓ 3-D MOLECULAR

NETWORK ↓ ↓

r+/r- Number of valence e- ↓ ↓

Co-ordination Lewis structure number (C.N.) ↓

↓ VSEPR theory

↓ CRYSTAL MOLECULAR STRUCTURE SHAPE


48 Lewis Structures

•  Single bond A:B denoted A-B •  Double bond A::B denonted A=B •  Trible Bond A:::B denoted A B •  Unshared Pair of valence electrons (A:) is known

as a lone pair (don’t contribute directly to bonding but do influence shape)


49 The Octet Rule

•  Each atom shares electrons with neighbouring atoms to achieve a total of eight valence electrons (an octet) [remember a closed shell noble gas configuration was achieved when 8 electrons occupy s and p shells, exception is hydrogen atom]

•  3 Steps –  Decide on the number of electrons in structure by adding

together the number of valence electrons provided by the atoms –  Write the Chemical Symbols of the atoms in the arrangement that

shows which atoms are bonded together –  Distribute the electron pairs to form single bonds, multiple bonds or lone

pairs


50 Examples

•  BH4- ion Step 1 Count Electrons Step 2 Organise Atoms Step 3 Distribute Electrons Step 4 Formal Charges Step 5 Structure


51 Step 1

B= [He] 2s22p1 à 3 Valence e- H = 1s1 à 1 Valence e-


52 Step 2

Classify the formula into a central atom A and a set of surrounding ligands B The Central atom most often is the least electronegative atom but a hydrogen atom is always a ligand and never a central atom


53 Step 3 Distribute Electrons Using the total no of electron pairs place a single bond between the central atom and the connecting ligand atom Place any extra electrons, in pairs, on the central atom. These are lone pairs. . If central atom has less than 8e-, form double (or triple bonds) to give central atom an octet (except group III).


54 Step 4 Formal Charge

Move electrons if possible to minimise the formal charge on the central atom by creating one or more double bonds where this is possible N.B. if the central atom is from the Second row and already has an octet it cannot be given any more electrons to minimise the formal charge.


55 Step 5 (Test structure )

If more than one equivalent structure can be written after minimising formal charges, write these as resonance structures (see slide 59) e.g. 03, NO2, SO4

2-


56 Step 6 ! Variants

1.  If there is an odd number of electrons, leave an odd electron on the

central atom e.g. NO 2.  Second row elements (> group 14(IV):C,N,O,F) always have an octet,

BUT heavier elements may have more than an octet where this is possible or necessary. (participation in bonding by d orbitals

3.  BF3 is an exception to the octet rule with all fluorine atoms attaining the octet rule with boron only 6*

* Bear in mind these are models for structures and not exact structures and although it fits in most cases other models such as hybridisation may be a better fit for BF3


57 Examples

•  BF4- ion


58 Examples

•  Lewis structures don’t give shape only pattern of bonds and lone pairs shape of BF4- is tetrahedral

•  Predicting shape will be carried out in VSEPR section


59 Resonance

•  In certain cases a single lewis structure is an inadequate description of a molecule

•  One resonance structure of NO3-


60 Formal Charge and Resonance

•  Formal charge can determine the lowest energy resonance structure as in NO2F

•  The structure with the highest charge has the highest energy such that structures with double bonds are most likely to be to most dominant resonant structures


61 Revision (Formal Charge and Oxidation Number are different)

•  Formal Charge is obtained by exaggerating the covalent character of a bond whereas oxidation numbers are obtained by exaggerating the ionic character of the bond


62 Tutorial Lewis Structures

•  Work out the lewis structures of •  H2, N2, C2H2, 03, S03, SO4 2-HCN, CHCl3,

NH4+,

•  SeF2, POCl3, XeO4, •  PO4

3-, PCl2-, OCN-, •  , NO2, SO2, •  XeOF4, PF5, SCl4-

•  (first atom is the central atom in nearly all all cases except those marked by underline)

ch4252 inorganic chemistry 1b spring 2013 semester...

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