Introduction to chemical bonding and Thermodynamic (6A4X0)
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Technische Universiteit Eindhoven (TUE)
Book
Principles of Modern Chemistry
This is a complete summary for the course "Introduction to Chemical Bonding and Thermodynamics" at TUe (Eindhoven University of Technology). The course code is 6A4X0.
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Technische Universiteit Eindhoven (TUE)
Chemical Engineering
Introduction to chemical bonding and Thermodynamic (6A4X0)
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(6A4X0) Introduction to
Chemical Bonding &
Thermodynamics
Full Notes!
1
,Contents INDEX Page number
-C1. Atomic Structure 3
-C2. Chemical Bonding 5
-C3. Quantum Mechanics 21
-C4. Quantum Mechanics and Atomic Structure 31
-C5. Quantum Mechanics and Molecular Structure 53
-C6. Bonding in Organic Molecules 69
-C7. Molecular Spectroscopy and Photochemistry 75
-C1. The Properties of Gases 94
-C2. The First Law of Thermodynamics 97
-C3. The Second and Third Law of Thermodynamics 122
2
, Chemical Bonding Part
Atomic Structure
Distribution of Mass and Charge in an Atom:
Most of the mass is concentrated in the centre of an atom, in the nucleus. The
nucleus is made up of protons and neutrons. While the electrons move around
in regions of space called orbitals. The number of protons and electrons are
the same as the number of positive charges must be equal to the negative
charges.
Proton number and nucleon number:
To calculate the number of protons, neutrons and electrons we used the
proton number and nucleon number.
Nucleon Number or
mass number:
The number of protons +
A electrons.
Symbol of the element.
X
Proton Number or atomic number:
Z
The number of protons which equals
the number of electrons.
Isotopes: Are atoms of the same element with the same number of protons
and electrons but different number of neutrons. The chemical properties of an
element depend on the number of electrons in the outer electron shell. As
isotopes of the same element have the same number of electrons, they have
the same chemical properties.
3
,Determining Percentage Abundance:
Given the atomic mass of an isotope mixture, we can find the percentage
abundance of one of the several isotopes by using the formula below.
𝐴𝑡𝑜𝑚𝑖𝑐 𝑚𝑎𝑠𝑠 𝑖𝑠𝑜𝑡𝑜𝑝𝑒 𝑚𝑖𝑥𝑡𝑢𝑟𝑒 = 𝑚1 × 𝑎𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒 𝑜𝑓 𝑚1 + 𝑚2 × 𝑎𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒 𝑜𝑓 𝑚2 + (…)
Where m1 and m2 are the atomic masses of each isotope.
If there are only 2 isotopes, then the formula is the following.
𝐴𝑡𝑜𝑚𝑖𝑐 𝑚𝑎𝑠𝑠 𝑖𝑠𝑜𝑡𝑜𝑝𝑒 𝑚𝑖𝑥𝑡𝑢𝑟𝑒 = 𝑚1 × 𝑎𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒 𝑜𝑓 𝑚1 + 𝑚2 × 𝑎𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒 𝑜𝑓 𝑚2
Which can finally be transformed into,
𝐴𝑡𝑜𝑚𝑖𝑐 𝑚𝑎𝑠𝑠 𝑖𝑠𝑜𝑡𝑜𝑝𝑒 𝑚𝑖𝑥𝑡𝑢𝑟𝑒 = 𝑚1 × 𝑎𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒 𝑜𝑓 𝑚1 + 𝑚2 × (1 − 𝑎𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒 𝑜𝑓 𝑚1 )
From here you calculate the abundance of the first isotope and finally the
abundance of the second isotope is determined by,
𝑎𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒 𝑜𝑓 𝑚2 = (1 − 𝑎𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒 𝑜𝑓 𝑚1 )
Determining Percentage Abundance Example:
To answer this question, we do the following:
4
, Chemical Bonding
First Ionization Energy: is the minimum energy necessary to remove an
electron from a neutral atom in the gas phase and form a positively charged
ion in the gas phase.
To achieve the ionization of X(g) to form the products X +(g) + e-, it is necessary
to supply energy to the neutral atom X(g). The energy added enables the
electron to escape from the potential energy well that holds it in the atom.
Hence, ionization energy is an endothermic process, as energy is required.
The ionization energy is a measure of the stability of the free atom. Those
atoms with larger ionization energies are more stable than those atoms with
smaller ionization energies because their electrons must be removed from
deeper potential energy wells.
In the periodic table, the values for the ionization energies generally increase
moving across a period (horizontal rows) from left to right as the nuclear
charge increases and is more difficult to remove an outer electron. However,
when we move down a group the nuclear charge also increases, but the effect
of the increase in size and the shielding of the other electrons is greater,
meaning that is easier to remove an electron and hence the ionisation energy
decreases.
Second Ionization Energy: is the minimum energy necessary to remove a
second electron.
The third, fourth, and higher ionization energies are defined in an analogous
fashion. Successive ionization energies always increase due to the greater
electrostatic attraction of the electron to the product ions, which have
increasingly greater positive charges. This is because when you are removing a
third electron for example, there is two electrons less in the ion than when a
first electron was going to be removed from the neutral atom. However, the
number of protons stays the same. Therefore, the attraction force the electron
being removed experiences is the same, however the repulsion force it
experiences is much less as there are already two electrons less. Hence, it is
more difficult to remove the third electron in comparison to removing the first
5
,electron, meaning that more energy is required leading to the third ionisation
energy having a higher value than the second or first ionization energies.
The first electron affinity: Is the enthalpy change when 1 mole of electrons is
added to 1 mole of gaseous atoms to form 1 mole of gaseous 1- ions under
standard conditions.
E.g.
The first electron affinity is exothermic; therefore, it has a negative value.
The second electron affinity: Is the enthalpy change when 1 mole of electrons
is added to 1 mole of a gaseous 1- ions to form 1 mole of gaseous 2- ions under
standard conditions.
E.g.
The second electron affinity has always a positive value as its endothermic, this
is because the electron is added to an ion which is already negative therefore it
must overcome the repulsion. To overcome this repulsion, we need to supply
energy.
Chemical Bonding:
Ionic bonds: This form when 2 or more atoms bond to get a full outer shell. To
do this they share electrons. Ionic bonds happen between metals and non-
metals. Ionic bonding is the electrostatic attraction between positive and
negative ions in an ionic crystal lattice.
Lattice:
Covalent bonds: They form when atoms of non-metals combine together and
become simple molecules. A covalent bond involves the sharing of electron
pairs between atoms. This electron pair are known as bonding pairs. Covalent
bonds can be single, double or triple. Molecules are held together by weak
intermolecular forces.
6
,Polar Covalent Bonds: Are bonds with a partial transfer of charge.
Electronegativity: Is a measure of the tendency to attract electrons.
Highly electronegative atoms readily accept electrons and negative ions. Highly
electropositive atoms readily donate electrons and form positive ions.
Electronegativity generally increases as you move from left to right across a
period and decreases as you move down a group.
Example Simple Problem:
As we can see, the reaction of Na + Cl to form the sodium chloride salt, consists
of the first ionization energy, to produce the Na + ion and the first electron
affinity to produce the Cl- ion. However, as we can see the total energy change
is hence, 147kJ, meaning that energy is required for the formation of NaCl and
hence, NaCl is expected to dissociate into Na+ and Cl-.
The ionic bond is formed due to the electrostatic Coulomb force of attraction
between the ions, this explains why NaCl is formed, and it does not dissociate
back into Na+ and Cl-.
Potential Energy between Charges (Coulomb Attraction):
𝑄1𝑄2 1
𝛥𝐸𝑐 = ⋅
4𝜋𝜀0 𝑟 1000
Where 𝛥𝐸𝑐 is the potential energy in kJ, 𝑄1 𝑄2 are the charges in C, r is the
separation between them in m and 𝜀0 is known as the permittivity of free
space.
To the formula above, if we wanted to find the potential energy for one mole
of NaCl, we should introduce the value for the constants above, including the
charges which will be 1.602 ⋅ 10−19 and −1.602 ⋅ 10−19 . Finally, we would
have to multiply the formula above by Avogadro’s Number as we want to find
the potential energy for 1 mole.
7
,Bond Dissociation Energy: Is the energy you must add to fully separate the
atoms in the molecule or break the bond.
When we determine the potential energy and bond dissociation energy, we
are making an approximation. This is because we neglect several things:
1. Other repulsive or attractive interactions
2. We assume that a salt such as NaCl is purely ionic, and it can be partially
ionic and partially covalent.
3. We assume they are point charges, ignoring the polarization.
Properties of Diatomic Molecules:
As we go down the group, the bond length increases and the bond energy
decreases, as less energy is required to break the bonds because long bonds
are weak bonds.
Lewis Structure:
First count the total amount of valence electrons (VE) you have. Then, draw
single bonds with neighbouring atoms to get a starting structure. Then, count
the remaining valence electrons you have. Place the remaining electrons by
using the octet rule and seeing if there are missing bonds and lone pairs.
Finally, check the formal charge. It is the difference between the number of
valence electrons an atom has when it is not bonded to other atoms and the
number it “owns” when it is bonded.
8
,Tricks for Lewis Structures:
First, we determine the maximum number of electrons (ME) by multiplying the
number of atoms by 8 or 2 if its hydrogen. Then, the number of bonds can be
𝑀𝐸−𝑉𝐸
determined by, 𝑛 𝑏𝑜𝑛𝑑𝑠 = . Sometimes, too little/many bonds appear
2
as a result, then, in this case, we must simply use the minimum amount of
bonds necessary. (If we have three atoms, at least three bonds are needed).
𝑉𝐸
Also, the number of lone pairs can be determined by, 𝑛 𝑙𝑜𝑛𝑒 𝑝𝑎𝑖𝑟𝑠 = −
2
𝑛 𝑏𝑜𝑛𝑑𝑠.
Resonance Contributor: Is the approximate structure with localized electrons.
Resonance Hybrid: Is the actual structure with delocalized electrons.
Resonance Contributors don’t exist in reality, it’s just a way of explaining the
resonance hybrid.
Rules For Drawing Resonance Contributors:
9
, Example 1:
Example 2:
Predicted Stabilities of Resonance Contributors:
All resonance contributors do not necessarily contribute equally to the
resonance hybrid. The degree to which each one contributes depends on its
predicted stability. Because resonance contributors are not real, their
stabilities can’t be measured. The stabilities of resonance contributors must be
predicted based on molecular features found in real molecules. The greater
the predicted stability the more it contributes to the resonance hybrid.
This are the main features that decrease the predicted stability of a resonance
contributor (in order of importance):
1. An atom with an incomplete octet
2. A negative charge that is not on the most electronegative atom
3. A positive charge that is on an electronegative atom
4. Charge separation
10
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