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Atomic Structure – Notes

2.1 Isotopes

(a) Isotopes are atoms of the same element which contain the same number of protons and electrons but different number of neutrons.

(b) Isotopes have the same number of electrons and therefore the same chemical properties.

(c) However, since they have different number of neutrons, they will have different masses and hence different physical properties such as melting point, density etc.

(d) They are found in all naturally occurring elements except F, Na, Al, P, Mn, Co, As.

Why do isotopes react similarly?
This is because in a chemical reaction, it is the electrons that are transferred between different atoms; atoms either gain, lose or share electrons

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Atomic Structure – Notes

Approach of writing electronic configurations

1. First, write down the electronic configurations for the neutral atom.

2. Remember to add electrons to the 4s orbitals before the 3d orbitals.

3. Remove or add the relevant number of electrons to the electronic configuration of the neutral atom to get that of the cation or anion
respectively.

4. Remember to remove electrons from the orbitals with the highest energy.Thus remove electrons from the 4s orbitals before the 3d orbitals.

Some Definitions

Isoelectronic: species containing the same number of electrons

Isotopic: species containing the same number of protons

Isotonic: species containing the same number of neutrons

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IONISATION ENERGIES

• Ionisation involves the removal of electron(s), forming a cation.

• When the electrons in the atom occupy the lowest energy levels, the atom is said to be in its
ground state.

• When an electron is promoted to one of the higher energy levels, the atom is unstable and is said to be in an excited state.

(a) M (g)—> M+ (g)+e– H = 1st I.E.

First ionisation energy (1st I.E.) is the energy required to remove 1 mole of electrons from one mole of gaseous atoms in the ground state to form one mole of gaseous singly- charged cations.

(b) M+ (g) —> M2+ (g)+e– H = 2nd I.E

Second ionisation energy (2nd I.E.) is the energy required to remove 1 mole of electrons from one mole of gaseous singly-charged cations to form one mole of gaseous doubly- charged cations.

• Ionisation energies affect the type of bond formed by the atom with other atoms. Elements with low ionisation energies will find it easy to lose an electron to form a cation, resulting in ionic bonds being formed.

• Ionisation energies are positive values (i.e. endothermic) since energy is absorbed during ionisation to overcome the attraction between electron and nucleus.

• The 2nd I.E. > 1st I.E. because more energy is required to remove an electron from a positive ion (compared to a neutral atom) due to greater electrostatic attraction between the positive ion and the valence electron.

Factors influencing ionisation energies

Ionisation energy (I.E.) of an atom is influenced mainly by two factors:

Effective nuclear charge, Zeff
• An electron in the atom faces two main electrostatic forces, namely the attractive force by the nucleus (nuclear charge) and the repulsive force by electrons closer than itself to the
nucleus (shielding effect).

• Effective nuclear charge is the combined effect of nuclear charge, Z and shielding effect, S, caused by inner electrons:

Zeff = nuclear charge (Z)  shielding effect (S)

• Higher Zeff ⇒ stronger forces of attraction between nucleus and valence electron,
⇒ higher ionisation energy

(a) Size of (positive) nuclear charge, Z
• indicates the electrostatic forces of attraction between the protons in the nucleus and the valence electrons
• nuclear charge increases with an increase in proton number
• stronger attraction between the positive nucleus and valence electrons
• more energy is required to remove the valence electron

(b) Shielding (or screening) effect of the inner electrons, S
• shielding of the valence electrons from the electrostatic attraction of the positively charged nucleus by the inner electrons
• electrons in the same subshell offer poor shielding for one another
• shielding effect increases with an increase in the number of inner electrons
• weaker attraction between the positive nucleus and valence electrons
• less energy is required to remove the valence electrons

Distance of the valence electron from the nucleus (i.e. the size/radius of the atom)
• attraction of the positive nucleus for the valence electron decreases as distance of electron from nucleus increases
• ionisation energy decreases as n increases

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Trends in First Ionisation Energies

(A) Across a period of the Periodic Table, the ionisation energies generally increase.

Reason:
• Nuclear charge increases as the proton number increases.
• Successive members of the period has one electron added to the same outermost shell, thus, the increase in screening effect is negligible.
• Effective nuclear charge increases.
• More energy is required to remove the more tightly held electrons, hence ionisation energy generally increases.

(B) Down a group of the Periodic Table, the ionisation energies generally decrease.

Reason:
• Nuclear charge increases as proton number increases.
• Electrons are added to a higher principal quantum shell which is further away from nucleus.
• Weaker electrostatic forces of attraction between the nucleus and valence electrons.
• Less energy is required to remove the valence electrons, hence ionisation energy generally decreases

(C)TRENDS IN THE ATOMIC RADIUS OF ELEMENTS

Across a period,
• nuclear charge increases
• shielding effect is relatively constant (number of inner core electrons is the same across a period)
• effective nuclear charge increases, resulting in electrons being pulled closer towards the nucleus
• atomic radius generally decreases.

Down a group,
• nuclear charge increases
• electrons are added to a higher principal quantum shall
• valence electrons are increasingly further away from the nucleus
• weaker electrostatic forces of attraction between valence electrons and nucleus
• atomic radius generally increases.

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2015 A level Chemistry Tuition Schedule

Level	Standard	Day	Time
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Chemical Bonding

Covalent Bond Strength
The strength of a covalent bond is measured by its bond energy (also called bond enthalpy) which is defined as the average energy absorbed when one mole of a particular bond is broken in the gaseous state.

The strength of a covalent bond is affected by the following factors:

(a) The effectiveness of overlap of the orbitals
In general, larger orbitals are more diffuse so that overlap is less effective in bonds formed by larger atoms compared to smaller atoms. For example the bond energies of chlorine to iodine are shown below.

BE(Cl—Cl): 244 kJ mol–1 BE(Br—Br): 193 kJ mol–1 BE(I—I):151 kJ mol–1

As the halogen increases in size, the valence orbital used in bonding is more diffuse so that the overlap of the orbitals is less effective from chlorine to iodine and bond energy decreases from chlorine to iodine.

Stop and Think

Question: Explain which bond is stronger, C—H or Si—H.
Answer: C—H bond is stronger since C is smaller than Si so that valence
orbital of C is less diffuse and overlap of its valence orbital with that
of H is more effective.

(b) The differences in electronegativities of the bonding atoms (bond polarity)

Electronegativity of an element measures the relative tendency of its atom to
attract the shared electron–pair in a covalent bond. Based on Pauling’s definition, fluorine, the most electronegative element, is given an arbitrary value of 4.0 and all values of the other elements are relative to it. The higher the value, the stronger the attraction.

Electronegativity decrease down a group.
Electronegativity increase­ across a period.

A polar covalent bond results if the bonded atoms have different electronegativities.

Partial charges (d+ and d–) arise on the two bonded atoms. The covalent bond is
described as possessing some ionic character.

For example in HF, F is more electronegative than H and hence attracts the
bonding electrons more strongly. The electron density distribution of the H—F
bond is asymmetrical and the H—F bond is polar, with F having a d– charge while
H having a d+ charge:

In addition to the existing covalent bond, there is now an increase in electrostatic attraction due to the two partial charges, which leads to increased bond strength.

In general, the greater the difference in electronegativity, the more polar is the covalent bond (i.e. greater the bond polarity) and stronger the covalent bond.

(c) Number of bonds between atoms (Single vs double vs triple bonds)

For the same bonding atoms, an increase in the number of bonds increases the
number of shared electrons between the two atoms i.e. there is increased
electrostatic attraction between the bond pairs and the two nuclei, hence bond
strength is increased.

Hence the strength of triple bond > double bond > single bond.

Bond Length
(i) The covalent bond length is the distance between the nuclei of the two
atoms in the bond.
(ii) Generally, the stronger the covalent bond, the shorter is the bond length.
(iii) The three factors of more effective overlap by less diffused orbitals, greater polarity of bond and increase in number of bonds between 2 atoms decrease bond length and increase bond strength.

A molecule with strong covalent bonds generally has less tendency to undergo
chemical change than does one with weak bonds. This is seen in nitrogen with its very large NºN bond energy (944 kJ mol–1). A very large amount of energy must be supplied to nitrogen to break the triple bond before nitrogen can react with other elements. Fluorine having a weak F–F bond (due to lone pair-lone pair repulsion in a very short bond) is very reactive.

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