Periodic trends
Atomic Radii
Atomic Radii is affected by two main factors :
(1) Nuclear charge (number of protons) : The stronger the ‘pull’ the protons have to the electrons with electrostatic attraction, then the smaller the size of the atom radii (2) Number of energy levels: The greater the number of energy levels, the larger the atomic radii. The internal energy levels “shield” and reduce electrostatic attraction of the valence electrons to the protons. Addition of another electron does not result in a fractional decrease in the electrostatic attraction to any given electron, but it does increase the electron-electron repulsion, so an overall decrease in Net attractive force. [The most significant factor is nuclear charge – atoms are neutral and as the number of protons increases, the number of electrons increases at the same rate. The increased numbers of electrons in the valence shell, also have an increased number of equivalent protons, and they will be pulled in tighter and therefore show a smaller atomic radii.] |
Atomic Radii Trends
Across the period the Atomic radii decreases
e.g. Li → Ne in period 2 As the nuclear charge increases across the period, so does the electrostatic attraction and so outer electrons are pulled closer to nucleus. The electron repulsion is balanced by the nuclear charge attractions, and as the nuclear charge gets larger, so the electrons get closer together. The net attractive electrostatic attraction is increased as the nuclear charge increases. |
Down the group the Atomic radii increases
e.g. Li → Fr in group one
Electrons are being added to successive energy levels and both charge on nucleus and electron repulsion increase in step to “cancel each other out”. However, successive energy levels are further from the nucleus therefore, there is a subtle increase in atomic radii and an overall decrease in Net electrostatic attraction, with an increase the electron-electron repulsion.
e.g. Li → Fr in group one
Electrons are being added to successive energy levels and both charge on nucleus and electron repulsion increase in step to “cancel each other out”. However, successive energy levels are further from the nucleus therefore, there is a subtle increase in atomic radii and an overall decrease in Net electrostatic attraction, with an increase the electron-electron repulsion.
Ionic Radii
Cations (metal ions) are smaller in radii than their atoms
The outside energy level of electrons are removed but the nuclear charge (number of protons) remains the same creating smaller radii than the atom. There is an increase in Net electrostatic attraction. Anions (non-metal ions) are larger in radii than their atoms Extra electrons are added to the outside valance shell that have to be accommodated for and there is no change to the nuclear charge. Electron-electron repulsion spreads the electrons out further creating larger radii than the atom. There is a decrease in Net electrostatic attraction. |
Ionic Radii trends
Cations i.e. Na to Na+
The inter-electron repulsion experienced by the electron cloud of the cation is less than the neutral atom (less electrons than protons) , and since both species have the same amount of nuclear charge, the net electrostatic attractive force on the electron cloud in the cation is greater than the neutral atom resulting in a smaller cation size. |
Anions i.e. Cl to Cl-
The inter-electron repulsion experienced by the electron cloud of the cation is greater than the neutral atom (more electrons than protons), and since both species have the same amount of nuclear charge, the net electrostatic attractive force on the electron cloud in the cation is less than the neutral atom resulting in a larger anion size. |
1st ionisation energy
An Ionisation equation can be written as: (M representing an atom)
M → M+ + e- ∆H = + kJ as this is always an endothermic reaction
If the ionisation energy is high, that means it requires a lot of energy to remove the outermost electron. If the ionisation energy is low, that means it takes only a small amount of energy to remove the outermost electron.
Ionisation energy is affected by two factors:
(1) Nuclear charge: As NC increases, there is a stronger pull to the electrons by electrostatic attraction → increased 1st I.E. with increased NC
(2) Number of energy levels: Electrons in a lower energy level are much closer to the nucleus and thus have much stronger net electrostatic attraction to it. Electrons in a lower energy level shell have electron repulsion but are closer together. Electrons in higher energy level shells experience less net electrostatic attraction to the nucleus, as they are further away → decreased 1st I.E with increased number energy levels
M → M+ + e- ∆H = + kJ as this is always an endothermic reaction
If the ionisation energy is high, that means it requires a lot of energy to remove the outermost electron. If the ionisation energy is low, that means it takes only a small amount of energy to remove the outermost electron.
Ionisation energy is affected by two factors:
(1) Nuclear charge: As NC increases, there is a stronger pull to the electrons by electrostatic attraction → increased 1st I.E. with increased NC
(2) Number of energy levels: Electrons in a lower energy level are much closer to the nucleus and thus have much stronger net electrostatic attraction to it. Electrons in a lower energy level shell have electron repulsion but are closer together. Electrons in higher energy level shells experience less net electrostatic attraction to the nucleus, as they are further away → decreased 1st I.E with increased number energy levels
1st ionisation energy Analogy
In order to remove an electron from an atom you need to overcome the nuclear attraction of its protons holding it around the nucleus. This can be shown by the mud that the car is stuck in. The more mud (nuclear attraction) the more energy to remove the car (electron). However, other electrons in the atom are repelling the electron to be removed – so the more people pushing the car (electron repulsion) the easier it is to extract the car (electron)
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1st ionisation energy trends
Across a period the 1st ionisation energy increases
As the nuclear charge increases, the attraction between the nucleus and the electrons increases and it requires more energy to remove an electron from the outermost energy level and that means there is a higher ionisation energy. As you go across the Periodic Table, nuclear charge is the most important consideration. Therefore, going across the Periodic Table, there will be a trend of increasing 1st ionisation energy, because of the increasing nuclear charge.
As the nuclear charge increases, the attraction between the nucleus and the electrons increases and it requires more energy to remove an electron from the outermost energy level and that means there is a higher ionisation energy. As you go across the Periodic Table, nuclear charge is the most important consideration. Therefore, going across the Periodic Table, there will be a trend of increasing 1st ionisation energy, because of the increasing nuclear charge.
Down a group the 1st ionisation energy decreases
Going down a group in the Periodic Table, the effect of increased nuclear charge is weighed against the effect of increased electron repulsion, and the number of energy levels becomes the predominant factor. With more energy levels, the outermost electrons (the valence electrons) are further from the nucleus and are not so strongly attracted to the nucleus, and therefore, there is a reduction in net electrostatic attraction. Thus, the ionisation energy of the elements decreases as you go down a group in the Periodic Table, because it is easier to remove the electrons. The more stable elements have higher ionisation energies. |
Electronegativity
Electronegativity is the tendency of an atom to attract bonding electrons from another atom. Higher electronegativity values mean a higher tendency to attract electrons. Atoms with high E.N. are strong oxidants (gain electrons).
Electronegativity is affected by two factors:
1. Nuclear charge: As an atom's nuclear charge increases, there is a stronger pull on electrons of another atom by electrostatic attraction.
2. Number of energy levels: the more energy levels an atom has, the lower the net electrostatic attraction and the radii of the atom is larger. Because this then creates a bigger distance between ‘neighbouring’ atoms, electrons from other atoms experience less electrostatic attraction to the nucleus of another atom. Therefore, an atom in the same group has less electronegativity than an atom above it with less energy levels( Even though it has more nuclear charge).
Electronegativity is affected by two factors:
1. Nuclear charge: As an atom's nuclear charge increases, there is a stronger pull on electrons of another atom by electrostatic attraction.
2. Number of energy levels: the more energy levels an atom has, the lower the net electrostatic attraction and the radii of the atom is larger. Because this then creates a bigger distance between ‘neighbouring’ atoms, electrons from other atoms experience less electrostatic attraction to the nucleus of another atom. Therefore, an atom in the same group has less electronegativity than an atom above it with less energy levels( Even though it has more nuclear charge).
Ionic – covalent bond continuum due to electronegativity
Bond types between atoms can depend on the electronegativity of the atoms. Rather than discrete categories, molecules fall along a continuum
If there is little difference in electronegativity between two atoms then they tend to form a covalent bond with no polarity difference. A greater electronegativity difference creates a polar bond with uneven “sharing” of valance electrons.
If the electronegativity is even greater then there will be a complete transfer of electron from one atom (Metal) to another atom (non-metal) and ions will form that are held together with an ionic bond.
If the electronegativity is even greater then there will be a complete transfer of electron from one atom (Metal) to another atom (non-metal) and ions will form that are held together with an ionic bond.
Electronegativity trends
Across a Period the electronegativity increases
e.g. Li → Ne: The atoms have increased “pulling power” as the nuclear charge is increasing. Electrons are held tighter to the nucleus and there is a greater net electrostatic attraction. This allows another atom to be closer and it has a stronger attraction to electrons from that atom, so electronegativity increases.
e.g. Li → Ne: The atoms have increased “pulling power” as the nuclear charge is increasing. Electrons are held tighter to the nucleus and there is a greater net electrostatic attraction. This allows another atom to be closer and it has a stronger attraction to electrons from that atom, so electronegativity increases.
Down a group the electronegativity decreases
e.g. Li → Fr Both nuclear attraction and electron repulsion increase in step, but with an overall decrease in Net electrostatic attraction. However, as successive shells increase atomic radii, then the electrostatic attraction of the nucleus to other atoms’ electrons decreases, so atoms have less electronegativity as you move down a group.
e.g. Li → Fr Both nuclear attraction and electron repulsion increase in step, but with an overall decrease in Net electrostatic attraction. However, as successive shells increase atomic radii, then the electrostatic attraction of the nucleus to other atoms’ electrons decreases, so atoms have less electronegativity as you move down a group.