effective nuclear charge (z eff )
DESCRIPTION
Effective Nuclear Charge (Z eff ) In a many-electron atom, each electron is attracted to the positively charged nucleus and repelled by the other negatively charged electrons. - PowerPoint PPT PresentationTRANSCRIPT
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Effective Nuclear Charge (Zeff)– In a many-electron atom, each electron is attracted to
the positively charged nucleus and repelled by the other negatively charged electrons.
– Zeff takes both of these factors into account and represents an estimate of the net electric field experienced by an electron.
– Zeff = Z – S – Z = number of protons in the nucleus.– S = screening constant; the number of core electrons
that screen the outer electrons from the positive charge in the nucleus.
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• Estimate the effective nuclear charge experienced by the outer electron in each of the following atoms:
– Potassium
– Fluorine
– Silicon
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Effective nuclear charge
Increases (b/c number of protons increases)
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Atomic Radius – represents one half the distance between the nuclei of atoms held together by a bond.
Decreases
Incr
ease
s
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As we move from left to right across a period, the atomic radius decreases.
Why does this occur?
• as we move from left to right, atomic number increases.
• Since there are more protons in the nucleus, the effective nuclear charge increases.
• This increases the attractive force felt by the outer e- making the atom smaller.
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As we move from the top of a group to the bottom of a group, the atomic radius increases.
Why does this occur?
• as we move down a group, the principal quantum number (n) increases.
• This means that the outer e- are more likely to be found in higher energy orbitals further from the nucleus.
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Ionic Radii• When atoms gain an electron(s), they become
negatively charged anions. Anions are larger than the corresponding neutral atom.
X– will be larger than X because:
• Negatively charged ions have more electrons in the electron cloud.
• Each electron repels the other negatively charged electrons, causing the electron cloud to “spread out”.
• More electron-electron repulsions leads to a larger radius than a neutral atom would have.
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Ionic Radii• When atoms lose an electron(s), they become
positively charged cations. Cations are smaller than the corresponding neutral atom.
M+ will be smaller than M because:
• Positively charged ions have less electrons in the electron cloud.
• Less electrons means there will be fewer electron-electron repulsions, making the ion smaller than a neutral atom would be.
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Ionization Energy – the energy required to remove an electron from the ground state of a gaseous atom or ion.
Increases
Dec
reas
es
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As we move from left to right across a period, the ionization energy increases.
Why does this occur?• as we move from left to right, atomic number
increases. • Since there are more protons in the nucleus, the
effective nuclear charge increases. • This increases the attractive force felt by the
outer e- making it harder to remove.
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As we move from the top of a group to the bottom of a group, the ionization energy decreases.
Why does this occur?
• as we move down a group, the principal quantum number (n) increases.
• This means that the outer e- are more likely to be found in higher energy orbitals further from the nucleus.
• This makes them less attracted to the nucleus so they take less energy to remove.
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Successive Ionization Energies• Removing a second (or third, etc.) electron
always requires even more energy than removing the first electron.
• This occurs because the atom gets smaller each time an electron is removed and this makes the outer electrons closer to the nucleus (and more strongly attracted to it).
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The ionization energies for aluminum are summarized in the table below.
Why is there such a large jump in the amount of energy required for I4?
• Once aluminum has lost 3e- it is Al3+ and has an octet of 8 valence electrons.
• According to the octet rule this is the most stable state for an atom, and removing one of it’s octet of valence electrons would require a very large amount of energy.
I1 I2 I3 I4
578 kJ 1817 kJ 2,745 kJ 11,577 kJ
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Electron Affinity – the energy change that occurs when an electron is added to a gaseous atom.
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• A negative affinity indicates a high attraction for an electron (energy is released)
• A positive affinity indicates that an electron is not likely to be gained by the atom.
• In general, electron affinity increases moving from left to right across a period…however it shows a lot more variation than the other trends.
• Therefore, it is more helpful to look at the trends for different groups of elements.
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Group Trends for Electron Affinity• Halogens have highly negative electron affinities
because they have the greatest attraction for electrons since they only need to gain one electron to reach a full outer shell.
• Noble Gases have positive electron affinities because they already have an octet of valence electrons and would have to add an electron to an unoccupied higher energy level.
• Group 5A sees a decrease in affinity because the p-sub-shell is half-full and adding another electron will result in increased electron-electron repulsions in the atom.
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Electronegativity – the ability of an atom in a molecule to attract electrons to itself.
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• Electronegativity tends to increase as we move from left to right across a period and tends to decrease as we move down a group.
• The most electronegative elements have highly negative electron affinities (attraction for additional electrons) and high ionization energies (tendency to hold onto their won electrons).