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A short class with Sarah and her students. The other ways are under More.
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Today, we'll discuss how atomic sizes change as we move across periods in the periodic table. Can anyone tell me what happens to atomic size as we go from left to right?
I think the atomic size decreases.
That's correct! As we move across a series, the atomic size typically decreases. This is due to the increase in effective nuclear charge that pulls electrons closer to the nucleus. A helpful way to remember this is to think of 'more protons, more pull!'
But why do the d-electrons not shield the other electrons effectively?
Great question! D-electrons have a poor shielding effect because they are in similar energy levels and do not effectively protect outer electrons from the nuclear charge. This is why we see that slight decrease in atomic size across a series.
So, does that mean that the atomic sizes of transition metals are generally smaller than those of s-block metals?
Yes, exactly! As we compare groups, transition metals tend to have smaller atomic sizes due to the reasons we've discussed.
Can you summarize that again?
Certainly! As we move across a period, the atomic size decreases due to increased nuclear charge and poor shielding by d-electrons. Remember the phrase 'more protons, more pull!' to help you. Any questions before we move on?
Now let's shift our focus to what happens when we move down a group in the periodic table. What trend do we expect in atomic size?
I think the atomic size increases down a group.
That's right! Atomic sizes typically increase as we move down a group. This is due to the addition of electron shells, which outweighs any increase in nuclear charge.
Does this trend apply to d-block elements as well?
Yes, it does! However, in the case of the 5d series, we have to consider the lanthanide contraction. Who can tell me what lanthanide contraction means?
Is it the fact that the atomic sizes don't increase as much due to the f-orbitals?
Exactly! The filling of the 4f orbitals leads to a strong effective nuclear charge that reduces the size of the elements in the 5d series. Think of it as a 'shrinkage' effect brought on by the d-orbitals being filled.
So, even though we're adding more electrons, why doesn't the size increase significantly?
Good question! The added electrons in the f-orbitals contribute to increased shielding, but the effective nuclear charge is still strong enough to keep the size smaller than expected. As a recap, we see an increase in atomic sizes down a group except in the case of 5d elements because of lanthanide contraction.
The distance from the nucleus to the outermost electrons, generally decreases across a period and increases down a group.
The size of ions compared to their neutral atoms, with potential increases or decreases based on the gain or loss of electrons.
The net positive charge experienced by outer electrons, increasing across a period.
The phenomenon that decreases the attraction between nucleus and valence electrons due to inner electrons.
The reduction of atomic and ionic size in 5d transition metal series due to strong nuclear pull from filled f-orbitals.
Across periods shorter we go, with protons pulling tightly in tow.
Imagine a group of kids holding hands. As they walk closer together, they get smaller and smaller — like atomic size across a period!
Remember 'BOSS' for Atomic Size: B - Bowers (decrease), O - Outer electrons, S - Shielding, S - Smaller.
Use 'SPEARS' for Ionic and Atomic Size influences: S - Size, P - Position, E - Effective nuclear charge, A - Anions/cations, R - Radius, S - Shielding.
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Atomic Size Trend
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Ionic Size of Cations
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Lanthanide Contraction
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The distance from the nucleus of an atom to the outermost shell of electrons.
The radius of an atom's ion, which can differ from the radius of its neutral atom.
The net positive charge experienced by electrons in an atom, accounting for shielding by inner electrons.
The phenomenon where the atomic and ionic radii decrease for the lanthanides due to the filling of 4f orbitals.
The reduction in effective nuclear charge experienced by outer electrons due to the presence of inner electron shells.
This section focuses on the trends in atomic and ionic sizes of the d-block elements, noting the slight decrease across periods and the increase down groups.
The atomic and ionic sizes of d-block elements exhibit distinct trends, showing a slight decrease across a series due to poor shielding by d-electrons and an increase down a group, influenced by lanthanide contraction in the 5d series.
In the d-block elements, atomic and ionic sizes exhibit trends that are essential for understanding their chemical behavior and properties. As we move across a series from left to right in the periodic table, there is a slight decrease in atomic and ionic sizes. This occurs because of the poor shielding effect provided by d-electrons, which leads to an increased effective nuclear charge felt by the outer electrons. Consequently, electrons are pulled closer to the nucleus, resulting in smaller atomic and ionic radii.
Conversely, when moving down a group in the periodic table, we observe an increase in atomic and ionic sizes. However, this increase is not uniform, particularly in the 5d series, where the phenomenon known as lanthanide contraction occurs. This contraction results in less than expected growth in size due to the strong effective nuclear charge caused by the filling of 4f orbitals, which counteracts the increase in size typically observed down a group.
Understanding these trends is crucial for predicting the properties and reactivities of transition metals, as size influences factors like electronegativity, ionization energy, and the ability to form complexes.