What are the trends for atomic radii? Why do these trends exist? 4.

1.9 – I can define and calculate the effective nuclear charge for an atom and explain how this impacts observed periodic trends – I can define atomic/ionic radius and explain how it relates to the effective nuclear charge. Furthermore, I can explain how this trend changes as you move throughout the Periodic Table and relate it to the elements quantum electron configuration.

The force felt by two charged objects is directly related to the size of the charge and inversely related to the distance between the charges. The force on the electron increases as the nuclear charge increases The force on the electron decreases as the nuclear charge decreases

Periodic trends depend on how well electrons “feel” the nucleus The effective nuclear charge (Z ef ) is the force felt between the nucleus and the electron of interest

In elements with many electrons, inner electrons shield the valence electrons. This occurs because the core electrons act to shield the valence electrons from the full positive charge of the nucleus.

Do you think the Z ef increases or decreases as we go down a group the Periodic Table? What about as we go across a period? Justify your answer with evidence and the idea of electron shielding.

The further away from the nucleus, the less the nuclear charge is “felt”, so Z ef decreases down a group. As you continue across a period, there are no more core electrons, but there is a stronger positive charge Therefore, as you go across a period Z ef increases.

Z ef decreases Z ef increases

What is the trend for atomic radii as you go throughout the Periodic Table? Justify your response

Atomic radii increase down a group and decrease across a period. They increase down a group because more orbitals are being added, making the element “bulkier” The decrease across a period because Z ef is increasing, thus causing the electron to be more attracted to the nucleus.

Removing electrons causes the radius to decrease Removing electrons causes orbitals to be removed and electron-electron repulsions to decrease Both of these lead to an increase in Z ef and a decrease in the ionic radius

Adding electrons causes the radius to increase Adding electrons causes an increase in electron – electron repulsions, so Z ef decreases. A decrease is Z ef causes the radius to be larger

Arrange Na +, K +, and K in order of increasing atomic radius

Arrange the ions K +, Cl -, Ca 2+, and S 2- in order of decreasing size.

Arrange the following atoms from smallest to largest: Rb +, Sr 2+, Y 3+

I am going to project up numerous statements You must decide whether you agree or disagree. Be prepared to defend your response!

1. Coulombs Law says that as you pull charges away from one another, the force increases. 2. The effective nuclear describes the attraction between the nucleus and the electron of interest. 3. Effective nuclear charge decreases down a group because there are more electrons. 4. Effective nuclear charge increases across a period because of the more positive nucleus. 5. The ionic radius is always larger than the atomic radius for a given atom.

5. Order the following from smallest to largest radius: Se 2-, Te 2-, Se - Answer: Se, Se 2-, Te 2 6. Order the following from largest to smallest radius: Co 3+, Fe 2+, Fe 3+ - Answer: Co 3+, Fe 2+, Fe The effective nuclear charge has littler overall impact on the atomic and ionic radius of an element

Pre – Lab for Lab 3 due at the start of lab on Thursday/Friday Final lab right up (results and conclusions) due Monday/Tuesday for Lab 2. To be on track, read sections 7.1 – 7.3 and answer the corresponding problems