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The Periodic Table and its Trends

The Periodic Table and its Trends. Effective Nuclear Charge ( Z eff ). e- are pulled to the nucleus because they are negative, and the nucleus is positive.

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The Periodic Table and its Trends

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  1. The Periodic Table and its Trends

  2. Effective Nuclear Charge (Zeff) • e- are pulled to the nucleus because they are negative, and the nucleus is positive. • In larger atoms the outer electrons are shielded from the positive nucleus because the core electrons block the “positive” charge of the nucleus. • Zeff is essentially the positive charge that a valence electron "sees".

  3. Effective Nuclear Charge Cont. • Z = nuclear charge • S = screening constant • S is usually close to the number of core electrons in the atom • Zeff = Z - S

  4. Trends in Zeff • As you go across a row of the PT, the Zeff increases. • WHY? • Because # of core electrons is the same, but the nuclear charge is increasing because of more protons (increasing atomic number) • As you go down a column of the PT, the Zeff only slightly increases. • We would calculate them to all be the same, however a larger nucleus is harder to shield

  5. Atomic Radii • There are two ways that scientists can look at the size of an atom: 1. The radius of an atom when it is not bonded (say in a gas phase) or 2. when it is bonded to something • Which do you think is a larger radius, nonbonding or bonding? • (Nonbonding) • WHY? • Nonbonding- the atoms don’t want to be together, so they will just bounce off each other. When they are bonded, they are attracted in some way, and therefore the radius will be smaller.

  6. Bonding Radii • What is most often used when talking about atom size • Found by looking at bond lengths between atoms. • Figure 7.7 page 261 shows the different atomic radii

  7. Trends in Atomic Radii • As you go down the rows, the atoms get ________ • bigger • WHY?? • Electrons move out to outer shells that are further away from the nucleus • As you go right across the rows, atomic radii get _________. • smaller • WHY? • the effective nuclear charge increases, pulling the e- in further. Even though e- are added too….they are still in the same shell

  8. Radii of IONS • Cations are ________ than their parent atom • Smaller • WHY? • Cations have given up electrons, and gotten rid of an electron shell • Anions are ________ than their parent atom • Larger • WHY? • More electons means more e-/e- repulsions, which results in the electrons spreading out more.

  9. Ionization Energy • The energy required to remove an electron from an atom. (to create a POSITVE ion) • A Lithium atom wants its electron to be removed, so that it has a full shell (Li+) • Flourine DOES NOT want an electron removed, it wants to gain another electron.

  10. 1st, 2nd, 3rdIonization Energies • The 1st Ionization energy is the energy required to remove the first electron from an atom, the 2nd ionization energy is the energy required to remove the second electron from an atom……etc. • Which is higher, the 1st or 2nd ionization energy for Lithium? • 2nd • The 1st electron wants to be removed, so it doesn’t take a lot of energy, but the 2nd is in a full outer shell (like a noble gas) so it doesn’t want to be removed.

  11. 1st Ionization Energy Trend • As you move ACROSS the rows, it become harder to pull off electrons (therefore the energy required to remove an electron is higher, ionization energy increases) • As you move DOWN the columns, it becomes easier to pull off electrons (they are further from the nucleus and are not held in as tightly)

  12. The orange color shows where a large change in ionization energy occurs. This large change occurs once the full octet (Noble Gas Configuration) has been achieved, because electrons really don’t want to leave that stable octet.

  13. Electron Affinity • The change in potential energy when an electron is added. • The ease at which an atom gains an electron • Electron affinities are thought of as negative (release of energy- exothermic process) • If there is a large energy change, then a stable ion is formed. If there is a small energy change, then the ion is not stable.

  14. Noble gases are not included in the trend. Halogens have the highest affinity. Noble gases don’t want any more electrons (their outer shells are full)

  15. Trends in Electron Affinity • Electron affinity increases across a row because the neutral atoms are more likely to gain an electron to have a full shell. • Electron affinity decrease only slightly down a column because the additional electron will be entering an orbital farther away from the nucleus, and thus would experience a lesser effective nuclear charge.

  16. Interesting trends in Electron Affinities • The electron affinity values are not what they are expected to be for atoms in group 15. • It seems as though this group doesn’t want that electron as much as one would think… • Why? • Because the p orbital is currently half filled….adding one more electron disrupts this partial stability. • Both half filled and fully filled orbitals create stability in the electronic structure of an atom

  17. Electronegativity • (very similar to electron affinity) • The tendency for an atom to attract an electron • Fluorine has the highest electronegativity • Trend is the same as electron affinity (and noble gases are excluded once again from the trend)

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