Chapter 6

Electron Configurations for Ions • Main group elements: – Cations – lose electrons from highest energy occupied orbital – Anions – gain electrons in lowest energy unoccupied orbital (Aufbau). Chapter 6 • Transition metals (d block) – Form cations – Lose electrons from the outer s orbital before losing electrons from the d orbital 1 Octet Rule Ionic Bonding – Octet Rule Many elements react to attain a lower energy level (more stable). Many elements want to attain a noble gas configuration. H+ Atoms lose/gain electrons to form ionic compounds. Atoms share electrons to form covalent compounds. In both cases, the “outer shell” is fully filled similar to a noble gas electron configuration. Lose electrons to gain noble gas configuration Gain electrons to gain noble gas configuration Transition (d orbital) metals do not attain noble gas configuration Ionic Bonding – Octet Rule Covalent Bonding – Octet Rule Share electrons to form an octet in the outer (valence) energy level 1s 2s 2p 1s 2s [He] 2p [Ne] Li+ has the electron configuration of He F- has the electron configuration of Ne Outer energy level has 8 (or 2) electrons by sharing electrons Actually it’s really more complicated for covalent bonds, as we’ll see in chapter 7 Effective nuclear charge (Zeff) Effective nuclear charge (Zeff) • Nuclear charge is the charge in the nucleus. – Number for protons • Effective nuclear charge is what an electron “feels” from the nucleus – Electrons affect each other within atoms – Electrons can shield one another from the nucleus – Multielectron atoms feel a charge less than the full nuclear charge Atomic Radii Atomic Radius If quantum mechanics tells us we can never know the exact location of an electron, how do we know where an atom “ends”? What trends do we see? Decrease radius across row (increase in p+ and e-) Atomic radius: one-half the distance between two nuclei in two adjacent atoms Increase radius down column More than one specific definition of radius depending on the type of bonding involved. So, the same element can actually have different radii. The ones we use in class are an average. (increase in p+ and e-) Van der Waals radius Covalent Radius d-block removed 9 Trends in Ionic Radii Atomic Radii Increase in Zeff Why do we see two different trends? How are the trends in ionic radii different/same as atomic radii? Increase in energy levels d-block removed d-block removed 12 Atomic and Ionic Radii Ionic Radius Cations: The radii of metals decrease when they become ions Anions: –Outer electrons in lower energy level (closer to the nucleus) –Increase in Zeff for remaining outer electrons –Outer electrons pulled in closer The radii of non-metals increase when they become ions Within the cation/anions, the trends till holds. –Outer electrons are in the same energy level as before –Decrease in the Zeff for outer electrons because of sheilding by more electrons in the same energy level Between the two groups, however, the relative sizes have changed Ionization Energy Ionization Energy •What is the correlation? • Energy to remove valence electron from neutral atom in the gas phase •Why? • X(g) → X+(g) + e- Atomic Radii and First Ionization Energies of the Alkali Metal (Group 1) Elements Ionozation Energy (kJ/mol) 550 – Positive because is required to remove atom – More energy to remove an electron from a noble gas configuration Li 500 Na 450 K 400 Rb Cs 350 300 140 190 240 290 Atomic Radius (pm) 19 Trends in First Ionization Energy Anomalous IE1 Notice an • Which member of each pair has the larger first ionization energy? Why? •F or Cl •O or F • Na or Mg •K or Na •N or unexpected trend in 1st ionization energies across rows In most cases, the smaller the radius, the higher the 1st ionization energy O (this one might be a trick) Why an anomaly? 20 Anomalous first ionization energies (IE1) • Successive ionization energies take….more energy. Why? • And certain, ionization energies for certain atoms take a lot more energy. Why? N O Larger atom: Smaller atom 75 pm 73 pm Larger IE1: Smaller IE1: 1402 kJ/mol 1314 kJ/mol 2s 2p 2s 1st ionization of N Successive Ionization Energies Does not fit the expected trend Takes more energy to remove an electron from a half-filled or completely filled subshell 2p 1st ionization of O Electron Affinity 2p 3s 2p 1st Ionization of Na 3s 2nd Ionization of Na Takes much more energy to remove an electron from a noble gas configuration • Energy released when electron is added to neutral atom in the gas phase • X(g) + e- → X-(g) – Generally negative because energy is released 2p 3s 1st Ionization of Mg 2p 3s 2nd Ionization of Mg 2p 3s 3rd Ionization of Mg Electron Affinity – More energy released, higher affinity for electron Anomalous electron affinities N Trend is not as smooth as others. Many exceptions Be Ne Why is nitrogen positive? 2s Positive or small negative number 2p 2s 2p Adding an electron Adding an electron to a half-filled p to a filled s subshell subshell sometimes destabilizes the destabilizes the atom. Or, the atom. Or, the atoms is less atom is less stable stable than than predicted predicted 2p 3s Adding an electron to a noble gas configuration greatly destabilizes the atom Lattice Energy: – First, look at the charges (magnitude) (magnitude) • Charges = Energy – Then, look at radii • Radii = Energy 29 Relative Lattice Energies • Which compound in each pair has the highest lattice energy? – NaCl or MgCl2 – MgO or Na2O – NaCl or KCl – NaBr or NaCl 30