Periodic Properties of the Elements Chapter 7 Periodic Properties of the Elements

Electron Configuration exceptions Energy cost to sticking two electrons in the same orbital Observation: sublevel with all half-filled orbitals energetically more stable Observation: sublevel with completely filled orbitals energetically more stable

Exceptions to electron configurations Observed: Cr: [Ar]4s13d5 Cu: [Ar]4s13d10 Only see these in transition metals Less energy difference between 4s and 3d

Exceptions to electron configurations Transition metals, when forming ions, lose s electrons first, NOT d!! Zn = [Ar]4s23d10 Zn+2 = [Ar]3d10 Why?

Charge interactions in an atom The closer + and – get, the stronger their attraction More + = stronger attraction for – Implications: behavior of electrons on outside of atom will depend on How close those electrons are, relatively, to the nucleus How strong a pull those electrons experience from the nucleus

Effective Nuclear Charge Effective nuclear charge: number of p+ in nucleus (Z) minus average number of e-(S) between nucleus and e- in question Zeff = Z – S Positive charge acting on outer e- is less than positive charge acting on closer e- As distance from nucleus increases, S increases and Zeff decreases The increase in electron shells (PELs) shields the nucleus, thus decreasing the attraction between the nucleus and the valence electrons

How orbitals differ More positive the nucleus, smaller the orbital Na 1s orbital = same shape as H 1s orbital, but is smaller b/c e- more strongly attracted to nucleus! “Greater effective nuclear charge!”

Atomic Size Have to measure distance between nuclei in diatomic – why? What pattern do you see?

Atomic Size (Trends) Within a group, atomic radius increases from top to bottom Adding energy levels of electrons, electrons pulling further away from nucleus Within a period, atomic radius decreases from left to right Greater effective nuclear charge within the same energy level – orbitals are shrinking in comparison to one another Stronger nucleus is pulling the electrons in closer

Ionic Radius For metals: For nonmetals: Positively charged ions (cations) formed when an atom of a metal loses one or more electrons Smaller than the corresponding atom Ex: Na+ is smaller than Na For nonmetals: Negatively charged ions (anions) formed when an atom of a nonmetal gains one or more electrons Larger than the corresponding atom Ex: Cl- is larger than Cl

1. Arrange the following atoms in order of increasing atomic radius: Na, Be, Mg. Practice Exercise Answer: Be < Mg < Na 2. Arrange these atoms and ions in order of decreasing size: Mg2+, Ca2+, and Ca. Answer: Ca > Ca+2 > Mg+2 3. Arrange the ions K+, Cl–, Ca2+, and S2– in order of decreasing size. Answer: S2– > Cl– > K+ > Ca2+.

Ionization Energy Energy required to remove an electron from a gaseous atom Highest energy electron removed first Furthest away First IE (I1) = removing first e- Second IE (I2) = removing second e- etc. etc.

IE for Mg I1 = 735 kJ/mole I2 = 1445 kJ/mole I3 = 7730 kJ/mole Effective nuclear charge increases as you remove electrons – increase proton to electron ratio Takes much more energy to remove core e- than valence e- b/c core e- feeling more pull from nucleus (greater effective nuclear charge)

IE

Explain this trend in IE For Element X – how many valence electrons? I1 = 580 kJ/mole I2 = 1815 kJ/mole I3 = 2740 kJ/mole I4 = 11,600 kJ/mole I5 = 16,430 kJ/mol

IE Trends Across a Period: Generally from left to right, I1 increases because there is a greater nuclear charge within the same energy level Down a Group: As you go down a group, I1 decreases because electrons are farther away from nucleus

IE Trend

BUT WAIT! Some exceptions! Explain the exceptions: Half filled and completely filled sublevels are harder to remove electrons from!

IE Examples Which will have the greater third ionization energy, Ca or S? Answer: Ca Arrange the following atoms in order of increasing first ionization energy: Ne, Na, P, Ar, K. Answer: K < Na < P < Ar < Ne

IE and PES PES is not in your book, pay careful attention to the notes!!!! Photoelectric effect Energy of an ejected electron: Etotal = KE(kinetic energy of e) + IE(ionization energy of e)

What happens when an electron is ejected? Vacuum level Vacuum level - I.E. Bound State Bound State + Metal (initial) - M+ K.E. Vacuum level I.E. Bound State + M+

Photoelectron Spectroscopy (PES) Machine irradiates metal with energy UV or X-ray Measures the kinetic energy of electrons when they “pop off” By measuring the total energy irradiating a metal (E) and the kinetic energy of the ejected electrons (KE), ionization energy can be determined for ANY e- (not just outermost!) Etotal = IE + KE Like mass spec, peaks in PES data

How to read PES data: B atom Energy (MJ/mol) Relative number of electrons 70 65 25 20 15 10 5 3 peaks = removing e- didn’t happen all at once Implies e- in 3 locations X-axis: energy required to remove the electrons often inverted (highest energy on left) BUT NOT NECESSARILY! ALWAYS CHECK AXIS! Y-axis: height of peak corresponds with ratio of electrons not precisely saying nof. e-, but gives ratio

How to read PES data: B atom Energy (MJ/mol) Relative number of electrons 70 65 25 20 15 10 5 Look at ratio of height of peaks on x axis 2:2:1 Which electrons hardest to remove? What does this mean? Hardest to remove = closest to nucleus

How to read PES data: B atom AND relative heights tell you 2e- in 1s, 2 e- in 2s, 1e- in 2p Relative number of electrons 1s22s22p1 70 65 25 20 15 10 5 Energy (MJ/mol) 2p sublevel (easiest to remove) 1s sublevel (hardest to remove) 2s sublevel (a bit easier to remove)

How to read PES data: B atom Relative number of electrons 70 65 25 20 15 10 5 Energy (MJ/mol) Therefore, it is Boron 5 electrons 1s22s22p1

Photoelectron Spectra http://www.chem.arizona.edu/chemt/Flash/photoelectron.html https://www.youtube.com/watch?v=NRIqXeY1R_I

Why PES? Look at what graph indicates: not all electrons in same energy level have same ionization energy What conclusions can be drawn from this? Sublevels! PES = evidence for quantum mechanical model of the atom Evidence for s,p,d,f sublevel

PES Example 4 peaks = 4 locations of electrons = 4 energy sublevels Relative number of electrons Energy (MJ/mol) 127 125 9 7 2 4 peaks = 4 locations of electrons = 4 energy sublevels Relative height of peaks tells you how many e- per sublevel Peak with highest energy = 1s, then 2s, 2p, 3s Which element is this? Mg: 1s22s22p63s2

PES Draw what you would expect the PES data for oxygen to look like Energy (MJ/mol) Relative number of electrons

Electronegativity and Electron Affinity Electronegativity: The ability of an atom to attract an electron to itself Scale of 0.0 – 4.0 Electron affinity: energy change associated when an electron is added to a gaseous atom Measured in kJ/mol The more negative, the greater the attraction of the atom for an electron If EA > 0, REALLY does not want an e-, negative ion higher in energy than separated atom and electron

Electronegativity: Trend EN increases left to right in a period Decreases from top to bottom F is most EN element on periodic table EA is different – does not change down a group, but does increase (gets more negative) across the p-block

IE vs. EA – the final showdown IE: energy required to take an electron AWAY! EA: energy associated with gaining an electron!

Metals vs. Nonmetals

Alkali Metals Group 1 As you move down the group Oxidize easily Decrease in IE Decrease in mp Increase in metallic character Oxidize easily React vigorously with water to produce metal hydroxide and hydrogen gas, highly exothermic Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)

Alkaline Earth Metals Group 2 Reactive, but not as reactive as group 1 metals

Halogens Group 17 (or 7A) High EA High IE High eneg F is the most reactive element on PT – high eneg, removes electrons from pretty much anything React with H2 to make acids H2(g) + X2 → 2HX(g)

Noble Gases Group 18 (8A) Stable, monatomic Completely filled s and p sublevels Were called “inert gases” until 1962 Have reacted under extreme circumstances Compounds of all noble gases have been created He was the last one – HeF2 in 1998 – why the last? Why F

Allotropes Different forms of the same element in the same physical state Oxygen: O2 (oxygen gas) and O3 (ozone) Carbon: Diamond and Graphite Sulfur: Have different structures and different physical and chemical properties