1) Configurations of ions 2) Trends in atom size (atomic radius) 3) Trends in ion size 4) Ionization Potential

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1 Chem 105 Friday 6 Nov ) Configurations of ions 2) Trends in atom size (atomic radius) 3) Trends in ion size 4) Ionization Potential 11/6/2009 1

2 Atomic Radius Measured in picometers (pm) 1 pm = m or Angstroms (Å) 1 Å = 100 pm = 10 8 cm Generally increase going down a group (down a column) and decrease going across a period (LtoR in a row) 11/6/2009 2

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4 Group 1 Alkali metals Group 8A Noble Gases 11/6/2009 4

5 One issue that arises with the term Atomic Radius is that the numbers may differ depending on how they are obtained! covalent radius = half distance between bonded atoms or calculated radius = distance out to arbitrary electron density based on quantum mechanics calculation (Schrödinger equation) or experimental based on crystal of metal atoms = ½ interatomic distance 11/6/2009 5

6 Comparing the electron distribution in a H atom vs. H 2 molecule The e /Å 3 contour The 0.01 e /Å 3 contour We define the calculated atomic radius = distance from nucleus out to electron density ~ e /Å 3 11/6/2009 6

7 Comparing the electron distribution in a H atom vs. H 2 molecule H atom H atom radius = 44 pm H covalent radius = 37 pm H 2 molecule HH dist = 74 pm Generally the covalent radius issmaller than the radius of free atombecause the electrons in the 11/6/ molecule are attracted by two or more nuclei. This shrinksthe whole electron cloud a bit.

8 Crystal structure (experimental) of metallic sodium. NaNa distance = 365 pm; so, Na radius = 365/2 = 183 pm 365 pm 11/6/2009 8

9 Group 1 Alkali metals Sodium: 184 pm 11/6/2009 9

10 Going from elementtoelement DOWN a group, you add a complete shell of electrons plus the same number of protons in the nucleus. For example, Group 2: Be, Mg, Ca e in 1s orbital e in 2s orbital 2e in 3s orbital 4+ 2e Berylium atom 12+ 2e 8e Magnesium atom 8e in 2s,2p orbitals Although the nuclear charge increases by 8+, adding a complete inner shell of 8 electrons shields the outer shell electrons from the increased positive charge. 11/6/

11 However, atoms get SMALLER going ACROSS a row LefttoRight. In this case, electrons are added to the same shell on the periphery of the atom, and the # of innershell electrons is constant. The outershell electrons DO NOT shield each otherfrom the increasing nuclear charge because they are spread out withapproximately same average distance from the nucleus. The nuclear charge increases by +1 for each electron added, and this added proton has a much larger effect on all the electrons compared to the effectof the added electron e 8e 13+ 2e 8e 14+ 2e 8e Magnesium atom radius = 145 pm Aluminum atom radius = 118 pm Silicon atom radius = 111 pm 11/6/

12 Transition metals decrease, then increase slightly at end of the series (Cu, Zn) 11/6/

13 21+ 2e 8e 8e 26+ 2e 8e 8e 30+ 2e 8e 8e Scandium atom radius = 144 pm Iron atom radius = 117 pm Zinc atom radius = 125 pm Decreases due to the increasing nuclear charge which is not shielded by outer electrons. Increase (or constant depending on which measurement). Now there are SO MANY ELECTRONS in outer shell they expand due to mutual repulsion. 11/6/ (covalent radii, Inorganic Chemistry, Miessler and Tarr, p 42)

14 Sizes of Ions Cations(remember ca + ion) always SMALLER than corresponding atom (you re removing electrons usually a whole shell without changing the nuclear charge) Anions Always LARGER than corresponding atom (you re adding electrons to complete a shell usually without changing the nuclear charge.) 11/6/

15 12+ 2e 8e 12+ 2e 8e Magnesium atom radius = 145 pm Magnesium 2 + ion radius = 72 pm Huge shrinkage you re stripping away the whole outer shell 17+ 2e 8e 17+ 2e 8e Outer shell expands (a lot) because you re adding an e without adding a proton in the nucleus. Chlorine atom radius = 99 pm Chloride ion (Cl) radius = 181 pm 11/6/

16 Atom and Common Anion Size Comparison These 3 anions have 10 e radii in picometers pm 11/6/

17 Cations radii in picometers pm 11/6/

18 Place the following atoms in order of increasing atomic radii: K, Mg, Ca, Rb K < Mg < Ca < Rb 2. K < Mg < Rb < Ca 3. Mg < Ca < K < Rb 4. K < Rb < Mg < Ca 5. Mg < K < Ca < Rb N = 75 1% 1% 67% 9% 21% K Mg Ca K < Mg < Ca < Rb K < Mg < Rb < Ca Mg < Ca < K < Rb K < Rb < Mg < Ca Mg < K < Ca < Rb Rb Excellent work, chem students! JK 11/6/

19 (OWL question on ion sizes) 11/6/

20 11/6/

21 Ionization energy (= ionization potential) definition A (g) > A + (g) + 1e E = 1 st ionization energy (kj/mol) A + (g) > A 2+ (g) + 1e E = 2 nd ionization energy (kj/mol) 11/6/

22 First Ionization Energies H(g) H + (g) + e ΔH = kj (For comparison, the thermite reaction gives off way less energy per mole of iron oxide consumed. Fe 2 O 3 (s) + 2 Al(s) Al 2 O 3 (s) + 2 Fe(l) ΔH = kj) He(g) He + (g) + e ΔH = kj The first ionization energy for helium is about twice the ionization energy for hydrogen because each electron in helium feels the attractive force of two protons, instead of one. Far less energy is required to remove an electron from a lithium atom, which has three protons in its nucleus. Li(g) Li + (g) + e ΔH = kj 11/6/

23 He The 1 st ionization energy (A > A + + e ) decreases going down a group. Ne Group 8A Ar Kr Xe Rn Li Na K Rb Cs Fr 11/6/

24 The 1 st ionization energy (A > A + + e ) decreases going down a group. Li Na Group 1A K Rb Cs Fr 11/6/

25 Li + Electron in 2s orbital Electron in 3s orbital Electron in 4s orbital Na + As you go down a group, the outermost electron(s) are further from nucleus, and are easier to remove. This is the same order as the chemical reactivity of these metals as reducing agents. ( Donate electrons ) K > Na > Li K + Another way to put this is that, as you go to larger atoms in the same group, the effective nuclear charge decreases due to shielding by inner electrons. 11/6/

26 He Ne The 1 st ionization energy (A > A + + e ) generally increases going (LtoR) across a row. Ar Kr You re adding a proton in nucleus and electron around the peripheryof atom. Electrons in the same shell do not shield each other from the nuclear charge too spread out. Xe Rn Li Na K Rb Cs Fr 11/6/

27 In Period 2, B,C, and N (and O, F, and Ne even more so) all have slightly lower ionization potentials than expected. He l tia n te o P n tio iza n Io st 1 H Li Be B C N O F Ne Na Mg Atomic Number > 11/6/

28 B and O have anomolouslylow 1 st ionization potentials, which means an outer electron is unexpectedly easy to remove. Boron, carbon, and nitrogen outermost electrons go into a porbital,which is less stable than an sorbital ( ). Oxygen, fluorine and neon outermost electrons go into a twoelectron porbital ( ). These are further destabilized by electronelectron repulsion within the orbital. 11/6/

29 The End 11/6/