2011 CHEM 120: CHEMICAL REACTIVITY

Transcription

1 2011 CHEM 120: CHEMICAL REACTIVITY INORGANIC CHEMISTRY SECTION Lecturer: Dr. M.D. Bala Textbook by Petrucci, Harwood, Herring and Madura 15 Lectures (4/10-29/10) 3 Tutorials 1 Quiz 1 Take-home test science/level1/chem120/lecturenotes.html

2 Alkali Metals The Periodic table Noble Gases p-block Alkaline Earth metals Transition metals d-block Halogens p-block Main Group p-block Main Group s- block Lanthanides and Actinides f- block

3 Metals, Nonmetals and Metalloids Metals Good conductors of heat and electricity Malleable and ductile Generally solids (apart from mercury) with moderate to high melting points Nonmetals Nonconductors of heat and electricity Brittle solids Some are gases at room temperature Metalloids Have some metallic and some nonmetallic properties Often behave as semiconductors Often form amphoteric oxides

4 Metals tend to Lose Electrons Form Cations e.g. M M n+ + ne - Loose electron Mg Mg e -

5 Nonmetals tend to Gain Electrons Form Anions X + ne - X n- gain electron e.g. S + 2e - S 2-

6 Main Group The s- and d block chemistry tends to be concerned with cations (metals) The p-block elements have much more diverse chemistry (metals, metalloids and nonmetals) The chemical character changes even within a group

7 PERIODIC TRENDS Lets look at some trends in experimentally determined properties, such as: ATOMIC RADII IONIC RADII IONIZATION ENERGIES ELECTRON AFFINITIES MAGNETIC PROPERTIES Lets start with atomic radii

8 Covalent radius The Sizes of Atoms and Ions Metallic radius One half the distance between the centres of two atoms that are bonded covalently. One half the distance between the centres of adjacent atoms in a solid metal. Ionic radius Radius of a spherical ion. It is atomic radius associated with an element in its ionic compound.

9 ATOMIC RADIUS ATOMIC RADIUS TRENDS IN ATOMIC RADII FOR GROUPS 1, 2 and 13 to 18 Atomic radii increase down the group Orbital's get bigger as n increases Why? Radial probability density extends further

10 ATOMIC RADIUS ATOMIC RADIUS TRENDS IN ATOMIC RADII FOR GROUPS 1, 2 and 13 to 18 Atomic radius Increase from R to L Increase Atomic radii decrease across (L to R) a period Why? The inner electrons do not screen the nucleus..

11 Atomic Radius

12 SCREENING AND PENETRATION SCREENING is to block the other outer electrons from the nucleus effect PENETRATION is to get close to the nucleus attraction nucleus electron repulsion attraction nucleus Core electrons repulsion Valence electron

13 Effective Nuclear Charge Net positive charge experienced by an electron (Z eff ) Z eff = Z - S where Z = is the number of protons in the nucleus S = average number of electrons responsible for screening In a many-electron atom, electrons are both attracted to the nucleus and repelled by other electrons. The nuclear charge that an electron experiences depends on both factors. e.g. Na = 1s 2, 2s 2, 2p 6, 3s 1 [Ne] 3s 1

14 SCREENING AND PENETRATION ATOMIC RADIUS DECREASES ACROSS PERIOD! WHY? As we add electrons in one by one (new element) The electrons go into same shell (orbital) and do not screen the nucleus any further than before. As nuclear charge goes up, electrons all held or attracted more closely to the nucleus. ATOMIC RADIUS INCREASES DOWN THE GROUP! WHY? When adding a full electronic orbital this gives more screening therefore electrons are held more loosely.

15 SCREENING AND PENETRATION Z eff Increases 1 1s s Radius decreases 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s La 5d 6p 7s Ac 6d 4f 5f 18 1s

16 SCREENING AND PENETRATION s s 2p 3s 4s 5s 6s La 7s Ac RADIUS INCREASES Z eff DECREASES 3d 4d 5d 6d 3p 4p 5p 6p 4f 5f 18 1s

17 Cationic Radii Cations are smaller than their parent atoms. Why? The outermost electron is removed (Radial probability density reduced) Repulsion forces between electrons are reduced.

18 Anionic Radii Anions are lager than their parent atoms. Why? Electrons are added (Radial probability density is increased) Electron repulsion forces between electrons are increased Down the group there is an increase in the value of n

19 Atomic and Ionic Radii of Elements Cations are smaller than their parent atoms: The outermost electron is removed and repulsions are reduced. Anions are larger than their parent atoms: Electrons are added and repulsions are increased.

20 Questions 1) Explain the following observations. a) The atomic radius of aluminium is 125 pm while the ionic radius of the same element is only 50 pm. b)the ionic radius of oxygen is 140 pm, 80 pm larger than its atomic radius. 2) The internuclear distance in solid magnesium oxide is 205 pm. The internuclear distance in solid magnesium chloride is 246 pm. Given that the ionic radius of O 2- is 140 pm, calculate the ionic radius of Cl -. 3) Predict which of the following ions has the larger ionic radius: Mg 2+ or Al 3+.

21 Prentice-Hall 2002 General Chemistry: Chapter 10 Slide 21 of 35

22 QUESTIONS. Which is bigger? EXAMPLES (i) Na or Rb (ii) K or Ca (iii) Ca or Ca 2+ (iv) Br or Br - What about ISOELECTRONIC SPECIES? (Same number of electrons) e.g. K +, Ca 2+ and Sc 3+ or Cl -, S 2- and P 3- Ar = [Ne] 3s 2 3p 6

23 Isoelectronic cations: Example K +, Ca 2+ and Sc 3+ Which is the smallest above? Answer The more positive the ionic charge, the smaller the ionic radius Isoelectronic Anions: Example Cl -, S 2- and P 3- Which is the largest above? Answer The more negative the ionic charge, the larger the ionic radius

24 QUESTION The species F -, Na +, Mg 2+ have relative sizes in the order 1. F - < Na + < Mg F - > Na + > Mg Na + > Mg 2+ > F - 4. Na + = Mg 2+ = F - 5. Mg 2+ > Na + > F - Which of the above statement true? ANSWER

25 Prentice-Hall 2002 General Chemistry: Chapter 10 Slide 25 of 35

26 Ionization Energy The energy required to remove an electron from an isolated atom or ion in the gas phase. It is considered a measure of the "reluctance" of an atom or ion to surrender an electron or the "strength" by which the electron is bound. The greater the ionization energy, the more difficult it is to remove an electron. e.g. Mg(g) Mg + (g) + e - I 1 = 738 kj Mg + (g) Mg 2+ (g) + e - I 2 = 1451 kj Z 2 eff I = R H n 2 E

27 First Ionization Energy

28 Trends in First Ionization Energies However, there are two apparent discontinuities in this trend.

29 Trends in First Ionization Energies The first occurs between Groups 2 and 13 Electron removed from p-orbital rather than s-orbital Electron farther from nucleus Small amount of repulsion by s electrons. e.g. Mg = [Ne] 3S 2 Al = [Ne] 3S 2 3P 1

30 Trends in First Ionization Energies The second occurs between Groups 15 and 16 Electron removed comes from doubly occupied orbital. Repulsion from other electron in orbital helps in its removal. e.g. P = [Ne] 3S 2 3P 3 S = [Ne] 3S 2 3P 4

31 Table 10.4 Ionization Energies of the Third-Period Elements (in kj/mol) I 2 (Mg) vs. I 3 (Mg) I 1 (Mg) vs. I 1 (Al) I 1 (P) vs. I 1 (S)

32 IONIZATION ENERGY IONIZATION ENERGY Zeff Decreases Trends in First Ionisation Energy First Ionization energies decrease down the group Why? Z eff INCREASES UP THE GROUP Electrons closer to nucleus more tightly held

33 IONISATION ENERGY Trends in First Ionisation Energy IONISATION ENERGY Z eff INCREASES Greater effective nuclear charge across period Poor shielding by electrons added

34 QUESTIONS 1) Which element in each of the following pairs has the higher I 1 value? (i) Na or Mg (ii) Mg or Al (iii) As or Sb 2) Explain why the first ionization energies (I 1 ) of Group 13 are smaller than those of Group 2 for a given value of n. 3) Explain why the first ionization energies (I 1 ) of nitrogen is higher than oxygen.

35 Prentice-Hall 2002 General Chemistry: Chapter 10 Slide 35 of 35

36 Electron Affinity The energy change associated with the addition of an electron to a gaseous atom. X (g) + e X (g) e.g. F(g) + e F (g) DE= -328 kjmol -1 exothermic process Electron affinity energy values can be either positive or negative. Energy is required to add an electron to an already stable configuration. e.g. Be and Mg (ns 2 ) noble gases (ns 2 np 6 )

37 First Electron Affinities

38 Second Electron Affinity Second electron affinities are all positive Why? e.g. It takes energy to overcome the repulsion between an electron and the already-negative ion. O (g) + e - O - (g) O - (g) + e - O 2- (g) EA 1 = -141 kj EA 2 = +744 kj

39 Questions 1. Predict which of the following elements has the more exothermic electron affinity: fluorine or sodium? 2. For the following group of elements, select the one with the most negative electron affinity: As, B, Cl, K, Mg, S?

40 Prentice-Hall 2002 General Chemistry: Chapter 10 Slide 40 of 35

41 Magnetic Properties Diamagnetic atoms or ions: All e - are paired. Weakly repelled by a magnetic field. Paramagnetic atoms or ions: Unpaired e -. Attracted to an external magnetic field.

42 Paramagnetism See example 9-4 Textbook by Petrucci Chapter 9 page 358

43 Diagonal Relationships Increase in Z eff Increase in size decrease in Z eff decrease in size Diagonal relationship exists because of size and charge effect. Moving from left to right in the periodic table, the atomic size decreases Moving down the group, the atoms and ions increase in size Thus on moving diagonally, the size remains nearly the same e.g. Li + = 0.76 Å and Mg 2+ = 0.72 Å

44 Diagonal Relationships These pairs (Li & Mg), (Be & Al), (B & Si ) etc. exhibit similar properties e.g. Boron and Silicon are both semiconductors Why? form halides that are hydrolysed in water and have acidic oxides The atom or their ion pairs have comparable sizes, ionisation potential, electron affinity, etc. e.g. (Li and Mg) or (Li + and Mg + )

45 Periodic properties of the elements