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Chapter 7. Periodic Properties of the Elements.. 7.1 Development of The Periodic Table. The number of known elements in 1800 was 31, but had increased.

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Presentation on theme: "Chapter 7. Periodic Properties of the Elements.. 7.1 Development of The Periodic Table. The number of known elements in 1800 was 31, but had increased."— Presentation transcript:

1 Chapter 7. Periodic Properties of the Elements.

2 7.1 Development of The Periodic Table. The number of known elements in 1800 was 31, but had increased to 63 by 1865. In 1869 Mendeleev and Lothar Meyer published almost identical classification schemes. Both noticed that similar chemical properties recurred when elements were placed in order of increasing atomic weight. Mendeleev advanced his ideas more vigorously, and had the insight to leave blank spaces and predict the existence of yet undiscovered elements. Dimitri Mendeleev (1834-1907)

3 Missing Element =eka-silicon Group 5 Group 1 Group 7 Group 4 Group 6 Group 3 Group 2 Group 2b Group 1b Group 4b Group 5b Group 6b

4 Thus, Mendeleev predicted the existence of what he called eka-silicon, which we today refer to as Germanium. Mendeleev predictions.Observed props. of Ge At. Wt.7272.59 g/mol Density:5.55.35 g/cc Sp. Heat:0.3050.309 J/g-K M.Pt.high947 o C Colordark graygrayish white Oxide:MO 2 GeO 2 Density oxide:4.74.70 Chloride:XCl 4 GeCl 4 B.Pt. chloride:<10084 o C

5 The shells and sub-shells being filled across the periodic table:

6 Appearance: Li Na Be Mg Group I A Group II A Because they have the same valence electron configuration, elements in the same group have similar chemistry:

7 Reactivity: Sodium in waterPotassium in water All the elements in group 1 react violently with water: 2 M (s) +2 H 2 O (l)  2 MOH (aq) + H 2 (g)

8 Atomic Numbers (Z): Henry Moseley (1887-1915) developed the concept of atomic numbers. Each element produces X-rays of a distinct frequency, which increased with atomic number. Atomic number = number of protons in the nucleus. The concept of At. No. allowed correct placement of e.g. Ar (Z = 18) before K (Z = 19) even though At. Wt. Ar > K. A similar problem occurs with Co (At. Wt = 58.9) and Ni (58.7), even though Z = 27 for Co and 28 for Ni. Henry Moseley (1887-1915)

9 The K x-ray spectrum of germanium: Energy (keV) (kilo-electron-volts)

10 Relationship between the frequency of the K x-ray peak and atomic number of the element (Z) Z Al Kr Ar K determining Z from energy of x-ray peak now places Ar and K in the right order

11 The core electrons and the valence electrons: Nucleus = 12+ core Electrons = Ne 10 e - valence electrons in 3s shell experience small net positive charge the core electrons shield the valence electrons from the nucleus e e The Mg atom

12 7.2 Effective Nuclear Charge. In a many-electron atom, each electron is simultaneously attracted to the nucleus, and repelled by the other electrons. Generally, the valence electrons are screened from the nucleus by the core electrons (~S, where S is the screening constant = number of core electrons). The effective nuclear charge is the approximate charge experienced by the valence electrons. Thus, the effective nuclear charge Zeff, experienced by the valence electrons is given by: Zeff=Z-S Thus, for Na, Zeff = 11 – 10 ~ 1+. However, because electrons in higher shells are less good at shielding the nucleus, the charge experience by Na + is actually about +2.5. number of core electrons

13 Which would you expect to experience a greater Z eff, a 2p electron of an F atom or a 3s electron of an Na atom? Z eff (F, 2p) = +9 – 2 = 7+ Z eff (Na, 3s) = +11 – 10 = 1+ much more tightly bound ! Which electron do you think would be easier to remove? The Na 3s electron is easily removed => metals easily lose electrons The F 2p electron is NOT easily removed => it is much harder to remove an electron from a halogen

14 7.3 Sizes of Atoms and Ions. We can picture atoms as having size in various ways. The bonding atomic radius, or covalent radius is half the distance in homonuclear diatomic molecules, e.g. H 2, Cl 2. The C-C distance in diamond gives us the covalent radius of C as 1.54/2 = 0.77 Å. Bond lengths are given quite closely as the sum of covalent radii, e.g. C-Cl = 0.77 + 0.99 = 1.76 Å. The non-bonding atomic radius = van der Waals radius.

15 Atomic non-bonding (van der Waals) and bonding (covalent) radii of atoms: iodine atom hydrogen atom Non-bonding atomic radii of I and H atoms 2 x I-I bonded Atomic radius I 2 molecule H 2 molecule HI molecule 2 x H-H bonded Atomic radius The bonding atomic radius (‘covalent radius’) of I is ½ the I-I distance in I 2. The H-I internuclear distance in HI is the sum of the bonding atomic radii of H and I.

16 Ionic radii: The ionic radii are the radii of the ions. For cations the ionic radius is smaller than the atomic radius, because there are fewer electrons around the nucleus, while for anions the ionic radii are larger than the atomic radii because there are now more electrons to repel each other: Lithium atom lithium cation radius = 1.34 Å ionic radius = 0.68 Å fluorine atom fluorine anion radius = 0.71 Å ionic radius = 1.33 Å Li + F-F- Li F

17 The sizes of atoms (gray) cations (pink) and anions (blue) (textbook p. 263):

18 Periodic trends in Atomic and ionic radii. 1) Atomic and ionic radius increases as we move down a group. 2) Atomic and ionic radius tends to decrease along a row from left to right

19 Variation in atomic radius with Z (atomic number) Atomic number (Z) He Ne Ar Kr Xe Rn Cs Rb K Na Li

20 Periodic trends in ionic radii. Ionic radii are based on the distance between ions in ionic compounds. In general, cations are smaller than the atoms from which they are derived, but anions are larger. O 2- F - Na + Mg 2+ Al 3+ 1.401.33 0.97 0.66 0.51Å

21 The sizes of atoms (gray) cations (pink) and anions (blue):

22 Ionization energy. The ionization energy is the minimum energy required to remove an electron from the ground state of the isolated gaseous atom or ion. E.g. First (I 1 )and second (I 2 ) ionization energies of Na + are given by ΔH for the processes: Na(g)→Na + (g)+e - Na + (g)→Na 2+ (g)+e -

23 Trends in 1 st Ionization Energies 1 st ionization energy

24 Variations in successive ionization energies: The value of I n increases as n increases for any one element. It increases sharply as we start to remove inner-shell electrons. Core (inner-shell) electrons are much harder to remove that valence electrons:

25 Why does the ionization energy increase sharply after removal of the 5 th electron for Nitrogen? What would a similar graph look like for beryllium (Be)? Nitrogen Beryllium core electrons are much harder to remove than valence electrons !

26 Periodic Trends in First Ionization Energies: If we look at how I 1 varies across the periodic table, we find that: 1. I 1 increases across a row with increasing Z. 2. I 1 decreases down a group, e.g. Li>Na>K>Rb>Cs, or He>Ne>Ar>Kr>Xe. 3. The representative elements (‘main group elements’) show a larger variation in I 1 than do transition elements. The f-block elements show only a small variation in I 1.

27 Variation in first ionization energy with Z along the periodic table: Atomic number (Z) He Ar Kr Xe Rn Ne CsRb K NaLi Noble gases occur at the peaks Alkali metals occur at the troughs

28 Electron Configuration of Ions: When electrons are lost from an atom to form a cation, they are always lost from the shell having the highest principal quantum number, i.e. the highest energy. e.g. Li(1s 2 2s 1 ) → Li + (1s 2 ) loses electron

29 Electronic configurations of transition metals (the d-block) For transition metal ions one finds that in e.g. Fe 2+, the two electrons lost are lost from the 4s 2 level, even though in the neutral atom this level is lower in energy than the 3d level. Fe ([Ar]3d 6 4s 2 ) → Fe 2+ ([Ar]3d 6 ).

30 Addition of electrons: When an electron is added, it goes into the empty or partially filled orbital having the lowest value of n. F (1s 2 2s 2 2p 5 )→F - (1s 2 2s 2 2p 6 )

31 7.5 Electron affinities: This is the energy for adding electrons to the atom. Cl(g) +e - →Cl - (g) ΔE = -349 kJ/mol The halogens have the most negative electron affinities, followed by the chalcogens. Group 2 has very small EA’s, because they have a filled s orbital. Similarly, group 5 has small EA’s because they have a half-filled p subshell, which is particularly stable. For the noble gases, EA’s are all positive, corresponding to the lack of chemistry, except for Xe.

32 Electron affinities (kJ/mol) for some elements:

33 7.6 Metals, non-metals, and metalloids. Metals are lustrous (shiny) and malleable (can be pounded into thin sheets) and ductile (can be drawn into wires), and conduct heat and electricity. All are solids at room temperature except Hg, and Cs and Ga melt at slightly above room temperature. Compounds of metals with non-metals tend to be ionic substances. Most metal oxides are basic, i.e. dissolve in water to give basic solutions. Na 2 O + H 2 O = 2 NaOH

34 Non-metals. These tend not to be lustrous, and usually are poor conductors of heat and electricity. Generally low-melting, although diamond melts at 3570 o C. Because of their electron affinities, non-metals tend to gain electrons when they react with metals, e.g. 2 Mg(s) +O 2 (g) → 2 MgO(s)

35 Compounds composed entirely of non- metals tend to be molecular substances. Most non-metal oxides dissolve in water to form acids: SO 3 (s) + H 2 O(l)→H 2 SO 4 (aq) CO 2 (g) +H 2 O(l)→H 2 CO 3 (aq) acid + base = salt + water H 2 SO 4 +CaO=CaSO 4 +H2O

36 7.7. Trends for the active metals. Students read 7.7 through 7.9 on own time.

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