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1 Chapter 7 Periodic Properties of the Elements. 2 The Periodic Table n Developed independently by German Julius Lothar Meyer and Russian Dmitri Mendeleev.

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Presentation on theme: "1 Chapter 7 Periodic Properties of the Elements. 2 The Periodic Table n Developed independently by German Julius Lothar Meyer and Russian Dmitri Mendeleev."— Presentation transcript:

1 1 Chapter 7 Periodic Properties of the Elements

2 2 The Periodic Table n Developed independently by German Julius Lothar Meyer and Russian Dmitri Mendeleev (1870s). n Didnt know much about atom. n Put in columns by similar properties. n Predicted properties of missing elements.

3 3 Details n Valence electrons- the electrons in the outermost energy levels (not d). n Core electrons- the inner electrons

4 4 7.2 Effective Nuclear Charge n Properties of atoms depend not only on their electron configurations but also on how strongly their outer electrons are attracted to the nucleus. n We view the net electric field of the nucleus as if it results from a single positive charge. This is known as the EFFECTIVE NUCLEAR CHARGE, Z eff.

5 5 Shielding n Valence electrons are attracted to the nucleus of the atom and are repelled by the other electrons in the atoms. –The inner electrons partially shield or screen the outer electrons from the attraction of the nucleus.

6 6 Periodic Trend of Z eff n The effective nuclear charge increases as we move across any row (period) of the table. –The number of core electrons stay the same as we move across the row, the actual nuclear charge increases.

7 7 Periodic Trend of Z eff n Going down a column, the effective nuclear charge experienced by valence electrons changes far less than it does across a row. –We would expect the effective nuclear charge for the outer electrons in lithium and sodium to be about the same, roughly 3-2 = +1 for Li and = +1 for Na.

8 8 7.3 Sizes of Atoms and Ions n Quantum Mechanical Model – no defined boundaries at which the electron distribution becomes zero. n We can define atomic size based on the distance between atoms in various situations. –Nonbonding atomic radius –Bonding atomic radius (shorter) n Allows us to estimate the bond lengths between different elements in molecules

9 9 Practice n Use page 231, Figure 7.5: –Natural gas used in home heating and cooking is odorless. Because natural gas leaks pose the danger of explosion or suffocation, various smelly substances are added to the gas to allow detection of a leak. One such substance is methyl mercaptan, CH 3 SH. Predict the lengths of the C-S, C-H, and S-H bonds in this molecule.

10 10 Practice n Use figure 7.5 to predict which will be greater, the P-Br bond length in PBr 3 or the As-Cl bond length in AsCl 3.

11 11 Periodic Trends in Atomic Radii n 1. Within each column, atomic radius tends to increase from top to bottom. –Results primarily from the increase in the principal quantum number of the outer electrons. n 2. Within each row, atomic radius tends to decrease from left to right. –The increase in the effective nuclear charge as we move across a row steadily draws the valence electrons closer to the nucleus.

12 12 You Try It… n Arrange the following atoms in order of increasing size: –P, S, As, Se n Arrange the following atoms in order of increasing atomic radius: –Na, Be, Mg

13 13 Periodic Trends in Ionic Radii n 1. Cations are smaller than their parent atoms n 2. Anions are larger than their parent atoms. n For ions carrying the same charge, size increases as we go down a column in the periodic table.

14 14 You Try It… n Arrange these atoms and ions in order of decreasing size: –Mg 2+, Ca 2+, Ca n Which of the following atoms and ions is the largest? –S -2, S, O -2

15 15 Isoelectronic series n A group of ions all containing the same number of electrons –O -2, F -, Na +, Mg 2+, Al 3+ all have 10 electrons. n We can list the members in order of increasing atomic number, and therefore nuclear charge increases as we move through the series. Because the number of electrons is constant, the radius of the ion decreases with increasing nuclear charge, as the electrons are more strongly attracted to the nucleus. –Oxide is the largest ion, smallest atomic number –aluminum the smallest ion, highest atomic number

16 16 You Try It… n Arrange the ions of potassium, chloride, calcium, and sulfide in order of decreasing size.

17 Ionization Energy n Ionization energy-the energy required to remove an electron from a gaseous atom n Highest energy electron removed first. First ionization energy ( I 1 ) is that required to remove the first electron. First ionization energy ( I 1 ) is that required to remove the first electron. Second ionization energy ( I 2 ) - the second electron Second ionization energy ( I 2 ) - the second electron n etc. etc.

18 18 Trends in ionization energy n for Mg I 1 = 735 kJ/moleI 1 = 735 kJ/mole I 2 = 1445 kJ/moleI 2 = 1445 kJ/mole I 3 = 7730 kJ/moleI 3 = 7730 kJ/mole n The effective nuclear charge increases as you remove electrons. n It takes much more energy to remove a core electron than a valence electron because there is less shielding.

19 19 Explain this trend n For Al I 1 = 580 kJ/moleI 1 = 580 kJ/mole I 2 = 1815 kJ/moleI 2 = 1815 kJ/mole I 3 = 2740 kJ/moleI 3 = 2740 kJ/mole I 4 = 11,600 kJ/moleI 4 = 11,600 kJ/mole

20 20 Across a Period Generally from left to right, I 1 increases because Generally from left to right, I 1 increases because –there is a greater nuclear charge with the same shielding. As you go down a group I 1 decreases because electrons are farther away. As you go down a group I 1 decreases because electrons are farther away.

21 21 It is not that simple Z eff changes as you go across a period, so will I 1 Z eff changes as you go across a period, so will I 1 n Half filled and filled orbitals are harder to remove electrons from. n heres what it looks like.

22 22 First Ionization energy Atomic number

23 23 First Ionization energy Atomic number

24 24 First Ionization energy Atomic number

25 25 Periodic Trends in 1 st Ionization Energies n 1. Within each row, I1 generally increases with increasing atomic number. –Alkali metals show the lowest, noble gases the highest. n 2. Within each column, the ionization energy generally decreases with increasing atomic number. –He>Ne>Ar>Kr>Xe

26 26 You Try It… n Referring to a periodic table, arrange the following atoms in order of increasing first ionization energy –Ne, Na, P, Ar, K n Write the electron configuration of: –Calcium Ion –Cobalt (III) Ion –Sulfide Ion

27 Electron Affinities n The energy change that occurs when an electron is added to a gaseous atom –Affinity = attraction n Cl(g) Cl + + e- Delta E = 1251 (first ionization energy) kJ/mol n Cl(g) + e- Cl - Delta E = -349 kJ/mol (electron affinity)

28 28 Remember n Ionization energy measures the ease with which an atom LOSES an electron. n Electron affinity measures the ease with which an atom GAINS an electron.

29 29 Affinity tends to increase across a period Affinity tends to decrease as you go down in a period Electrons farther from the nucleus experience less nuclear attraction Some irregularities due to repulsive forces in the relatively small p orbitals Electron Affinity - the energy change associated with the addition of an electron Values of Electron affinity are always negative, the more negative the value the more stable the atom becomes. Therefore, the more negative the value, the more that atom wants an electron.

30 30 Table of Electron Affinities

31 31 A measure of the ability of an atom in a chemical compound to attract electrons Electronegativities tend to increase across a period Electronegativities tend to decrease down a group or remain the same Electronegativity

32 32 Periodic Table of Electronegativities

33 33 Summary of Periodic Trends


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