Final Note on Nodes The H atom wave functions tell us that there are only radial nodes for the 1s, 2s, 3s.. orbitals. Angular or planar nodes become important as we move to p orbitals (one planar node) and d orbitals (2 planar nodes). This is seen in the slide on the next table and graphical representations of d orbitals. (Orbital shapes impt in chemical bonding.)

Hydrogen Atom Wavefunctions – Number of Nodes Orbital Designation Total # Nodes (n-1) Planar Nodes ( l) Radial Nodes 1s00 (s orbital)0 2s10 (s orbital)1 2p11 (p orbital)0 3s20 (s orbital)2 3p21 (p orbital)1 3d22 (d orbital)0

Representations of the five d orbitals FIGURE 8-30 Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 8Slide 3 of 50

The Periodic Table In studying the electronic structure of atoms we mentioned that chemical families of elements have similar valence shell electron configurations. Historically, however, a detailed knowledge of atomic structure came subsequent to the observation that groups of elements had similar chemical properties.

The Periodic Table In the modern Periodic Table elements are arranged in order of increasing atomic number so that groups of elements with similar chemical properties appear in columns. The 100+ known elements are commonly divided into three groups – the metals, non-metals and the metalloids. These three sets of elements have rather different physical properties.

Slide 6 of 35 Metals and Nonmetals and Their Ions Metals – Good conductors of heat and electricity. – Malleable and ductile. – Moderate to high melting points. Nonmetals – Nonconductors of heat and electricity. – Brittle solids. – Some are gases at room temperature. Metalloids – Metallic and non-metallic properties Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 9 Slide 6 of 35

Electrical and Thermal Conductivity The next two slides present electrical and thermal conductivity data (room temperature) for a number of elements. The periodic trends are more important than the units. For illustrative purposes relative thermal conductivities are shown on the second slide. Again, in the modern world, silicon is rather important!

Electrical Conductivity Elements ElementConductivity (S/m) Cost ($/kg) Element Type Silver6.2 x ,198Metal Copper5.9 x Metal Gold4.5 x ,000Metal Nickel1.4 x 10 7 Metal Iron1.0 x 10 7 InexpensiveMetal Germanium2.0 x 10 3 ExpensiveMetalloid Silicon1.0 x 10 3 VariableMetalloid Bromine1.0 x Nonmetal Sulfur1.0 x InexpensiveNonmetal

Thermal Conductivity Elements ElementRelative Thermal Conductivity Element Type Silver1.000Metal Copper0.93Metal Gold0.74Metal Nickel0.21Metal Iron0.19Metal Germanium0.14 (Suprise?)Metalloid Silicon0.35(Surprise?)Metalloid Bromine Nonmetal Sulfur Nonmetal

The Periodic Table Much early work on the Periodic Table was done by the Russian chemist Dimitri Mendeleev and the German chemist Lothar Meyer. (The English chemist John Newlands noticed that elements, when ordered according increasing atomic weight showed “recurring/similar” chemical and physical properties at intervals of eight elements.) It goes without saying that all famous chemists eventually receive philatelic recognition.

Slide 11 of Classifying the Elements: The Periodic Law and the Periodic Table 1869Dimitri Mendeleev Lothar Meyer Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 9 Slide 11 of 35 When the elements are arranged in order of increasing atomic mass, certain sets of properties recur periodically.

Periodic Table - Many Subdivisions! The Periodic Table may be subdivided in many ways. One finds metals on the left and non- metals on the right. Elements with similar chemical properties appear in columns. As one moves across a row or period in the Periodic Table a range of chemical and physical properties is seen. Periods generally end with a chemically unreactive Noble Gas ( a Noble Gas compound was first synthesized in Canada).

Periodic Table - Many Subdivisions! Further subdivisions of the Periodic Table are illustrated on the next slide. One can, for example, divide the elements into Main Group elements and Transition Metals. From high school chemistry you will recall that transition metals have relatively complex chemistry – forming, for example, a variety of ions in ionic compounds (eg. Fe 2+ and Fe 3+ ).

Metal and Non-metal Monatomic Ions Another familiar result/fact from high school chemistry is that in ionic compounds metals form a range of positive ions and non-metals from negative ions. Main Group metals typically lose one or more electrons to form a monatomic ion with a Noble Gas electron configuration. Main Group non-metals gain one or more electrons and also form monatomic ions with a Noble Gas electron configuration.

Slide 15 of 35 Main-Group Metal Ions Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 9 Slide 15 of 35 Metals tend to lose electrons to attain noble gas electron configurations.

Slide 16 of 35 Main-Group Nonmetal Ions Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 9 Slide 16 of 35 Nonmetals tend to gain electrons to attain noble-gas electron configurations

Slide 17 of 35 Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 9 Slide 17 of 35

Atomic and Ionic Radii The “size” of an individual H atom or He atom (for example) in the gas phase in difficult to define exactly. Why? (Recall the probabilistic description of electronic structure considered earlier and in Chapter 8 of the text).For molecules we can use spectroscopic methods to determine internuclear distances in gas phase molecules. X-ray diffraction gives internuclear distances in crystalline solids.

Atomic and Ionic Radii We will use spectroscopic and X-ray diffraction data to give us working estimates of atom and ionic radii (sizes). Atoms are small. We will use picometers (pm = m) and Angstroms ( m) to describe atom/ion sizes. From a homonuclear diatomic molecule (such as H 2, Cl 2 or Na 2 ) the covalent radius of an atom can be obtained as half of the internuclear separation.

Slide 20 of Sizes of Atoms and Ions Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 9 Slide 20 of 35 FIGURE 9-3 Covalent, metallic, and ionic radii compared

Atomic and Ionic Radii - Surprises? The results presented on the previous slide are perhaps a little startling. Let’s use the metallic radius for Na as an estimate for the size of a neutral Na atom. Na→ Na + + e - (11e - ) (10e - ) Both the Na atom and the Na + ion have 11 protons. If removing an e - from an atom were a process analogous to peeling an orange we might expect the Na atom to be about 10% larger than the Na + ion. 9/10/2015Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 9Slide 21 of 35

Atomic and Ionic Radii - Surprises? In fact, the Na atom (radius ~ 186 pm) and the Na + ion (radius ~ 99 pm) have very different sizes. On a volume basis the relative sizes are given by (186pm/99pm) 3 = 6.6! We can gain some insight into this large difference by looking first at the electron configurations of the Na atom and the Na + ion. Na atom: 1s 2 2s 2 2p 6 3s 1 Na + ion:1s 2 2s 2 2p 6 9/10/2015Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 9Slide 22 of 35

Atomic and Ionic Radii - Continued We see that the Na atom has electrons in two principal energy levels but the Na + ion has only two principal energy levels with electrons. Moving from the Na atom to the Na + ion the number of (repulsive) coulombic interactions between electrons also drops. Why? This effect should also cause the size of the ion to be smaller than that of the neutral atom.

Atomic and Ionic Radii - Continued A concept of screening electrons and effective nuclear charge is very useful in discussing atomic radii, ionization energies and other properties that show periodic trends. The key idea is that the outermost electrons in an atom experience an effective nuclear charge that is smaller than the actual nuclear charge due to the presence of the “inner” or screening electrons that lie closer (on average!) to the nucleus than the valence electrons.

Atomic and Ionic Radii - Continued The concepts of screening electrons and effective nuclear charge are illustrated on the next slide for the magnesium (Mg) atom. The atomic number of Mg is 12 (12 protons). The condensed electron configuration for the Mg atom is 1s 2 2s 2 2p 6 3s 2. The number of core or inner electrons lying “between” the valence electrons and the nucleus is 10. We say that the valence electrons see an effective nuclear charge of +2. Z effective = Z – S = = 2

Slide 26 of 35 The shielding effect and effective nuclear charge, Z eff FIGURE 9-6 Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 9 Slide 26 of 35 Screen of electron charge from 10 core electrons (-) (12+)

Slide 27 of 35 Ionic Radius Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 9 Slide 27 of 35 FIGURE 9-7 A comparison of atomic and ionic sizes

Slide 28 of 35 Ionic Radius Copyright © 2011 Pearson Canada Inc. Slide 28 of 35 General Chemistry: Chapter 9 Cations are smaller than the atoms from which they are formed. For isoelectronic cations, the more positive the ionic charge, the smaller the ionic radius. Anions are larger than the atoms from which they are formed. For isoelectronic anions, the more negative the charge, the larger the ionic radius.

Slide 29 of 35 Covalent and anionic radii compared FIGURE 9-8 Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 9 Slide 29 of 35

Slide 30 of 35 A comparison of some atomic and ionic radii FIGURE 9-9 Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 9 Slide 30 of 35

Sodium Atom and Sodium Ion The startling contraction in size when the sodium ion forms from the neutral atom reflects the fact that the third principal energy level has been emptied completely, the outermost electrons are now more strongly attracted to the nucleus and electron-electron repulsions have been reduced. Why?

Periodic Table – Binary Ionic Compounds Class Examples: We’ll look at several examples of writing chemical reactions for binary ionic (metal-nonmetal) compounds from the constituent elements. As well, we want now to be able to recognize basic oxides and consider their reactions with water. We’ll look at the more comples examples of binary covalent compounds in a few days.