Properties of Materials

1. Properties of Materials • The observable properties of pure substances are not those of individual atoms but are those of large collections of particles whether they be atoms, molecules, or ions. • Physical and chemical properties of materials can be explained by gaining an understanding of the structural units of these materials and the forces involved in holding these units together.

2. Sructures of Crystals

3. Properties of Materials • Substances can be classified according to the structural units which make up the aggregates and the type of bond or force that exists between the units. • 1. Ionic Substances - made up of ions which are held together by electrostatic forces called ionic bonds. • These are relatively strong forces.

4. 2. Molecular Substances - structural units are polar or non-polar molecules. Attractive forces are called Van der Waals forces. • These forces are generated by periodic shifting in electron clouds which result in attractions between nuclei and electrons. • They are relatively weak forces.

5. 3. Network or macromolecular solids - structural units are atoms held together by covalent bonds. Very strong forces. • 4. Metallic substances - structural units are atoms which are held together by metallic bonds

6. Ionic Substances • These are characterized by the following properties: • 1. as liquids they are good conductors of electricity due to the presence of ions • 2. as solids they are poor conductors because there are no freely mobile IONS 3. forces between the ions are strong so they have high melting points and boiling points and low volatilities and vapour pressures

7. Ionic Substances • 4. They are brittle and easily broken when stress is exerted because ions of the same charge will encounter one another resulting in increased repulsion. View here 5. When dissolved in water they are good conductors.

8. Na1+

9. F = kq1q2 r2 Factors Affecting Bond Strength and Crystal Stability • The force of attraction between ions is determined by the size and charge of the ions. This force can be mathematically represented by • Where: k = a constant • q = charge on the ions • r = the distance between the ions

10. Electrostatic Forces • This tells us that as ion charge increases "f" increases and bond strength increases. It also tells us that as "r" increases "f" decreases and bond stength decreases. • This force of attraction can be used to predict the melting point, boiling point and solubility of ionic solids. • Example - compare the melting points and solubilities of CsCl and NaCl.

11. Ionic Crystals • The stronger the forces between the ions the higher the melting and boiling points and the lower the solubility. • Since the chloride ions are the same size lets compare the Cs ion with the Na ion. • Cs is in the 6th PERIOD of the periodic table. It has 6 electron shells. Na is in the 3rd PERIOD of the periodic table so it has 3 electron shells.

12. Ionic Crystals • This means that Cs ions are larger than Na ions. Therefore the distance between Cs ions and Cl ions in a crystal of CsCl is greater than the distance between Na ions and Cl ions in a NaCl crystal. If the distance is greater then the force must be weaker. That is why CsCl melts and boils at lower temperatures than NaCl and is also the reason it is more soluble in water. Keep in mind this generalization holds because Cs ions and Na ions are the same charge, 1+.

13. Let's compare compounds with elements in the same PERIOD but different groups. • MgO NaCl • Mg2+ O2- Na1+ Cl1- • Since the force of attraction between ions is proportional to the magnitude of the charges the force of attraction between Mg2+ and O2- is greater than the force of attraction between Na1+ and Cl1-. That is why the melting point of MgO is higher than NaCl and the solubility of MgO in water is lower. (Forces are stronger so its harder to dissolve) • Questions - Page 324 # 2 - 6.

14. Molecular Crystals • In general these substances have these properties: • 1. They do not conduct electricity as solids or liquids because there are no free electrons in the solids and the liquids do not contain charged particles. • 2. Many exist as gases as room temperature and are relatively volatile because the forces holding the molecules together (Van der Waals forces) are quite weak

15. Molecular Crystals • 3. The melting points and boiling points are relatively low. • 4. The solids are generally soft and have a waxy consistency. • 5. Large amounts of energy are required for chemical decomposition because the forces within the molecules are stronger than the forces between the molecules.

16. Molecular Crystals • The basic structural units in molecular crystals are either polar or non-polar molecules and there are wide variations in the properties of individual molecular substances. In general, polar molecular crystals have higher melting points than non-polar molecular crystals because the forces between the molecules are stronger.

17. Molecular shape will influence properties. Sulfur, for instance, as a solid exists as ring-shaped molecules with the formula S8 packed into tight crystals. When it melts it first changes into a yellow liquid that flows easily. As its temperature increases, however, these rings break and the sulfur atoms align in long chains, which don't flow easily so the viscosity increases and the colour darkens. As temperature further increases the viscosity decreases as the chains break apart and flow more easily.

18. The forces of attraction between molecules in molecular crystals can be separated into three categories: • 1. Dipole forces of attraction (in polar molecular crystals) • 2. Van der Waals forces (predominate in non-polar molecular crystals) • 3. Hydrogen bonds (occur between N and H, O and H, and F and H)

19. Van der Waals forces, in general, get stronger as the molecules in the crystal get larger. This, presumably, is due to the increased numbers of protons and electrons. This trend can be used to predict trends in boiling points of similar compounds. • Example - compare the boiling points of SbH3, AsH3, and PH3. • pg. 311 in textbook

20. As the number of electrons increases the boiling point increases because the van der Waals forces between the molecules get stronger. • Look at the boiling point of NH3. Why does it deviate from this trend? • H bonding is responsible. • H atoms exert an electrostatic force of attraction on electron pairs of other highly electronegative atoms. (N, O, and F)

21. H bonding is responsible for many of the unique properties of water. • 1. The density of the solid being less than the liquid. (thats why ice floats) • 2. The relatively high boiling point of such a small molecule. (thats why water exists as a liquid between the temperatures 0 and 100 at 100 kPa) • 3. Its ability to dissolve many solids. • Questions - pg. 324 # 8 - 12

22. Network (Macromolecular) Solids • These crystals consist of networks of atoms connected by strong covalent bonds. These may be atoms of the same (diamond), or different (silicon carbide) elements. • In general they are: • 1. highly stable, extremely hard, and insoluble in most solvents because of the strength of the bonds between the atoms • 2. poor electrical conductors because the electrons are very tightly held

23. Properties of these network crystals depend on the arrangement of the atoms in the crystals. • Example - Graphite and Diamond - both made of C atoms

24. Graphite is a two dimensional crystal with C atoms arranged in hexagonal rings. These rings are interconnected to form large sheets. The C atoms in the rings are held together by strong covalent bonds so graphite has a high melting and boiling point.

25. The individual sheets are held together by weak bonds so graphite is soft. Graphite is also a conductor of electricity so there must be loose electrons moving between the C atom sheets. Its luster is also due to the ability of the electrons to absorb and emit light.

26. Diamond crystals are poor electrical conductors and transparent because there are no loose electrons to absorb energy. The C atoms are joined by strong tetrahedrally oriented covalent bonds which accounts for the extreme hardness. • Questions - pg. 324 # 14, 15

27. Metallic Crystals open-pit copper mine • Most metals are characterized by these properties: 1. Metallic surfaces are good reflectors of light because of the abundance of mobile electrons which absorb and re-emit light. 2. Excellent conductors of electricity and heat because of the abundance of mobile electrons.

28. 3. Ductility (ability to be drawn into wires) and malleability (ability to be shaped into sheets) because of the bonding flexibility between the metallic ions and the sea of mobile electrons. • 4. Electron emmission caused by heat and light. • A metal consists of positive metallic ions arranged in a "sea" of highly mobile electrons. The electrons are free to move throughout the crystal rather like gas molecules confined in a container.

29. The metallic bond is the result of the attraction between the positively charged "kernels" and the surrounding electrons. • The strength of a metallic bond depends on the nuclear charge and the number of electrons in the outer energy level. As the number of outer electrons increases, the strength of the bond increases, and the melting and boiling points, and hardness increases.

30. Copper Crystal 2+ - - - - - - 2+ - 1+ - 1+ - 2+ - 1+ - - - - - “sea” of electrons - - 2+ - - 1+ 2+ - - - - 1+ - - - - - 2+ 2+

31. Summary • Compare the physical properties of the elements in period 3 of the periodic table. (pg. 323 in the text) These properties can be understood if you compare the forces of attraction between the structural units. Remember, in general, the stronger the forces of attraction the higher the melting point, boiling point and the lower the volatility.

32. The forces of attraction between the particles in the metallic crystals increase as you go from Na to Al due to an increase in the number of "mobile electrons" from the valence shell. • Silicon crystals have much higher attractions between the atoms because of the covalent bonds in this network solid. • The attractive forces in phosphorus are much smaller because this material forms molecular crystals P4.

33. S8 molecules in sulphur crystals are larger than P4 molecules so the van der Waals forces increase. • Cl2 molecules are smaller still so van der Waals forces continue to decrease. • Ar is monatomic so the van der Waals forces continue to decrease. • Questions - pg 324 - 325 # 16 - 20, 22, 24, 27, 35 - 37.

35. Type of Solid Types Of Particles Intermolec ular Forces Factors Charge –Bigger- Bigger, Ion Size Bigger-smaller + and – ions Ionic Ionic bonds Polar Molecular Polar Molecules Dipole Dipole • EN Bigger-bigger London Dispersion (LD or VdW) Non-Polar Molecular Non-polar Molecules Molecular Size Bigger-bigger Ion charge # e1- in sea + ions Metallic bonds Metallic • EN Bigger- smaller • # bonds/atom Covalent Network Neutral atoms Covalent bonds

36. Is the substance an element or a compound? Is it a metal or nonmetal? Element Compound metal+ nonmetal nonmetals Covalent Network Crystal Group III + V SiO2 strong covalent bonds the more bonds per atom the stronger the crystal Nonmetal Molecular Crystals Metal Metallic crystals metallic bonds depend on charge of positive ion and # of e1- in the sea of e1- the greater the charge the stronger the force ionic crystals ionic bonds beween ions of opposite charge depend on magnitude of charge and distance between ions closer the ions the stronger the force higher the product of charge the stronger the force Nonpolar Molecular Crystals VdW forces are weak and get bigger as molecules get bigger Covalent Network Crystal Group 1V strong covalent bonds the more bonds per atom the stronger the crystal Polar or Nonpolar Molecules (use Lewis Structure and Shape) Polar Molecular Crystals dipoles depend on difference in En H bonds if H and N,O,F Non-Polar Molecular Crystals VdW forces depend on size of molecule