Task 001
- 2. Q1. All solids are crystalline. Is it true? Discuss your answer. • No, not all solids are crystalline. • This is because solid can be divided into crystalline and amorphous. They have highly ordered three-dimensional arrangement of particles. The presence of such long-range order is the defining property of crystals. Amorphous are randomly arranged solid particles in three- dimension. • Examples of amorphous are glass, plastic, rubber.
- 3. Q2. What is a molecule? • Molecule is a group of atoms that are held together tightly enough by covalent bonds to behave as a single particle. • For example is water. • A molecule always has a definite composition and structure and also has little tendency to gain or lose atoms. • Molecule representing the smallest fundamental unit of a chemical compound that can take part in a chemical reaction.
- 4. Q3. How are the crystal structures of solids usually determined? • The structure of a crystal are usually determine by the interference patterns produced when an X-ray beam passes through it. • A crystal consists of a regular array of atoms, each of which is able to scatter an electromagnetic wave that happen to strike it. • A beam of X-rays, all with the same wavelength, that falls upon a crystal will be scattered in all directions within it; but owing to the regular arrangement of the atoms, in certain directions, the scattered waves will constructively interfere with one another while in others they will destructively interfere. The phenomenon is known as X-ray diffraction. • The resulting pattern of high and low X-ray intensities can be analyzed to yield an arrangement in space of the scattering centers, which are atoms of the crystal.
- 5. Q4. When a compound is formed, atoms of the elements present are linked by chemical bonds. Describe the two types of chemical bonds which are ionic and covalent. Ionic bonds Covalent bonds The complete transfer of valence electron(s) between atoms. The sharing of electrons between atoms. A type of chemical bond that generates two oppositely charged ions. Occurs between two atoms of the same element or of elements close to each other in the periodic table. Occurs primarily between non-metals; however, it can also be observed between non- metals and metals. If atoms have similar electro-negativities (the same affinity for electrons), covalent bonds are most likely to occur. the metal loses electrons to become a positively charged cation, whereas the non-metal accepts those electrons to become a negatively charged anion. Because both atoms have the same affinity for electrons and neither has a tendency to donate them, they share electrons in order to achieve octet configuration and become more stable. require an electron donor, often a metal, and an electron acceptor, a non-metal. Covalent bonds include interactions of the sigma and pi orbitals; therefore, covalent bonds lead to formation of single, double, triple, and quadruple bonds. By losing those electrons, these metals can achieve noble gas configuration and satisfy the octet rule. Similarly, non-metals that have close to 8 electrons in their valence shells tend to readily accept electrons to achieve noble gas configuration. In ionic bonding, more than 1 electron can be donated or received to satisfy the octet rule. The charges on the anion and cation correspond to the number of electrons donated or received. In ionic bonds, the net charge of the compound must be zero. the ionization energy of the atom is too large and the electron affinity of the atom is too small for ionic bonding to occur. For example: carbon does not form ionic bonds because it has 4 valence electrons, half of an octet. To form ionic bonds, Carbon molecules must either gain or lose 4 electrons. This is highly un-favourable; therefore, carbon molecules share their 4 valence electrons through single, double, and triple bonds so that each atom can achieve noble gas configurations.
- 6. Q5. Why the ionic bonds usually result in the form of crystalline solids, not in molecules? • Ionic compounds form crystals that are composed of oppositely charged ions: a positively charged cation and a negatively charged anion. • It takes a lot of energy to overcome ionic bonds due to the strong attraction between opposite charges. • Ionic compounds have very high melting points, often between 300 and 1,000 degrees Celsius (572 to 1,832 degrees Fahrenheit). • Most ionic compounds can be dissolved in water, forming a solution of free ions that will conduct electricity. • They may be simple binary salts like sodium chloride (NaCl), or table salt, where one atom of a metallic element (sodium) is bonded to one atom of a non-metallic element (chlorine). • They may also be composed of polyatomic ions such as NH4NO3 (ammonium nitrate). Polyatomic ions are groups of atoms that share electrons (called covalent bonding) and function in a compound as if they constituted a single charged ion.
- 7. Q6. Label the following diagrams • Crystal name: Salt • Bonding type: Ionic bond • Crystal name: Diamond • Bonding type: Covalent bond
- 8. Q7. Complete the following table to differentiate the four types of crystalline solids. TYPE NATURE OF THE BOND EXAMPLE PROPERTIES Ionic - Made up of positive and negative ions - Held together by electrostatic forces of attractions Table salt, NaCl - very high melting points and brittleness - poor conductors in the solid state Covalent - Sharing of electron pairs between atoms. CO2 - carbon dioxide - Low melting and boiling points - Poor electrical conductors in all phases - Many soluble in nonpolar liquids but not in water Metallic - Interaction between the delocalised electrons and the metal nuclei. Transition metal - conduct heat - conduct electricity - generally high melting and boiling points - Strong - malleable (can be hammered or pressed out of shape without breaking) - ductile (able to be drawn into a wire) - metallic lustre - opaque (reflect light) Molecular - held together by the van der Waals forces hydrocarbons, ice, sugar, fullerenes, sulphur and solid carbon dioxide. - Soft - have relatively low melting temperature - Pure molecular solids are electrical insulators but they can be made conductive by doping.
- 9. Q8. You are given two solids of almost identical appearance, one of which is held together by ionic bonds and the other by van der Waals bonds. How could you tell them apart? 1. Melting and boiling point • The attraction between opposite ions is very strong. A lot of kinetic energy is thus required to overcome them and the melting point and boiling point of ionic compounds is very high. • In the liquid state, the ions still retain their charge and the attraction between the ions is still strong. Much more energy is required to separate the ions completely and the difference between the melting and boiling point is thus large.The higher the charge on the ions, and the smaller they are, the stronger the attraction between them will be and the higher the melting and boiling points. In MgO, the ions have a 2+ and 2- charge and thus the attraction between them is stronger than in NaCl, so the melting and boiling points are higher.
- 10. 2. Electrical Conductivity • Since ionic solids contain ions, they are attracted by electric fields and will, if possible, move towards the electrodes and thus conduct electricity. In the solid state, however, the ions are not free to move since they are tightly held in place by each other. Thus ionic compounds do not conduct electricity in the solid state. Ionic solids are thus good insulators. • In the liquid state, the ions are free to move and so can move towards their respective electrodes. Thus ionic compounds can conduct electricity in the liquid state. 3. Mechanical properties • Since ions are held strongly in place by the other ions, they cannot move or slip over each other easily and are hence hard and brittle.
- 11. 4. Van der Waals • The number of electrons in the molecule • The greater the number of electrons in a molecule, the greater the likelihood of a distortion and thus the greater the frequency and magnitude of the temporary dipoles. Thus the Van der Waal's forces between the molecules are stronger and the melting and boiling points are larger. • Surface area of the molecules: The larger the surface area of a molecule, the more contact it will have with adjacent molecules. Thus the greater its ability to induce a dipole in an adjacent molecule and the greater the Van der Waal's forces and melting and boiling points. • van der waals solid will be softer and will melt at a much lower temperature
- 12. Q9. Why can metals be deformed with relative ease whereas covalent and ionic solids are quite brittle? • Metals are characterized by their ability to reflect light, called lustre, their high electrical and thermal conductivity, their high heat capacity, and their malleability and ductility. Every lattice point in a pure metallic element is occupied by an atom of the same metal. The packing efficiency in metallic crystals tends to be high, so the resulting metallic solids are dense, with each atom having as many as 12 nearest neighbours. • Bonding in metallic solids is quite different from the bonding in the other kinds of solids we have discussed. Because all the atoms are the same, there can be no ionic bonding, yet metals always contain too few electrons or valence orbitals to form covalent bonds with each of their neighbours. Instead, the valence electrons are delocalized throughout the crystal, providing a strong cohesive force that holds the metal atoms together.
- 13. • The strength of metallic bonds varies dramatically. For example, cesium melts at 28.4°C, and mercury is a liquid at room temperature, whereas tungsten melts at 3680°C. • Metallic bonds tend to be weakest for elements that have nearly empty (as in Cs) or nearly full (Hg) valence subshells, and strongest for elements with approximately half-filled valence shells (as in W). • As a result, the melting points of the metals increase to a maximum around group 6 and then decrease again from left to right across the d block. • Other properties related to the strength of metallic bonds, such as enthalpies of fusion, boiling points, and hardness, have similar periodic trends.Valence electrons in a metallic solid are delocalized, providing a strong cohesive force that holds the atoms together. • Atoms in metal can be readily rearranged in position because bonding occurs by means of a sea of freely moving electrons
- 14. Q10. Van der Waals forces hold inert-gas atoms together to form solids, but they cannot hold such atoms together to form molecules in the gaseous state. Why not? • Van der Waals forces hold inert gas atoms together to form solids but they cannot hold such atoms together forming molecule in gaseous state because such forces are too weak to hold inert gas molecules together against forces exerted during collision in gaseous state.
- 15. Q11. Do the silicon and diamond are transparent to the visible light? Why? State their forbidden band gap. • Silicon is opaque because its electrons readily absorb photons of visible light and enter the upper energy band. • Diamond is transparent because photons of visible light do not have enough energy for electrons to absorb them and enter the upper band. • The forbidden band in silicon is 1.1 eV wide and in diamond it is 6 eV.