Structure of Matter PDF

Chapter One The structure of matter Prepared by Dr. Mohamed Hasabelnaby Lecturer of physics, School of Allied Health Sciences, Badr University. 1.1 The Structure of Matter All the objects around us whether living or non-living are matt...

Chapter One The structure of matter Prepared by Dr. Mohamed Hasabelnaby Lecturer of physics, School of Allied Health Sciences, Badr University. 1.1 The Structure of Matter All the objects around us whether living or non-living are matter. Water we drink, food we eat, air we breathe, chair we sit on, are all examples of matter. Matter is anything that has mass and takes up space. Matter appears in a huge variety of forms such as rocks, trees, computer, clouds, people, etc. Matter embraces each and everything around us. Therefore, in order to understand the world, it would be necessary to understand the matter. Solids, liquids and gases are made of atoms; in recent year’s direct visual evidence of atoms have been obtained by scientists using powerful microscopes (e.g. atomic force microscope). For details of the internal structure of atoms these microscopes are not yet sufficiently powerful. Atom is the smallest particle of an element that has all the properties of that element. Atoms are the basic building blocks of matter that make up everyday objects. -10 As we know from physics, the radii of atoms are approximately 1–5 ×10 m, and the -24 -21 mass of a single atom of an element is in the range 10 to 10 g, depending upon the element chosen. At the center of each atom we find the nucleus, a small object which radius -15 about (1–10 ×10 m) that contains almost all the mass of the atom (Fig. 1.1). All atoms are found to contain three basic particles, protons, electrons and neutrons. Protons have a positive electrical charge; electrons have a negative charge of identical magnitude to the proton, since the overall charge of the atom is zero. The mass of a proton is slightly less than that of the neutron, while the mass of the electron is negligible compared with the proton mass. Neutrons are uncharged particles with masses approximately equal to the mass of a proton; the table (1.1) gives the relative charge and mass for each particle. Why study structure and properties of matter? The characteristics of materials are functions of their atomic and molecular structures. Knowledge enables appropriate choice of materials to perform certain functions. The bulk properties of matter depend on structure at the atomic, molecular, microscopic, macroscopic levels.  Figure (1.1): The structure of atom. 1.1.1 Volume of Atoms Accounting for the sizes of protons, neutrons, and electrons, most of the volume of an atom greater than 99% is in fact empty space. Despite all this empty space, solid objects do not just pass through one another. The electrons that surround all atoms are negatively charged and cause atoms to repel one another, preventing atoms from occupying the same space. These intermolecular forces prevent you from falling through an object like your chair. Table (1.1): The relative charge and mass for atomic particles. 1.1.2 Atomic Particles Protons  Protons are positively charged subatomic particles. The charge of a proton is -19 1e, which corresponds to approximately 1.602 × 10 C. -27  The mass of a proton is approximately 1.672 × 10 g.  Protons are over 1800 times heavier than electrons.  The total number of protons in the atoms of an element is always equal to the atomic number of the element. Neutrons  The mass of a neutron is almost the same as that of a proton i.e. 1.674×10-27 g.  Neutrons are electrically neutral particles and carry no charge.  Different isotopes of an element have the same number of protons but vary in the number of neutrons present in their respective nuclei. Electrons -19  The charge of an electron is -1e, which approximates to -1.602 × 10 C. -31  The mass of an electron is approximately 9.1 × 10 g.  Due to the relatively negligible mass of electrons, they are ignored when calculating the mass of an atom. Atomic Mass -27 Protons and neutrons have approximately the same mass, about 1.67 × 10 grams. Scientists define this amount of mass as one atomic mass unit (amu). Although similar in mass, protons are positively charged, while neutrons have no charge. Therefore, the number of neutrons in an atom contributes significantly to its mass, but not to its charge. -31 Electrons are much smaller in mass than protons, weighing only 9.11 × 10 grams, or about 1/1800 of an atomic mass unit. Therefore, they do not contribute much to an element’s overall atomic mass. When considering atomic mass, it is customary to ignore the mass of any electrons and calculate the atoms mass based on the number of protons and neutrons alone (is the number of protons plus the number of neutrons). Atomic Number The atomic number is the number of protons in an element. Chemical properties of an element are determined by the atomic number Z of the nucleus. 1.2 Chemical Bonding and Molecular Structure A chemical bond is an attraction between atoms that allows the formation of chemical substances that contain two or more atoms. The bond is caused by the electrostatic force of attraction between opposite charges, (between electrons and nuclei). When two atoms approach each other closely enough for their electron clouds to interpenetrate, the electrons of one atom repel the electrons of the other, and same happens with the nuclei of the two atoms. At that time each atom’s electron are attracted to each other atom’s nucleus, if the atoms come closer, the attractive forces can offset the repulsive forces , the energy of the two atoms decreases and a bond is formed. That is the atoms are stick together. When the energy decreases is small, we have bonds called van der Waals bonds, when the energy decreases is larger, then we speak about chemical bonds. Thus, we can define chemical bonding as follows: “The attractive force which holds various constituents (atom, ions, etc.) together and stabilizes them by the overall loss of energy.” The stronger the chemical bonding between the constituents, the more stable the resulting compound would be. The opposite also holds true; if the chemical bonding between the constituents is weak, the resulting compound would lack stability and would easily undergo another reaction to give a more stable chemical compound (containing stronger bonds). That’s the atoms combine together to lose their energy and find stability. Why form chemical bonds? The basic answer is that atoms are trying to reach the most stable (lowest-energy) state that they can. Many atoms become stable when their valence shell is filled with electrons or when they satisfy the octet rule (by having eight valence electrons). If atoms don’t have this arrangement, they’ll “want” to reach it by gaining, losing, or sharing electrons via bonds. Why do certain atoms combine while others do not? This is mainly because a compound forms only when there is an attractive force leading to the lowering of energy. On the other hand, in case of a repulsive force, we find an increase in overall energy of the system. Thus, we do not see the formation of any compounds. 1.2.1 Types of Chemical Bonds When atoms participate in chemical bonding and yield compounds, the stability of the resulting compound can be gauged by the type of chemical bonds it contains. The types of chemical bonds formed vary in strength and properties. There are 3 primary types of chemical bonds which are formed by atoms or molecules to yield compounds. These types of chemical bonds include:  Ionic Bonds  Covalent Bonds  Metallic Bonds These types of bonds in chemical bonding are formed from the loss, gain, or sharing of electrons between two atoms/molecules. 1.2.1.1 Ionic Bonding Ionic bonding is the complete transfer of valence electron between atoms. In ionic bonds, the metal loses electron to become a positively charged cation, whereas the nonmetal accepts those electron to become a negatively charged anion. Ionic bonds require an electron donor (often a metal) and an electron acceptor (a nonmetal). Ionic bonding is observed because metals have few electrons in their outer- most orbitals. By losing those electrons, these metals can achieve most stable configuration and satisfy the octet rule. Similarly, nonmetals that have close to 8 electrons in their valence shells tend to readily accept electrons to achieve, when the transfer of electrons occurs, an electrostatic attraction between the two ions of opposite charge takes place and an ionic bond is formed. the ionic compound are represented by common (NaCl) as shown in figure (1.2):. 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. Figure (1.2): Ionic Bonds (NaCl). Characteristic properties of ionic bonds are:  Heat resistant and insulators as solid.  In ionic solution, they dissociate into ions and conduct electricity.  Insoluble in organic solvent.  Basic bond for glasses and ceramic. 1.2.1.2 Covalent Bonding As the name suggests, covalent bonding involves the sharing (co, meaning joint) of valence (outer shell) electrons. The atoms involved in covalent bonding arrange themselves in order to achieve the greatest energetic stability. And the valence electrons are shared sometimes equally, and sometimes unequally – between neighboring atoms. This bonding occurs primarily between nonmetals; however, it can also be observed between nonmetals and metals. Elements having very high ionization energies are incapable of transferring electrons and elements having very low electron affinity cannot take up electrons. The atoms of such elements tend to share their electrons with the atoms of other elements or with other atoms of the same element in a way that both the atoms obtain octet configuration in their respective valence shell and thus achieve stability. One, two, or three pairs of electrons may be shared between atoms, resulting in single, double, or triple bonds, respectively. The more electrons that are shared between two atoms, the stronger their bond will be. The kinds of molecules of which we are built (fats, protein and other), the food we eat (carbohydrates), and the clothes we wear (cotton, wool and synthetics fibers) all consist of covalent bonding molecules. As an example of covalent bonding, let’s look at water. A single water molecule, H2O, consists of two hydrogen atoms bonded to one oxygen atom. Each hydrogen shares an electron with oxygen, and oxygen shares one of its electrons with each hydrogen as shown in figure (1.3):  Figure (1.3): Covalent bond (H2O). Characteristic properties of covalent bonds:  Very strong.  Insulators.  Basic bond for polymers (fat, fatty acid).  Resist inorganic solvent. 1.2.1.3 Metallic Bonding Metallic bonding is a type of chemical bonding that rises from the electrostatic attractive force between conduction electrons (in the form of an electron cloud of delocalized electrons) and positively charged metal ions. The force that holds atoms together in a metallic substance. Such a solid consists of closely packed atoms. In most cases, the outermost electron shell of each of the metal atoms overlaps with a large number of neighboring atoms. As a consequence, the valence electrons continually move from one atom to another and are not associated with any specific pair of atoms. In short, the valence electrons in metals are nonlocalized, capable of wandering relatively freely throughout the entire crystal. The atoms that the electrons leave behind become positive ions, and the interaction between such ions and valence electrons gives rise to binding force that holds the metallic crystal together. Characteristic properties of metallic bond  High electric conductivity.  High thermal conductivity.  Opaque because the free electrons may absorb light.  It leads to crystalline arrangement in metals. Figure (1.4): Metallic bonds. 1.3 The Structure of Solids Solid is one of the three main states of matter, along with liquid and gas. In a solid, the particles are packed closely together and are not free to move about within the substance. Molecular motion for the particles in a solid is confined to very small vibrations of the atoms around their fixed positions; therefore, solids have a fixed shape that is difficult to change. Solids also have a definite volume; that is, they keep their size no matter how you try to change them. A large majority of substances around us are solids. The distinctive features of solids are:  They have a definite shape.  They are rigid and hard.  They have fixed volume. These characteristics can be explained on the basis of following facts:  The constituent units of solids are held very close to each other so that the packing of the constituents is very efficient. Consequently, solids have high densities.   Since the constituents of solids are closely packed, it imparts rigidity and hardness to solids.  The constituents of solids are held together by strong forces of attraction. This results in their having definite shape and fixed volume. 1.3.1 Classification of Solids Solids are classified into categories on the basis of how the constituent particles that forms the solid are arranged:  Amorphous solids.  Crystalline solids. The two types of solids have different characteristics. 1.3.1.1 Amorphous Solids An amorphous solid is a substance whose constituents do not possess a regularity arrangement. The molecules are distributed at random without regularity or repetition in their arrangement so there is no specific form or shape in their structure. Important examples of amorphous solids are glass, wax and plastics. An amorphous solid does not have a definite melting point; instead, it melts gradually over a range of temperatures, because the bonds do not break all at once. This means an amorphous solid will melt into a soft, malleable state (think candle wax or molten glass) before turning completely into a liquid. 1.3.1.2 Crystalline Solids A crystalline solid is a substance whose constituents possess a regularity arrangement in a definite geometric pattern in three dimensions in what is called a crystal lattice or space lattice. The atoms may be held together by ionic bonds as in sodium chloride, covalent bonds as in diamond, or metallic bond as in metals. They have a sharp and definite melting point. 1.3.2 Types of Crystal Systems or crystal lattice Since solid materials are of many different shapes, it may appear at first sight that may be an infinite number of interfacial combinations. But this is not true. A careful examination of several thousand crystals of various substances reveals that there are only seven possible crystal symmetries exhibited by solids. Different solids which exhibit the same symmetry elements are all classified as belonging to the same system. The simplest way to study is to consider a unit cell; a unit cell is the smallest repeat unit in a crystal lattice (the building block of a crystal). The three dimensions arrangement of atoms, molecules or ions inside a crystal is called a crystal lattice. 1. Cubic system. All three axes are of equal length and are mutually perpendicular to each other. Some typical examples are diamond, silver and cesium chloride. 2. Tetragonal system. To produce tetragonal symmetry, the cube is elongated in one direction. The axes are still at right angle to each other, but one axis is longer (or shorter) than the other two axes. One example is lead tungsten. 3. Orthorhombic system. It is also called rhombic system. It consists of three mutually perpendicular axes of unequal lengths. For examples, Alpha-sulphur, ammonium sulphate. 4. Rhombohedral system. It consists of three equal axes which are inclined to each other at same angle but it is not 90°. Some common examples are arsenic, antimony, bismuth and calcite. 5. Monoclinic system. In monoclinic system the three axes are of unequal length and are no longer perpendicular to each other. Monoclinic system differs from rhombic symmetry in that one of axes does not make 90° with the plane of other two axes. Examples include potassium chloride and high temperature form of sulphur. 6. Triclinic system. It consists of three unequal axes and none is perpendicular to any of the others. This system has the lowest symmetry. There is no simple axis or plane of symmetry. Boric acid is an example of this class. 7. Hexagonal system. In this type of symmetry the atoms are arranged in the form of hexagons and these unit cell has two edges of one above length (a=b). The symmetry axis (c) is at 90° to these two axes which make an angle of 120° with one another. Graphite is a common example of hexagonal symmetry. 1.4 Atomic Packing Factor (APF) Some of the space of the structural is not occupied by the atoms. The fraction of space occupied by the atoms is called the atomic packing factor. Atomic Packing Factor (APF) tells you what percent of an object is made of atoms vs. empty space. You can think of this as a volume density, or as an indication of how tightly-packed the atoms are. Calculating the atomic packing factor for a crystal is simple: for some repeating volume, calculate the volume of the atoms inside and divide by the total volume. APF = Usually, this “repeating volume” is just the volume of the unit cell. The unit cell is defined as the simplest repeating unit in a crystal. For simple cubic system, it equals 0.54 which indicates that nearly 50% of the space is free. Materials having higher atomic packing factor usually have higher densities, stability and strength properties. 1.5 Correlation between Atomic Structure and Properties:- The properties of materials depend basically on the type of bonds which dominate the structure:- 1. Density is controlled by atomic weight, atomic radius and atomic packing factor. 2. Melting and boiling temperature can correlate with the strength of bond (inversely proportional). Stronger bonds need higher temperature to impart the energy for melting. 3. Thermal expansion of materials is inversely with their melting temperature i.e. higher melting temperature, the less the coefficient of thermal expansion? 4. The electrical and thermal conductivity are very dependent on the nature of atomic bonds e.g. thermal conductivity is higher in material with metallic bonds. 5. Strength can be primarily governed by the type of bond, although the arrangement of atoms controls the deformation and resistance to stress. 6. Crystalline structures have lower energy level while amorphous structures have higher energy due to irregular arrangement. 7. Atomic structures are generally stronger than molecular structure because primary bonds control the properties.