Chapter 2: Atomic Structure and Interatomic Bonding Lecture Notes PDF

EMJ 16302 Principles of Engineering Materials Chapter 2: Atomic Structure and Interatomic Bonding SEM 1 SA 2024/2025...

EMJ 16302 Principles of Engineering Materials Chapter 2: Atomic Structure and Interatomic Bonding SEM 1 SA 2024/2025 Prof. Ir. Ts. Dr. Hazry bin Hj. Desa Centre of Excellence for Unmanned Aerial Systems (COE-UAS) W2 – 23/10/2024 1. ATOMIC STRUCTURE  History of the Atom Democritus was a pioneering thinker who introduced the concept of atoms. Dalton’s Model: Dalton’s model tells us that atoms are the building blocks of matter, and they combine in specific ways to create everything around us! John Dalton introduced the idea that atoms have different masses and the concept of relative atomic masses. History of the Atom Thompson’s The plum pudding model was proposed by J.J. Thomson in 1904 after he discovered the electron. This model was an attempt to describe the structure of the atom, which, at that time, was not well understood.  History of the Atom The plum pudding model was replaced after Ernest Rutherford’s gold foil experiment (also known as the Rutherford experiment) in 1909 The Atom Atoms are made up of particles called protons, neutrons and electrons which are responsible for the mass and charge of atoms. The Structure of the Atom An atom is the smallest unit of matter that retains all of the chemical properties of an element. Atoms combine to form molecules, which then interact to form solids, gases, or liquids. For example, water is composed of hydrogen and oxygen atoms that have combined to form water molecules.  Atomic Particles Atom consist of three basic particles: protons, electrons and neutrons. The nucleus (center) of the atom contains the protons (positively charged) and neutrons (no charge). The outermost regions of the atom are called electron shells and contain the electrons (negatively charged). Atoms have different properties based on the arrangement and number of their basic particles. The hydrogen atom (H) contains only 1 proton, 1 electron and no neutrons. This can be determined using the atomic number and the mass number of the element. Properties are dictated by atomic structure Many properties of materials are determined by atomic structure particularly valence electrons Strong vs. week Brittle vs. ductile Electrical conductor vs. insulator Transparent vs. opaque We need to know: Types of atom present Types of bonds between atoms Packing of atoms Atom  Basics of Atomic Structure Atom consists of: ❑ Nucleus → proton (+) and neutrons. ❑Electrons (-) (m=9.11 X10-31 Kg) Atomic number, 𝒁 = # of protons or # of electrons of neutral species. Mass number, A = protons + neutrons in isotope. 1 proton or 1 neutron weighs = 1 amu Atomic weight (Atomic Mass) = This is the weighted average mass of all the isotopes of an element as they occur naturally, measured in atomic mass units (amu). 1 amu/atom = 1g/mol -Relative atomic mass- 1 atomic mass unit (amu) = 1 Dalton (Da) = 1.66053906660 × 10⁻²⁷ kg (1.66 X 10-24g) This value represents the mass of one-twelfth of a carbon-12 atom, which is the standard used for atomic mass units and Daltons. Atomic Number and Mass Numbers The atomic number is the number of protons in an element, while the mass number is the number of protons plus the number of neutrons.  Atomic Numbers and Mass Numbers Symbol of element Carbon as example Mass number = Protons + Neutrons Atomic number (Number of protons) Each chemical element: Characterized by number of protons (atomic number) (or Atomic Weight) Atomic Mass Protons and neutrons have approximately the same mass, about 1.67 X 10-24 grams. Scientists define this amount of mass as one atomic mass unit (amu) or 1 Dalton. 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. Electrons are much smaller in mass than protons, weighing only 9.11 X 10-28 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 atom’s mass based on the number of proton and neutrons alone. Electrons contribute greatly to the atom’s charge, as each electron has a negative charge equal to the positive charge of a proton. Scientist define this charges as “+1” and “-1”. In an uncharged, neutral atom, the number of electrons orbiting the nucleus is equal the the number of protons inside the nucleus. In these atoms, the positive and negative charges cancel each other out, leading to an atom with no net charge. Proton = 1.67 X 10-24 g (or 1 amu) Electron = 9.11 X 10-28 g ≈ 0 amu  Protons, Neutrons and Electrons Particles Charge Mass (amu) Location Proton +1 1 nucleus Neutron 0 1 nucleus Electron -1 0 orbitals Protons, neutrons and electrons: Both protons and neutrons have a mass of 1 amu and are found in the nucleus. However, protons have a charge of +1, and neutrons are uncharged. Electrons have a mass of approximately 0 amu, orbit the nucleus, and have a charge of -1. Atomic Numbers Neutral atoms of an element contain an equal number of protons and electrons. The number of protons determines and element’s atomic number (Z) and distinguishes one element from another. For example, carbon’s atomic number (Z) is 6 because it has 6 protons. The number of neutrons can vary to produce isotopes, which are atoms of the same element that have different numbers of neutrons. The number of electrons can also be different in atoms of the same element, thus producing ions (charged atoms). For instance, iron, Fe, can exist in its neutral state, or in the +2 and +3 ionic states.  Isotopes: atoms of some elements have two or more different atomic masses. (Atoms with the same number of protons but different number of neutrons) The atomic mass unit (amu): computations of atomic weight. Key Terms Atom: The smallest possible amount of matter which still retains its identity as a chemical element, consisting of a nucleus surrounded by electrons. Proton: Positively charged subatomic particle forming part of the nucleus of an atom and determining the atomic number of an element. It weighs 1 amu. Neutron: A subatomic particle forming part of the nucleus of an atom. It has no charge. It is equal in mass to a proton, or it weighs 1 amu. Key Points An atom is composed of 2 regions: the nucleus, which is in the center of the atom and contains protons and neutrons, and the outer region of the atom, which holds its electrons in orbit around the nucleus. Protons and neutrons have approximately the same mass, about 1.67 X 10-24 grams, which scientists define as one atomic mass unit (amu) or 1 Delton. Each electron has a negative charge (-1) equal to the positive charge of a proton (+1). Neutrons are uncharged particles found within the nucleus. 2. ATOMIC MODEL Bohr Atomic Model Niels Bohr made significant contributions to our understanding of atomic structure and quantum mechanics. His most notable achievements include the development of the Bohr model of the atom. Properties of Electrons under the Bohr Model In 1913, Bohr suggested that electrons could only have certain classical motions: Electrons in atoms orbit the nucleus. The electrons can only orbit stably, without radiating, in certain orbits (called by Bohr the “stationary orbits”) at a certain discrete set of distances from the nucleus. These orbits are associated with definite energies and are also called energy shells or energy levels. Electrons can only gain or lose energy by jumping from one allowed orbit to another, absorbing or emitting electromagnetic radiation with a frequency (𝜈) determined by the energy difference of the levels according to the Planck relation. Bohr Atomic Model The Bohr model represents an early attempt to describe electrons in atoms, in terms of both position (electron orbitals) and energy (quantized energy levels). electrons travel around the nucleus in specific circular orbits or shells, each orbit corresponding to a fixed energy level. Schematic representation of the Bohr atom These orbits are quantized, meaning the electron can only occupy certain allowed energy levels.  Wave-mechanical model Electron is to exhibit both wave-like 1926, by Erwin Schrödinger and particle-like characteristic. An electron is no longer treated as a particle moving in a discreate There is only a probability that orbital; an electron will be at a given rather, position is considered to be position. Orbitals are more closely the probability of an electron’s modeled by a electron cloud. being at various locations around the nucleus. Position is described the probability distribution or electron cloud. Instead, in the Bohr model, electrons are located at very discrete levels. However, in the wave-mechanical model, electrons are Comparison of the (a) Bohr and (b) wave-mechanical atom scattered within a range represented by the models in terms of electron distribution. electron cloud.  Quantum Numbers In wave mechanics, every electron in an atom is characterized by four parameters, called quantum numbers (𝒏, 𝒍, 𝒎𝒍 and 𝒎𝒔 ) to identify the energy & shape of boundary space/electron cloud, and the spin for any electron in any atom. The size, shape, and spatial orientation of an electron’s probability density are specified by three of these quantum numbers (𝒍, 𝒎𝒍 and 𝒎𝒔 ). Shells (𝒏) are specified by a principal quantum number n, which may take integral values beginning with unity (no. 1). Sometimes shells denoted as: K, L ,N , O …(n = 1,2,3,4,…) -Shell (𝑛)- Number of energy level  Electrons are both particles and waves. An orbital is a region of probability where an electron can be found. There are 𝒔, 𝒑, 𝒅 and 𝒇 sub-shells (orbital) with different shape. The s, p, d, and f subshells refer to the different types of Quantum number orbitals that electrons can occupy within each energy level or electron shell of an atom. 4 types of subshells Atomic Energy Level Bohr energy levels separate into electron subshells. 3rd energy level, 𝑛 = 3 𝒏=3 2nd energy level, 𝑛 = 2 𝒏=2 1st energy level, 𝑛 = 1 𝒏=1 Electron that occupies the 3rd energy level has more energy than electron that occupies in the 2nd energy level. + Electrons that closer to the nucleus have lower energy levels. Electrons that are farther from the nucleus have higher energy - levels. -  The terms s, p, d, and f refer to the sublevels or subshells of an atom where electrons are found. These sublevels are part of the structure of electron orbitals, which describe the regions around the nucleus where electrons are likely to be located. The letters correspond to specific types of orbitals, each with different shapes and energy levels. Number of energy level is equal to number of sub-shells 𝑛 = 1, 𝑠 subshell 1𝑠 𝑛 = 2, 𝑠 and 𝑝 subshells 2𝑠 2𝑝 𝑛 = 3, 𝑠, 𝑝 and 𝑑 subshells 3𝑠 3𝑝 3𝑑 𝑛 = 4, 𝑠, 𝑝, 𝑑 and 𝑓 subshells 4𝑠 4𝑝 4𝑑 4𝑓 s : The s sublevel is spherical in shape and can hold a maximum of 2 electrons. It is present in every energy level (n = 1, 2, 3,...). p : The p sublevel is dumbbell-shaped and can hold a maximum of 6 electrons. It appears starting from the second energy level (n = 2). d : The d sublevel has more complex shapes (clover-like) and can hold a maximum of 10 electrons. It appears starting from the third energy level (n = 3). f : The f sublevel has even more complex shapes and can hold a maximum of 14 electrons. It appears starting from the fourth energy level (n = 4). These sublevels are arranged in order of increasing energy and are filled according to the Aufbau principle, which governs how electrons occupy these orbitals in an atom.  Quantum Numbers Symbol Number Possible Values 𝑛 Principle quantum number – represent energy level of an electron. 1,2,3,4.... 𝑙 Angular momentum quantum number – represent the shape of orbital which 0,1,2,3.. (𝑛-1) is associated with 𝒔,𝒑,𝒅 and 𝒇. 𝑚𝑙 Magnetic quantum number – represent how many orbitals they are of a type −𝑙,…,-1,0,1,…, 𝑙 per energy level. (2𝑙 + 1) 𝑚𝑠 Spin quantum number – represent electron spin inside an orbital. +1/2, -1/2 No. of No. of 4 parameters of quantum numbers Sub-shells orbitals electrons 𝒏=3 𝒏=2 𝑙=0 𝒔 Spherical shape 1 2 𝒏=1 𝑙=1 𝒑 Dumbbell shape 3 6 + - 𝑙=2 𝒅 Clover leave shape 5 10 𝑙=3 𝒇 Unusual/complex shape 7 14  Quantum Numbers Symbol Number Possible Values 𝑛 Principle quantum number – represent energy level of an electron. 1,2,3,4.... 𝑙 Angular momentum quantum number – represent the shape of orbital which 0,1,2,3.. (𝑛-1) is associated with 𝒔,𝒑,𝒅 and 𝒇. 𝑚𝑙 Magnetic quantum number – represent how many orbitals they are of a type −𝑙,…,-1,0,1,…, 𝑙 per energy level. (2𝑙 + 1) 𝑚𝑠 Spin quantum number – represent electron spin inside an orbital. +1/2, -1/2 Each arrow in the box represents an electron, Orbital as and the direction of the arrow (up or down) Box indicates the electron's spin (clockwise or Sub-shells counterclockwise). Orbital Diagrams 𝒔 𝑙=0 𝑚𝑙 = 0 (1 - 𝑠 orbital per energy level) 𝒑 𝑙=1 𝑚𝑙 = -1, 0, 1 (3 - 𝑝 orbitals per energy level) 𝒅 𝑙=2 𝑚𝑙 = -2, -1, 0, 1, 2 (5 - 𝑑 orbitals per energy level) 𝒇 𝑙=3 𝑚𝑙 = -3, -2, -1, 0, 1, 2, 3 (7 – 𝑓 orbitals per energy level)  Table: Summary of Relationships among the Quantum Numbers 𝑛, 𝑙, 𝑚𝑙 , and Numbers of Orbitals and Electrons 𝒔 sublevel has 1 orbital, 𝒑 sublevel has 3 orbitals, 𝒅 sublevel has 5 orbitals, 𝒇 sublevel has 7 orbitals. Each "orbital" can hold a maximum of 2 electrons, so the number of orbitals in a sublevel determines the maximum number of electrons it can contain. Electron Configurations/𝑠,𝑝,𝑑,𝑓 notation Electron states: value of energy that are permitted for electrons. Pauli exclusion principle: (1) Each orbital can only hold 2 electrons maximum and (2) two electrons in the same orbital must have opposite spins. Thus, 𝑠, 𝑝, 𝑑 and 𝑓 subshells may each accommodate, respectively, a total of 2, 6, 10, 14 electrons. There is a hierarchy, i.e 𝑠 orbital will be filled before 𝑝 orbitals which will be filled before 𝑑 orbitals and so on. (𝑠 < 𝑝 < 𝑑 < 𝑓) (note, this is a general rule but there are exceptions) Valance electron: Electrons that occupy the outmost shell. Ground state: When all the electrons occupy the lowest possible energies. Table: Expected Electron Configurations for Some Common Elements (𝑠, 𝑝, 𝑑, 𝑓 notations) The Periodic Table All the elements has been classified according to electron configuration in the periodic table. The elements are situated with increasing atomic number, in 7 horizontal rows called periods. Most of the elements really come under the metal classification and sometimes termed electropositive elements. Electropositive elements: they are capable of giving up their few valence electrons to become positively charged ions. The elements situated on the right-hand side of the table are electronegative; readily accept electrons to form negatively charged ions, or sometimes they share electrons with other atoms.  The alkali and the alkaline earth Group Group 0: The metal, Groups IA and IIA Groups IIIA, IVA, and VA Group VIIA: Two inert gases, respectively, having one and two (B, Si, Ge, As etc.) display VIA: Two deficient which have electrons in outmost shell. characteristic that are electrons (halogen) filled electron intermediate between the deficient shells and metals and non-metals by stable electron virtue of their valence configuration. electron structures. Groups IIIB – Groups IIB: The transition metals, partially filled 𝑑 electron states and in some cases one or two electrons in the next higher energy shell. 3. ATOMIC BONDING IN SOLIDS 1. Bonding Forces and Energies. When two atoms are brought close together, several interactions occur between their nuclei and electrons. These interactions can result in the formation of a bond (attraction) or repulsion, depending on the distance between the atoms and the forces involved. At long distances: Weak attraction occurs between the atoms. When they are at long distance between 2 atoms..they can't see each other, so they are not attracted to each other. At intermediate distances: Attraction is stronger, and bonding may occur, resulting in a stable, lower-energy state. But, as you start to get them closer together, they start to interact, and they want to be closer and there is a force that pull them together. At very close distances: Repulsive forces dominate, increasing potential energy and preventing the atoms from getting any closer. They don’t want to get too close, if they get too close then they start repelling each other.  Atomic Bonding in Solids Atomic Bonding Forces and Atomic Bonding Energies. Why are atoms attracted to each other? It is due to Coulombic force. Coulombic force: the force created between positively charged & negatively charged entities. 𝐹𝑁 = 𝐹𝐴 + 𝐹𝑅 A force is attractive if the objects interacting exert a pull on each Net force Attractive force Repulsive force other, creating a tendency to move closer. On the other hand, a force is repulsive if the interacting objects push each other away.  -Interatomic Forces & Energy Curves- 𝐸𝑅 Repulsive Energy The distance between the 2 atoms center 𝐸 Equilibrium distance between 2 Repulsive (kJ/mol) atoms in a substance 𝒓𝟎 r (nm) 𝒓 𝑬𝑵 Attractive Net Energy 𝑬𝑨 Attractive Energy If change the 𝑟 , there are 2 different energies: Minimum energy level 𝑬𝟎 i) 𝑬𝑨 Attractive Energy When the energy is ii) 𝑬𝑹 Repulsive Energy +ve, the net force is repulsive, the atoms are Constant Number repelling each other. 𝐴 𝐹 When the energy is – 𝑬𝑨 = Attractive Force 𝑟 𝑭𝑨 ve, the net force is Attractive 𝐵 attractive, the atoms 𝑬𝑹 = 𝑛 value is different 𝑭𝑵 attract each other. 𝑟𝑛 based on type of 𝒓𝟎 bonding Net Force r 𝑬𝑵 = 𝑬𝑨 + 𝑬𝑹 Repulsive 𝐹𝑅 Repulsive Force At 𝑟0 = Net force (𝐹𝑁 ) equal to ZERO. Means that atom want to be in location where there is no pushing on each other and no pulling on each other.  Interatomic Forces & Energy Curves Bonding Forces and Energies. Considering the interaction between two isolated atoms as they are brought into close proximity from an infinite separation. At larger distance, the interaction are negligible. As the atom approach, each exerts forces on the other. ▪ Attractive ▪ Repulsive Ultimately, the outer electron shells of the two atoms begin overlap, and a strong repulsive force comes into play. Figure shows attractive, repulsive and net potential energies and forces as a function of interatomic separation of two atoms. Bonding Models Bonding holds atoms together to form solids materials. In solids, atoms are held at preferred distance from each other (equilibrium distances). Distance larger or smaller than equilibrium distances are not preferred. Equilibrium spacing 𝑟0 is approximately 0.3nm. Consequently, as atomic bonds are stretched, atoms tend to attract each other, and as the bond are compressed, atoms repel each other. Simple bonding models assume that the total bonding results from sum of two forces: an attractive force 𝐹𝐴 and repulsive force 𝐹𝑅. The net force 𝑭𝑵 = 𝑭𝑨 + 𝑭𝑹 The repulsive force dominates at small distance, and the attractive force at larger distance. At equilibrium they are just equal. When 𝑭𝑨 and 𝑭𝑹 balance, or become equal, there is no net force: 𝑭𝑨 + 𝑭𝑹 = 0 𝑬𝟎 correspond energy at minimum point, it represents the energy that would be required to separate two atoms to an infinite separation. 2. Primary Interatomic Bonds. Primary bonding: i. Ionics (transfer of valence electrons) ii. Covalent (sharing of valence electrons, directional) iii. Metallic (delocalization of valence electrons) i) Ionic Bonding Ionic bonding is always found in compounds that are composed of both metallic and non-metallic elements. Atoms of a metallic element easily give up their valence electrons to the non-metallic atoms. Chloride ion Sodium ion Example: Sodium Chloride (NaCl) Sodium loses 1 electron to form sodium ion (transfer to chlorine) They are oppositely charged ions and will be electrostatically attracted to one another. This electrostatic attraction is an ionic bond. Na (1s2 2s2 2p6 3s1) + Cl (1s2 2s2 2p6 3s2 3p5) Na+ (1s2 2s2 2p6) + Cl- (1s2 2s2 2p6 3s2 3p6) Na (1s2 2s2 2p6 3s1) + Cl (1s2 2s2 2p6 3s2 3p5) Na+ (1s2 2s2 2p6) + Cl- (1s2 2s2 2p6 3s2 3p6) It transfer one valence 3s electron to a chlorine atom. Then, the chlorine atom has a net negative ion. The attractive bonding force are coulombic; that is, positive and negative ions, by virtue of their net electrical charge, attract one another. Ionic bonding in sodium chloride (NaCl) ii) Covalent Bonding In covalent bonding, stable electron configurations are assumed by the sharing of electrons between adjacent atoms. Two atoms that are covalently bonded will each contribute at least one electron to the bond, and the shared electrons may be considered to belong to both atoms. The carbon atom has four valence electrons The four hydrogen atoms has a single valence electron Covalent bonding in a molecule of methane (CH4) Covalent Bonding The covalent bonding is directional; usually only few patterns of overlap are possible, consequently, only few spatial arrangements of atom are possible. Many nonmetallic elemental molecules (H2, Cl2, F2, etc.) as well as molecules containing dissimilar atoms, such as CH4, H2O, HNO3 and HF are covalently bonded. This type of bounding is found in elemental solids: Diamonds Silicon Germanium and Gallium arsenide (GaAs), indium antimonide (InSb), and silicon carbide (SiC)  iii) Metallic Bonding Metallic bonding is found in metals and their alloys. Metallic materials have one, two or at most three valance electrons. With this model, these valence electrons are not bound to any particular atom in the solid and are more or less free to drift throughout the entire metal. They may be thought of as belonging to the metal as a whole or forming a ”sea of electrons” or an “electron cloud”. The remaining non-valence electrons and atomic nuclei form called ion cores, which possess a net positive charge (Ion cores) equal in magnitude to the total valence electrons charge per atom. Each atom is surrounded by electrons in its outermost shell. These electrons are free to move. (Sea of valence electrons) Metallic Bonding  There is a strong electrostatic attraction between the positive ions and the negative electrons. These forces of attraction hold everything together in a regular structure to form a metallic bond. This gives the metal its overall strength. 3. Secondary Bonding or Van Der Waals Bonding Secondary, van der Waals, or physical bonds are weak; bonding energies are typically on the order of only 10kJ/mol (0.1 eV/atom). Secondary bonding is evidenced for the inert gases, which have stable electron structures, and, in addition, between molecules in molecular structures that are covalently bonded. Secondary bonding forces arise from atomic or molecular dipoles. In essence, an electric dipole exists whenever there is some separation of positive and negative portions of an atom or molecule. Dipole: Two equal and opposite charges separated by some distance Van Der Waals Bonding between two dipoles  The bonding results from the coulombic attraction between the positive end of one dipole and the negative region of an adjacent one.  SUMMARY Electrons in ▪ The two atomic models are Bohr and wave mechanical. Whereas the Bohr model assumes electrons to be Atoms particles orbiting the nucleus in discrete paths, in wave mechanics we consider them to be wavelike and treat electron position in terms of a probability distribution. ▪ The energies of electrons are quantized—that is, only specific values of energy are allowed. ▪ The four electron quantum numbers are 𝑛, 𝑙, 𝑚𝑙 and 𝑚𝑠. They specify, respectively, electron orbital size, orbital shape, number of electron orbitals, and spin moment. ▪ According to the Pauli exclusion principle, each electron state can accommodate no more than two electrons, which must have opposite spins. Bonding Force ▪ Bonding force and bonding energy are related to one another. and Energy ▪ Attractive, repulsive, and net energies for two atoms or ions depend on interatomic separation per the schematic plot of Figure page 38 & 39. Primary ▪ For ionic bonds, electrically charged ions are formed by the transference of valence electrons from one atom Interatomic type to another. Bonds ▪ There is a sharing of valence electrons between adjacent atoms when bonding is covalent. ▪ With metallic bonding, the valence electrons form a “sea of electrons” that is uniformly dispersed around the metal iron cores and acts as a form of glue for them. Secondary ▪ Relatively weak van der Waals bonds result from attractive forces between electric dipoles, which may be Bonding or van induced or permanent. der Waals Bonding END of Chapter 2