Lecture 2 (Atomic Structure) - PDF

Second Lecture Atomic Structure and Interatomic Bonding DR. MOHAMED EBEED WHY STUDY Atomic Structure and Interatomic Bonding? to explain a material’s properties.  1. Atomic Structure Each atom consists of a very small nucleus (protons and neutrons), which is encircled by moving electrons....

Second Lecture Atomic Structure and Interatomic Bonding DR. MOHAMED EBEED WHY STUDY Atomic Structure and Interatomic Bonding? to explain a material’s properties.  1. Atomic Structure Each atom consists of a very small nucleus (protons and neutrons), which is encircled by moving electrons. The protons have a positive electric charge (1.602 10-19 C)  The electrons have a negative electric charge The neutrons have no electric charge More than 99.94% of an atom's mass is in the nucleus The protons and neutrons have approximately The same mass is 1.67 10-27 kg The electron mass is 9.11 10-31 kg  1. Atomic Structure Atomic number (Z) the number of protons in the nucleus. This atomic number ranges in integral units from 1 for hydrogen to 92 for uranium. The atomic mass (A) of a specific atom may be expressed as the sum of the masses of protons and neutrons within the nucleus. Isotopes (‫ )النظائر‬are variants of a particular chemical element which differ in neutron number. For example, carbon-12, carbon-13 and carbon-14 are three isotopes of the element carbon with mass numbers 12, 13 and 14 respectively. The atomic number of carbon is 6, which means that every carbon atom has 6 protons, so that the neutron numbers of these isotopes are 6, 7 and 8 respectively 1. Atomic Structure Although the number of protons (Z) is the same for all atoms of a given element, the number of neutrons (N) may be variable. Atomic mass A ≈ Z + N The atomic mass unit (amu) may be used to compute atomic weight. atomic mass unit = amu = 1/12 mass of 12C  1 amu / atom = 1 g / mol 1mol of substance = 6.022 x 1023 molecules or atoms For example, the atomic weight of iron is 55.85 amu/atom, or 55.85 g/mol  2. Bohr atomic model The Bohr model shows that the electrons in atoms are in orbits of differing energy around the nucleus (think of planets orbiting around the sun). Bohr used the term energy levels (or shells) to describe these orbits of differing energy. He said that the energy of an electron is quantized, meaning electrons can have one energy level or another but nothing in between. The energy level an electron normally occupies is called its ground state. But it can move to a higher-energy, less-stable level, or shell, by absorbing energy. This higher- energy, less-stable state is called the electron’s excited state. After it’s done being excited, the electron can return to its original ground state by releasing the energy it has absorbed, as shown in the diagram below. 2. Bohr atomic model  2. Bohr atomic model Bohr's theory was successful in that: 1. It provided a physical model of the atom, whose internal energy levels matched those of the observed hydrogen spectrum. 2. It accounted for the stability of atoms. 3. It applied equally well to other one electron atoms such as a singly ionized helium ion. Bohr's theory failed in that: 1. It broke down when applied to many electron atoms, because it took no account of the interactions between electrons in orbit. 2. With the development of more precise spectroscope techniques, it became apparent that each of the excited states was not a unique single energy level, but a group of finely separated levels 3. Wave-mechanical model Quantum Mechanical model 1. -Electrons are NOT in circular orbits around nucleus. 2. -Electrons are in a 3-D region around the nucleus called atomic orbitals. 3. -The atomic orbital describes the probable location of the electron 4. The quantum mechanical model describes the probable location of electrons in atoms by describing: - Principal energy level [n] -------------- (1,2,3,4,5,6,7) - Energy sublevel (angular moment quantum number) [L]----(0 =s,1 =p,2= d,3 =f) - Orbital in each sublevel (ml) (magnetic quantum number)---- (-L,L) L= 0 (ml =0) L= 1 (ml =-1,0,1) ) L= 2 (ml =-2,-1,0,1,2) L= 3 (ml =-3,-2,- 1,0,1,2,3) - Spin (ms) = (-1/2,1/2)  4. ELECTRONS IN ATOMS Principal Energy Level (n) (principal quantum number) "shells" Indicates the relative size and energy of atomic orbitals. n=integers: n= 1, 2, 3, etc. As n increases: > orbital becomes larger > electron spends more time farther away from nucleus  atom's energy level increases The general formula is that the nth shell can in principle hold up to 2(n2) 2(n2) electrons The shells are labeled K, L, M, N, O, P, and Q; or 1, 2, 3, 4, 5, 6, and 7  4. ELECTRONS IN ATOMS Energy sublevel (l) (The second quantum number) Principal energy levels are broken down into sublevels. Sublevels define the orbital shape (s, p, d, f) > n=1, 1 sublevel (s) > n=2, 2 sublevels (s, p) > n=3, 3 sublevels (s, p, d) > n=4, 4 sublevels (s, p, d, f)  4. ELECTRONS IN ATOMS Orbitals (in each sublevel) (third quantum number) ml Each sublevel has a different number of orbitals. s: 1 orbital p: 3 orbitals d: 5 orbitals f: 7 orbitals Forth quantum number ms: spin moment of an electron Associated with each electron is a spin moment, which must be oriented either up or down 4. ELECTRONS IN ATOMS 4. Electron configurations Electron configurations are a simple way of writing down the locations of all of the electrons in an atom. Aufbau Principle– Electrons enter orbit alsof lowest energy first.  4. ELECTRONS IN ATOMS  the valence electrons are those that occupy the outermost shell. These electrons are extremely important; as will be seen, they participate in the bonding between atoms to form atomic and molecular aggregates. Furthermore, many of the physical and chemical properties of solids are based on these valence electrons.  some atoms have what are termed stable electron configurations; that is, the states within the outermost or valence electron shell are completely.  Normally this corresponds to the occupation of just the s and p states for the outermost shell by a total of eight electrons, as in neon, argon, and krypton; one exception is helium, which contains only two 1s electrons.These elements (Ne, Ar, Kr, and He) are the inert, or noble, gases, which are virtually unreactive chemically. 4. ELECTRONS IN ATOMS have discrete energy states tend to occupy lowest available energy state 4. ELECTRONS IN ATOMS Stable electron configurations... have complete s and p subshells tend to be unreactive. 4 SURVEY OF ELEMENTS Most elements: Electron configuration not stable. Electron configuration 1s1 1s2 (stable) 1s22s1 1s22s2 1s22s22p1 Adapted from Table 2.2, 1s22s22p2 Callister 6e.... 1s22s22p6 (stable) 1s22s22p63s1 1s22s22p63s2 1s22s22p63s23p1... 1s22s22p63s23p6 (stable)... 1s22s22p63s23p63d104s246 (stable) Why? Valence (outer) shell usually not filled completely. 5  5. THE PERIODIC TABLE  All the elements have been classified according to electron configuration in the periodic table.  the elements are situated, with increasing atomic number, in seven horizontal rows called periods. The arrangement is such that all elements arrayed in a given column or group have similar valence electron structures.  The elements positioned in Group 0, the rightmost group, are the inert gases, noble gases  electropositive elements, indicating that they are capable of giving up their few valence electrons to become positively charged ions  electronegative elements, they readily accept electrons to form negatively charged ions, or sometimes they share electrons with other atoms ELECTRONEGATIVITY Ranges from 0.7 to 4.0, Large values: tendency to acquire electrons. fr is the least in negativity with 0.7 f is the largest with 4 Smaller electronegativity Larger electronegativity Adapted from Fig. 2.7, Callister 6e. (Fig. 2.7 is adapted from Linus Pauling, The Nature of the Chemical Bond, 3rd edition, Copyright 1939 and 1940, 3rd edition. Copyright 1960 by Cornell University. 7 it's the amount of energy released when an atom gains an electron. ionization electron affinity always UP and RIGHT atomic radius always DOWN and LEFT 6. MOLECULES Many of the common molecules are composed of groups of atoms that are bound together by strong covalent bonds; these include elemental diatomic molecules (F2, O2, H2, etc.) as well as a host of compounds (H2O, CO2, HNO3, C6H6, CH4, etc.). In the condensed liquid and solid states, bonds between molecules are weak secondary ones. Consequently, molecular materials have relatively low melting and boiling temperatures. Most of those that have small molecules composed of a few atoms are gases at ordinary, or ambient, temperatures and pressures. On the other hand, many of the modern polymers, being molecular materials composed of extremely large molecules, exist as solids 7. BONDING FORCES AND ENERGIES Attractive force and repulsive force FN = FA + FR Equilibrium State FN = FA + FR = 0 Potential Energy E = ∫F drEN = EA + ER  7. BONDING FORCES AND ENERGIES A. Primary interatomic bonds (1) Ionic Bonding  It is always found in compounds that are composed of both metallic and nonmetallic elements.  Atoms of a metallic element easily give up their valence electrons to the nonmetallic atoms  In the process all the atoms acquire stable or inert gas configurations and, in addition, an electrical charge; that is, they become ions  Large difference in electronegativity required (>2) 7. BONDING FORCES AND ENERGIES (1) Ionic Bonding Example Example 1: NaCl 7. BONDING FORCES AND ENERGIES (1) Ionic Bonding Example Example 2: CaCl2 7. BONDING FORCES AND ENERGIES the attractive energy EA and the repulsion energy ER are function of the interatomic distance according to 𝐴 𝐵 1 𝐸𝐴 = 𝐸𝑅 = 𝑤ℎ𝑒𝑟𝑒, 𝐴 = 𝑍1 𝑒 𝑍2 𝑒 𝑟 𝑟𝑛 4𝜋𝜖0 where is the permittivity of a vacuum (8.85 10-12 F/m), Z1 and Z2 are the valences of the two ion types, and e is the electronic charge (1.602 10-19 C). In these expressions, A, B, and n are constants whose values depend on the particular ionic system. The value of n is approximately 8. https://sciencenotes.org/valences-of-the-elements/  7. BONDING FORCES AND ENERGIES (2) Covalent Bonding  A covalent bond, also called a molecular bond, is a chemical bond that involves the sharing of electron pairs between atoms. These electron pairs are known as shared pairs or bonding pairs.  A covalent bond is constructed If the electronegativity is less than 1.7.  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. Furthermore, this type of bonding is found in elemental solids such as diamond (carbon), silicon, and germanium and other solid compounds composed of elements that are located on the right-hand side of the periodic table, such as gallium arsenide (GaAs), indium antimonide (InSb), and silicon carbide (SiC).  7. BONDING FORCES AND ENERGIES (2) Covalent Bonding  A covalent bond, also called a molecular bond, is a chemical bond that involves the sharing of electron pairs between atoms. These electron pairs are known as shared pairs or bonding pairs.  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. Furthermore, this type of bonding is found in elemental solids such as diamond (carbon), silicon, and germanium and other solid compounds composed of elements that are located on the right-hand side of the periodic table, such as gallium arsenide (GaAs), indium antimonide (InSb), and silicon carbide (SiC). 7. BONDING FORCES AND ENERGIES (2) Covalent Bonding Examples  7. BONDING FORCES AND ENERGIES (2) Covalent Bonding There are three types of covalent bond depending upon the number of shared electron pairs. (1) Single covalent bond A covalent bond formed by the mutual sharing of one electron pair between two atoms is called a "single covalent bond.“ (2) double covalent bond A covalent bond formed between two atoms by the mutual sharing of two electron pairs is called a "double covalent bond“ (3) triple covalent bond A covalent bond formed by the mutual sharing of three electron pairs is called a "Triple covalent bond".  7. BONDING FORCES AND ENERGIES (4) Polar covalent bond A covalent bond formed between two different atoms is known as Polar covalent bond. For example when a Covalent bond is formed between H and Cl , it is polar in nature because Cl is more electronegative than H atom. Therefore, electron cloud is shifted towards Cl atom. Due to this reason a partial -ve charge appeared on Cl atom and an equal +ve charge on H ------() (5) Non-polar bond A covalent bond formed between two like atoms is known as Non-polar bond. Since difference of electro negativity is zero therefore, both atoms attract electron pair equally and no charge appears on any atom and the whole molecule becomes neutral. 7. BONDING FORCES AND ENERGIES (3) Metallic Bonding  Metallic bonding, the final primary bonding type, is found in metals and their alloys.  Metallic bonding is a type of chemical bonding that arises from the electrostatic attractive force between conduction electrons (in the form of an electron cloud of delocalized electrons) and positively charged metal ions. It may be described as the sharing of free electrons among a lattice of positively charged ions (cations).  Metallic bonding is found in the periodic table for Group IA and IIA elementsand, in fact, for all elemental metals  7. BONDING FORCES AND ENERGIES B. Secondary bonding or van der waals bonding  Secondary, van der Waals, or physical bonds are weak in comparison to the primary or chemical ones; bonding energies are typically on the order of only 10 kJ/mol (0.1 eV/atom).  Secondary bonds are weak in comparison to primary bonds  They are found in most materials, but their effects are often overshadowed by the strength of the primary bonding.  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.  7. BONDING FORCES AND ENERGIES Hydrogen bonding  Hydrogen bonding, a special type of secondary bonding, is found to exist between some molecules that have hydrogen as one of the constituents  hydrogen bond is a partially electrostatic attraction between a hydrogen (H) which is bound to a more electronegative atom such as nitrogen (N), oxygen (O), or fluorine (F), and another adjacent atom bearing a lone pair of electrons 7. BONDING FORCES AND ENERGIES 7. Important problems Problem 1 7. Important problems Problem 2