CHAPTER 2 1/1/2016. Atomic Structure. The Periodic Table Columns: Similar Valence Structure. Atomic Structure and Bonding

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1 inert gases 1/1/016 APTER Periodic Table Atomic Structure and Bonding 1 Source: Davis, M. and Davis, R., Fundamentals of hemical Reaction Engineering, McGraw-ill, 003. The Periodic Table olumns: Similar Valence Structure Atomic Structure give up 1e - give up e - Li Be Na Mg give up 3e - K a Sc Rb Sr Y accept e - accept 1e - O Te I e F Ne Adapted from S l Ar Fig..6, allister & Se Br Kr Rethwisch 8e. Xe Valence electrons determine all of the following properties 1) hemical ) Electrical 3) Thermal 4) Optical s Ba Po At Rn Fr Ra Electropositive elements: Readily give up electrons to become + ions. Electronegative elements: Readily acquire electrons to become - ions. hapter - 3 hapter - 4 opyright The The McGraw-ill ompanies, Inc. Permission required for reproduction or display or display opyright The The McGraw-ill ompanies, Inc. Permission required for reproduction or display or display istory of Atom Rutherford Experiment 17 th century: Robert Boyle asserted that elements are made up of simple bodies which themselves are not made up of any other bodies. 19 th century: John Dalton stated that matter is made up of small particles called atoms. 19 th century: enri Becquerel and Marie and Pierre urie in France, introduced the concept of radioactivity. Joseph J. Thompson found electrons thru cathode ray tube. In 1910: Ernest Rutherford found protons. In 193: James hadwick found neutrons. Rutherford experiment: Discovery of proton

2 Nucleus Diameter : m Accounts for almost all mass Positive harge Proton Mass : x 10 4 g harge : 1.60 x Structure of Atoms ATOM Basic Unit of an Element Diameter : m. Neutrally harged Neutron Mass : x 10 4 g Neutral harge Electron loud Mass : x 10 8 g harge : x 10 9 Accounts for all volume Atomic Structure (Freshman hem.) atom electrons 9.11 x kg protons neutrons } 1.67 x 10-7 kg atomic number = # of protons in nucleus of atom = # of electrons of neutral species A [=] atomic mass unit = amu = 1/1 mass of 1 Atomic wt = wt of 6.0 x 10 3 molecules or atoms 1 amu/atom = 1g/mol e etc. 7 hapter - 8 Atomic Number and Atomic Mass Example Problem 9 Atomic Number = Number of Protons in the nucleus Unique to an element Example :- ydrogen = 1, Uranium = 9 Relative atomic mass = Mass in grams of 6.03 x 10 3 (Avogadro Number) Atoms. The mass number (A) is the sum of protons and neutrons in a nucleus of an atom. Example :- arbon has 6 Protons and 6 Neutrons. A= 1. One Atomic Mass unit is 1/1 th of mass of carbon atom. One gram mole = Gram atomic mass of an element. Isotope: Variations of element with same atomic number but different mass number. One gram Mole of arbon 1 Grams Of arbon 6.03 x 10 3 arbon Atoms 10 A 100 gram alloy of nickel and copper consists of 75 wt% u and 5 wt% Ni. What are percentage of u and Ni atoms in this alloy? Given:- 75g u Atomic Weight g Ni Atomic Weight g Number of gram moles of u = mol g/mol 5g Number of gram moles of Ni = mol g/mol Atomic Percentage of u = % ( ) Atomic Percentage of Ni = % ( ) 11 Example problem An intermetallic compound has the chemical formula NixAly, where x and y are simple integers, and consists of 4.04 wt% nickel and wt% aluminum. What is the simplest formula of this nickel aluminide? No. of moles of Ni = 4.04 g Ni / 1 mol Ni /58.71 g Ni = mol No. of moles of Al = g Al / 1 mol Al /6.98 g Al =.148 mol total =.864 mol mole fraction of Ni = /.864 = 0.5 mole fraction of Al =.148 /.864 = 0.75 The simplest formula is Ni 0.5 Al or NiAl 3. 1 Planck s Quantum Theory Max Planck, discovered that atoms and molecules emit energy only in certain discrete quantities, called quanta. James lerk Maxwell proposed that the nature of visible light is atom emits energy in the form of electromagnetic radiation. E = hυ = hc/λ is always released in integer multiples of hυ

3 Electron Structure of Atoms: Bohr s Theory (1913) Electron rotates at definite energy levels. is absorbed to move to higher energy level. is emitted during transition to lower level. change due to transition = ΔE = hc Absorb (Photon) Emit (Photon) h = Planks onstant = 6.63 x J.s c= Speed of light λ = Wavelength of light Neils enrik Davis Bohr s Atom Model (1913) fig_0_01 levels 13 in ydrogen Atom Emission Spectrum of ydrogen ydrogen atom has one proton and one electron of hydrogen atoms for different energy levels is given by (n=1,..) principal quantum E ev n numbers Example:- If an electron undergoes transition from n=3 state to n= state, the energy of photon emitted is E 1. 89ev 3 required to completely remove an electron from hydrogen atom is known as ionization energy Emission Spectra Emission spectra of hydrogen: animation. lick the figure below to view the animation (this animation has voice). x mu h 4 Uncertainty Principle and Schrodinger s Wave Functions Bohr s model fails to explain complex atoms. Louis de Broglie: Particles of matter such as electrons could be treated in terms of both particle and wave. λ = h / mv Werner eisenberg (uncertainty principle): It is impossible to simultaneously determine the exact position and the exact momentum of a body. Δx mδu h/4π Δx is the uncertainty in the position, and Δu is the uncertainty in speed. We can only provide the probability of finding an electron with a given energy within a given space (electron density)

4 Electron Density Quantum Numbers of Electrons of Atoms Solution of the wave equation is in terms of a wave function, ψ (orbitals). The square of the wave function represents electron density. Boundary surface representation. 0.1 nm Total probability 0.05 nm Principal Quantum Number (n) Represents main energy levels. Range 1 to 7. Larger the n higher the energy. n=1 n= n=3 Subsidiary Quantum Number l Represents sub energy levels (orbital). Range 0 n-1. Represented by letters s,p,d and f. n=1 n= s orbital (l=0) p Orbital (l=1) 19 0 Electronic Structure Electrons have wavelike and particulate properties. This means that electrons are in orbitals defined by a probability. Each orbital at discrete energy level is determined by quantum numbers. Electron States Electrons have discrete energy states tend to occupy lowest available energy state. 4d 4p 3d 4s N-shell n = 4 Adapted from Fig..4, allister & Rethwisch 8e. Quantum # Designation n = principal (energy level-shell) K, L, M, N, O (1,, 3, etc.) l = subsidiary (orbitals) s, p, d, f (0, 1,, 3,, n -1) m l = magnetic 1, 3, 5, 7 (-l to +l) m s = spin ½, -½ 3p M-shell n = 3 3s p s 1s L-shell n = K-shell n = 1 hapter - 1 hapter - SURVEY OF ELEMENTS Most elements: Electron configuration not stable. Element ydrogen elium Lithium Beryllium Boron arbon Neon Sodium Magnesium Aluminum Argon Krypton Atomic # Electron configuration 1 1s 1 1s (stable) 3 1s s 1 4 1s s 5 1s s p 1 6 1s s p s s p 6 (stable) 1s s p 6 3s 1 1s s p 6 3s 1s s p 6 3s 3p 1 Adapted from Table., allister & Rethwisch 8e. 1s s p 6 3s 3p 6 (stable) 1s s p 6 3s 3p 6 3d 10 4s 4p 6 (stable) Why? Valence (outer) shell usually not filled completely. fig_0_04 hapter - hapter - 4 4

5 Electron onfigurations Valence electrons those in unfilled shells Filled shells more stable Valence electrons are most available for bonding and tend to control the chemical properties example: (atomic number = 6) 1s s p valence electrons Electronic onfigurations ex: Fe - atomic # = 6 1s s p 6 3s 3p 6 3d 6 4s 4d 4p 3d 4s 3p M-shell n = 3 3s p s 1s N-shell n = 4 L-shell n = K-shell n = 1 valence electrons hapter - 5 hapter - 6 s, p and d Orbitals ybridization ybridization: Animation. lick the figure below to view the animation (this animation has voice). 7 8 Quantum Numbers of Electrons of Atoms Electron Structure of Multielectron Atom Magnetic Quantum Number m l. Represents spatial orientation of single atomic orbital. Permissible values are l to +l. Example:- if l=1, m l = -1,0,+1. I.e. l+1 allowed values. No effect on energy. Electron spin quantum number m s. Specifies two directions of electron spin. Directions are clockwise or anticlockwise. Values are +1/ or 1/. Two electrons on same orbital have opposite spins. No effect on energy. Maximum number of electrons in each atomic shell is given by n. Atomic size (radius) increases with addition of shells. Electron onfiguration lists the arrangement of electrons in orbital. Example :- Number of Electrons Orbital letters 1s s p 6 3s Principal Quantum Numbers For Iron, (Z=6), Electronic configuration is 1s s p 6 3s 3p 6 3d 6 4s

6 Multielectron Atoms Nucleus charge effect: The higher the charge of the nucleus, the higher is the attraction force on an electron and the lower the energy of the electron. Shielding effect: Electrons shield each other from the full force of the nucleus. The inner electrons shield the outer electrons and do so more effectively. In a given principal shell, n, the lower the value of l, the lower will be the energy of the subshell; s < p < d <f. The Quantum-Mechanical Model and the Periodic Table Elements are classified according to their ground state electron configuration. Figure Periodic Variations in Atomic Size Atomic Radius Atomic size: half the distance between the nuclei of two adjacent atoms (metallic radius) OR identical (covalent radius). Affected by principal quantum number and size of the nucleus. Atomic radius: animation. lick the figure below to view the animation (this animation has voice) Trends in Ionization Oxidation Number is required to remove an electron from its atom. First ionization energy plays the key role in the chemical reactivity. As the atomic size decreases it takes more energy to remove an electron. As the first outer core electron is removed, it takes more energy to remove a second outer core electron Positive oxidation number: The number of outer electrons that an atom can give up through the ionization process

7 Electron Structure and hemical Activity Electron Structure and hemical Activity Except elium, most noble gasses (Ne, Ar, Kr, Xe, Rn) are chemically very stable All have s p 6 configuration for outermost shell. elium has 1s configuration Electropositive elements give electrons during chemical reactions to form cations. ations are indicated by positive oxidation numbers Example:- Fe : 1s s p 6 3s 3p 6 3d 6 4s Fe + : 1s s p 6 3s 3p 6 3d 6 Fe 3+ : 1s s p 6 3s 3p 6 3d 5 Electronegative elements accept electrons during chemical reaction. Some elements behave as both electronegative and electropositive. Electronegativity is the degree to which the atom attracts electrons to itself Measured on a scale of 0 to 4.1 Example :- Electronegativity of Fluorine is 4.1 Electronegativity of Sodium is 1. Electropositive Na Te N O Fl 0 K 1 W Se 3 4 Electronegative Trends in Electron Affinity Metals, Metalloids, and Nonmetals Electron affinity: Tendency to accept one or more electrons and release energy. Electron affinity increases (more energy is released after accepting an electron) as we move to the right across a period and decreases as we move down in a group. Groups 6A and 7A have in general the highest electron affinities. Reactive metals: (or simply metals): Electro positive materials, have the natural tendency of losing electrons and in the process form cations. Reactive nonmetals (or simply nonmetals): Electronegative, they have the natural tendency of accepting electrons and in the process form anions. Metalloids: an behave either in a metallic or a nonmetallic manner. Examples: In group 4A, the carbon and the next two members, silicon and germanium, are metalloids while tin and lead, are metals. In group 5A, nitrogen and phosphorous are nonmetals, arsenic and antimony are metalloids, and finally bismuth is a metal Pop Quiz Primary Bonds 41 (Level: Knowledge and omprehension) 4 Bonding with other atoms, the potential energy of each bonding atom is lowered resulting in a more stable state. Three primary bonding combinations : 1) metal-nonmetal, ) nonmetal-nonmetal, and 3) metal-metal. Ionic bonds :- Strong atomic bonds due to transfer of electrons ovalent bonds :- Large interactive force due to sharing of electrons Metallic bonds :- Non-directional bonds formed by sharing of electrons Permanent Dipole bonds :- Weak intermolecular bonds due to attraction between the ends of permanent dipoles. Fluctuating Dipole bonds :- Very weak electric dipole bonds due to asymmetric distribution of electron densities. 7

8 Ionic Bonding Ionic bonding is due to electrostatic force of attraction between cations and anions. It can form between metallic and nonmetallic elements. Electrons are transferred from electropositive to electronegative atoms Ionic Bonds Large difference in electronegativity. When a metal forms a cation, its radius reduces and When a nonmetal forms an anion, its radius increases. Electropositive Element Electron Transfer Electronegative Atom ation +ve charge Electrostatic Attraction Anion -ve charge 43 IONI BOND The electronegativity variations 44 Ionic bonding in Nal Sodium Atom Na Ionic Bonding - Example 3s 1 3p 6 hlorine Atom l Ionic Force for Ion Pair Nucleus of one ion attracts electron of another ion. The electron clouds of ion repulse each other when they are sufficiently close. These two forces will balance each other when the equilibrium interionic distance, a 0, is reached and a bond is formed 45 Sodium Ion Na + I O N I B O N D hlorine Ion l - 46 Force versus separation distance for a pair of oppositely charged ions Figure.16 Ion Force for Ion Pair Interionic Force - Example F Z ez e 1 0a 1 4 Z1,Z = Number of electrons removed or added during ion formation e = Electron harge, a = Interionic seperation distance ε = Permeability of free space (8.85 x 10-1 c /Nm ) (n and b are constants) Z Z e 0a 4 attractive Force of attraction between Na+ and l - ions Z 1 = +1 for Na +, Z = -1 for l - e = 1.60 x 10-19, ε 0 = 8.85 x 10-1 /Nm a 0 = Sum of Radii of Na + and l - ions = nm nm =.76 x m 47 F F net repulsive nb Z 1Z e a n1 nb n 4 a 0 a 1 48 F attraction Z1Z e 4 0a 19 ( 1)( 1)( ) N (8.85 x 10 /Nm)(.76 x 10 m) Na + l- a 0 8

9 Interionic Energies for Ion Pairs Net potential energy for a pair of oppositely charged ions = Z Z e b 1 E net n 4 a 0 a Attraction Released Repulsion Absorbed E net is minimum when ions are at equilibrium seperation distance a 0 Ion Arrangements in Ionic Solids Ionic bonds are Non Directional (i.e. no orientation) Geometric arrangements are present in solids to maintain electric neutrality. Example:- in Nal, six l- ions pack around central Na+ Ions sl Nal Ionic packing In Nal and sl Figure.18 As the ratio of cation to anion radius decreases, fewer anion surround central cation Bonding Energies Bonding Lattice energies and melting points of ionically bonded solids are high. Lattice energy decreases when size of ion increases. Multiple bonding electrons increase lattice energy. Example :- Nal sl BaO Lattice energy = 766 kj/mol Melting point = 801 o Lattice energy = 649 kj/mol Melting Point = 646 o Lattice energy = 317 kj/mol Melting point = 193 o onsider production of LiF: result in the release of about 617 kj/mole. Step 1. onverting solid Li to gaseous Li (1s s 1 ): 161 kj/mole of energy. Step. onverting the F molecule to F atoms: 79.5 kj/mole. Step 3. Removing the s 1 electron of Li to form a cation, Li + : 50 kj/mole. Step 4. Transferring or adding an electron to the F atom to form an anion, F - : -38 kj/mole. Step 5. Formation of an ionic solid from gaseous ions: lattice energy, unknown=-617 kj [161 kj kj + 50 kj 38 kj] = kj 51 5 Lattice, Material Properties ovalent Bonding Ionic solids are hard, rigid and strong and brittle. Excellent Insulators. In ovalent bonding, outer s and p electrons are shared between two atoms to obtain noble gas configuration. Takes place between elements with small differences in electronegativity and close by in periodic table. In ydrogen, a bond is formed between atoms by sharing their 1s 1 electrons + Electron Pair Overlapping Electron louds s 1 Electrons ydrogen Molecule 9

10 ovalent Bonding - Examples In case of F, O and N, covalent bonding is formed by sharing p electrons Fluorine gas (Outer orbital s p 5 ) share one p electron to attain noble gas configuration. F + F F F Oxygen (Outer orbital - s p 4 ) atoms share two p electrons F F Bond =160KJ/mol O + O O O O = O Bond =8KJ/mol Bond Length, Bond order and Bond For a given pair of atoms, with higher bond order, the bond length will decrease; as bond length decreases, bond energy will increase (, F, N ) Nonpolar bonds: sharing of the bonding electrons is equal between the atoms and the bonds. Polar covalent bond: Sharing of the bonding electrons is unequal (F, NaF). 55 Nitrogen (Outer orbital - s p 3 ) atoms share three p electrons N + N N N N N Bond =54KJ/mol 56 Pop Quiz ovalent Bonding in arbon (Level: Knowledge and omprehension) arbon has electronic configuration 1s s p Ground State arrangement 1s s p Two ½ filed p orbitals Indicates carbon Forms two ovalent bonds ybridization causes one of the s orbitals promoted to p orbital. Result four sp3 orbitals. 1s p Four ½ filled sp 3 orbitals Indicates four covalent bonds are formed Structure of Diamond arbon ontaining Molecules Four sp 3 orbitals are directed symmetrically toward corners of regular tetrahedron. This structure gives high hardness, high bonding strength (711kJ/mol) and high melting temperature (3550 o ). arbon Atom Tetrahedral arrangement in diamond In Methane, arbon forms four covalent bonds with ydrogen. Methane Molecules are very weakly molecule bonded together resulting in low melting temperature (-183 o ). arbon also forms bonds with itself. Molecules with multiple carbon bonds are more reactive. Examples:- Ethylene Acetylene

11 ovalent Bonding in Benzene hemical composition of Benzene is 6 6. The arbon atoms are arranged in hexagonal ring. Single and double bonds alternate between the atoms. Ionic Versus ovalent Bonding Ionic versus ovalent bonds: Animation. lick the figure below to view the animation (this animation has voice). 61 Structure of Benzene Simplified Notations 6 Metallic Bonding Metallic Bonds (ont..) Atoms in metals are closely packed in crystal structure. Loosely bounded valence electrons are attracted towards nucleus of other atoms. Electrons spread out among atoms forming electron clouds. These free electrons are Positive Ion reason for electric conductivity and ductility Since outer electrons are shared by many atoms, metallic bonds are Non-directional Overall energy of individual atoms are lowered by metallic bonds Minimum energy between atoms exist at equilibrium distance a 0 Fewer the number of valence electrons involved, more metallic the bond is. Example:- Na Bonding energy 108KJ/mol, Melting temperature 97.7 o igher the number of valence electrons involved, higher is the bonding energy. Example:- a Bonding energy 177KJ/mol, Melting temperature 851 o 63 Valence electron charge cloud 64 Metallic Bonds and Material Properties Metallic Bonds and Material Properties The bond energies and the melting point of metals vary greatly depending on the number of valence electrons and the percent metallic bonding. Pure metals are significantly more malleable than ionic or covalent networked materials. Strength of a pure metal can be significantly increased through alloying. Pure metals are excellent conductors of heat and electricity

12 Secondary or van der Waals Bonding Secondary bonds are due to attractions of electric dipoles in atoms or molecules. Dipoles are created when positive and negative charge centers exist. Fluctuating Dipoles Weak secondary bonds in noble gasses. Dipoles are created due to asymmetrical distribution of electron charges. Electron cloud charge changes with time. +q -q Dipole moment=μ =q.d q= Electric charge d = separation distance Figure.6 d There two types of bonds i. permanent and ii. fluctuating. Symmetrical distribution of electron charge Asymmetrical Distribution (hanges with time) Permanent Dipoles ydrogen Bonds Dipoles that do not fluctuate with time are called Permanent dipoles. Examples:- 4 methane Symmetrical Arrangement of 4 - bonds No Dipole moment ydrogen bonds are Dipole-Dipole interaction between polar bonds containing hydrogen atom. Example :- In water, dipole is created due to asymmetrical arrangement of hydrogen atoms. Attraction between positive oxygen pole and negative hydrogen pole. 3 l chloromethane Asymmetrical Tetrahedral arrangement reates Dipole O ydrogen Bond Short Video The origin of water? The biggest star in universe? SUN

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