Chapter 2: Atomic Structure and Interatomic Bonding

Transcription

1 Chapter 2: Atomic Structure and Interatomic Bonding Atomic Structure Electron Configuration Periodic Table Primary Bonding Ionic Covalent Metallic Secondary Bonding or van der Waals Bonding Three types of Dipole Bonding Molecules

2 Why Study Atomic Structure and Interatomic Bonding? Two Allotropes of CARBON Graphite Diamond - Relatively soft - Greasy feel to it - Reasonably good conductor of electricity - The hardest known material - Poor conductor of electricity The disparities in properties are attributed to a type of interatomic bonding found in graphite that does not exist in diamond

3 Atomic Structure (Freshman Chem.) atom electrons 9.11 x kg protons } neutrons 1.67 x kg atomic number = # of protons in nucleus of atom = # of electrons of neutral species A [=] atomic mass unit = amu = 1/12 mass of 12 C Atomic wt = wt of x molecules or atoms 1 amu/atom = 1g/mol C H etc. 3

4 Atomic Models

5 Timeline of Atomic Theory Democritus atomos Aristotle and Plato Dalton J.J. Thomson 460 BC 360 BC E. Rutherford N. Bohr Wave Mechanical 5

6 ~ 400 BC - Democritus Ancient Greek philosopher Democritus coined the term átomos which means "uncuttable" or "the smallest indivisible particle of matter". Structure of Matter Physical world VOID + BEING

7 Democritus Democritus postulated that atoms were completely solid, hard and small particles with no internal structure and has an infinite variety of shapes and sizes. This theory was ignored for more than 2000 years! 7

8 The Atomic Theory of Matter Aristotle and Plato There can be no ultimately indivisible particles. Believed that fire, earth, air and water were the four main elements that world was made up of. 8

9 1803 John Dalton English instructor and natural philosopher Each element consists of atoms of single unique type and can join to form chemical compounds. Beginning the modern atomic theory

10 The Atomic Theory of Matter Dalton s Atomic Theory Postulates 1. Elements are composed of extremely small particles called atoms. 2. All atoms of a given element are identical, having the same size, mass and chemical properties. The atoms of different elements are different. Different elements have different atomic properties such as atomic mass. 10

11 Assist. Prof. Dr. İlkay KALAY The Atomic Theory of Matter Dalton s Atomic Theory 3. Atoms of an element are neither created nor destroyed by any chemical reactions. Chemical reactions only involve the combination, rearrangement or separation of atoms. 4. Compounds are formed from the combination of atoms of more than one element. A compound always has the same relative number and kind of atoms. 11

12 Mendeleev Building upon earlier discoveries by scientists, Mendeleev published the first functional PERIODIC TABLE. Certain chemical properties of elements repeat periodically when arranged by atomic number. The periodic table was first developed by Mendeleev and Meyer on the bases of similarity in chemical and physical properties exhibited on certain elements.

13 Periodic Table Draft of the first periodic table, Mendeleev, 1869

14 1869

15 Today: Periodic Table of the Elements

16 1897 Sir J. J. Thomson Although Dalton had postulated that atoms were indivisible, the studies have shown a more complex structure for an atom. J. J. Thomson conducted a series of experiments which showed that the atoms were not indivisible Discovered the electron (1906 Nobel Prize in Physics).

17 Atom is composed of even smaller particles but how the particles fit together? J. J. Thomson proposed the atom consisted of a uniform positive sphere of matter in which the electrons were embedded. Plum Pudding (1904): The atom as being made up of electrons swarming in a sea of positive charge. J. J. Thomson's "plum-pudding" model of the atom 17

18 1909 E. Rutherford Tested and disproved the Plum Pudding Model. Rutherford's experiment on the scattering of α particles by metal foil Results: Majority of a particles transmitted (pass through) or deflected through small angles Tiny fraction deflected through large angles

19 1909 E. Rutherford Conclusion: Disproved the Plum-Pudding Model Large amount of the atom's charge and mass is concentrated into a small region Atom was mostly empty space Objections to Rutherford model The laws of classical mechanics predict that the electron will release electromagnetic radiation while orbiting a nucleus. Because the electron would lose energy, it would gradually spiral inwards, collapsing into the nucleus. This atom model is unsuccessful, because it predicts that all atoms are unstable.

20 The Modern View of Atomic Structure The structure of an atom Electron Nucleus Particle Charge Mass (amu) Electron Negative (-1) x 10-4 Proton Positive (+1) Neutron Neutral (0) Neutron Proton Particle Charge (C) Electron x Proton x Neutron 0 20

21 Assist. Prof. Dr. İlkay KALAY The Modern View of Atomic Structure The structure of an atom The atoms are small and the atomic dimensions are expressed in terms of Angstrom (Å) unit. 1 Å = m Nucleus ~ 10-4 Å 1-5 Å Schematic view of an 21

22 Assist. Prof. Dr. İlkay KALAY The Modern View of Atomic Structure The structure of an atom ü All atoms of an element have the same number of protons in the nucleus. ü The number of protons in the nucleus of an atom is called atomic number (Z). ü Because the atoms have no net electrical charge, # of protons = # of electrons. ü The total number of protons and neutrons in a nucleus is called mass number (A). Mass number = A Atomic number = Z 12 C 6 22

23 Assist. Prof. Dr. İlkay KALAY The Modern View of Atomic Structure ü All nuclei of the atoms of a particular element have the same atomic number, but the the number of neutrons and so mass numbers may be different. Atoms of a given element with a different mass numbers are called isotopes. ü Most elements occur in nature as mixtures of isotopes. Symbol C 6C 13 6C 14 6C # of electrons # of protons # of neutrons An atom of a specific isotope is called a nuclide. 23

24 The Periodic Table The periodic table, developed in 1869, shows the arrangement of elements with similar chemical and physical properties in order of increasing atomic number. Periodic table presents atomic number, atomic symbol and atomic weight for each element. 40 Atomic number Zr Atomic symbol Atomic weight 24

25 1912 N. Bohr Many phenomena involving electrons in solids could not be explained in terms of CLASSICAL MECHANICS. We need QUANTUM MECHANICS

26 Line Spectra and The Bohr Model Bohr s Model Bohr proposed a model of the hydrogen atom that explains its line spectrum. Bohr s postulates Rutherford atom is correct Classical EM theory not applicable to orbiting e- Newtonian mechanics applicable to orbiting e- E electron = E kinetic + E potential Only orbits of certain radii, corresponding to certain definite energies are permitted for electrons in an atom. An electron in a permitted orbit has a specific energy and is in allowed energy state. An electron in an allowed energy state will not radiate energy and therefore will not spiral into the nucleus. During the transition of an electron from one allowed energy state to another, energy is only emitted or absorbed by an electron. This energy is emitted or absorbed as a photon, ΔE = Ef-Ei= hν = hc/λ where c = νλ 26

27 Line Spectra and The Bohr Model Energy States of the Hydrogen Atom Energy levels in the hydrogen atom from the Bohr model. - The arrows refer to the transi7ons of the electron from one allowed energy state to another. - The states shown are those for which n = 1 through n = 6, and the state for n =, for which the energy, E, equals zero. 27

28 Bohr s Model The lower the value of n, the smaller the radius of the orbit, and the lower the energy level. n=1 ground state (lowest energy state) (orbit closest to the nucleus) n=2, 3, or higher excited state The arrows refer to the transi7ons of the electron from one allowed energy state to another. Energy levels in the hydrogen atom from the Bohr model. 28

29 BOHR ATOM orbital electrons: n = principal quantum number n=3 2 1 Adapted from Fig. 2.1, Callister 6e. Nucleus: Z = # protons N = # neutrons Atomic mass A Z + N 2

30 Sommerfeld German theoretical physicist Modified the Bohr Model suppose we have plurality of orbits a shell containing multiple orbits: ORBITALS How to capture these new ideas quantitatively? We need new quantum numbers: n, l, m, s n principal quantum number, distance of an electron from the nucleus l subshell, describes the shape of the subshell m number of energy states in a subshell s spin moment

31 Wave mechanics to arrive at same place: E=E(n,l,m,s) The Bohr model significant limitations Resolution: Wave-mechanical model (electron is considered to exhibit both wave-like and particle-like characteristics). De Broglie: If a photon which has no mass, can behave as a particle, does an electron which has mass can behave as a wave (1920)? = h/p = h/mv Heisenberg: Uncertainty Principle I don t know where any of one of electrons is, but I can tell you an average where any of one of them is likely to be Schrodinger

32 De Broglie wavelength:! = h m" Heisenberg Uncertainty Principle: Beyond Bohr s Model De Broglie wavelength states that electrons have wavelike mo7on It is impossible to know simultaneously the exact posi7on and!x.!m! " h 4" momentum of a par7cle. The more accurately we know the posi7on of the par7cle (smaller Δx), the less accurately we know its speed (larger Δu) and vice versa. Wave mo>on of objects on the atomic scale Schrödinger s wave equacon incorporates both the wave- like and par.cle- like behavior of the electron. Wave function, ψ, Ψ 2 provides information about an electron s location when it s in an allowed energy state

33 Quantum Mechanics and Atomic Orbitals Electron density: probability of the electron being at a point Higher density of dots region: larger values of Ψ 2 Electron- density diagram in the ground state of the hydrogen atom

34 Bohr Model vs. Wave Mechanical Atom Model With WM model electron no longer treated as a particle moving in a discrete orbital, rather position is described by a probability distribution Bohr WM

35 Quantum Numbers According to quantum mechanics, each electron in an atom is described by four different quantum numbers, three of which (n, l, m l ) specify the wave function that gives the probability of finding the electron at various points in space. n. The principal quantum number (K, L, M, N, O and so on that correspond to 1, 2, 3,4, 5 As n increases: - The orbital becomes larger - The electron spends more time farther from the nucleus - The electron has a higher energy and is therefore less tightly bound to the nucleus. l. The second quantum number l can have integral values from 0 to n-1 for each value of n defines the shape of the orbital The value of l of a particular orbital is designated by s, p, d and f, corresponding l values of 0, 1, 2 and 3 respectively s: sharp p:principal d:diffuse f: fundamental m l. The magnetic quantum number Have integral values between l and l including zero. Describes the orientation of the orbital in space

36 What is the filling sequence of electrons in orbitals by n, l, m, s is not adequate? AUFBAU PRINCIPLE 3 principles: 1. Pauli Exclusion Principle:only one electron can have a given set of four quantum numbers. 2. Electrons -have discrete energy states -fill orbitals from lowest en. to highest en. 3. Hund s rule

37 Electron Configura>ons Hund s rule states that the lowest energy is attained by maximizing the number of electrons with the same electron spin. For example, for a carbon atom to achieve its lowest energy, the two 2p electrons will have the same spin. 37

38 Quantum Numbers l m l m s = ±½

39 Quantum Numbers Electrons fill quantum levels in order of increasing energy ( only n and l make significant differences in energy configurations). 1s, 2s, 2p, 3s,3p,4s,3d,4p,5s,4d,5p,6s,4f,5d,. When all electrons are at the lowest possible energy levels => ground state Excited states do exist such as in glow discharges etc Valence electrons occupy the outermost filled shell. Valence electrons are responsible for all bonding!

40 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 determined by quantum numbers. Quantum # Designation n = principal (energy level-shell) K, L, M, N, O (1, 2, 3, etc.) l = subsidiary (orbitals) s, p, d, f (0, 1, 2, 3,, n -1) m l = magnetic 1, 3, 5, 7 (-l to +l) m s = spin ½, -½ 40

41 Electron Energy States Electrons... have discrete energy states tend to occupy lowest available energy state. 4d 4p N-shell n = 4 3d 4s Energy 3p M-shell n = 3 3s 2p 2s Adapted from Fig. 2.4, Callister 7e. L-shell n = 2 1s K-shell n = 1 41

42 SURVEY OF ELEMENTS Most elements: Electron configuration not stable. Element Hydrogen Helium Lithium Beryllium Boron Carbon... Neon Sodium Magnesium Aluminum... Argon... Krypton Atomic # Electron configuration 1s 1 1s 2 (stable) 1s 2 2s 1 1s 2 2s 2 1s 2 2s 2 2p 1 1s 2 2s 2 2p s 2 2s 2 2p 6 (stable) 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 3s 2 1s 2 2s 2 2p 6 3s 2 3p 1... Adapted from Table 2.2, Callister 7e. 1s 2 2s 2 2p 6 3s 2 3p 6 (stable)... 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 (stable) Why? Valence (outer) shell usually not filled completely. 42

43 Assist. Prof. Dr. İlkay KALAY Electron Configura>ons Elements in any given group in the periodic table have the same type of electron arrangements in their outermost shells. The outer shell electrons those that lie outside the orbitals occupied in the next lowest noble gas element are called its valence electrons, whereas the electrons in the inner shells are called the core electrons. 43

44 Electron Configurations 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: C (atomic number = 6) 1s 2 2s 2 2p 2 valence electrons 44

45 Assist. Prof. Dr. İlkay KALAY Electron Configura>ons and the Periodic Table 45

46 Atomic Structure Valence electrons determine all of the following properties 1) Chemical 2) Electrical 3) Thermal 4) Optical 46

47 Electronic Configurations ex: Fe - atomic # = 26 1s 2 2s 2 2p 6 3s 2 3p 6 3d 6 4s 2 4d 4p 3d 4s N-shell n = 4 valence electrons Energy 3p M-shell n = 3 3s 2p 2s 1s L-shell n = 2 K-shell n = 1 47

48 The Periodic Table Columns: Similar Valence Structure give up 1e give up 2e give up 3e H Li Be Na Mg K Ca Sc accept 2e accept 1e inert gases O S Se F Cl Br He Ne Ar Kr Adapted from Fig. 2.6, Callister 7e. Rb Sr Y Te I Xe Cs Ba Po At Rn Fr Ra Electropositive elements: Readily give up electrons to become + ions. Electronegative elements: Readily acquire electrons to become - ions. 48

49 Electronegativity - Tells us whether a given bond will be nonpolar covalent, polar covalent or ionic. - The ability of an atom in a molecule to attract electrons to itself. - Ranges from 0.7 to 4.0. Smaller electronegativity Larger electronegativity 49

50 REVIEW OF ATOMIC STRUCTURE (FRESHMAN CHEMISTRY) ATOMS = (PROTONS+NEUTRONS) + ELECTRONS NUCLEUS BONDING Mass of an atom: Proton and Neutron: ~ 1.67 x kg Electron: 9.11 x kg Charge: Electrons and protons: (±) 1.60 x C Neutrons are neutral The atomic mass (A): total mass of protons + total mass of neutrons Atomic weight ~ Atomic mass # of protons are used to identify elements (Z) # of neutron are used to identify isotopes ( e.g. 14 C 6 and 12 C 6 ) Isotopes are written as follows: A X Z, i.e. 1 H 1, 2 H 1, 3 H 1

51 Atomic bonding in solids Things are made of atoms little particles that move around, attracting each other when they are a little distance apart, but repelling upon being squeezed into one another. In that one sentence... there is an enormous amount of information about the world. Richard P. Feynman

52 Why Study Atomic Structure and Interatomic Bonding? Two Allotropes of CARBON Graphite Diamond - Relatively soft - Greasy feel to it - Reasonably good conductor of electricity - The hardest known material - Poor conductor of electricity The disparities in properties are attributed to a type of interatomic bonding found in graphite that does not exist in diamond

53 Atomic Bonding in Solids Interatomic forces that bind the atoms together are important to understand many properties of materials. Start with two atoms infinitely separated At large distances, interactions are negligible At small distances, each atom exerts forces on the other. Two types of forces: attractive, F A, and repulsive, F R Interatomic separation, r Attractive component is due to type of the bonding (minimize energy thru electronic configuration) Repulsive component is due to negatively charged electron clouds for two atoms and important only at small values of r

54 Bonding Forces and Energies F N =F A +F R At equilibrium F A +F R =0 E N = E A +E R 54

55 Atomic Bonding Essentially atoms either want to give up (transfer) or acquire (share) electrons to complete electron configurations; minimize their energy Transfer of electrons => ionic bond Sharing of electrons => covalent Metallic bond => sea of electrons Sulfur Magnesium oxide Gold Bromine COVALENT BONDING Sucrose Sharing of electrons between two atoms Magnesium Copper METALLIC BONDING Bonding of metal atoms to neighbor atoms Potassium dichromate IONIC BONDING Electrostatic forces between ions Nickel (II) oxide

56 Ionic Bonding Occurs between + and ions (anion and cation). Requires electron transfer. Large difference in electronegativity required. Example: NaCl Na (metal) unstable electron Cl (nonmetal) unstable Na (cation) stable + - Coulombic Attraction Cl (anion) stable 56

57 Ionic bond metal + nonmetal donates electrons accepts electrons Dissimilar electronegativities ex: MgO Mg 1s 2 2s 2 2p 6 3s 2 O 1s 2 2s 2 2p 4 [Ne] 3s 2 Mg 2+ 1s 2 2s 2 2p 6 O 2-1s 2 2s 2 2p 6 [Ne] [Ne] 57

58 IONIC BONDING Oppositely charged ions attract, attractive force is coulombic. Ionic bond is non-directional, ions get attracted to one another in any direction. Bonding energies are high => 2 to 5 ev/atom,molecule,ion Hard materials, brittle, high melting temperature, electrically and thermally insulating 8

59 Ionic Bonding Energy minimum energy most stable Energy balance of attractive and repulsive terms E N = E A + E R = A r B r n Repulsive energy E R Interatomic separation r Net energy E N Adapted from Fig. 2.8(b), Callister 7e. Attractive energy E A 59

60 Examples: Ionic Bonding Predominant bonding in Ceramics NaCl MgO CaF 2 CsCl Give up electrons Acquire electrons Adapted from Fig. 2.7, Callister 7e. (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. 60

61 Covalent Bonding Requires shared electrons Example: CH 4 C: has 4 valence e -, needs 4 more CH 4 H shared electrons from carbon atom H: has 1 valence e -, needs 1 more Electronegativities are comparable. H C H H shared electrons from hydrogen atoms Adapted from Fig. 2.10, Callister 7e. 61

62 COVALENT BONDING Covalent bonds are formed by sharing of the valence electrons Covalent bonds are very directional Covalent bond model: an atom can have at most 8-N covalent bonds, where N = number of valence electrons Diamond, sp3 Covalent bonds can be very strong, eg diamond, SiC, Si, etc, also can be very weak, eg Bismuth Polymeric materials do exhibit covalent type bonding. 10

63 H 2.1 Li 1.0 Na 0.9 K 0.8 Rb 0.8 Cs 0.7 Fr 0.7 EXAMPLES: COVALENT BONDING Be 1.5 Mg 1.2 Ca 1.0 Sr 1.0 Ba 0.9 Ra 0.9 H2 Ti 1.5 Cr 1.6 Fe 1.8 H2O C(diamond) SiC Ni 1.8 Zn 1.8 Ga 1.6 column IVA C 2.5 Si 1.8 As 2.0 GaAs 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. Ge 1.8 Sn 1.8 Pb 1.8 O 2.0 F 4.0 Cl 3.0 Br 2.8 I 2.5 At 2.2 He - Ne - Ar - Kr - Xe - Rn - F2 Cl2 Molecules with nonmetals Molecules with metals and nonmetals Elemental solids (RHS of Periodic Table) Compound solids (about column IVA) 11

64 Primary Bonding Ionic-Covalent Mixed Bonding % ionic character = # 1" e " (X A"X B ) % 4 % % $ 2 & ( ( ( ' x ( 100 %) where X A & X B are Pauling electronegativities of the corresponding elements Ex: MgO X Mg = 1.3 X O = 3.5 (3.5 % ionic character = 1 e 4 1.3) 2 x (100%) = 70.2% ionic 64

65 METALLIC BONDING Arises from a sea of donated valence electrons (1, 2, or 3 from each atom). Non valence and atomic nuclei form ion cores. Ion cores in the sea of electrons. Valance electrons belong no one particular atom but drift throughout the entire metal. Free electrons shield + ly charged ions from repelling each other Adapted from Fig. 2.11, Callister 6e. Primary bond for metals and their alloys 12

66 Arises from interaction between dipoles Fluctuating dipoles Permanent dipoles-molecule induced -general case: -ex: liquid HCl SECONDARY BONDING asymmetric electron clouds secondary bonding H H Adapted from Fig. 2.13, Callister 7e. secondary bonding H Cl secondary bonding ex: liquid H 2 H 2 H 2 H Cl H secondary bonding H Adapted from Fig. 2.14, Callister 7e. -ex: polymer secondary bonding secondary bonding 66

67 Type Ionic Summary: Bonding Bond Energy Comments Large! Nondirectional (ceramics) Covalent Metallic Secondary Variable large-diamond small-bismuth Variable large-tungsten small-mercury smallest Directional (semiconductors, ceramics polymer chains) Nondirectional (metals) Directional inter-chain (polymer) inter-molecular 67

68 Bonding Energies

69 Bonding in Solids The physical properties of crystalline solids, such as melting point and hardness, depend both on the arrangements of particles and on the attractive forces between them. 69

70 Properties From Bonding: T m Bond length, r r Melting Temperature, T m Energy Bond energy, E o Energy r o smaller T m r unstretched length r o r E o = bond energy larger T m T m is larger if E o is larger. 70

71 PROPERTIES FROM BONDING: E Elastic modulus, E F ΔL A = E o L o Elastic modulus E ~ curvature at ro Energy unstretched length r o r E is larger if Eo is larger. smaller Elastic Modulus larger Elastic Modulus 16

72 Summary: Primary Bonds Ceramics (Ionic & covalent bonding): Metals (Metallic bonding): Large bond energy large T m large E small α Variable bond energy moderate T m moderate E moderate α Polymers (Covalent & Secondary): secondary bonding Directional Properties Secondary bonding dominates small T m small E large α 72