ATOMIC STRUCTURE AND BONDING. IE-114 Materials Science and General Chemistry Lecture-2

Transcription

1 ATOMIC STRUCTURE AND BONDING IE-114 Materials Science and General Chemistry Lecture-2

2 Outline Atomic Structure (Fundamental concepts, Atomic models (Bohr and Wave-Mechanical Atomic Model), Electron configurations) Periodic Table (Classification of elements, their characteristics) Atomic Bonding in Solid Materials (Primary Bonding:Ionic, covalent, metallic bonds) (Secondary bonding: Fluactuating Induced Dipole, Polar Molecule-Induced Dipole, Permanent Dipole Bonds)

3 Why atomic structure is important? Some properties of solid materials depend on atomic nature and its arrangement. Type of atomic arrangements : crytalline or amorphous Type of bonding (interactions among the atoms) determines the melting temperature (T m ), coefficient of thermal expansion(α), mechanical properties, i.e. Elastic modulus, E Graphite and Diamond graphite diamond Graphite: Soft and greasy feel Diamond: Hardest known material different type of interatomic bonding in graphite and diamond.

4 Atomic Structure Atom = nucleus (protons+neutron) and electrons Charge =1.60x10-19 C Electrons are negatively(-) charged, Protons are positively(+) charged Neutrons are electrically neutral particles. Atomic Number (Z): Number of protons in the nucleus Electrically neutral atom; # protons = #electrons

5 Mass of an Atom Atomic mass (A) for an atom = masses of protons(z)+masses of neutrons(n) (electrons are not considered, because...?) Mass of proton = mass of neutrons=1.67x10-27 kg Mass of electron= 9.11x10-31 kg Atoms with two or more atomic masses (ISOTOPES) For all atoms of an element the number of protons are the same, but the number of neutrons may vary, which vary the atomic mass, Example: 12 C, 13 C, 14 C Atomic Weight: Weighted average of the atomic masses of the atom s naturally occuring isotopes

6 The atomic mass unit (amu) is used for the computation of atomic weight. Scale: 1 amu=1/12 of the atomic mass of the Carbon (C) (A= for carbon 12 isotope) 1 amu/atom= 1 g/mol ( 1 mol of a substance=6.023x10 23 atoms) For example: Fe A=55.85 amu/atom or g/mol (this is most commonly used form)

7 Structure of Atom 1) Bohr Atomic Model 2) Wave-Mechanical Model BOHR ATOMIC MODEL: (Used hydrogen atom) 1) Electrons are assumed to be positioned around the nucleus in discrete orbitals 2) Position of the electron is more or less well defined in its orbital. Nucleus: Z = # protons N = # neutrons

8 Energy of Electrons Bohr Atomic model describe the electrons in terms of their positions (orbitals) and energy (quantized energy levels by Rydberg equation). E= - (2π 2 me 4 /n 2 h 2 ) = - (13.6/n 2 ) ev e: electron charge m: electron mass n: principal quantum number or principal energy levels(1,2,3,.) An electron can change its energy level To a higher level by absorbing energy, to a lower level by emitting energy Example: n=1 E 1 = ev n=2 E 2 = ev E 2 >E 1 If electron changes its energy level from 1 to 2 (from lower to higher energy level) It must absorb energy, and the amount of energy absorbed; E =E 2 - E 1 = -3.4 (-13.4) = 10.0 ev

9 Allowed energy levels for hydrogen electron in Bohr Model Ionization energy: Energy required to remove the electron completely from the atom Ionization energy for hydrogen electron is 13,6 ev Figure.The first three electron energy states for the Bohr hydrogen aton * Bohr s model was not able to explain quantitatively the spectra of the atoms more complex than hydrogen and the model could not have been modified.

10 WAVE-MECHANICAL MODEL: Limitations of Bohr model was resolved by this model and electrons are considered to behave both wave-like and particle-like. Electrons are no longer treated as a particle moving in discrite orbitals. Position of electron is described by a probability distribution or electron cloud Heisenberg s uncertainty principle; Position and momentum of a small particle such as an electron can not be determined simultaneously. BOHR MODEL WAVE-MECHANICAL MODEL Since the position of an electron can not be precisely determined, an electron charge cloud density distribution is used

11 Motion of electron around its nucleus and its energy is characterized by 4 QUANTUM NUMBERS (n, l, m l, m s ) 1) Principal quantum number,n: Represents main energy levels for the electrons or shells (n=1 to 7 ) n=1 (first shell, K) n= 2 second shell(l), so forth... or 2) Secondary quantum number, l: Specifies subenergy levels within the main energy levels(subshells) and related to shape of the electron subshell l=n-1 Number designation of l: Letter designation of l : s p d f g h Principle quantum number, n Shell Designation Subshells 1 K s 2 L s,p 3 M s,p,d 4 N s,p,d,f

12 3) Magnetic quantum number, m l : Number of orbitals or energy states for each subshell m l =2l+1 Example: For a given l, m l can range from +l to l l=0 (s subshell) m l = 1 energy state (0) l=1 (p subshell) m l = 3 energy states (+1, 0, -1) l=2 (d subshell) m l = 5 energy states (+2,+1, 0, -1,-2) Pauli s Exclusion Principle: No two electrons can have the identical values for all four of their quantum numbers 4) Electron spin quantum number, m s : Specifies two allowed spin directions for an electron. m s = +1/2 and -1/2

13 Comparison of electron energy states in Bohr and Wave-mechanical Models BOHR MODEL WAVE-MECHANICAL MODEL

14 The maximum number of electrons in each shell in an atom is 2n 2 ***Electrons fill up the the lowest possible energy states in the electron shells and subshells

15 Electron Configurations Represents the manner in which the states are occupied The number of electrons in each subshell is indicated by a superscript after the shell-subshell designation. Example: H 1s 1 He 1s 2 Na 1s 2 2s 2 2p 6 3s 1 Na 1s 2 2s 2 2p 6 3s1 Principal quantum number,n (SHELL K,L,M,..) p subshell has three orbitals: m l =-1,0,+1, each of these orbitals contains 2 electrons Secondary quantum number,l (SUBSHELL; s,p,d,f) s subshell has one orbital: m l =0 Orbital contains 2 electrons

16 Valance Electrons:The electrons occupying the outermost shell These electrons participate in bonding. Many of the physical and chemical properties of solids are based on these valence electrons. Stable electron configurations have complete s and p subshells tend to be unreactive. (inert, or noble, gases)

17 Survey of Elements * Most of the elements are not stable. Electron configuration 1s 1 1s 2 1s 2 2s 1 1s 2 2s 2 1s 2 2s 2 2p 1 1s 2 2s 2 2p 2... (stable) 1s 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 s 2 2s 2 2p 6 3s 2 3p 6 (stable)... 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s (stable)

18 Periodic Table Elements are classified according to electron configuration in this table. Same column, or group, have similar chemical and physical properties due to similar valence electron configurations These properties change gradually and systematically across each period moving horizontally. Atomic Number(Z) +

19 Groups are designated at the top by the numbers 0-7 and by the letters A and B. A group elements- Representative or main group elements B group elements- Transition elements *Group 0: inert gases (filled electron shells) *Group IA and IIA are alkali (except H) and alkaline earth metals *Group IIIA, IVA and VA elements have characteristics between metal and nonmetals because of their valence electron configurations. *Group VIIA (halogens) and VIA elements= one and two electrons deficient respectively from having stable configurations. *Groups from IIIB to IIB are transition metals, with partially filled d electron states and in some cases one or two electrons in the next higher shell.

20 Electropositive and Electronegative Elements Electropositive elements: Elements capable of giving up their electrons to become positively charged ions (located on the left of the table.) Electronegative elements: Elements ready to accept electrons to form negatively charged ions or to share their electrons. Electropositive elements: Readily give up electrons to become + ions. Electronegative elements: Readily acquire electrons to become - ions.

21 Electronegativity The degree to which an atom attracts electrons to itself Ranges from 0.7 to 4.0 Large values: tendency to acquire electrons. Smaller electronegativity Larger electronegativity

22 Atomic Bonding in Solids Atomic bonding can be explained by interaction of two isolated atoms Net force is zero (equilibrium state) r 0 =equilibrium spacing Physical properties are related to interatomic forces that bind the atoms together

23 Bonding energy, E o Energy required to seperate these two atoms to an infinite seperation Bonding energy (E o ); Solids >Liquids > Gases 1) Bonding arises from the tendency of the atoms to assume stable electron structures 2) Valance electrons are involved 3) Nature of bond depends on the electron structure

24 Properties from bonding Bond energy, Eo 1) Melting Temperature, Tm Tm is larger if Eo is larger.

25 2) Coefficient of thermal expansion, α coeff. thermal expansion L Lo = (T 2 -T 1 ) α ~ symmetry at ro is larger if Eo is smaller.

26 3) Elastic Modulus, E Elastic modulus F L A = E o L o E ~ curvature at ro Energy E is larger if Eo is larger. unstretched length r o smaller Elastic Modulus larger Elastic Modulus r

27 Types of Bondings 1)Primary Bonding: Ionic bonding Covalent bonding Metallic bonding 2)Secondary bonding: Fluactuating Induced Dipole Polar Molecule-Induced Dipole Permanent Dipole Bonds

28 Found in compounds formed by metallic and nonmetallic elements (occurs between + and ions) Requires electron transfer. Ionic Bonding Large difference in electronegativity is required. Example: NaCl E net = E att. + E rep. E net = - (A/r) + B/r n B, and n are constants. n is approximately 8. (A = (Z 1 Z 2 e 2 /4π o ) + B/r n ) The ionic bonding is nondirectional, that is the magnitude of the bond is equal in all directions. The predominant bonding in ceramics is ionic.

29 Give up electrons Acquire electrons

30 Covalent Bonding Stable electron configurations are assumed by sharing of electrons between adjacent atoms. Example: CH4(methane) H feels like helium electron configuration, while C feels like neon electron configuration. C: has 4 valence e, needs 4 more Electronegativities are comparable. H: has 1 valence e, needs 1 more Covalent bonding is directional It forms between two specific atoms and may exist only in the direction between one atom and another.

31 Molecules with nonmetals Molecules with metals and nonmetals Elemental solids (RHS of Periodic Table) Compound solids (about column IVA)

32 # = 8-N Number of covalent bonds N = number of valence electrons Example: Cl atom 7 valence electrons, an atom can have maximum 1 more bond (completing the valence orbital electron number to eight) Covalent bonds may be extremely strong (like in diamonds) or may be weak (like in Bismuth).Polymeric materials are covalently bonded materials. %Ionic character Some bonds are partially ionic and partially covalent. The degree of either bond is controlled by the electronegativities of the composing atoms. %ionic character = (1-e -(0.25)(X A-X B )2 )x100 (XA and XB are the electronegativities of the respective elements) As the electronegativity difference gets higher, the bonding becomes more ionic.

33 Metallic Bonding Valence electrons are not bound to any particular atom in the solid and they are more or less free to move throughout the entire metal. Primary bond for metals and their alloys Metallic bond is nondirectional. Metalling bonding explains the heat and electric conductivity of the metallic materials as well as their ductility.

34 Secondary Bonding Arises from interaction between dipoles Fluctuating dipoles Permanent dipoles-molecule induced -general case: -ex: liquid HCl -ex: polymer