Module1
- 1. Module 1 We will cover the electronic structure of atoms and the types of bonds that can form between atoms. Module 1 -Atomic number and Atomic mass -Electrons -Electronic configuration -Periodic Table -Bonding between atoms 1
- 2. Atom: Basic Element of Matter Atomic Number # of PROTONS in nucleus of atom = # of ELECTRONS of neutral species Atomic mass Mass of PROTON + Mass of NEUTRON 2 Mass Charge Electron 9.11x10-31 kg -1.69x10-19 C Proton 1.67x 10-27 kg 1.69x10-19 C Neutron 1.67x 10-27 kg 0
- 3. 3 Atomic mass unit (amu), 1 amu = 1/12 mass of 12C =1.67x10-27kg 1 atom 12C=12 amu Mole The amount of substance that has as many atoms/molecules as in 12 grams of Carbon. Avogadro’s number 1 mole of a substance has 6.023 x 1023 number of particles Atomic Mass Unit This is defined in units of gram mole. 1 kg-mole will contain atoms6.023x10 26
- 4. 4 Isotopes Variants of the same element having different number of neutrons. Atomic Weight Weighted average of the atomic masses of the naturally occurring isotopes of a particular element in atomic mass unit. As elements have isotopes atomic weight is not an integer for example : C 12.001 or H 1.008 For example, atomic weight of copper is 63.54 amu/atom or 63.54 g/mole
- 5. 5 How many Fe atoms are there in one gram of Fe ? Atomic mass of Fe = 55.85 g/mol How many grams in one amu of a material?
- 6. So, where are the electrons? • Electrons (negatively charged particles) orbit the nucleus. • We will learn about two different models describing the placement of the electrons around the nucleus. 6
- 7. Structure of Atoms- Two Models Quantum Mechanics- set of principles and laws that govern systems of atomic and sub atomic entities. 1) Bohr Model Electrons revolve around nucleus in discrete orbitals. Deals with position (electron orbitals) and energy (quantized energy). 2) Wave Mechanical Model Electrons exhibit both wave-like and particle-like characteristics Probability of electron being in a location - Energy quantized in shells and subshells. 7
- 8. Bohr Model • Protons and neutrons in the nucleus and are held together by the strong force (overcomes electrostatic repulsion). 8 Electrons revolve around a positively charged nucleus in discrete orbits (K, L, M or n=1, 2, 3 respectively) with specific levels of energy. Electrons positions are fixed as such, however, an electron can jump to higher or lower energy level by absorption or emission of energy respectively
- 9. Wave Mechanical Model Electrons are in orbitals defined by a probability. Each orbital at discrete energy level determined by 4 quantum numbers. Quantum numbers : Four parameters that describe the distribution and position of the electrons Principal quantum number, n = principal (energy level-shell) K, L, M, N, O (1, 2, 3, etc.) • Determines the size of atom or distance of shell from nucleus • Number of electrons in a shell is 2n2 9
- 10. 10 Angular quantum number, l = subsidiary (orbitals) It signifies the subshells/electron orbital s, p, d, f The values of l ranges from (0, 1,…, n -1) For example , M shell n=3, so l is 0,1,2 s, p and d Magnetic quantum number, ml It signifies the number of energy states in each orbital. 1, 3, 5, 7 (-l to +l) determines number of energy states in an orbital For example, s orbital l=0 , so ml = 0 One energy state p orbital l=1, so ml = -1,0,+1 Three energy states
- 11. 11 Pauli’s exclusion Principle: It states that no more than two electrons having opposite spin can occupy the same energy state. For example, s orbital has one energy state and can accommodate only two electrons having opposite spin(ms = ±1/2) Spin quantum number, ms It signifies the spin moment associated with the electrons. It can have only two values, +1/2 and -1/2
- 12. 12 n l subshells ml Energy states ms Maximum number of electrons in shell Electronic configuration K shell 1 0 0 ±1/2 2 1s2 L shell 2 0 0 ±1/2 8 electrons 2s2 2p6 One 2s state and Three 2p states 1 +1 ±1/2 0 ±1/2 -1 ±1/2 M shell 3 0 0 ±1/2 18 electrons 3s2 3p6 3d10 One 3s state Three 3p state Five 3d state 1 +1 ±1/2 0 ±1/2 -1 ±1/2 2 +2 ±1/2 +1 ±1/2 0 ±1/2 -1 ±1/2 -2 ±1/2 N shell 4 32 electrons 4s2 4p6 4d10 4f14 It denotes the manner in which the electrons are distributed in the orbital shells. Electronic Configuration
- 13. Determining Electron Configuration: Fill Lowest Energy Sites First Aufbau principle: Electrons will sit in the lowest energy positions first. 13 Madelung’s rule: Orbitals fill in the order of increasing (n+l) value. For elements with same values of (n+l) the one with lower value of n will be filled first.
- 14. What is the electronic configuration for Carbon and Iron? 14
- 15. What is the electronic configuration for Carbon and Iron? • Carbon – Symbol ‘C’ on the periodic chart – Atomic number 6 – Had 6 electrons • Iron – Symbol ‘Fe’ on the periodic chart – Atomic number 26 – Had 26 electrons 15
- 17. Determining Electron Configuration for Carbon 17 1s 2s 2p K-shell n = 1 L-shell n = 2 3s 3p M-shell n = 3 3d 4s 4p 4d Energy N-shell n = 4 L-shell with two subshells K-shell with one subshells
- 18. Example: Iron 18 Lowest energy orbitals are filled first .
- 19. Adapted from Fig. 2.4, Callister 7e. 1s 2s 2p K-shell n = 1 L-shell n = 2 3s 3p M-shell n = 3 3d 4s 4p 4d Energy N-shell n = 4 Determining Electron Configuration for Iron 19
- 20. 20 Electron configuration (stable) ... ... 1s22s22p 63s23p 6 (stable) ... 1s22s22p 63s23p 63d 10 4s24p6 (stable) Atomic # 18 ... 36 Element 1s11Hydrogen 1s22Helium 1s22s13Lithium 1s22s24Beryllium 1s22s22p 15Boron 1s22s22p 26Carbon ... 1s22s22p 6 (stable)10Neon 1s22s22p 63s111Sodium 1s22s22p 63s212Magnesium 1s22s22p 63s23p 113Aluminum ... Argon ... Krypton Valence electrons are most available for bonding and tend to control the chemical properties Filled shells more stable Adapted from Table 2.2, Callister & Rethwisch 3e. Stable configurations are lowest energy and do not take part in any chemical reaction.
- 21. Periodic Table 21 18 numbered groups (columns) and 7 periods (horizontal rows)
- 22. 22 Electropositive elements: Readily give up electrons to become + ions. Electronegative elements: Readily acquire electrons to become - ions. giveup1e- giveup2e- giveup3e- inertgases accept1e- accept2e- O Se Te Po At I Br He Ne Ar Kr Xe Rn F ClS Li Be H Na Mg BaCs RaFr CaK Sc SrRb Y Periodic Table Determine properties from valence electrons
- 23. Electronegativity is a measure of how readily an atom accepts an electron to form a negatively charged ion. Electronegativity 23 Smaller electronegativity Larger electronegativity
- 24. 24 • Atomic radii describes the distance of the outermost energy level from the nucleus. • Atomic radii decrease as move from left to right across of period and increase down a group. • For ions , For example Fe, Fe2+ and Fe3+ , Fe3+ < Fe2+ < Fe Atomic Radii n is constant ‘n’ increases
- 25. Bonding Between Atoms Source: R.E. Smallman, Ch1, F1.3 25 • Type of bond between atoms will depend on each atom’s electronic structure. • The valence electrons take part in bonding. • The atoms – ‘acquire’, ‘loose’or ‘share’ valence electrons to achieve lowest energy configuration.
- 26. 26 Primary Classification Primary Bonds- High amount of energy needed to break bond Secondary Bonds- Lower amount of energy needed to break bonds
- 27. Na (metal) unstable Cl (nonmetal) unstable electron + - Coulombic Attraction Na (cation) stable Cl (anion) stable 27 • Requires electron transfer. • Large difference in electronegativity required. Example: NaCl
- 28. Examples: Ionic Bonding Give up electrons Acquire electrons NaCl MgO CaF2 CsCl 28
- 29. Example: MgO Mg 1s2 2s2 2p6 3s2 O 1s2 2s2 2p4 Mg2+ 1s2 2s2 2p6 O2- 1s2 2s2 2p6 29 Donates electrons Accepts electrons Ionic bonding found between atoms with large difference in electronegativity ---in compounds composed of a metal and non-metal and require the transfer of valence electrons Bonds are non-directional in nature.
- 30. Covalent Bonding C: has 4 valence e-, needs 4 more H: has 1 valence e-, needs 1 more Electronegativities are comparable. Adapted from Fig. 2.10, Callister 7e. Share valence electrons Formed between atoms of similar electronegativity s & p orbitals dominate bonding shared electrons from carbon atom shared electrons from hydrogen atoms H H H H C CH4 30 Example: CH4 These bonds are directional in nature.
- 31. 31 Covalent bonding form through sharing of valence electrons between atoms of similar electronegativity.
- 32. Metallic Bonding • Valence electrons sea/ cloud of electrons which are shared by all atoms. • Rest of electrons + Nuclei ion core. • Sea of electrons shields ion cores. 32 Metallic bonding arises from interaction between sea of electrons and ion Core
- 33. 33 Secondary Bonding – Van der Waals bonding Dipole-Dipole Dipole induced Dipole Hydrogen Bonding Fluctuating dipole (weakest) H Cl H Clsecondary bonding secondary bonding + - + - Example: Liquid HCl Dipole-Dipole interaction: Between molecules with permanent dipole moments. Adapted from Fig. 2.13, Callister 7e. ~3.3kJ/mole
- 34. 34 Hydrogen Bonding Bond strength: 10-50 kJ/mole A form of dipole-dipole interaction. Interaction between hydrogen atom and another atom of high electronegativity such as Fluorine or Oxygen or Nitrogen.
- 35. 35 Water- Anomalous expansion on cooling --- decrease in density
- 36. 36 Polarity vs. Dipole Asymmetric distribution of electrical charge over the atoms joined by a bond or a molecule http://sscchemistry.weebly.com/polaritydipole-and-hybridization.html Non-polar Polar Bond between two non-metal atoms with greater difference in electronegativity – Polar covalent bond. H 𝛿 + −𝐶l 𝛿 − Oxygen – 1s2 2s2 2p4
- 37. 37 Dipole Induced Dipole: Polar molecules (with asymmetric arrangement of positively and negatively charged regions) can induce dipoles in adjacent nonpolar molecules. Example: Dipole in water molecule (polar) induces dipole in oxygen molecule (non- polar)
- 38. 38 asymmetric electron clouds + - + - secondary bonding HH HH H 2 H 2 secondary bonding ex: liquid H2 Fluctuating Dipole Adapted from Fig. 2.14, Callister 7e. Weak attraction forces between dipoles due to natural oscillation of atoms leading to momentary breakdown of charge symmetry and temporary dipoles.
- 39. Bonding of Atoms Types Ionic Covalent Metallic Secondary Bond Energy Large! large-Diamond small-Bismuth Variable large-Tungsten 850 kJ/mole small-Mercury- 68 kJ/mole smallest 39 Large Bond energy Lack of free electrons --Poor conductors of heat and electricity. Brittle in nature. Moderate Bond Energy Valence electrons are free to move -- Good thermal and electrical conductivities. -- Bond strength increases with atomic number—increase in melting point, hardness and strength. -- Agile electrons- Ductile and malleable 4-50 kJ/mole 600-1500 kJ/mole 100-1200kJ/mole
- 40. • Energy – minimum energy most stable Energy balance of attractive and repulsive terms Attractive energy EA Net energy EN Repulsive energy ER Interatomic separation r r A n r B EN = EA + ER = +- Adapted from Fig. 2.8(b), Callister 7e. 40
- 41. Interactions Between Atoms Relationship between interaction energy E and force F Force Total energy: 41 r A n r B E = EA + ER = +- rF(r)dE dr dE(r) F
- 42. 42 How do we find forces and bond length?
- 43. Find Attractive Force: 43 Find Repulsive Force! Find Bond Length!
- 44. Secondary Bonding Example Consider the van der Waals bonding in solid Argon. The potential energy as a function of interatomic separation can be modeled as Calculate the equilibrium bond length and bond energy. A=10.37x10-78J-m6 B=16.16x10-135J-m12 E r( ) = -Ar-6 + Br-12 44
- 45. 45 The Potential we just looked at is called Lennard-Lones potential - 612 4 rr ULG ε is the depth of the potential well σ is the finite distance at which the inter-particle potential is zero Find equilibrium bond length
- 46. 46 Indicate the Equilibrium bond distance and Bond Energy in the following Energy-interatomic distance plots: Eo “bond energy” Energy r o r
- 47. • Bond length, r • Bond energy, Eo • Melting Temperature, Tm Tm is larger if Eo is larger. r o r Energy r larger Tm smaller Tm Eo = “bond energy” Energy r o r unstretched length 47
- 48. 48 A steep slope at the equilibrium bond distance indicates high modulus of elasticity of stiffness. Force vs. Interatomic separation distance plots of Material A and Material B. r0 r0
- 49. 49 Ceramics (Ionic & covalent bonding): Metals (Metallic bonding): Polymers (Covalent & Secondary): Large bond energy large Tm large E small α Variable bond energy moderate Tm moderate E • moderate α Directional Properties Secondary bonding dominates small Tm small E • large α
- 50. Example 50 Calculate the force of attraction between a Ca2+ and an O2- ion the centers of which are separated by a distance of 1.25 nm. Charge on electron =1.67 x10-19 Coulomb
- 51. 51 • Bonding energies between 600 and 1500 kJ/mol (or 3 to 8 eV/atom) are considered to be relatively large with high Tm