Atomic structure & interatomic bonding Chapter two 1
Atomic Structure Mass Charge Proton 1.67 х 10-27 kg + 1.60 х 10-19 C Neutron 1.67 х 10-27 kg Neutral Electron 9.11 х 10-31 kg - 1.60 х 10-19 C Electron mass = 1/1836 that of a proton Radius of an atom= 0.1 nm = 0.1 x10-9 m (1Angstrom) 50,000,000 atoms lined up measure 10mm!!! Nucleus takes up 10-14 of the total volume of atom and has diameter of 4-15 fm (femtometer = 10-15 m) Precision: How closely measurements of the same quantity come to each other. Accuracy: How close an experimental observation lies to the true value. 2
General Notes: # of protons gives chemical identification of the element # of neutrons (N) defines isotope number # of protons = atomic number (Z) Atomic mass (A) = mass of protons + mass of neutrons The atomic mass unit (amu) is often used to express atomic weight. 1 amu is defined as 1/12 of the atomic mass of the most common isotope of carbon atom that has 6 protons (Z=6) and six neutrons (N=6). m proton m neutron = 1.67 х 10-27 kg = 1 amu. The atomic mass of the 12 C atom is 12 amu. Atomic mass (A) atomic number (Z) + # of neutrons (N) 3
Atomic weight = weighted average of the atomic masses of the atoms naturally occurring isotopes. Atomic weight of carbon is 12.011 amu. The atomic weight is often specified in mass per mole. A mole is the amount of matter that has a mass in grams equal to the atomic mass in amu of the atoms (A mole of carbon has a mass of 12 grams). The number of atoms in a mole is called the Avogadro number, N av = 6.023 10 23. N av = 1 gram/1 amu. Atomic weight of Fe = 55.85 amu/atom = 55.85 g/mol The number of atoms per cm 3, n, for material of densityρ(g/cm 3 ) and atomic mass A (g/mol): n = N av ρ/ A 4
Examples: Graphite (carbon): ρ = 2.3 g/cm 3, A = 12 g/mol n = 6.023 10 23 atoms/mol 2.3 g/cm 3 / 12 g/mol = 11.5 10 22 atoms/cm 3 Diamond (carbon): ρ = 3.5 g/cm 3, A= 12 g/mol n = 6.023 10 23 atoms/mol 3.5 g/cm 3 / 12 g/mol = 17.5 10 22 atoms/cm 3 Water (H 2 O) ρ = 1 g/cm 3, A w = 18 g/mol n = 6.023 10 23 molecules/mol 1 g/cm 3 / 18 g/mol = 3.3 10 22 molecules/cm 3 For material with n = 6 10 22 atoms/cm 3 calculate mean distance between atoms L. 5
Atomic Models: The electrons form a cloud around the nucleus. This picture looks like a mini planetary system. But quantum mechanics tells us that this analogy is not correct!! Electrons move not in circular orbits, but in 'fuzzy orbits. Actually, we cannot tell how it moves, but only can say what is the probability of finding it at some distance from the nucleus. Only certain orbits or shells of electron probability densities are allowed. 6
The shells & electrons are identified by four quantum number, n, l, m l and m s The quantum numbers arise from solution of Schrodinger s equation. Pauli Exclusion Principle: only one electron can have a given set of the four quantum numbers. Now we can give a short description for the quantum numbers. 7
Primary Quantum Number n Can have values from 1 to infinity, but they can only be integers K, L, M, N Represents the energy of the orbital, which is also related to the size of the orbital An orbital is the region of space where you are likely to find the electron 8
Angular Momentum Quantum Number l Shape of the orbital Can have values from 0 to n-1 s, p, d, f, g, h.. If there is more than one electron present, the angular momentum quantum number also affects the orbital energy (also called the azimuthal quantum number) 9
Magnetic Quantum Number m l Can have integer values from l to +l Thus, if n=1, l =0, and ml must equal 0 In other words, it can only have one value If n=2, then l can equal either 0 or 1 If it equals 1, then m l can equal 1, 0 or +1 It can have three values 10
Remember s orbitals correspond to l = 0 p orbitals correspond to l = 1 1 3 d orbitals correspond to l = 2 f orbitals correspond to l = 3 How many orbitals are possible for each of these types? 5 7 11
Spin Quantum Number ms +1/2-1/2 Two electrons of opposite spin fill each orbital 12
The first three quantum numbers define an orbital You need all four to define an electron Shorthand Notation Germanium has 32 protons and 32 electrons 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 2 13
Electron Shells Bonding occurs only with the electrons in the outer most shells called the valence electrons Inner electrons are called the core electrons The valence electrons are those in the outer s and p orbitals, and any unfilled d and p orbitals. 14
1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 2 Core electrons Valence electrons 1s 2 2s 2 2p 6 3s 2 3p 6 3d 2 4s 2 Chemistry happens in the valence shell 15
Fundamental Concepts First ionization energy (IE): it s also called ionization potential, it is the energy required to remove the most weakly bound electron from an isolated gaseous atom Atom (g) + IE = positive ion (g) + e - and can be calculated from the equation: IE = 13.6 Z 2 / n 2 16
Fundamental Concepts Electron Affinity (EA): the reverse process to the ionization energy, it is the energy change associated with an isolated gaseous atom accepting one electron Atom (g) + e - = negative ion (g) EA : positive if energy released. : negative if energy required. 17
Fundamental Concepts Atomic and ionic radii : in general, positive ions are smaller than neutral atoms, while negative ions are larger. 18
Fundamental Concepts Electronegativity (χ) : independent measure for atom attraction to electrons from another atom in a bond forming. It can be calculated from: or from: (χ) = (IE + EA) / 2 (χ) = [ {0.31 (n + 1 ± c) } / r ] + 0.5 n: # of valence electrons. c: any formal valence charge on the atom. r: covalent radius. 19
Periodic Table S-block d-block p-block f-block 20
Trends in The Periodic Table (IE) (EA) Atomic & ionic radii Electronegativity 21
Atomic Bonding: The bond is an electrostatic force that bind atoms or molecules together. Binding Energy: There are both attractive and repulsive forces acting on atoms When they are balanced a bond is formed When the total energy of a pair of atoms is minimized, a bond is formed 22
Attraction Attraction force F A Force Repulsion r o Interatomic separation Net force F N ( F A +F R ) Repulsion force F R Attraction Potential energy; E Repulsion E o Attraction energy Interatomic separation Net energy (E= min) Repulsion energy 23
E o : Bonding energy. r 0 : Bond length. Types of bonds: Primary Strong Chemical secondary weak physical Primary bond is created when there is direct interaction of electrons between two or more atoms. Secondary bond occur due to indirect interaction of electrons in adjacent atoms or molecules. 24
Primary bonds: Electronegativity controls how elements combine (bond) with each other because it provides a measure of the excess binding energy between atoms A and B, A-B (in kj/mol) : A-B = 96.5 ( χ A χ B ) 2 The excess binding energy is related to the energy required to separate two bonded atoms, bond dissociation energy, DE AB : A-B = DE AB [ (DE AA ) (DE BB )] 1/2 25
Types of Primary bonds: Electronegativity difference > 2.0 Ionic bond Electronegativity difference < 0.4 covalent bond 0.4 < Electronegativity difference < 2.0 Polar covalent bond Special types of primary bonds is metallic bond 26
Ionic Bonds Metal-Nonmetal Cation-anion Non-directional Poor electrical conductivity Poor thermal conductivity Ceramics are formed from ionic bonds What is a molecule? 27
Covalent Bonds Nonmetal nonmetal Directional bonds Poor electrical conductivity Poor thermal conductivity Polymers are covalently bonded Compounds 28
Polar covalent bond: A bond neither truly ionic nor totally covalent. (HF) Partial Ionic character: Ionic character % = 100 [ 1- exp { -0.25 ( χ A χ B ) 2 }] 29
Metallic Bonds Metal-Metal Non-directional Electrons are free to move around Good electrical conductivity Good thermal conductivity What is a molecule? 30
Mixed bonding 2 or more metals may form an intermetallic compound A mixture of ionic and metallic bonds Ceramics (nonmetal metal) are usually a mixture of ionic and covalent As the electronegativity difference increases, the bond becomes more ionic 31
Secondary Bonding Metallic compounds and ionic compounds form crystals. But how do molecules of covalently bonded elements stick together? Secondary Bonds 32
Types of Secondary Forces Van der Waal s: Dipole-dipole forces. Dipole-induced dipole London dispersion forces Hydrogen bonding What is dipole? Atom or molecule that have some separation of positive and negative portions 33
Van der Waal s Interaction between permanent dipoles The interaction of permanent dipoles (analogous to magnets but having an electrostatic dipole moment). CH 3 Cl CH 3 Cl 34
Dipole induced dipole A permanent dipole moment can induce a dipole in a neighboring molecule in which the unperturbed centers of positive and negative charge are otherwise coincident CH 4 CH 3 Cl 35
London dispersion forces: No permanent dipole It does have an instantaneous dipole moment These instantaneous dipoles orient themselves with their neighbors to give an overall force of attraction. 36
H-bonding Hydrogen bonding: can be viewed as the interaction between very strong dipoles in - OH or -NH2 groups. It is reflected in the very high boiling point of water compared with molecules of similar size. H 2 O H 2 O 37
Bonding Energy How does bonding energy relate to melting point? Modulus of Elasticity? Coefficient of Thermal Expansion? Hint: The higher the bonding energy the more tightly the atoms are held together. 38