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Chapter 5 Periodicity & Atomic Structure 1 The Periodic Table 01 • The periodic table is the most important organizing principle in chemistry. • Chemical and physical properties of elements in the same group are similar. • All chemical and physical properties vary in a periodic manner, hence the name periodic table. 2 Development of the Periodic Table Mendeleev’s Periodic Table (1871) Until the discovery of the proton, the elements were typically organized by increasing atomic weight. The modern organization is by increasing atomic number. 3 The Periodic Table 04 4 Ar = 39.95 amu, would appear on the right of K = 39.10 amu Development of Modern Periodic Table • In 1913 Moseley discovered that when elements were irradiated with high energy radiation y emitted X-ray. He used the frequency of the emitted radiation to calculate the atomic number. √ υ = a( Z-b) The Periodic Table 03 6 What is light made of ? 1) Made of Waves? Waves interfere, 2) Made of particles? 3) Made of both? 4) What are electron’s made of? 8 Electromagnetic Radiation 02 9 Electromagnetic Radiation 01 • Frequency (, Greek nu): Number of peaks that pass a given point per unit time. • Wavelength (, Greek lambda): Distance from one wave peak to the next. • Amplitude: Height measured from the center of the wave. The square of the amplitude gives intensity. 10 Electromagnetic Radiation 03 11 Electromagnetic Radiation 04 • Speed of a wave is the wavelength (in meters) multiplied by its frequency (in reciprocal seconds, (s–1) ). – – Wavelength x Frequency = Speed (m) x (s–1) = c (m/s) 12 Learning Check • The red light in a laser pointer comes from a diode laser that has a wavelength of about 630 nm. What is the frequency of the light? C = 2.9979 x 108 m.s–1 13 A photon has a frequency of 6.0 x 104 Hz (s-1). Convert this frequency into wavelength (nm). Does this frequency fall in the visible region? c = 2.9979 x 108 m/s–1 x=c = c/ = 2.9979 x 108 m/s / 6.0 x 104 Hz = 5.0 x 103 m = 5.0 x 1012 nm Radio wave 14 Atomic Spectra 01 • Atomic spectra: Result from excited atoms emitting light. • Line spectra: Result from electron transitions between specific energy levels. 15 Maxwell (1873), proposed that visible light consists of electromagnetic waves. Electromagnetic radiation is the emission and transmission of energy in the form of electromagnetic waves. c = 2.9979 x 108 m/s–1 Speed of light (c) in vacuum = 3.00 x 108 m/s All electromagnetic radiation x=c 16 7.1 Line Emission Spectrum of Hydrogen Atoms 17 7.3 18 “Photoelectric Effect” h Depending on metal used only lights of certain minimum frequency could cause ejection of electron KE e- Light must be composed of particles called photon Planks and Einstein Light has both: 1. wave nature 2. particle nature 19 Particle like Properties of Electromagnetic Energy 20 Planck in 1900 Energy (light) is emitted or absorbed in discrete units (quantum). E=hx Planck’s constant (h) h = 6.63 x 10-34 J•s 21 “Photoelectric Effect” 22 “Photoelectric Effect” Solved by Einstein in 1905 Light has both: 1. wave nature 2. particle nature h KE e- Photon is a “particle” of light 23 Bohr’s Model of the Atom (1913) What happens in a hydrogen lamp? Line Emission Spectrum of Hydrogen Atoms 1. e- can only have specific (quantized) energy values 2. light is emitted as emoves from one energy level to a lower energy level En = -RH ( 1 n2 ) n (principal quantum number) = 1,2,3,… RH (constant) = 2.18 x 10-18J 24 E = h E = h 25 Hydrogen Atomic Spectra Line Emission Spectrum of Hydrogen Atoms 26 ni = 3 ni = 3 ni = 2 nf = 2 Ephoton = DE = Ef - Ei 1 Ef = -RH ( 2 nf 1 Ei = -RH ( 2 ni 1 DE = RH( 2 ni ) ) 1 n2f ) RH (constant) = 2.18 x 10-18J nnf f==11 m = initial n = final 1 1 1 =R 2 E = h x C/λ m n2 27 R = 1.097 X 10-2 nm-1 Why is e- energy quantized? De Broglie (1924) reasoned that e- is both particle and wave. = h/mv v = velocity of em = mass of eh = 6.63 x 10-34 J•s 28 What is the de Broglie wavelength (in nm) associated with a 2.5 g Ping-Pong ball traveling at 15.6 m/s? = h/mv h in J•s m(mass) in kg V( velocity) in (m/s), J = Kg.m2/S2) = 6.63 x 10-34 J.S ((Kg.m2/S2)/J) / ((2.5 x 10-3 Kg) x (15.6 m/s)) = 1.7 x 10-32 m = 1.7 x 10-23 nm 29 Uncertainty Principle W. Heisenberg 1901-1976 Problem of defining nature of electrons in atoms solved by W. Heisenberg. Cannot simultaneously define the position and momentum (= m•v) of an electron. We define e- energy exactly but accept limitation that we do not know exact position. 30 http://www.youtube.com/watch?v=DfPeprQ7oGc Schrodinger Wave Equation In 1926 Schrodinger wrote an equation that described both the particle and wave nature of the eWave function (Y, psi) describes: 1. energy of e- with a given Y 2. probability of finding e- in a volume of space Schrodinger’s equation can only be solved exactly for the hydrogen atom. Must approximate its solution for multi-electron systems. 31 Schrodinger Wave Equation Y = fn(n, l, ml, ms) Existence (and energy) of electron in atom is described by its unique wave function Y. Pauli exclusion principle - no two electrons in an atom can have the same four quantum numbers. Each seat is uniquely identified (E, R12, S8) Each seat can hold only one individual at a time 32 Schrodinger Wave Equation Y = fn(n, l, ml, ms) principal quantum number n n = 1, 2, 3, 4, …. distance of e- from the nucleus n=1 n=2 n=3 33 Electron Radial Distribution 01 34 Electron Radial Distribution 02 • s Orbital Shapes: 35 1s Orbital 36 2s Orbital 37 3s Orbital 38 Schrodinger Wave Equation Y = fn(n, l, ml, ms) angular momentum quantum number l for a given value of n, l = 0, 1, 2, 3, … n-1 n = 1, l = 0 n = 2, l = 0 or 1 n = 3, l = 0, 1, or 2 l=0 l=1 l=2 l=3 s orbital p orbital d orbital f orbital Shape of the “volume” of space that the e- occupies 39 QUANTUM NUMBERS Y = fn(n, l, ml, ms) The shape, size, and energy of each orbital is a function of 3 quantum numbers: n (major) ---> shell l (angular) ---> subshell ml (magnetic) ---> designates an orbital within a subshell 40 QUANTUM NUMBERS Symbol Values Description n (major) 1, 2, 3, .. Orbital size and energy l (angular) 0, 1, 2, .. n-1 Orbital shape and energy (subshell) ml (magnetic) -l..0..+l Orbital orientation # of orbitals in subshell = 2 l + 1 41 Schrodinger Wave Equation Y = fn(n, l, ml, ms) angular momentum quantum number l for a given value of n, l = 0, 1, 2, 3, … n-1 n = 1, l = 0 n = 2, l = 0 or 1 n = 3, l = 0, 1, or 2 l=0 l=1 l=2 l=3 s orbital p orbital d orbital f orbital Shape of the “volume” of space that the e- occupies 42 l = 0 (s orbitals) l = 1 (p orbitals) 43 7.6 Electron Radial Distribution 03 • p Orbital Shapes: 44 l = 2 (d orbitals) 45 7.6 Schrodinger Wave Equation Y = fn(n, l, ml, ms) magnetic quantum number ml for a given value of l ml = -l, …., 0, …. +l if l = 1 (p orbital), ml = -1, 0, or 1 if l = 2 (d orbital), ml = -2, -1, 0, 1, or 2 orientation of the orbital in space 46 7.6 ml = -1 ml = -2 ml = 0 ml = -1 ml = 0 ml = 1 ml = 1 ml = 2 47 Schrodinger Wave Equation Y = fn(n, l, ml, ms) spin quantum number ms ms = +½ or -½ ms = +½ ms = -½ 48 Schrodinger Wave Equation Y = fn(n, l, ml, ms) Existence (and energy) of electron in atom is described by its unique wave function Y. Pauli exclusion principle - no two electrons in an atom can have the same four quantum numbers. Each seat is uniquely identified (E, R12, S8) Each seat can hold only one individual at a time 49 Schrodinger Wave Equation Y = fn(n, l, ml, ms) angular momentum quantum number l for a given value of n, l = 0, 1, 2, 3, … n-1 n = 1, l = 0 n = 2, l = 0 or 1 n = 3, l = 0, 1, or 2 l=0 l=1 l=2 l=3 s orbital p orbital d orbital f orbital Shape of the “volume” of space that the e- occupies 50 7.6 f-Orbitals • 51 Schrodinger Wave Equation Y = fn(n, l, ml, ms) Shell – electrons with the same value of n Subshell – electrons with the same values of n and l Orbital – electrons with the same values of n, l, and ml How many electrons can an orbital hold? If n, l, and ml are fixed, then ms = ½ or - ½ Y = (n, l, ml, ½) or Y = (n, l, ml, -½) An orbital can hold 2 electrons 52 How many 2p orbitals are there in an atom? n=2 If l = 1, then ml = -1, 0, or +1 2p 3 orbitals l=1 How many electrons can be placed in the 3d subshell? n=3 3d l=2 If l = 2, then ml = -2, -1, 0, +1, or +2 5 orbitals which can hold a total of 10 e53 Effective Nuclear Charge 01 • Electron shielding leads to energy differences among orbitals within a shell. • Net nuclear charge felt by an electron is called the effective nuclear charge (Zeff). 54 Effective Nuclear Charge 02 • Zeff is lower than actual nuclear charge. • Zeff increases toward nucleus • Energy of electron 1)The higher the main shell the more the energy of electron. 2)Subshell also contribute to the energy of electron: ns > np > nd > nf • This explains certain periodic changes observed. 55 Effective Nuclear Charge 03 56 Energy of orbitals in a multi-electron atom Energy depends on n and l n=3 l = 2 n=3 l = 0 n=2 l = 0 n=1 l = 0 n=3 l = 1 n=2 l = 1 57 “Fill up” electrons in lowest energy orbitals (Aufbau principle) ?? Be Li B5 C 3 64electrons electrons 22s 222s 22p 12 1 BBe Li1s1s 1s 2s H He12electron electrons He H 1s 1s12 58 The most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins (Hund’s rule). Ne97 C N O F 6 810 electrons electrons electrons 22s 222p 22p 5 246 3 Ne C N O F 1s 1s222s 59 Order of orbitals (filling) in multi-electron atom 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s 60 < 4f< 5d< 6p< 7s< 5f< 6d< 7p Electron Configuration of Atoms • Rules of Aufbau Principle: 1. Lower energy orbitals fill first. 2. Each orbital holds two electrons; each with different ms. 3. Half-fill degenerate orbitals before pairing electrons. 61 What is the electron configuration of Mg? Mg 12 electrons 1s < 2s < 2p < 3s < 3p < 4s 1s22s22p63s2 2 + 2 + 6 + 2 = 12 electrons Abbreviated as [Ne]3s2 What are the possible quantum numbers for the last (outermost) electron in Cl? Cl 17 electrons 1s22s22p63s23p5 1s < 2s < 2p < 3s < 3p < 4s 2 + 2 + 6 + 2 + 5 = 17 electrons Last electron added to 3p orbital n=3 l=1 ml = -1, 0, or +1 ms = ½ or -½ 62 4f 5f 63 ns2np6 ns2np5 ns2np4 ns2np3 ns2np2 ns2np1 d10 d5 d1 ns2 ns1 Electron Configuration and the Periodic Table Electron Configuration of Atoms 06 64 65 Using periodic table write Noble gas notation for the following elements: a)S b)Fe [Ne]3s23p4 [Ar] 4s23d6 c)Se [Ar] 4s23d104p4 d)Gd [Xe]6s24f75d1 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s 66 < 4f< 5d< 6p< 7s< 5f< 6d< 7p 67 Electron Configuration of Atoms 08 • Anomalous Electron Configurations: Result from unusual stability of half-filled & full-filled subshells. • Chromium should be [Ar] 4s2 3d4, but is [Ar] 4s1 3d5 • Copper should be [Ar] 4s2 3d9, but is [Ar] 4s1 3d10 68 Electron Spin Quantum Number Diamagnetic: NOT attracted to a magnetic field Paramagnetic: substance is attracted to a magnetic field. Substance has unpaired electrons. 69 * See page 261 of your book for molecular structure of N2 vs O2 Molecular Orbital Theory: Other Diatomic Molecules Paramagnetic unpaired electrons 2p Diamagnetic all electrons paired 2p 71 Periodic Properties 01 72 Effective nuclear charge (Zeff) is the “positive charge” felt by an electron. Zeff = Z - s 0 < s < Z (s = shielding constant) (Sigma) Zeff Z – number of inner or core electrons Z Core Na 11 10 1 186 Mg 12 10 2 160 ~ Zeff Radius/pm Al 13 10 3 143 Si 14 10 4 132 Within a Period as Zeff increases radius decreases 73 Size of the atoms 74 (pm) 75