Ch 7 Atomic Structure & Periodicity Summary (PDF)

Chapter 7 –Atomic structure & periodicity 7.1 Electromagnetic Radiation 7.2 The Nature of Matter 7.3 The Atomic Spectrum of Hydrogen 7.4 The Bohr Model 7.5 The Quantum Mechanical Description of the Atom 7.7 Orbitals shapes & energy 7.8 Electron Spin and the Pauli Principle 7.9-7.11The Aufba...

Chapter 7 –Atomic structure & periodicity 7.1 Electromagnetic Radiation 7.2 The Nature of Matter 7.3 The Atomic Spectrum of Hydrogen 7.4 The Bohr Model 7.5 The Quantum Mechanical Description of the Atom 7.7 Orbitals shapes & energy 7.8 Electron Spin and the Pauli Principle 7.9-7.11The Aufbau Principle and the Periodic Table 7.12 Periodic Trends in Atomic Properties Chapter 8 Bonding: General Concepts 8.4 Ions: Electron Configurations and Sizes  Electromagnetic Radiation One of the ways that energy travels through space is by electromagnetic radiation. Named so because it has electrical and magnetic fields that simultaneously oscillate in planes mutually perpendicular to each other and to the direction of propagation through space. Figure 7.1 - Electromagnetic radiation has oscillating electric (E) and magnetic (H) fields in planes perpendicular to each other and to the direction of propagation.  Electromagnetic Radiation Waves are periodic oscillations that transmit energy through space They are characterized by wavelength, frequency and speed. Wavelength () Distance between 2 consecutive peaks or troughs in a wave Described by any appropriate unit of distance (m, m=10-6 m, nm=10-9m or pm=10-12 m) Frequency () Number of waves per second that pass a given point in space. Units are : oscillations per second or 1/s = s-1, which is called the hertz (Hz) Since all types of electromagnetic radiation travel through vacuum at the speed of light, short-wavelength must have a high frequency.  = c c = speed of light = 2.99792458 x 108 m/s  Electromagnetic Radiation Figure 7.2 - The nature of waves. Note that the radiation with the shortest wavelength has the highest frequency. This diagram can represent the oscillating electric or magnetic field of the wave.  Electromagnetic Radiation Electromagnetic radiation provide an important means of energy transfer Figure 7.3 - Classification of electromagnetic radiation.  The Nature of Matter At the end of the 19th century, it was thought that matter and energy were distinct. Matter was thought to consist of particles (specific mass and position in space) and energy was described as a wave (massless and delocalized). It was assumed that energy is continuous and that any amount of energy could be released in any radiation process.  The Nature of Matter In 1901, Max Planck discovered that energy is not continuous. It can be gained or lost in whole-number multiples of the quantity h h (Planck’s constant) = 6.626 10-34 J.s h=hc/ λ The energy is quantized and can be transferred only in discrete units of size h. Each of these small “packets” of energy is called quantum. A system can transfer energy only in whole quanta. Thus energy seems to have particulate properties.  The Nature of Matter Example 7.1 The blue color in fireworks is often achieved by heating copper (I) chloride (CuCl) to about 1200°C. The hot compound emits blue light having a wavelength of 450 nm. What is the increment of energy (the quantum) that is emitted at 4.5 x 102 nm by CuCl?  The Nature of Matter SOLUTION ΔE = h = c = 2.9979 x 108 m/s = 6.66 x 1014 s-1 λ 450 x 10-9 m So ΔE = h = (6.626 x 10-34 J s)(6.66 x 1014 s-1) = 4.41 x 10-19 J  The Nature of Matter Dual nature of light: wave and particle properties Figure 7.4 - Electromagnetic radiation exhibits wave properties and particulate properties. The energy of each photon of the radiation is related to the wavelength and frequency by the equation Ephoton = h = hc/.  The Nature of Matter Louis de Broglie - 1923: - The particle and the wave properties are related by the expression: λ is the wavelength (m) h λ= h is the Planck’s constant (J.s or Kg m2 /s) mv m is the mass (Kg) v is the velocity (m/s) - The wave properties become observable only for submicroscopic objects due to the smallness of Planck’s constant.  The Nature of Matter Example 7.2 Compare the wavelength for an electron (m= 9.11 x 10-31 kg) traveling at a speed of 1 x 107 m/s with that for a ball (mass = 0.1 Kg) traveling at 35 m/s.  The Atomic Spectrum of Hydrogen When a high energy discharge is passed through a sample of hydrogen gas, the H2 molecules absorb energy, causing some of the H-H bonds to break. The resulting hydrogen atoms are excited; that is, they contain excess energy, which they release by emitting light of various wavelengths to produce what is called the emission spectrum of the hydrogen atom. White light – prism – continuous spectrum- all wavelengths of visible light Hydrogen emission spectrum in the visible – prism – few lines (discrete wavelengths) – line spectrum  The Atomic Spectrum of Hydrogen Figure 7.5 - (a) A continuous spectrum containing all wavelengths of visible light (indicated by the first letters of the colors of the rainbow). (b) The hydrogen line spectrum contains only a few discrete wavelengths.  The Bohr Model In 1913, Bohr developed a model for the hydrogen atom. - e- in a hydrogen atom moves around the nucleus only in certain allowed circular orbits. - by using the theories of classical physics and by making some new assumptions, he gave the expression for the energy levels available to the electron in the hydrogen atom E = -2.178 x 10-18 J (1 /n2) - n is an integer, the larger its value, the larger is the orbit radius - The negative sign means that the energy of the e- bound to the nucleus is lower than it would be if the electron were at an infinite distance ( n= ∞). -1 is the atomic number Z of Hydrogen - This equation can be used to calculate the change in energy when the electron changes orbits.  The Bohr Model Calculate the change in energy when the electron in an excited hydrogen atom falls back from n=6 to n=1. E6= -2.178 x 10-18 J (1/62) = -6.05 x 10-20 J E1= -2.178 x 10-18 J (1/12) = -2.178 x 10-18 J ΔE = E1 – E6 = (-2.178 x 10-18 ) - (- 6.05 x 10-20 ) = -2.118 x 10-18 J The negative sign for the change in energy indicates that the atom has lost energy and is now more stable. The absolute value of ΔE must be used for the calculation of λ. h c (6.626 x 10-34 J s) (2.9979 x 108 m/s) = 9.379 x 108 m λ= = ΔE 2.118 x 10-18 J  The Bohr Model Summary: A general equation for the electron moving from one level (ninitial) to another level (nfinal): ΔE = Efinal – Einitial = (-2.178 x 10-18 J) (1/nfinal2) - (-2.178 x 10-18 J) (1/ninitial2) = (-2.178 x 10-18 J) (1/nfinal2 - 1/ninitial2)  The Bohr Model Example 7.3 Calculate the energy required to excite the hydrogen electron from level n=1 to level n=2. Also calculate the wavelength of light that must be absorbed by a hydrogen atom in its ground state to reach this excited state.  The Bohr Model SOLUTION E1= -2.178 x 10-18 J (1/12) = -2.178 x 10-18 J E2= -2.178 x 10-18 J (1/22) = -5.445 x 10-19 J ΔE = E2 – E1 = (-5.445 x 10-19) - (-2.178 x 10-18) = 1.634 x 10-18 J The positive value of indicates that the system has gained energy. h c = (6.626 x 10-34 J s) (2.9979 x 108 m/s) = 1.216 x 10-7 m λ= ΔE 1.634 x 10-18 J  The Bohr Model Limitations of the Bohr Model - The energies calculated by Bohr agreed with the values obtained from the hydrogen emission spectrum. - However, when applied to atoms other than hydrogen, it did not work at all.