Chapter 3B Modern Atomic Theory 1 CHAPTER OUTLINE Waves Electromagnetic Radiation Dual Nature of Light Bohr Model of Atom Quantum Mechanical Model of Atom Electron Configuration Electron Configuration & Periodic Table Abbreviated Electron Configuration 2 1
WAVES q All Wavelength waves are (λ) characterized is the distance by between wavelength, any 2 frequency successive and crests speed. or troughs. wavelength (measured wavelength from peak to (measured peak) from trough to trough) 3 10.1 WAVES q Frequency (nu,ν) is the number of waves produced per unit time. q Wavelength and frequency are inversely proportional. As wavelength of a wave increases its frequency decreases inversely q proportional Speed tells how fast waves travel through space. 4 10.1 2
ELECTROMAGNETIC RADIATION q Energy travels through space as electromagnetic radiation. This radiation takes many forms, such as sunlight, microwaves, radio waves, etc. q In vacuum, all electromagnetic waves travel at the speed of light (3.00 x 10 8 m/s), and differ from each other in their frequency and wavelength. 5 ELECTROMAGNETIC RADIATION q These classification waves range of from electromagnetic γ-rays (short waves λ, high f) according to radio Short waves to their (long frequency λ, low Long f). is called electromagnetic wavelength spectrum. wavelength High frequency Low frequency 6 3
ELECTROMAGNETIC RADIATION Visible light is a X-rays have longer Infrared waves have small part of the EM λ but lower υ than longer λ but lower υ spectrum γ-rays than visible light 7 10.2 DUAL NATURE OF LIGHT q When white light is passed through a glass prism, it is dispersed into a spectrum of colors. q This is evidence of the wave nature of light. 8 4
DUAL NATURE OF LIGHT q Red Scientists light has also longer have wavelength much evidence and less that energy light than beams blue light act as a stream of tiny particles, called photons. A photon of red light A photon of blue light 9 DUAL NATURE OF LIGHT q Scientists, therefore, use both the wave and particle models for explaining light. This is referred to as the wave-particle nature of light. q Flatland Video q Scientists also discovered that when atoms are energized at high temperatures or by high voltage, they can radiate light. Neon lights are an example of this property of atoms. 10 5
ATOMIC LINE SPECTRUM q When These Each element lines the light indicate possesses from that the a atom light unique is formed placed line spectrum only at through certain that can wavelengths a be prism, used to a series identify and frequencies of it. brightly colored that lights, correspond called to a specific line spectrum colors. is formed. Each line represents a particular λ and 11 BOHR MODEL OF ATOM q Neils Bohr s In this Bohr, model, a Danish of the the physicist, studied the hydrogen atom electrons consisted could atom of only extensively, and developed a model electrons occupy for particular orbiting the atom the energy that was able to explain the line nucleus levels, spectrum. and at different could jump distances to higher levels from the by nucleus, absorbing called energy. energy levels. 12 6
BOHR MODEL OF ATOM q The lowest energy level is called ground state, and the higher energy levels are called excited states. q When electrons absorb energy through heating or electricity, they move to higher energy levels and become excited. energy 13 BOHR MODEL OF ATOM q When excited electrons return to Lower the ground energy state, energy is emitted transition as a photon of light give off is released. red q The color (wavelength) light of the light Higher emitted energy is determined transition by the give difference off in blue energy light between the two states (excited and ground). 14 7
BOHR MODEL OF ATOM q q q The line spectrum is produced by many of these transitions between excited and ground states. Bohr s model was able to successfully explain the hydrogen atom, but could not be applied to larger atoms. Quantum Mechanics & Structure of Atom 15 QUANTUM MECHANICAL MODEL OF ATOM q q q q In 1926 Erwin Shrödinger created a mathematical model that showed electrons as both particles and waves. This model was called the quantum mechanical model. Double-Slit Experiment This model predicted electrons to be located in a probability region called orbitals. An orbital is defined as a region around the nucleus where there is a high probability of finding an electron. 16 8
QUANTUM MECHANICAL MODEL OF ATOM q Based on this model, there are discrete principal energy levels within the atom. q Principal As n energy increases, levels the are designated energy by n. of the q The electrons in increases an atom can exist in any principal energy level. 17 QUANTUM MECHANICAL MODEL OF ATOM q Each principal energy level is subdivided into sublevels. q The sublevels are designated by the letters s, p, d and f. q As n increases, the number of sublevels increases. 18 10.7, 10.8 9
QUANTUM MECHANICAL MODEL OF ATOM q Within The number the sublevels, of orbitals the within electrons the sublevels are located in orbitals. vary with The their orbitals type. are also designated by the letters s, p, d and f. s sublevel = 1 orbital = 2 electrons p sublevel = 3 orbitals d sublevel = 5 orbitals f sublevel = 7 orbitals = 6 electrons = 10 electrons = 14 electrons An orbital can hold a maximum of 2 electrons 19 ELECTRON CONFIGURATION q Similarities of behavior in the periodic table are due to the similarities in the electron arrangement of the atoms. This arrangement is called electron configuration. q The modern model of the atom describes the electron cloud consisting of separate energy levels, each containing a fixed number of electrons. q Each orbital can be occupied by no more than 2 electrons, each with opposite spins. 20 10
ELECTRON CONFIGURATION q q q q The electrons occupy the orbitals from the lowest energy level to the highest level. The energy of the orbitals on any level are in the following order: s < p < d < f. Each orbital on a sublevel must be occupied by a single electron before a second electron enters. For example, all three p orbitals must contain one electron before a second electron enters a p orbital (Hund s Rule). Visualizing Orbitals 21 ELECTRON CONFIGURATION q Electron configurations can be written as: 2 p 6 Number of electrons in orbitals Principal energy level Type of orbital 22 11
ELECTRON CONFIGURATION q Another notation, called the orbital notation is shown below: Electrons in orbital with opposing spins Principal energy level 1 s Type of orbital 23 ELECTRON CONFIGURATION H 1s 1 1s Hydrogen has 1 electron. It will occupy the orbital of lowest energy which is the 1s. He 1s 2 1s Helium has two electrons. Both helium electrons occupy the 1s orbital with opposite spins. 24 12
ELECTRON CONFIGURATION Li 1s 2 2s 1 1s 2s The 1s orbital is filled. Lithium s third electron will enter the 2s orbital. Be 1s 2s 1s 2 2s 2 The 2s orbital fills upon the addition of beryllium s third and fourth electrons. 25 B 1s 2 2s 2 2p 1 1s Boron has the first p electron. The three 2p orbitals have the same energy. It does not matter which orbital fills first. C 1s 2s ELECTRON CONFIGURATION 2s 2p 2p 1s 2 2s 2 2p 2 The second p electron of carbon enters a different p orbital than the first p due to Hund s Rule. 26 13
N 1s 2 2s 2 2p 3 1s The third p electron of nitrogen enters a different p orbital than its first two p electrons due to Hund s Rule. O 1s 2s ELECTRON CONFIGURATION 2s 2p 2p 1s 2 2s 2 2p 4 The last p electron of oxygen pairs opposite of another since each orbital has an electron in it and Hund s Rule is satisfied. 27 F 1s 2 2s 2 2p 5 1s Two of the p electrons for fluorine pair up with other electrons in the p orbitals. Ne 1s 2s ELECTRON CONFIGURATION 2s 2p 2p 1s 2 2s 2 2p 6 The last p electron for neon pairs up with the last lone electron and completely fills the 2 nd energy level. 28 14
ELECTRON CONFIGURATION q As electrons occupy the 3 rd energy level and higher, some anomalies occur in the order of the energy of the orbitals. q Knowledge of these anomalies is important in order to determine the correct electron configuration for the atoms. q The following study aid is used by beginning students to remember these exceptions to the order of orbital energies. 29 ELECTRON CONFIGURATION 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 30 15
ELECTRON CONFIG. & PERIODIC TABLE q The horizontal rows in the periodic table are called periods. The period number corresponds to the number of energy levels that are occupied in that atom. q The vertical columns in the periodic table are called groups or families. For the main-group elements, the group number corresponds to the number of electrons in the outermost filled energy level (valence electrons). 31 ELECTRON CONFIG. & PERIODIC TABLE 4 One energy energy 3 energy levels level levels 32 16
ELECTRON CONFIG. & PERIODIC TABLE 3 valence 1 valence 5 valence electrons electron electrons 33 ELECTRON CONFIG. & PERIODIC TABLE q The Note valence that elements electrons in the configuration same group for have the elements similar electron in periods configurations. 1-3 are shown below. 34 10.15 17
ELECTRON CONFIG. & PERIODIC TABLE 10.16 Arrangement of orbitals in the periodic table 35 ELECTRON CONFIG. & PERIODIC TABLE d orbital numbers are 1 less than the period number 10.16 36 18
ELECTRON CONFIG. & PERIODIC TABLE f orbital numbers are 2 less than the period number 37 10.16 ELECTRON CONFIG. & PERIODIC TABLE q The electrons in an atom fill from the lowest to the highest orbitals. q The knowledge of the location of the orbitals on the periodic table can greatly help the writing of electron configurations for large atoms. q The energy order of the sublevels is shown next. Note that some anomalies occur in the energy level of d and f sublevels. 38 10.15 19
ELECTRON CONFIG. & PERIODIC TABLE 39 10.15 Example 1: Use the periodic table to write complete electron configuration for phosphorus. P Z = 15 10 5 electrons remaining used 1s 2 2s 2 2p 6 3s 2 3p 3 Core electrons Valence electrons 40 20
Example 2: Draw an orbital notation diagram for the last incomplete level of chlorine and determine the number of unpaired electrons. 3p 3s 41 Example 2: Cl 3s 3p One unpaired electron 42 21
ABBREVIATED ELECTRON CONFIG. q When writing electron configurations for larger atoms, an abbreviated configuration is used. q In writing this configuration, the non-valence (core) electrons are summarized by writing the symbol of the noble gas prior to the element in brackets followed by configuration of the valence electrons. 43 ABBREVIATED ELECTRON CONFIG. K Z = 19 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 Previous noble gas core [Ar] 4s 1 electrons valence electron 44 22
ABBREVIATED ELECTRON CONFIG. Br Z = 35 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 core [Ar] 4s 2 3d 10 4p 5 electrons valence electrons 45 Example 3: Write abbreviated electron configurations for each element listed below: Fe Z = 26 18 8 electrons 620 electrons remaining used remaining used [Ar] 4s 2 3d 6 46 23
Example 3: Write abbreviated electron configurations for each element listed below: Sb Z = 51 [Kr] 5 valence electrons 5s 2 4d 10 15 36 electrons 38 13 electrons 48 3 electrons remaining used remaining used remaining used 5p 3 47 TRENDS IN PERIODIC PROPERTIES q The electron configuration of atoms are an important factor in the physical and chemical properties of the elements. q Some of these properties include: atomic size, ionization energy and metallic character. q These properties are commonly known as periodic properties and increase or decrease across a period or group, and are repeated in each successive period or group. 48 24
ATOMIC SIZE q The size of the atom is determined by its atomic radius, which is the distance of the valence electron from the nucleus. q For each group of the representative elements, the atomic size increases going down the group, because the valence electrons from each energy level are further from the nucleus. 49 ATOMIC SIZE 50 25
ATOMIC SIZE q The atomic radius of the representative elements are affected by the number of protons in the nucleus (nuclear charge). q For elements going across a period, the atomic size decreases because the increased nuclear charge of each atom pulls the electrons closer to the nucleus, making it smaller. 51 ATOMIC SIZE 52 26
IONIZATION ENERGY q The ionization energy is the energy required to remove a valence electron from the atom in a gaseous state. q When an electron is removed from an atom, a cation (+ ion) with a 1+ charge is formed. Na (g) + IE Na + + e - 53 IONIZATION ENERGY q The ionization energy decreases going down a group, because less energy is required to Larger remove atom an electron from the Less outer IE shell since it is further from the nucleus. 54 27
q Going across a period, the ionization energy increases because the increased nuclear charge of the atom holds the valence electrons more tightly and therefore it is more difficult to remove. IONIZATION ENERGY 55 IONIZATION ENERGY q In general, the ionization energy is low for metals and high for non-metals. q Review of ionization energies of elements in periods 2-4 indicate some anomalies to the general increasing trend. 56 28
IONIZATION ENERGY q These anomalies are caused by more stable electron configurations More More of stable stable the atoms in groups 2 (complete s sublevel) Higher (1/2 and filled) IE group 5 (half-filled p sublevels) that cause an Higher increase IE in their ionization energy compared to the next element. Be 1s 2 2s 2 B 1s 2 2s 2 2p 1 N 1s 2 2s 2 2p 3 O 1s 2 2s 2 2p 4 57 METALLIC CHARACTER q Metallic character is the ability of an atom to lose electrons easily. q This character is more prevalent in the elements on the left side of the periodic table (metals), and decreases going across a period and increases for elements going down a group. 58 29
METALLIC CHARACTER Most metallic elements Least metallic elements 59 Example 1: Select the element in each pair with the larger atomic radius: Li or K Larger due to more energy levels 60 30
Example 1: Select the element in each pair with the larger atomic radius: K or Br Larger due to less nuclear charge 61 Example 1: Select the element in each pair with the larger atomic radius: P or Cl Larger due to less nuclear charge 62 31
Example 2: Indicate the element in each set that has the higher ionization energy and explain your choice: K or Na Higher IE due to less energy levels 63 Example 2: Indicate the element in each set that has the higher ionization energy and explain your choice: Mg or Cl Higher IE due to more nuclear charge 64 32
Example 2: Indicate the element in each set that has the higher ionization energy and explain your choice: F N or C Highest IE due to most nuclear charge 65 THE END 66 33