9/30/2011 Chapter 11 Atomic Theory: The Quantum Model of the Atom Electromagnetic Radiation Electromagnetic radiation A form of energy that occurs as waves, consists of both electric and magnetic fields, and includes X-rays, ultraviolet, visible light, infrared, microwaves, and radio and television waves. Electromagnetic spectrum The range of possible electromagnetic radiation wavelengths and frequencies. Figure 11.1-The electromagnetic radiation. Visible light is only a small portion of the electromagnetic spectrum, covering a wavelength range of about 400 to 700 nm. There seems to be no upper or lower limit to the theoretical possible lengths of electromagnetic waves, although experimentally measured waves vary between about 1011 to 10-12 m. 1
Electromagnetic Radiation Speed of light, c 3.00 10 8 m/s For short distances, the time needed for light to travel from an object to your eye is imperceptible. Velocity of an electromagnetic wave c = λν λ(lambda) is wavelength ν(nu) is frequency Waves Terms Velocityis the linear speed of a point on a wave. Wavelength(λ) is the distance between corresponding points on a wave. Frequency(ν)is the number of complete waves passing a point per second. c = λν λ and ν have an inverse relationship. Figure 11.2-Wave properties. Water waves are similar to electromagnetic waves in that they can be described by properties such as velocity, wavelength, and frequency. 2
9/30/2011 Electromagnetic Radiation (cont d) Continuous spectrum A spectrum having no distinct lines; the range of wavelengths is uninterrupted. Line spectrum A spectrum that has lines at certain wavelengths and nothing in between those lines. Discrete lines The individually distinct lines in a line spectrum. Figure 11.3-Dispersion of white light by a prism. White light is passed through slits and then with a prism. It is separated into a continuous spectrum of all wavelengths of visible light. A visible spectrum is observed like in figure 11.1. Figure 11.4-Dispersion of light from a gas discharge tube filled with hydrogen. This is like a neon light, except that neon given red light. The magenta light from hydrogen is passed through slits and then through a prism. It is separated into a line spectrum made up of four wavelengths of visible light. 3
Figure 11.5-Line spectra of hydrogen, mercury, and neon. Each element produces a unique spectrum that can be used to identify the element. The hydrogen spectrum corresponds with that shown in Figure 11.4. The neon spectrum is a combination of the colors we see as the red light of a neon sign. Electromagnetic Radiation (cont d) The energy of a wave and its frequency are proportional: E ν. Energy released by electrons is in the form of a masslesspacket of electromagnetic radiation known as a photon. Wave-particle duality Light has both wavelike properties and particle-like properties. Bohr s Model In 1913, NielsBohr, a Danish scientist, suggested that an atom consists of a dense nucleus containing most of the atom s positive charge and mass and surrounding by electrons which travel in regions called orbits. He noted that the orbits are large in size in compared to the atom s nucleus. Most of the space for an atom is empty. 4
9/30/2011 Bohr s Model (cont d) This model would be useful for a hydrogen like atom (1 electron) and he stated that the amount of energy possessed by the electron in a hydrogen atoms and the radius of its orbit are quantized. Quantized means it s a fixed amount and in our case, each orbit possesses a fixed amount of energy. Electrons must reside on the orbit, not between orbits in its transition. Each energy level is an integer (i.e. 1, 2, 3, etc.) represented by n. Figure 11.6-The quantum concept. A woman on a ramp can stop at any level above ground. Her elevation is not quantized. A woman on stairs can stop only on a step. Her elevation is quantized by h1, h2, h3, and h4. Figure 11.7-The Bohr model of the hydrogen atom. The electrons is allowed to circle the nucleus only at certain radii and with certain energies, the first four are shown. Electrons in the ground state level, n = 1, can absorb the exact amount of energy to raise to any over level, such as n = 2, 3, or 4. The electrons at such an excited state is unstable and drops back to the n = 1 level in one or more steps. 5
Bohr s Model (cont d) Quantum jump or leap The process by which an electron moves between orbits. Ground state Electron in lowest-energy orbit available (H atom); all electrons in the lowest possible energy levels (atoms with multiple electrons). Excited state Electron in orbit with energy higher than ground state (H atom); one or more electrons has an energy level above ground state (atoms with multiple electrons). Bohr s Model (cont d) When an electron falls from an excited state to a lower energy state, it releases the energy by emitting a photon of electromagnetic radiation. The energy of this photon is quantized, and thus corresponds to a line in the spectrum of the element. Orbits in the Bohr model of the hydrogen atom are called Principal Energy Levels or Principal Quantum Numbers (n). Figure 11.8-Ground and excited state. When an electron in the ground state absorbs energy, it is promoted to an excited state. When the electron returns to the ground state, energy can be released in the form of a photon of electromagnetic radiation that corresponds to one of the lines in that element s like spectrum. E is the amount of energy absorbed or emitted. 6
Quantum-Mechanical Model Louie de Broglie stated that matter in motion in association to waves, specifically for electrons. Additional studies done by Erwin Schrödinger used wave mechanics to describe the behavior of electrons in an atom. Quantum-Mechanical Model In this model, there are 4 specific quantum numbers which describe the electron energy. General Specific Principal Energy Levels, n Sublevels Electron Orbitals Orbital Occupancy Principle Energy Levels This term is described by (n). It is usually an integer (1, 2, 3, 4, 5, 6, 7, etc.). Generally, energy increases with increasing nand distance of the electron from the nucleus increases with increasing n. n= 1 < n= 2 < n= 3. < n= 7 7
Sublevels At each energy level, there are various sublevels associated in to each level. There are 4 sublevels (s, p, d, and f). Each sublevel has a different amount of energy depending on how close it is to the nucleus. The number of sublevels present equals the principle quantum number (n). Sublevels At n = 1, there is only one sublevel (s) At n = 2, there are two sublevels (s and p) and in order of increasing energy: 2s < 2p At n = 3, there are three sublevels (s, p and d) and in order of increasing energy: 3s < 3p < 3d At n = 4, there are four sublevels (s, p, d and f) and in order of increasing energy: 4s < 4p < 4d < 4f Orbitals Recall, Bohr s model viewed the behavior of electrons through regions called orbits. By the studies, primarily that Schrödinger did, one can predict the best area an electron can be found is through regions called orbitals. Locating the electron is based on probability where depending on the shape of the orbital, one can predict the likelihood of finding the electron. ORBITALS are a region in space around a nucleus in which there is a high probability of finding the electron. 8
9/30/2011 Figure 11.10-The Bohr model of the atom compared with the quantum model. Sublevels Sublevels Each sublevel has a number of orbitals that depends on the sublevel quantum number: s sublevels have 1 orbital p sublevels have 3 orbitals d sublevels have 5 orbitals f sublevels have 7 orbitals Figure 11.11-Shapes of electron orbitals accounting to the quantum mechanical model of the atom 9
Pauli Exclusion Principle This statement states that no orbital can only have two total electrons of opposite spin. At any time, the orbital can be (1) unoccupied, (2) occupied by one electron, or (3) occupied by two electrons. Summary The Quantum Mechanical Model of the Atom Principal Energy Levels n= 1, 2, 3, 4, 5, 6, 7 Sublevels s, p, d, f Orbitals and Orbital Occupancy Sublevel Orbitals Maximum Electrons per Sublevel s 1 1 2 = 2 p 3 3 2 = 6 d 5 5 2 = 10 f 7 7 2 = 14 Electronic Configuration An electronic configurationdisplays all on the electrons in their respective orbitalsin its ground state configuration. Each orbital is identified by its principle quantum number (n) and its specific orbital. Key rules in the assignments of electrons to orbitals: 1. At ground state, the electrons fill the lowest-energy orbitals available. 2. No orbital can have more than 2 electrons. 10
Diagonal Rule When writing the electronic configuration of an atom or ion it is written as followed: 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f By its energy (lowest to highest): 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d < 7p Figure 11.12-Arrangement of periodic table according to atomic sublevels. Figure 11.13-Sublevel energy diagram. All sublevels are positioned vertically according to a general energy-level scale shown at the left. Each box represents an orbital. Orbitals are filled from lowest energy level which generally can hold two electrons. 11
Ground-State Configurations EXAMPLES: Hydrogen: 1s 1 Nitrogen: 1s 2 2s 2 2p 3 [He] 2s 2 2p 3 Zinc 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 [Ar] 4s 2 3d 10 Figure 11.15-Ground-state electron configurations of neutral atoms. Procedure Writing Electronic Configurations 1. Locate the element in the periodic table. From its position in the table, identify and write the electron configuration of its highest occupied energy sublevel. (Leave room for writing lower-energy sublevels to the left.) 2. To the left of what has already been written, list all lower-energy sublevels in order of increasing energy. 3. For each filled lower-energy sublevel, write as a superscript the number of electrons that fill that sublevel. (There are two selectrons, ns 2, six pelectrons, np 6, and ten delectrons, nd 10.) Exceptions: for chromium and copper, the 4ssublevel has only one electron, 4s 1.) 4. Confirm that the total number of electrons is the same as the atomic number. Electronic Configuration Notation Principle Quantum Number ENERGY Level Orbital Type (Sublevel type) 4s 1 Number of electrons in the orbital 12
Writing Electronic Configuration There are two expressions for writing the electron configuration for an atom or ion. EXAMPLE: Li atom 1. Complete/Ground-State Write out all orbitalsand its corresponding electrons per subshells. NOTATION: 1s 2 2s 1 2. Condensed/Noble Gas Write out the noble gas configuration for the most complete filled period and write the uncompleted period by writing it out by their appropriate subshells. NOTATION: [He] 2s 1 PRACTICE Problem Aluminum Write the complete ground-state configuration with an aluminum atom. Z (atomic number) = 13 For an atom, it has 13 electrons. Therefore, the electronic configuration of this atom is as follows: 1s 2 2s 2 2p 6 3s 2 3p 1 CHECK: 2 + 2 + 6 + 2 + 1 = 13 (Yes, it s equal to Z) PRACTICE Problem Aluminum w/neon Write the noble gas configuration for an aluminum atom. Al: 1s 2 2s 2 2p 6 3s 2 3p 1 Ne: 1s 2 2s 2 2p 6 Since both have the same ten electrons at the beginning, one can replace these coreelectrons with its element symbol in brackets. [Ne] 3s 2 3p 1 13
PRACTICE Problem Copper and Chromium Write the complete ground-state configuration with an chromium. Z (atomic number) = 24 For an atom, it has 24 electrons. Therefore, the electronic configuration of this atom is as follows: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 CHECK: 2 + 2 + 6 + 2 + 6 + 2 + 4 = 24 (Yes, it s equal to Z) However, since the an empty box for a d orbital, it would like to obtain an electron so each orbital has 1 electron each so by acquiring it for a higher energy level orbital (4s), the final, complete electronic configuration is written as 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 PRACTICE Problem Copper and Chromium Write the complete ground-state configuration with an copper. Z (atomic number) = 29 For an atom, it has 29 electrons. Therefore, the electronic configuration of this atom is as follows: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 9 CHECK: 2 + 2 + 6 + 2 + 6 + 2 + 9 = 29 (Yes, it s equal to Z) However, since the an half-filled box for a d orbital, it would like to obtain an electron so each orbital has 2 electrons each so by acquiringit for a higher energy level orbital (4s), the final, complete electronicconfiguration is written as 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 Valence Electrons From the elements electronic configuration, one can also determine the valence electrons of the atom. When one determines the elements valence electrons, we determine it by the electrons located on the outermost shell (by n). 14
Calcium Valence Electrons EXAMPLE Noble Gas Configuration: [Ar] 4s 2 Electronic Configuration: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 Ca Electronic Dot Diagram for Ca Valence Electron, n is the highest. For this atom, it has 2 valence electrons. Electron Dot Symbols When drawing the electron dot diagram for these elements, we are only interested in the atom s valence electrons, not the total number of electrons. In a calcium atom, there are two electrons in n=4 level, therefore, it has 2 valence electrons in its dot diagram. Valence Electrons Main Group Families Group 1A 1 valence electron ns 1 Group 2A 2 valence electrons ns 2 Group 3A 3 valence electrons ns 2 np 1 Group 4A 4 valence electrons ns 2 np 2 Group 5A 5 valence electrons ns 2 np 3 Group 6A 6 valence electrons ns 2 np 4 Group 7A 7 valence electrons ns 2 np 5 Group 8A 8 valence electrons ns 2 np 6 15
Periodic TABLE Trends Mendeleev and Meyer when they made the final periodic table, the elements are organized them based on reoccurring physical and chemical properties observed. 1. Atomic Radius 2. Ionization Energy 3. Chemical Families 4. Metallic Character Atomic Radius By definition, the atomic size deals with the diameter of an ion or atom. 16
Figure 11.16-Sizes of atoms of main group elements, expressed in picometers (1 pm = 10-12 m). Atomic Radius Trends The two key trends are as follows: 1. Highest occupied principal energy level. As the period increases, the atomic size of the atom increases. REASON: As n (principle energy level) increases, the electrons will occupy higher energy levels which results to an increase in the size of the atom. 2. Nuclear Charge. As the atomic number (or protons) increases, the atomic size of the atom decreases. REASON: The positive charge of the atom increases which pulls electrons closer to the nucleus and thus the size of the atom is smaller. Practice Question Rank the atoms from smallest to largest radius and explain why. Pb, Si, C, P 17
Ionization Energy By definition, ionization energy is an amount of energy required to remove one electron from a neutral gaseous atom of an element. Figure 11.17-The formation of a sodium ion from a sodium atom. Figure 11.18-First ionization energy plotted as a function of atomic number, to show periodic properties of elements. Ionization Energy Trends The two key trends are as follows: 1. Highest occupied principal energy level. As the period increases, the ionization energy decreases. REASON: As n (principle energy level) increases, the attractive force between the electrons and the nucleus decreases as the distance increases. 2. Nuclear Charge. As the atomic number (or protons) increases, the increases. REASON: As the number of electrons are removed, the positive charge increases which leads to a stronger attractive force which will require a larger amount of energy. 18
Practice Question Rank the atoms from smallest to largest amount of ionization energy and explain why. B,Tl, In Na, S, Al Chemical Families Trends Chemical families have similar chemical properties exhibited between the elements in the group. REVIEW: Group 1A-Alkali Metals Group 2A-Alkaline Earth Metals Group 7A-Halogens Group 8A-Noble Gases Chemical Families Alkali Metals and Alkaline Earth Metals Alkali Metals Group 1A Has 1 valence electron (ns 1 ) Chemical Reaction: M M + + e - Alkaline Earth Metals Group 2A Has 2 valence electron (ns 2 ) Chemical Reaction: M M 2+ + 2e - 19
9/30/2011 Figure 11.21-The alkali metal and alkaline earth metal families. Alkali metals have 1 valence electron (ns1) and alkaline eath metals have two valence electrons (ns2) Chemical Families Halogens, Noble Gases and Hydrogen Halogens Group 7A Has 7 valence electrons (ns7) Chemical Reaction: M + e- M- Noble Gases Group 8A Has 8 valence electrons (ns8) Chemical Reaction: Unreactive Hydrogen Has 1 valence electron Chemical Reaction: M + e- MM M+ + e - Figure 11.22-The halogen family. 20
Metallic Character Metallic character deals with the likelihood of losing an electron. At a particulate level, the lower the amount of energy needed to remove its valence electron(s), the higher the metallic character will be for the element. Let s look at the properties for metals and nonmetals. Metals Lose electrons easily to form cations Nonmetals Tend to gain electrons to form anions 1, 2, or 3 valence electrons 4 or more valence electrons Low ionization energies Form compounds with nonmetals, but not with other metals High electrical conductivity High thermal conductivity Malleable (can be hammered into sheets) Ductile (can be drawn into wires) High ionization energies Forms compounds with metals and with other nonmetals (molecules) Poor electrical conductivity (Carbon in the form of graphite is an exception) Poor thermal conductivity; good insulator Brittle Nonductile Figure 11.23-Metals and nonmetals. Green identifies elements that are metalloids (semimetals or semiconductors), which have properties that are intermediate between those of metals and nonmetals. 21
Ionization Energy Trends The two key trends are as follows: 1. As one goes from left to right within a period, the metallic character decreases. REASON: As the nuclear charge increases, the electrons are held more tighter and thus its metallic character will decrease within a period of similar energy level. 2. As one goes from top to bottom within a group, the metallic character increases. REASON: As n (principle energy level) increases, the valence electrons in the outermost shell are held less tightly and thus will easily be more removable. 22