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Transcript
Chapter 9 Electronic Structure and Periodic Trends Homework Assigned Problems (odd numbers only) “Questions and Problems” 9.1 to 9.71 (begins on page 258) “Additional Questions and Problems” 9.81 to 9.115 (page 284-286) “Challenge Questions” 9.119, 9.121, 9.123 (page 286) Electromagnetic Radiation Matter is anything that has mass and occupies space. Nearly all changes that matter undergoes requires the absorption or release of energy Energy is the capacity to do work The process of moving matter against an opposing force. Forms of energy include heat, electrical, and light One way energy is transmitted through space is by Electromagnetic Radiation Transmits from one place to another in the form of a wave Given off by atoms when they have been excited by any form of energy Electromagnetic Radiation Light (radiant) energy, which is visible and invisible Classified into types according to the frequency of the wave Sunlight, visible light, radio waves, microwaves (ovens), X-rays, and heat from a fire (infrared), are all forms of this radiant energy These forms of radiant energy exhibit the same wavelike characteristics Wavelength and Frequency Electromagnetic radiation is radiant (light) energy that travels in waves at the speed of light The waves have three basic characteristics: wavelength, frequency, and speed The highest point on the wave is a peak Wavelength (l = distance between neighboring peaks) generally measured in nanometers (1 nm = 10-9 m) Velocity (v = how fast the wave is moving) c = speed of light 3.00 x 108 m/s Amplitude (how tall the waves are) Frequency (u = the number of waves that pass a point in a given time) generally measured in Hertz (Hz), 1 Hz = 1 wave/sec = 1 sec-1 c = u x l Waves frequency wavelength frequency wavelength C = speed of light Electromagnetic Spectrum Classified by wavelength: Lower energy (longer wavelength, lower frequency) Higher energy (shorter wavelength, higher frequency) Radiowaves: AM/FM/TV signals, cell phones, low frequency and energy Microwaves: Microwave ovens and radar Infrared (IR): Heat from sunlight, infrared lamps for heating Visible: The only EM radiation detected by the human eye ROYGBIV Ultraviolet: Shorter in wavelength than visible violet light, sunlight X-rays: Higher in energy than UV Gamma rays: Highest in energy, harmful to cells Wavelengths of EM Radiation Atomic Spectra and Energy Levels When white light passes through a prism it produces a continuous rainbow of colors from (red to violet) From red to violet the wavelength becomes shorter Atomic Spectra and Energy Levels When an element is heated (strontium and barium) light is produced If this light is passed through a prism, it does not produce a continuous rainbow, only certain colors Atomic Spectra and Energy Levels Only specific colors are produced in the visible region. This is called a “bright-line spectrum” Each line produced is a specific color, and thus has a specific energy Each element produces a unique set of lines (colors) which represents energy associated with a specific process in the atom Light Energy and Photons Scientists associated the lines of an atomic spectrum with changes in an electrons energy (“Bohr Model”) An electron in a higher energy state will return to a lower energy state The energy that is given off (emitted) corresponds to the energy difference between the higher and lower energy states The light emitted behaves like a stream of small particles called “photons” Electron Energy Levels Electrons possess energy; they are in constant motion in the large empty space of the atom The arrangement of electrons in an atom corresponds to an electron’s energy The electron resides outside the nucleus in one of seven fixed energy levels Energy levels are quantized: Only certain energy values are allowed Light Energy and Photons The energy of a photon is related by the equation E = hν “The energy of a photon is directly proportional to its frequency” “The energy of a photon is inversely proportional to its wavelength” Electron Energy Levels The different lines in an atomic spectrum are associated with changes in an electrons energy Each electron resides in a specific E level called it’s principal quantum number (n, where n=1, n=2…) Electrons closer to nucleus have lower energy (lower n values) Electrons farther from the nucleus have higher energy (higher n values) Electron Energy Levels Electrons can be “excited” to a higher E level with the absorption of E The energy absorbed is equal to the difference between the two E states When an electron loses E and falls to a lower E level, it emits EM radiation (photon) Electron Energy Levels If the EM radiation wavelength is in the visible spectrum a color is seen Energy Levels of Hydrogen: The Bohr Model In 1913 Bohr developed a quantum model based on the emission spectrum for hydrogen The proposal was based on the electron in hydrogen moving around the nucleus in a circular orbit Energy Levels of Hydrogen/ The Bohr Model The Bohr atom nucleus Energy Levels of Hydrogen/ The Bohr Model The Bohr atom has several orbits with a specific radius and specific energy Each orbit or energy level is identified by “n” the principal quantum number Electrons can be “excited” to a higher energy level with nucleus absorption of energy The energy absorbed and released is equal to the energy difference between the two states Energy Levels of Hydrogen/ The Bohr Model The energy levels calculated by the Bohr model closely agreed with the values obtained from the hydrogen emission spectrum The Bohr model did not work for other atoms Energy levels were OK but another model was needed to describe the location of the electron about the nucleus Shrodinger in 1926 (DeBroglie, Heisenberg) developed the more precise quantum mechanical model The quantum (wave) mechanical model is the current theory of atomic structure Quantum Mechanical Model The electron is treated not as a particle but as a wave bound to the nucleus The electron does not move around the nucleus in a circular path (orbit) Instead, the electron is found in orbitals. It is not an circular path for the electron An orbital indicates the probability of finding an electron near a particular point in space An orbital is a map of electron density in 3-D space Each orbital is characterized by a series of numbers called quantum numbers Electron Energy Levels The energy of an electron and its distances from the nucleus can be grouped into levels Principal quantum number “n” is the major energy level in the atom It has values of n =1, 2, 3, etc. As “n” increases the size of the principal energy level (shell) increases Principal E level electron capacity = 2n2 Electron Sublevels All electrons in a principal E level (shell) do not have the same energy Each principal level is divided into 1, 2, 3, or 4 sublevels (subshells) An E level contains the same number of sublevels (s, p, d, and f) as its own pr. energy level number # of sublevels in a principal E level = n Electron Sublevels The order of the increasing energy for sublevels (within an E level) The sublevels with the lowest to highest energy: s sublevel (holds up to 2 electrons) p sublevel (holds up to 6 electrons) d sublevel (holds up to 10 electrons) f sublevel (holds up to 14 electrons) Lowest energy s<p<d<f Highest energy Orbitals The third term used to describe electron arrangement about the atomic nucleus (shells, subshells) is the orbital Since the electron location cannot be known exactly, the location of the electron is described in term of probability, not exact paths Region in space around the nucleus where there is a high (90%) probability of finding an electron of a specific energy Orbitals Orbital shapes are 3-D regions where the highest probability exists Each orbital is represented by four quantum numbers Orbitals within the same sublevel differ mainly in orientation Orbitals of the same type, but in different E levels (i.e. 1s, 2s, 3s) have the same general shape, but differ in size s-Orbitals Only one type of orbital Spherical in shape The larger the energy level, the larger the sphere Holds two electrons s-Orbitals 1s Fig10_23 2s 3s p-Orbitals Can only occur in n=2 or higher levels Are dumb-bell in shape Three sub-orbitals (px, py and pz) each holding 2 electrons for a total of 6 electrons in a porbital pOrbitals z y x (a) z z y y x x (b) Fig10_21 (c) d-Orbitals Five possible d-orbitals Odd shapes Only possible in n=3 and larger energy levels Holds a total of 10 electrons z z y z y y dOrbitals x x x dyz dxz dxy z z y y x x dx2 - y2 Fig10_24 dz2 f-Orbitals Seven possible types of forbital Shapes very difficult, so don’t have to know Can hold a total of 14 electrons Only possible for energy levels n=4 and higher Writing Orbital Diagrams and Electron Configurations To show how the electrons are distributed in the E levels within an atom Orbital diagrams Electron configurations The most stable arrangement of electrons is one where the electrons are in the lowest energy sublevels possible Writing Orbital Diagrams and Electron Configurations The most stable arrangement of electrons is called “ground-state electronic configuration” The most stable, lowest E arrangement of the electrons The GS configuration for an element with many electrons is determined by a building-up process Writing Orbital Diagrams and Electron Configurations For the building-up process, begin by adding electrons to specific E levels beginning with the 1s sublevel Continue in the order of increasing sublevel energies: 1s→2s →2p →3s →3p →4s →3d →4p →5s →4d →etc. Orbital Diagram The notation illustrating the electron arrangement in terms of which energy levels and sublevels are occupied Uses the building-up principal Hund’s Rule: When electrons are placed in a set of orbitals of equal energy, the orbitals will be occupied by one electron each before pairing together Notation Draw a box for each orbital Use an arrow up or down to represent an electron Only one up and one down arrow is allowed in a box 1s 2s 2p Filling of Orbitals In General: Begin filling from the lowest to the highest energy level If there are more than one sub-orbital possible, electrons will spread out first instead of doubling up Once each sub-orbital is filled with one electron, they will double up, but MUST have opposite spins (Hund’s Rule) Orbitals Review s-orbitals Only one per n Can hold two electrons for a total of 2 electrons in an s-orbital p-orbitals Three per n Can each hold two electrons for a total of 6 electrons in a p-orbital Orbitals Review d-orbitals Five per n Can each hold two electrons for a total of 10 electrons in a d-orbital f-orbitals Seven per n Can each hold two electrons for a total of 14 electrons in an f-orbital Orbital Diagram hydrogen Only one electron Occupies the 1s orbital helium Two electrons Both occupy the 1s orbital lithium Three electrons 1s 1s 1s 2s Two occupy the 1s orbital, one occupies the 2s orbital Electron Configurations and the Periodic Table No need to memorize the filling order of the electron, use the periodic table The atomic numbers are in order of increasing sublevel Can “build-up” atoms by reading across the periods from left to right By following a path of increasing atomic number and note the various subshells as they are encountered Each box in the table (across a period) is an increase in one electron Electron Configurations and the Periodic Table The elements are arranged by increasing atomic number The periodic table is divided into sections based on the type of subshell (s, p, d, or f) which receives the last electron in the build up process Different blocks on the periodic table correspond to the s, p, d, or f sublevels Electron Configurations and the Periodic Table s-block elements (Groups 1A and 2A) gain their last electron in an s-sublevel p-block elements (Groups 3A to 8A) gain their last electron in a p-sublevel d-block elements (transition metals) gain their last electron in a d-sublevel. First appear after calcium (element 20) d-sublevel is (n-1) less than the period number f-block elements are in the two bottom rows of the periodic table f-sublevel is (n-2) less than the period number Subshell Filling Order 1 2 3 4 5 6 7 (n-1)d np ns (n-2) f Writing Electronic Configurations Using Sublevel Blocks Locate the element, the number of electrons is equal to the atomic number Lowest energy sublevel fills first, then the next lowest following a path across each period The configuration of each element builds on the previous element The p, d, or f sublevels must completely fill with electrons before moving to the next higher sublevel Electron Configuration Example #1 Write the complete electron configuration for chlorine Chlorine is atomic number 17 (on the periodic table) so the neutral atom has 17 electrons Writing sublevel blocks in order up to chlorine gives: 1s22s22p63s23px Electron Configuration Example #1 1 2 3 4 5 6 7 (n-1) d np ns (n-2) f Electron Configuration Example #1 2 2 6 2 5 Cl : 1s 2s 2p 3s 3p 2 5 or [Ne] 3s 3p 1s 2s 2p 3s 3p Electron Configuration Example #2 Write the complete electron configuration for calcium Calcium is atomic number 20 (on the periodic table) so the neutral atom has 20 electrons Writing sublevel blocks in order up to calcium gives: 1s22s22p63s23p64sx Electron Configuration Example #2 1 2 3 4 5 6 7 (n-1) d np ns (n-2) f Electron Configuration Example #2 2 2 6 2 6 Ca : 1s 2s 2p 3s 3p 4s or [Ar] 4s 1s 2s 2p 2 2 3s 3p 4s Electron Configs Examples 2 2 6 2 6 2 Ca : 1s 2s 2p 3s 3p 4s or [Ar] 4s 2 Periodic Trends of the Elements Per. Table: Graphically represents the behavior of the elements Elements are arranged by increasing atomic number In the periodic table, elements with similar properties occur at regular intervals The arrangement of electrons and not the mass that determines chemical properties of the elements Periodic Trends of the Elements/Valence Electrons Valence electrons: The electrons in the outermost energy level “n” (where n = 1, 2, 3 …) The most important (chemically) Always found in the outermost s or p sublevels Group number equals the valence electrons for each element in that group Applies only to the groups 1A-8A Periodic Trends of the Elements/Valence Electrons Group IA elements have one valence electron: ns1 Group IIA elements have two valence electron: ns2 Group VIIA elements have seven valence electron: ns2np6 Periodic Trends of the Elements/Valence Electrons Write the electron configuration for lithium Li: 1s22s1 Write the electron configuration for sodium Na: 1s22s22p63s1 Each group 1A element has a single electron in an s-sublevel. This is the (one) valence electron Atomic Size For representative (main group) elements only Describes the volume of the electron cloud in the atoms Dependent upon the electron configuration of the atoms Atomic Size Within groups: The atomic radius increases from top to bottom Increase in the period number Principal E level (n) increases Valence electron is further from the nucleus Atomic Size Across periods: The atomic radius decreases from L to R with increasing atomic number Each element increases in proton and electron number Increase in + nuclear charge Valence electrons pulled closer to the nucleus Size of Atoms and Their Ions The formation of a positive ion requires the loss of one or more valence electrons Loss of the outermost (valence) causes a reduction in atomic size Positive ions are always smaller than their parent ions Size of Atoms and Their Ions The formation of a - ion requires the addition of one or more electrons to the valence shell of an atom There is no increase in + nuclear charge to offset the added electron’s - charge Increase in size due to repulsion between electrons Ionization Energy The minimum energy required to remove one electron from an atom of an element The more tightly an electron is held, the higher the ionization energy Ionization Energy In the same group (top to bottom) Ionization Energy decreases Energy required to remove an electron decreases Due to larger principal energy level (larger n value) This puts outer electron farther from nucleus As n increases, ionization energy decreases Across same period (left to right) Ionization Energy increases Metals (left end) have lower ionization E Tend to lose electrons to form + ions Nonmetals (right end) have higher ionization E Tend to gain electrons in chemical reactions End