Survey
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
Electronic Structure of Atoms Chapter 6 Light • Made up of electromagnetic radiation. • Waves of electric and magnetic fields at right angles to each other. Parts of a wave Wavelength l Frequency (n = number of cycles in 1 second Measured in hertz 1 hertz = 1cycle/second Frequency = n Kinds of EM waves • There are many different EM waves • different l and n • Visible Light is only the part our eyes can detect. (colors of the rainbow) • Greater wavelength means, smaller frequency Gamma Rays X-Rays UV Infrared Microwave Radio Visible Spectrum The speed of light, c • in a vacuum is 2.998 x 108 m/s • c = 3.0 x 108 m/s • c = ln Examples What is the wavelength of light with a frequency 5.89 x 1014 Hz? l = c v = 3.0 x 108 m/s 5.89 x 1014 Hz = 5.09 x 10-7 m = 509 nm (green light) What is the frequency of blue light with a wavelength of 484 nm? v= c l = 3.0 x 108 m/s 484 x 109 m = 6.20 x 1014 Hz Planck and the Quantum Theory • Energy is gained or lost in whole number multiples (n) of the quantity hv. • Similar to energy required to go up stairs (opposed to going up a ramp) • Planck found that Energy is transferred to matter in “energy packets” called a quantum (hv) • Frequency = v • Planck’s constant = h = 6.63 x 10-34 J-s DE = nhn Einstein, the Photoelectric Effect, and Photons • EM radiation is quantized a stream of particles -- “photons” • Ephoton = hn = hc/l • Combine this with E = mc2 • You get the apparent mass of a photon. m = h / (lc) Is light a Wave or does it consist of particles? • Both… • Macroscopically like a wave, • But consists of a collection of photons that we only see at the atomic level. • called The Wave-Particle Duality (Like describing an entire beach and then beginning to examine the grains of sand.) Examples • Calculate the energy of one photon of yellow light whose wavelength is 589nm 1. Find the frequency • 5.09 x 1014 s-1 2. Then use Plank’s equation to find E • 3.37 x 10-19 J Matter as a wave • Using the velocity (v) instead of the frequency (n we get: • De Broglie’s equation l = h/mv • Can calculate the wavelength of an object. Line Spectra • Spectrum = the range of frequencies present in light • Continuous Spectrum = contains all wavelengths of light. (white light… can be broken down into “rainbow”) • Line Spectrum = contains only specific wavelengths of light. Hydrogen spectrum • Emission spectrum because these are the colors it gives off or emits. • Called a bright line emission spectrum. • There are just a few discrete lines showing 656 nm 434 nm 410 nm 486 nm Visible Spectrum Bright Line Spectra • Excited electrons return to lower NRG states • NRG is emitted in the form of a photon of definite wavelength. • Definite change in energy corresponds to: – Definite frequency – Definite wavelength • Use DE = hn = hc / l • Only certain energies are possible within any atom. Niels Bohr • Developed the Quantum Model • Described the atom like a solar system • Electrons attracted to (+) nucleus because of their (-) charge • Electrons didn’t fall into nucleus because they were moving around Bohr’s atom • Found only certain NRGs were allowed; called them NRG levels. • Putting NRG into atom moves electron away from the nucleus (ground state excited state) • When e- returns to ground state, it gives off light of a certain NRG The Bohr Atom n=4 n=3 n=2 n=1 Available NRG levels E = -2.178 x 10-18 J (Z2 / n2 ) • n = quantum number (NRG level) • Z = nuclear charge (+1 for Hydrogen) • J = energy in joules • The more negative the NRG is, the more stable the atom will be. change in Energy • When the electron moves from one energy level to another: • DE = Efinal - Einitial DE = -2.178 x 10-18J [(1/ nf2)–(1/ ni2)] l = hc / DE Shortcomings of Bohr Model • Only works for Hydrogen atoms • Electrons don’t move in circular orbits • The quantization of energy is right, but not because they are circling like planets • Questions Bohr couldn’t answer: Why are e- confined to only certain energy levels? Why don’t e- eventually spiral and crash into the nucleus? The Quantum Mechanical Model • New approach that viewed electron as a standing wave of NRG • Standing waves don’t propagate through space • Standing waves are fixed at both ends (similar to vibrations of a stringed instrument) What’s possible? • You can only have a standing wave if you have complete waves. • There are only certain allowed waves. • In the atom there are certain allowed waves called electrons. • 1925 Erwin Schroedinger described the wave function of the electron. “The Schroedinger Equation” • Much math but what is important are the solutions. Schroedinger’s Equation 2x2 22 • • • • • • 2y2 22 2z2 22 82m h2 (E V) = 0 The wave function, is a F(x, y, z) Solutions to the equation are called orbitals. These are not Bohr orbits. Each solution is tied to a certain energy. These are the energy levels. Many strange and seemingly impossible behaviors occur when the electron is treated as a wave! Orbitals • Orbitals are not circular orbits for electrons • Orbitals are areas of probability for locating electrons There is a limit to what we can know… • about how the electron is moving or how it gets from one energy level to another. • about both the position and the momentum of an object. • The Heisenberg Uncertainty Principle - “we cannot know the exact location and exact momentum of an electron at the same time.” Quantum Mechanical Model and Quantum Numbers • Note: A quantum mechanical orbital is not the same as a Bohr orbit because the motion of the electron in an atom cannot be precisely measured or tracked. (Heisenberg uncertainty Principle) • There are 4 quantum numbers to describe the “location” of an electron. (sort of like how a zip code works) Principal Quantum Number (n) • Indicates probable distance from the nucleus (old Bohr orbitals) • Gives the size and energy of the orbital • Has integer values >0 • According to the periodic table, what would the highest principal quantum number be? Angular Momentum Quantum (l ) • Gives the shape of the orbital (more detail to come) • Integral values from 0 to (n-1) for each principal quantum number (n) Value of l 0 1 2 3 4 Letter used for shape* s p d f g *letters s, p, d, f come from the words sharp, principal, diffuse, and fundamental, which were used to describe certain features of spectra before quantum mechanics was developed. Magnetic Quantum Number (ml ) • Relates to the orientation of the orbital in space relative to the other orbitals. (It tells you if the orbital will be on the x, y or z axis.) • Integral values from l to –l including 0. n l 1 2 3 4 0 0 1 0 1 Orbital designation 1s 2s 2p 3s 3p ml 0 0 -1, 0, 1 0 -1, 0, 1 # of orbitals 1 1 3 1 3 2 0 1 2 3 3d 4s 4p 4d 4f -2, -1, 0, 1, 2 0 -1, 0, 1 -2, -1, 0, 1, 2 -3, -2, -1, 0, 1, 2, 3 5 1 3 5 7 Important Observations 1. The shell w/ quantum #n will have exactly n subshells. 2. Each subshell has a specific number of orbitals. Each orbital corresponds to a different allowed value of ml. For a given value of l, there are 2l + 1 allowed values of ml. 3. The total number of orbitals in a shell is n2. The resulting number of orbitals for the shells – 1, 4, 9, 16 – is related to a pattern seen in the periodic table… We see the number of elements in the table – 2, 8, 18, 32 – equal twice these numbers… S orbitals n=1 n=2 n=3 P orbitals At another energy level the solutions are “dumbell” shaped. There are 3 possible solutions for this energy leve P Orbitals All 3 p orbitals may exist at the same time. d orbitals At another energy we get “flower” shaped orbitals for a solution. All 5 may exist at the same time F orbitals And finally, at another energy, 7 f orbitals are the solution. Orbital Energies • All orbitals with the same value of n have the same energy • The lowest energy state is called the “ground state” • When the atom absorbs energy, electrons may move to higher energy orbitals – “excited state” Electron Spin Quantum Number (ms ) • An individual orbital can hold only 2 electrons • Electrons must have opposite spins (why important?) • Spin can have two values +½ or –½ Pauli Exclusion Principle “in a given atom, no two electrons can have the same set of four quantum numbers” What this means for the atom? • Each atomic sub-orbital may contain a maximum of 2 electrons • Those electrons must have opposite spins Increasing energy 7s 6s 5s 4s 3s 2s 1s 7p 6p 6d 5p 4d 4p 5d 3d 3p 2p Helium with 2 electrons 5f 4f Increasing energy 7s 6s 5s 4s 3s 2s 1s 7p 6p 6d 5p 4d 4p 5d 3d 3p 2p Li with 3 electrons 5f 4f Increasing energy 7s 6s 5s 4s 3s 2s 1s 7p 6p 6d 5p 4d 4p 5d 3d 3p 2p Boron with 5 electrons 5f 4f 2 more important rules: • Aufbau Principle – electrons enter orbitals of lowest energy first. • Hund’s Rule -- When electrons occupy orbitals of equal energy, one electron enters each orbital before they pair. For Example: 2s 2p After the s sublevel gets two electrons, three electrons enter the p orbitals before they pair. Increasing energy 7s 6s 5s 4s 3s 2s 1s 7p 6p 6d 5p 4d 4p 3p 2p 5d 3d 5f 4f Electron Configuratoin p s d f 3 QUESTIONS TO ASK • What Row? –(principle energy level) • What section? –(type of sub-orbital) • What seat? –(how many electrons in that suborbital) Example 1: Write the electron configuration for nitrogen. 7N 2 2 3 1s 2s 2p Example Write the electron 2: configuration for Fe. 26Fe 2 2 6 2 6 2 6 Condensed Electron Configurations • Put the symbol for the Noble gas from the previous principal energy level, then add the electron configuration after that point. • Example 1 for Nitrogen: [He] 2s22p3 • Example 2 for Iron: • [Ar] 4s23d6 The History of the Modern Periodic Table See separate slide show for Periodic Table History Periodic Law • When elements are arranged in order of increasing atomic #, elements with similar properties appear at regular intervals. Atomic Radius (pm) 250 200 150 100 50 0 0 5 10 Atomic Number 15 20 Chemical Reactivity Families Similar valence e- within a group result in similar chemical properties 1 2 3 4 5 6 7 •Alkali Metals •Alkaline Earth Metals •Transition Metals •Halogens •Noble Gases Periodic Table Reveals Periodic Trends • Effective Nuclear charge • Reactivity • atomic size or radius • bonding characteristics • ionization energy • crystal configurations • electron affinity • acidic properties • electronegativity • densities • metallic character • Melting/Boiling points Electron screening or shielding • Electrons are attracted to the nucleus • Electrons are repulsed by other electrons • Electrons would be bound more tightly if other electrons weren’t present. • The net nuclear charge felt by an electron is called the effective nuclear charge ( Zeff ). Quantum Mechanical Model Zeff is lower than actual nuclear charge. Zeff increases toward nucleus ns > np > nd > nf This explains certain periodic changes observed. Effective Nuclear Charge ( Zeff) • The effective nuclear charge acting on an electron equals the number of protons in the nucleus, Z, minus the average number of electrons, S that are between the nucleus and the electron in question. Zeff = Z S Zeff = attractive forces repulsive forces Zeff = # protons # shielding electrons For Example, Lithium vs. Carbon Li Zeff = 3 2 = 1 C Zeff = 6 2 = 4 When moving across a row: The greater the Zeff value, the smaller the atom’s radius. So, carbon has a much smaller atomic radius compared to lithium: Rcarbon =77 pm Rlithium = 152 pm Trend #1 Atomic Radii Increases to Left and Down 1 2 3 4 5 6 7 •Why larger going down? •Higher energy levels have larger orbitals •Shielding - core e- block the attraction between the nucleus and the valence e•Why smaller to the right? • Increased nuclear charge without additional shielding pulls e- in tighter Practice… • Referring to a periodic table, arrange the following atoms in order of increasing size: – Phosphorus – Sulfur – Arsenic – Selenium • S < P < Se < As Atomic radii The Periodic Table & Radii Periodic Trend is Due to Effective Nuclear Charge Atomic Radii vs. Zeff: Trends in Ionic Radii • Using your knowledge of Zeff, how would the size of a cation compare to neutral atom? Anion? Trends in Ionic Radii • The cation of an atom decreases in size. • The more positive an ion is, the smaller it is because Zeff increases • The anion of an atom increases in size. • The more negative an ion, the larger it is because Zeff decreases. Cations lose electrons, become smaller Anions gain electrons, become bigger Ion Radii Increases down 1 2 3 4 5 6 7 Increases moving across, but depends if cation OR anion +3 +4 -3 -2 -1 Ions and Ionic Radii Practice… • Arrange the following atoms and ions in order of decreasing size: – Mg2+ – Ca2+ – Ca • Which of the following ions is the largest: – S2–S – O2- Practice… • Arrange the following ions in order of decreasing size: – S2– Cl– K+ – Ca2+ • Which of the following ions is the largest? – Rb+ – Sr2+ – Y3+ Trend in Ionization Energy • Ionization NRG is the NRG required to remove an electron from an atom Successive Ionization NRG • Ionization energy increases for successive electrons from the same atom. Why do you think there is such a big jump for Mg3+? *Notice the large jump in ionization energy when a core e is removed. • The smaller the atom, the higher the ionization energy due to Zeff • Bigger atoms have lower ionization NRG due to the fact that the electrons are further away from the nucleus and therefore easier to remove. Decreases Increases Practice… • Which of the following elements would have the highest second ionization energy? Justify your answer. –Sodium, Sulfur, or Calcium • Which will have the greater third ionization energy, Ca or S? Justify your answer. Practice… • Referring to a periodic table, arrange the following atoms in order of increasing first ionization energy (Ne, Na, P, Ar, K) Justify your answer. • Based on the trends discussed in this section, predict which of the following atoms (B, Al, C or Si) has the lowest first ionization energy and which has the highest first ionization energy. Electron Affinity • The energy change associated with the addition of an electron • Tends to increase across a period • Tends to decrease as you go down a group • Abbreviation is Eea, it has units of kJ/mol. Values are generally negative because energy is released. • Value of Eea results from interplay of nucleus electron attraction, and electron–electron repulsion. Ionization NRG vs. Electron Affinity • Ionization energy measures the ease with which an atom loses an electron • Electron affinity measures the ease with which an atom gains an electron Electron Affinity Trends in Electronegativity • tendency for an atom to attract electrons when it is chemically combined with another atom. • decreases as you move down a group • increases as you go across a period from left to right. Trend #5 Metallic Character • The metallic character of atoms can be related to the desire to lose electrons. • The lower an atom’s ionizatoin energy, the greater its metallic character will be. • On the periodic table, the metallic character of the atoms increase down a family and decreases from left to right across a period. Metals Nonmetals • Shiny Luster • Various colors (most silvery) • Solids are malleable and ductile • Good conductors of heat and electricity • Most metal oxides are ionic solids that are basic • Tend to form cations in aqueous solution • • • • No luster Various colors Brittle solids Poor conductors of heat and electricity • Most nonmetal oxides are molecular substances that form acidic solutions • Tend to form anions or oxyanions in aqueous solution Metallic Character Increases moving down and across to the left 1 2 3 4 5 6 7 Rb Cs Ba Fr Ra Lower left corner -- elements most likely to lose their valence electrons Metals and Nonmetals • Low ionization energies of metals means they tend to form cations (positive ions) relatively easily • Due to their electron affinities, nonmetals tend to gain electrons when they react with metals. # 6 Melting/Boiling Points • Highest in the middle of a period (generally). 1 2 3 4 5 6 7 Some Important Properties of Alkali Metals • Soft metallic solids • Easily lose valence electrons (Reducing Agents) – React with halogens to form salts – React violently with water • Large Hydration NRG – Positive ionic charge makes ions attractive to polar water molecules Alkaline Earth Metals… • Harder and more dense than Alkali Metals • Less reactive than alkali metals (lower first ionization energies) • Reactivity increases as you move down the periodic table. The Halogens… • “Salt Formers” • Melting and Boiling Points increase with atomic number. • Highly negative electron affinities • Tendency to gain electrons and form halide ions Noble Gases … • • • • Monoatomic ions Gases at room temperature Large 1st ionization energies “Exceptionally” unreactive Practice… • Look at Sample Integrative Exercise 7 on page 264