Why Patterns for Charges of Common Cations and Anions? Electrons in Atoms

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1 Electrons in Atoms From Light to Energy of Electrons in Atom Quantum mechanical description of Atom 1. Principal quantum number: Shell 2. Orientation (shape) of : Subshell 3. Orbitals hold electrons with opposite spin Electron configuration and Orbital Diagram Periodic patterns relate to Electron configuration 1 Why Patterns for Charges of Common Cations and Anions? Why metals always tend to lose electron(s) whereas nonmetals tend to gain electron(s)? Why Group IA metals (e.g., Na) always form cation with +1 charge, Group IIA cations always with +2? Why Group VIIA (halogens) always form anions with 1 charge, Group VIA monatomic anions always 2 charge? 3 Blimp Chemistry: Hydrogen vs. Helium Blimps float: filled with a gas that is less dense than the air Early blimps used the gas Hydrogen: flammability led to the Hindenburg disaster Blimps now use Helium gas: not flammable nor any chemical reactions Why Hydrogen and Helium are so different in Chemical Reaction? Electromagnetic Radiation Light: one of the forms of energy Electromagnetic radiation electromagnetic radiation travels in Waves Wave properties Wave speed Height (amplitude) Wave length Frequency 2 4 1

2 Electromagnetic Waves How fast Light travels? Velocity c = speed of light = x 10 8 m/s in vacuum all types of light energy travel at the same speed Frequency = #peaks pass a point in a second generally measured in Hertz (Hz), Low frequency High Frequency Visible Light Colors in visible light: Different frequency of electromagnetic radiation: Frequency for visible light: Red < Yellow < Green < Blue < Violet Photons with different frequency have a different amount of energy frequency Energy of the photons 5 Energy: Red < Yellow < Green < Blue < Violet 7 Energy of Electromagnetic Radiation Max Planck: German Physicist Light consists of numerous, individual packets of electromagnetic energy, called photon. The energy of each photon is proportional to its frequency E photon = hv 6 Man-made Rainbow? Prism: An optical element that refract ( bend ) light. Sunlight passed through a prism is separated into all its colors - this is called a continuous spectrum Nowadays, a silver-colored CD disc can generate your homemade spectrum! 8 2

3 Electromagnetic Spectrum Everyday spectra Common street light (containing mainly Na vapor) Indoor fluorescent light (containing Hg vapor) 9 11 Types of Electromagnetic Radiation By the frequency from low to high Radiowaves : low frequency and energy Light s Relationship to Matter Atoms can acquire extra energy, but they must eventually release it Microwaves Infrared (IR) Visible: ROYGBIV Ultraviolet (UV) X-rays Gamma rays high frequency and energy 10 When atoms emit energy, it always is released in the form of light However, atoms don t emit all colors, only very specific wavelengths in fact, the spectrum of wavelengths can be used to identify the element 12 3

4 Emission Spectrum of Hydrogen 13 Emission spectra: Fingerprint of atoms White light Sample Sample Emission light Absorption Spectrum Emission Spectrum Absorption Spectrum of H Emission Spectrum of H

5 Why Line Spectra? Another way for the same question: Why atoms can only emit or absorb certain amount of energies? An simple guess would be that an atom could only have very specific amounts of energy; When they absorb or release energies (photon), the change in the energies they possess would be certain amount. Electron Orbits Electrons travel in orbits around the nucleus more like shells than planet orbits the farther the electron is from the nucleus, the more energy it has Bohr Model of the Atom The energy of the atom was quantized The amount of energy in the atom was related to the electron s position in the atom (Electron Orbit) The atom could only have very specific amounts of energy (n = 1, 2, 3, ) 18 Orbits and Energy each Orbit has a specific amount of Energy Energy of each orbit is characterized by an integer n The larger n, the more energy an electron in that orbit has, the farther it is from the nucleus n: quantum number 20 5

6 Energy Transitions Quantum-Mechanical Orbitals when the atom gains energy, the electron leaps from a lower energy orbit to higher energy orbit: Excitation ( spectra) Quantum Physicists including Schrödinger: Electrons show up with a particular probability at certain location of the atom when the electron leaps from a higher energy orbit to lower energy orbit, energy is emitted as a photon of light: Relaxation ( spectra) 21 Orbital: A region where the electrons show up a very high probability when it has a particular amount of energy generally set at 90 or 95% 23 Bohr Model of the Atom Hydrogen Spectrum Quantum-Mechanical Model: Quantum Numbers Three quantum numbers: quantize the energy Principal quantum number, n, specifies the main energy level for the orbital the higher n value, the higher energy of the electrons, the further away electrons are located from the nucleus

7 Quantum-Mechanical: Quantum Numbers f orbitals Principal energy shell has one or more Subshells the number of subshells = the Principal quantum number n = 1, one subshell; n = 2, two subshells; n = 3, three subshells Subshell Quantum numbers: s, p, d, f each Subshell has orbitals with a particular shape the shape represents the probability map 90% probability of finding electron in that region 25 Tro: Chemistry: A Molecular Approach, 2/e s Orbital Shapes of Subshells p Orbitals: p x, p y, p z How does the 1s Subshell Differ from the 2s Subshell? (colors: signs of wavefunction) d Orbitals

8 Shells & Subshells Energy 7s 6s 5s 4s 6p 5p 4p 3p 6 d 5d 4d 3d 5f 4f 3s 2p 2s 29 1s 31 Subshells and Orbitals Among the subshells of a principal shell, slightly different energies: s < p < d < f each subshell contains one or more Orbitals s : 1 orbital p : 3 orbitals d : 5 orbitals f : 7 orbitals within one subshell, different orbitals have the same energy. Example: 2p x, 2p y and 2p z Order of Subshell Filling in Ground State Electron Configurations 1. Diagram putting each energy shell on a row and listing the subshells, (s, p, d, f), for that shell in order of energy, (left-to-right) 2. draw arrows through the diagonals, looping back to the next diagonal each time 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s

9 Electron Configurations Definition: The distribution of electrons into the various energy shells (n = 1,2,3, ) and subshells (s, p, d, f) in an atom in its ground state Each energy shell and subshell has a maximum number of electrons it can hold Subshell s =, p =, d =, f = Shell n: 1 = 2e, 2 = 8e, 3 = 18e, 4 = 32e Electrons fill in the energy shells and subshells in order of energy, from low energy up Aufbau Principal ( Construction in German) 33 Orbital Diagrams often an orbital as a square the electrons in that orbital as arrows the direction of the arrow represents the spin of the electron unoccupied orbital orbital with 1 electron orbital with 2 electrons 35 Spinning Electron(s) in Orbital Experiments (Stern and Gerlach) showed Electrons spin on an axis generating their own magnetic field Pauli Exclusion Principle each Orbital may have a maximum of 2 electrons, with opposite spin Two electrons sharing the same orbital must have Opposite spins so their magnetic fields will cancel analogous to two bar magnets in parallel: only opposite alignment could stabilize each other. 34 How electrons in an atom are filled into orbitals 1. How Electrons fill subshells with multiple orbitals 2. How Electrons fill subshells with higher n number first 36 9

10 Filling the Orbitals in a Subshell with Electrons Energy shells fill from lowest energy to high Subshells fill from lowest energy to high s p d f Orbitals of the same subshell have the same energy. Three 2p orbitals; Five 3d orbitals Electrons prefer spreading out in orbitals of same subshell before they pair up in orbitals. Hund s Rule Example: 2p 3 _ _ _ instead of 37 Example: Ground State Orbital Diagram and Electron Configuration of Magnesium 1. Determine the number of electrons: Atomic number = #protons = #electrons = 2. Draw boxes to represent the subshells 3. Add one electron to each box in a set, then pair the electrons before filling the next subshell When pair, put in opposite arrows: 4. Use the diagram to write the electron configuration (1s 2 2s 2 ) 39 Electron Configuration of Atoms in their Ground State Electron configuration: a listing of the subshells in order of filling with the number of electrons in that subshell written as a superscript Kr = 36 electrons = 1s 2 2s 2 2p 6 More example: Write Electron Configuration and Orbital Diagram for a chlorine atom chlorine: electrons a shorthand way : use the symbol of the previous noble gas in [] for the inner electrons, then just write the last set Rb = 37 electrons = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 =

11 Valence Electrons Definition: the electrons in all the subshells with the highest principal energy shell Example: electrons in bold Mg = [Ne]3s 2 O = [He]2s 2 2p 4 Br = [Ar]4s 2 3d 10 4p 5 Electrons Configurations and the Periodic Table Core electrons: electrons in lower energy shells Chemists have observed that one of the most important factors in the way an atom behaves, both chemically and physically, is the Number of Valence electrons Valence Electrons Rb = 37 electrons = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 the highest principal energy shell is the 5 th : valence electron + core electrons Kr = 36 electrons = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 the highest principal energy shell is the 4 th : valence electrons + core electrons Electron Configurations from the Periodic Table Example: Be 2s 2 B2s 2 2p 1 C2s 2 2p 2 N2s 2 2p 3 O2s 2 2p 4 Elements in the same period (row) have Valence Electrons in the same principal energy shell. #Valence electrons increases by 1 from to Example: IIA: Be 2s 2 Ca 3s 2 Sr 4s 2 Ba 5s 2 VIIA: F 2s 2 2p 5 Cl 3s 2 3p 5 Br 4s 2 4p 5 I5s 2 5p 5 Elements in the same group have the same and same kind of subshell

12 Electron Configuration & the Periodic Table Elements in the same Group have similar chemical and physical properties their valence shell electron configuration is the same No. Valence electrons for the main group elements is the same as the Group Number Example: Group IA: ns 1 ; Group IIIA: ns 2 np 1 Group VIIA: ns 2 np 5 Electron Configuration from the Periodic Table Inner electron configuration = Noble gas of the preceding period Outer electron configuration: from the preceding Noble gas the next period (Subshells) Element the valence energy shell = the period number the d block is always one energy shell below the period number and the f is two energy shells below s 1 s 2 Electron Configuration & the Periodic Table d 1 d 2 d 3 d 4 d 5 d 6 d 7 d 8 d 9 d 10 p 1 p 2 p 3 p 4 p 5 s 2 p Electron Configuration from the Periodic Table 1A 8A 2A 3A 4A 5A 6A 7A Ne 3s 2 P 3p 3 f 1 f 2 f 3 f 4 f 5 f 6 f 7 f 8 f 9 f 10 f 11 f 12 f 13 f P = [Ne]3s 2 3p 3 P has 5 valence electrons 48 12

13 Electron Configuration from the Periodic Table 1A 2A 4s 2 3d 10 As = [Ar]4s 2 3d 10 4p 3 As has 5 valence electrons 3A 4A 5A 6A 7A As 4p 3 8A Ar 49 Electron Configuration: Noble Gas Noble gases have 8 valence electrons except for He, which has only 2 electrons Noble gases are especially nonreactive He and Ne are practically inert The reason: the electron configuration of the noble gases is especially stable 51 Electron configuration & Chemical Reactivity Chemical properties of the elements are largely determined by No. Valence electrons Why elements in groups? Since elements in the same column have the same #valence electrons, they show similar properties Everyone Wants to Be Like a Noble Gas! Alkali Metals (Group 1A) have one more electron than the previous noble gas, [NG]ns 1 tend to lose their extra ONE electron, resulting in the same electron configuration as a noble gas forming a cation with a 1+ charge Na Na + Li Li

14 Everyone Wants to Be Like a Noble Gas! Halogens (Group 7A) one fewer electron than the next noble gas: [NG]ns 2 np 5 Reactions with Metals: tend to gain an electron and attain the electron configuration of the next noble gas: [NG]ns 2 np 5 + 1e [NG]ns 2 np 6 forming an anion with charge 1-: Cl Cl - Reactions with Nonmetals: tend to share electrons so that each attains the electron configuration of a noble gas 53 Stable Electron Configuration And Ion Charge Metals: Cations by losing enough electrons to get the same electron configuration as the previous noble gas Nonmetals: Anions by gaining enough electrons to get the same electron configuration as the next noble gas 55 Everyone Wants to Be Like a Noble Gas! Summary Alkali Metals as a group are the most reactive metals they react with many things and do so rapidly Halogens are the most reactive group of nonmetals one reason for their high reactivity: they are only ONE electron away from having a very stable electron configuration the same as a noble gas Example: Write Electron Configuration for the following ions Sulfide ion: charge =, #electrons = Aluminum ion: charge =, #electrons =

15 Trends in Atomic Size Metallic Character Metals malleable & ductile shiny, lusterous, reflect light conduct heat and electricity most oxides basic and ionic form cations in solution lose electrons in reactions oxidized Nonmetals brittle in solid state dull electrical and thermal insulators most oxides are acidic and molecular form anions and polyatomic anions gain electrons in reactions - reduced Trends in Atomic Size Down a group: crease valence shell farther from nucleus, weaker attraction Trends in Metallic Character Across a period (left to right): crease More protons to attract valence shell electrons Electrons added to same valence shell valence shell held closer

16 Electron Configuration Affects the Size of Atoms and Metallic Character: Within a Group Within the same Group, from top to bottom: As valence shell number n increases valence electron(s) further away from the nucleus Atomic Radius weaker Coulombic force (electrostatic force) withholding valence electrons electrons easier to be lost metallic character 61 Electron Configuration Affects the Size of Atoms and Metallic Character: Over the Period Within the same Period (row), from left to right: Same valence shell number n As Nucleus has increasing number of protons (p + ) Stronger Coulombic force (electrostatic force) withholding valence electrons Valence Electrons closer the nuclues Atomic Radius Valence electrons harder to be lost metallic character 63 Example: Group IIA Be (4p + & 4e - ) Mg (12p + & 12e - ) 2e - 2e- 4 p + 2e - 8e - 2e - 12 p + Example: Period 2 From Li (3 protons) to Ne (10 protons), attraction increases 1e - 2e - 3+ Li (3p + & 3e - ) 2e - 2e e - 3e - 5+ Be (4p + & 4e - ) B (5p + & 5e - ) Ca (20p + & 20e - ) 2e - 8e - 8e - 2e - 4e e- C (6p + & 6e - ) 6e e- O (8p + & 8e - ) 8e e- Ne (10p + & 10e - ) 20 p

17 Practice Choose the Larger Atom in Each Pair Cor O Li or K C or Al Se or I? 65 Practice Choose the More Metallic Element in Each Pair Snor Te Si or Sn Br or Te Se or I? 66 17