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1 Electrons and Periodicity • Adapted from www.chemistrygeek.com Rutherford Model of the Atom (1911) What is wrong with this model of the atom? 2 Bohr Model of the Atom (1913) 3 •Electrons are found in circular paths around the nucleus •The paths are called Energy Levels (n) •Each level has a specific energy •The lowest level is called Ground State •Electrons can jump to higher levels by absorbing energy (Excited State) •Electrons can return to a lower level by releasing energy (light) •Bohr’s model only works for atoms of Hydrogen Wave-Particle Duality • Photons of light and most subatomic particles (i.e. electrons) behave as both PARTICLES and WAVES. https://www.youtube.com/watch?v=MB03zWIQO1Q 4 Heisenberg Uncertainty Principle (1927) W. Heisenberg 1901-1976 Problem of defining nature of electrons in atoms solved by W. Heisenberg. Cannot simultaneously define the position and momentum (= m•v) of an electron. We can define electron energy exactly but accept limitation that we do not know exact position. 5 Heisenberg Uncertainty Principle (2015) 6 Electron Cloud Model (1926) • • • • 7 Proposed by Erwin Schrödinger Applied idea of electrons behaving as a wave Electrons are no longer in defined paths/orbits Uses Electron Clouds (Orbitals) – defined areas where an electron is likely to be found • Based on probability 8 Electromagnetic Spectrum In increasing energy, ROY G BIV 9 Wave Behavior wavelength Visible light Amplitude wavelength Node Ultraviolet radiation Visible Spectrum Long wavelength --> small frequency Short wavelength --> high frequency What colors have --long wavelengths? --high frequencies? --short wavelengths? --small frequencies? increasing frequency increasing wavelength 10 Wave Behavior • Wavelength – distance between 2 crests of a wave – Abbreviated with Greek letter lambda ( – Units are meters (m) λ) • Frequency – speed at which crests pass a given point – Abbreviated with Greek letter nu ( – Units are 1/s, s-1 or Hertz ν) • Wavelength and Frequency are related by the equation: • = c – c =speed of light = 3.00 x 108 m/sec – Inverse relationship (one goes up, the other goes down) 11 Wave Behavior • Energy and frequency are directly related • Amount of energy released by an excited electron corresponds to a specific frequency of light •E=hxν –(E) Energy measured in Joules (J) –(h) Planck’s constant = 6.626 x 10-34 J s –(ν) Frequency measured in 1/s or hertz 12 13 Wave Behavior • What is the frequency of a wave that has a wavelength of 2.56 x 10-7 m? c = λ x ν 2.8 x 108 m/s = (2.56 x 10-7 m) 2.8 x 108 m/s 2.56 x 10-7 m = ν 1.1 x 1015 1/s = ν • What is the Energy of this wave? E = h x ν x ν E = (6.626 x 10-34 Js) x (1.1 x 1015 1/s) E = 7.2 x 10-19 J Learning Check • What is the frequency of EMR having a wavelength of 5.55 x 10-7 m? • What is the Energy of this same wave? 14 Continuous Spectrum of White Light •If you separate white light with a prism you will see ALL colors of the rainbow •Continuous because there are no gaps between colors 15 Creating an Atomic Spectrum 16 • Discharge tube – tube containing atoms of a gas connected to an electricity source • When electricity is turned in the gas will emit light Atomic Spectrum of Excited Hydrogen Gas 17 •If the light from a discharge tube is separated by a prism you will only see very specific lines of color separated by black regions •Called atomic emission spectra or line spectra Atomic Emission Spectra of Excited Atoms 18 • When electricity is supplied to the gas atoms the electrons absorb Energy and become EXCITED • Electrons eventually have to release this energy in the form of light • Excited atoms emit light of only certain wavelengths (lines of color) • The wavelengths of emitted light depend on the element. 19 Line Spectra of Other Elements Arrangement of Electrons in Atoms Electrons in atoms are arranged as LEVELS (n) SUBLEVELS (l) ORBITALS (ml) 20 Quantum Numbers •Every electron has a unique set of 4 quantum numbers that defines it’s position in the atom •Think of it as the electron’s address! 21 Quantum Numbers 22 23 Learning Check • Write a set of quantum numbers for the 4f orbital • Which orbital is represented by the following sets of Quantum numbers a. b. c. d. e. n = 3, l = 0, ml = 0 n = 2, l = 1, ml = 1 n = 4, l = 2, ml = -1 n = 3, l = 3, ml = 2 n = 3, l = 1, ml = 2 24 Energy Levels • Each energy level has a number called the PRINCIPAL QUANTUM NUMBER (n) • Currently n can be 1 thru 7, because there are 7 periods on the periodic table 25 Energy Levels n=1 n=2 n=3 n=4 Types of Orbitals • Orbital - most probable area to find an electron • Orbitals can take on different shapes • So far, we have 4 shapes. They are named s, p, d, and f. • No more than 2 e- assigned to an orbital – one spins clockwise, one spins counterclockwise 26 Types of Orbitals 27 S Orbital (sharp) • The s orbital is the simplest orbital • The shape is a sphere • An electron can move anywhere within the sphere • L=0 • Every principal level (n) has an s orbital 28 Relative sizes of the spherical 1s, 2s, and 3s orbitals of hydrogen. 29 p Orbitals (proximal) • A p orbital has a dumb-bell shape • There are three different directions for p-orbitals • These are called x, y, and z • The electron can move anywhere in the nodes • L=1 • Every level above n = 1 has a p orbital 30 There is a PLANAR NODE thru the nucleus, which is an area of zero probability of finding an electron p Orbitals • The three p orbitals lie 90o apart in space 31 d Orbitals (diffuse) • A d orbital has a clover shape • There are 5 different directions for a d orbital • L= 2 • Every level above n=2 has a d orbital 32 33 The shapes and labels of the five 3d orbitals. 34 f Orbitals (fundamental) • f orbitals have a flower shape • There are 7 directions for d orbitals • L= 3 • Every level above n = 3 has an f orbital Electron Configurations • The arrangement of electrons in an atom into atomic orbital • Electron configurations are filled according to three rules (Aufbau, Pauli, Hund) • We need electron configurations so that we can determine the number of electrons in the outermost energy level. These are called valence electrons. • The number of valence electrons determines the number and types of bonds and atom can make 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14… etc. 35 Electron Configurations 4 2p Energy Level Number of electrons in the sublevel Sublevel 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14… etc. 36 Electron Configuration Rules 37 • Aufbau Principle – electrons are filled starting from the lowest energy levels first • Pauli Exclusion Principle – no more than 2 electrons can fill into orbitals at one time and they must have opposite spin (clock-wise/counter-clockwise) • Hund Rule – In p, d and f sublevels, electrons will have to fill one clockwise electron into all available orbitals before filling in counter-clockwise electrons Energy Level (n) Sublevel (l) Orbital (ml) Aufbau Principle • Aufbau Principle states that electrons fill from the lowest possible energy to the highest energy • The Aufbau Diagram is a memory device that helps you remember the order of the filling of the orbitals from lowest energy to highest energy 38 Aufbau Diagram Why are d and f orbitals always in lower energy levels? •d and f orbitals require LARGE amounts of energy to fill •It’s easier to fill an s orbital before a d or f •This is the reason for the diagram •BE SURE TO FOLLOW THE ARROWS IN ORDER! 39 Pauli Exclusion Principle 40 How many total electrons can be in a sublevel? Remember: A maximum of two electrons can be placed in an orbital. s orbitals p orbitals d orbitals f orbitals Number of orbitals Number of electrons Lithium 41 Group 1A Atomic number = 3 1s22s1 ---> 3 total electrons 3p 3s 2p 2s 1s Carbon 3p 3s 2p 2s 1s Group 4A Atomic number = 6 1s2 2s2 2p2 ---> 6 total electrons Here we see HUND’S RULE. 42 Let’s Try It! • Write the electron configuration for the following elements: H He B N F Ne Na 43 Orbitals and the Periodic Table • Orbitals grouped in s, p, d, and f orbitals (sharp, proximal, diffuse, and fundamental) • The d orbitals are n-1 for Rows 4-7 • The f orbitals are n-2 for Rows 6-7 s orbitals f orbitals d orbitals p orbitals 44 Shorthand Notation • A way of abbreviating long electron configurations • Use Noble gases as placeholders for completely filled levels • Only write out the configuration for the Valence (outermost) Level He Ne Ar Kr Xe Rn 45 Shorthand Notation • Step 1: Find the closest noble gas to the atom (or ion) that has a lower atomic number. • Step 2: Write the symbol for the Noble Gas in brackets [ ] • Step 3: Find the energy level below your Noble Gas • Step 4: Write the configuration for that energy level and stop when you reach your target element 46 Shorthand Notation • Chlorine – Longhand is 1s2 2s2 2p6 3s2 3p5 You can abbreviate the first 10 electrons with a noble gas, Neon. [Ne] replaces 1s2 2s2 2p6 The next energy level after Neon is 3 Write out the configuration for level 3 until you reach Cl [Ne] 3s2 3p5 47 48 Practice Shorthand Notation • Write the shorthand notation for each of the following atoms: Br K Sb Zn Sn Np Valence Electrons Electrons are divided between core and valence electrons B 1s2 2s2 2p1 Core = [He] , valence = 2s2 2p1 Br [Ar] 3d10 4s2 4p5 Core = [Ar] 3d10 , valence = 4s2 4p5 49 Electron Configurations for Ions • Remember… – Negative ions have gained electrons – Positive ions have lost electrons • Electrons that are lost or gained should be added/removed from the highest energy level (not the highest orbital in energy!) 50 51 Electron Configurations for Ions Tin (Sn) Atom: [Kr] 5s2 4d10 5p2 Sn+4 ion: [Kr] 4d10 Sn+2 ion: [Kr] 5s2 4d10 Note that the electrons came out of the highest energy level, not the highest energy orbital! 52 Electron Configurations for Ions Bromine Atom: [Ar] 4s2 3d10 4p5 Br- ion: [Ar] 4s2 3d10 4p6 Note that the electrons went into the highest energy level, not the highest energy orbital! Ion Configurations To form anions from elements, add 1 or more e- from the highest sublevel. P [Ne] 3s2 3p3 + 3e- ---> P3- [Ne] 3s2 3p6 or [Ar] 3p 3p 3s 3s 2p 2p 2s 2s 1s 1s 53 Try Some Ions! • Write the longhand notation for these: FLi+ Mg+2 • Write the shorthand notation for these: O-2 Ba+2 Al+3 54 (HONORS only) 55 Exceptions to the Aufbau Principle • Remember d and f orbitals require LARGE amounts of energy • If we can’t fill these sublevels, then the next best thing is to be HALF full (one electron in each orbital in the sublevel) • There are many exceptions, but the most common ones are the elements whose configurations end in d4 – Cr, Mo, W d9 – Cu, Ag, Au (HONORS only) Exceptions to the Aufbau Principle 56 d4 is one electron short of being HALF full In order to become more stable (require less energy), one of the closest s electrons will actually go into the d, making it d5 instead of d4. Example: Cr (theoretical) [Ar] 4s2 3d4 Cr (actual) [Ar] 4s1 3d5. Procedure: Find the closest s orbital. Steal one electron from it, and add it to the d. (HONORS only) 57 Exceptions to the Aufbau Principle OK, so this helps the d, but what about the poor s orbital that loses an electron? Remember, half full is good… and when an s loses 1, it too becomes half full! So… having the s half full and the d half full is usually lower in energy than having the s full and the d to have one empty orbital. (HONORS only) Exceptions to the Aufbau Principle 58 d9 is one electron short of being full Just like d4, one of the closest s electrons will go into the d, this time making it d10 instead of d9. Example: Au (theoretical) [Xe] 6s2 4f14 5d9 Au (actual) [Xe] 6s1 4f14 5d10. Procedure: Same as before! Find the closest s orbital. Steal one electron from it, and add it to the d. (HONORS only) 59 Try These! • Write the shorthand notation for: Cu W Au General Periodic Trends • Atomic and ionic size • Ionization energy • Electronegativity Higher effective nuclear charge Electrons held more tightly Larger orbitals. Electrons held less tightly. 60 61 Atomic Size • Size goes UP on going down a group. • Because electrons are added further from the nucleus, there is less attraction. This is due to additional energy levels and the shielding effect. Each additional energy level “shields” the electrons from being pulled in toward the nucleus. • Size goes DOWN on going across a period. 62 Atomic Size Size decreases across a period owing to increase in the positive charge from the protons. Each added electron feels a greater and greater + charge because the protons are pulling in the same direction, where the electrons are scattered. Large Small 63 Which is Bigger? • Na or K ? • Na or Mg ? • Al or I ? Ion Sizes Li,152 pm 3e and 3p Does+ the size go up+ or down Li , 60 pm when losing an 2e and 3 p electron to form a cation? 64 65 Ion Sizes + Li,152 pm 3e and 3p Li + , 78 pm 2e and 3 p Forming a cation. • CATIONS are SMALLER than the atoms from which they come. • The electron/proton attraction has gone UP and so size DECREASES. Ion Sizes Does the size go up or down when gaining an electron to form an anion? 66 67 Ion Sizes F, 71 pm 9e and 9p F- , 133 pm 10 e and 9 p Forming an anion. • ANIONS are LARGER than the atoms from which they come. • The electron/proton attraction has gone DOWN and so size INCREASES. • Trends in ion sizes are the same as atom sizes. Trends in Ion Sizes Figure 8.13 68 69 Which is Bigger? • Cl or Cl- ? • K+ or K ? • Ca or Ca+2 ? • I- or Br- ? Ionization Energy 70 IE = energy required to remove an electron from an atom (in the gas phase). Mg (g) + 738 kJ ---> Mg+ (g) + e- This is called the FIRST ionization energy because we removed only the OUTERMOST electron Mg+ (g) + 1451 kJ ---> Mg2+ (g) + eThis is the SECOND IE. Trends in Ionization Energy • IE increases across a period because the positive charge increases. • Metals lose electrons more easily than nonmetals. • Nonmetals lose electrons with difficulty (they like to GAIN electrons). 71 72 Trends in Ionization Energy • IE increases UP a group • Because size increases (Shielding Effect) 73 Which has a higher 1st ionization energy? • Mg or Ca ? • Al or S ? • Cs or Ba ? 74 Electronegativity, is a measure of the ability of an atom in a molecule to attract electrons to itself. Concept proposed by Linus Pauling 1901-1994 Periodic Trends: Electronegativity • In a group: Atoms with fewer energy levels can attract electrons better (less shielding). So, electronegativity increases UP a group of elements. • In a period: More protons, while the energy levels are the same, means atoms can better attract electrons. So, electronegativity increases RIGHT in a period of elements. 75 Electronegativity 76 77 Which is more electronegative? • F or Cl ? • Na or K ? • Sn or I ?