CEM 141 - Chapter 2 Slides PDF

CEM 141 – Chapter 2 Note: This is a condensed set of notes that does not contain everything your instructor will present in class. You should add to these notes during the lecture! 1. Make sure to record answers to questions/polls, explanations, and drawings. 2. Write down any ques...

CEM 141 – Chapter 2 Note: This is a condensed set of notes that does not contain everything your instructor will present in class. You should add to these notes during the lecture! 1. Make sure to record answers to questions/polls, explanations, and drawings. 2. Write down any questions you have so that you can ask at the end of the lecture, during office hours, or in the chemistry help room. 3. These notes are not divided by lecture and do not contain announcements or homework answers. You may wish to add slide numbers as you go through each lecture. 4. Your instructor may make changes to these slides or the order of these slides before presenting them in class.  Chapter 2 Electrons and Orbitals Model of Atom How did we get to this model? In this chapter, we are going to look at where the electrons are and how they affect some properties of elements. – For example, why He atoms interact through LDFs whereas H atoms form covalent bonds.  To understand atomic structure we need to understand electromagnetic radiation. What is electromagnetic radiation? What are some examples that you encounter? Examples of Electromagnetic Radiation How do we characterize electromagnetic radiation? What is different about radio waves and X-rays? Two models are used to describe the behavior of light (electromagnetic radiation). Which is correct? What is light? A) A wave B) A particle C) Both D) Neither  Light is a Wave Watch this video for a primer on “waves” Characterization of Waves - Wavelength Characterization of Waves – Frequency  Electromagnetic Radiation Amplitude is the height of peaks (________) λ = wavelength (m): ν = frequency (Hz = s-1): Wavelength and frequency are related by the speed of light: c = λ × ν – Speed of light = c = 3.00 × 108 m s-1 Energy of light – increases as frequency increases (and wavelength decreases) – Not related to amplitude! A 1.Which has higher intensity? 2.Which has higher frequency? B 3.Which has a longer wavelength? 4.Which one is more likely to be red light (as opposed to blue)?  The Electromagnetic Spectrum Short Long wavelength wavelength High Low frequency frequency Low energy High energy Note: Scale on the Spectrum is a log Scale Let's say we have 2 A. 1m objects 1 m apart (100 m) …. B. 10 m C. 20 m How far apart would 2 objects be that were 101 D. 100 m m apart? E. 1000 m How far apart would 2 objects be that were 102 m apart? How far apart would 2 objects be that were 103 m apart?  Scale Which e/m A. Radiowaves ( ≈ radiation has m) wavelengths on B. Microwaves ( ≈ the order of the cm) sizes of atoms? C. X-rays ( ≈ nm) D. Gamma rays ( ≈ pm) How big are these waveleng ths? Determine the wavelength of an X-ray with a frequency of 3.0 × 1018 Hz in meters. (c = 3.0 × 108 m/s). A. 1.0 × 10–10 m B. 9.0 × 1026 m C. 9.0 × 10–19 m D. 1.0 × 10–1 m We saw that an X-ray with a frequency of 3.0 × 1018 Hz has a wavelength of 1.0 × 10–10 m. What is this in nanometers? Remember: A. 1.0 × 10–10 nm 1 nm = 1 × 10–9 m B. 9.0 × 1026 nm This is the same as C. 9.0 × 10–19 nm saying: D. 1.0 × 10–1 nm 1 m = 1 × 109 nm We’re making a claim that light is a wave. What is the evidence? Properties of Waves: Diffraction Particles don’t diffract  Properties of Waves: Interference If waves are in- phase: If waves are out-of- phase:  Properties of Waves: Interference Patterns double slit experiment (Watch the first 2:15 only) Properties of Waves: Diffraction Properties of Waves: Interference  So, light is a wave? With a frequency, a wavelength, and a speed? Problem: Wave nature of light does not explain _____________. Which has the longest wavelength? A. X-rays B. visible C. infrared Which has the highest frequency? A. X-rays B. visible C. infrared Which has the highest energy? A. X-rays B. visible C. infrared  Energy of Light According to the wave model of light, energy should increase with intensity, but it doesn’t! Instead, higher frequency (and shorter wavelength) light has higher energy. If light is a wave, this doesn’t make sense!  Light is a Particle What is our evidence?  Photoelectric Effect Description Many metals emit _________ when electromagnetic radiation shines on the surface. Uses: photomultipliers, photocells, garage door openers How does this work? – The light is transferring ______ to the electrons at the metal surface where it is transformed into kinetic energy that give the electrons enough energy to “leave” the atoms in the metal. Why does adding energy in the form of light (e/m radiation) allow electrons to “escape”? Why are electrons stuck on atoms in the first place?  The Photoelectric Effect Photoelectric effect simulation http://phet.colorado.edu/en/simulation/photoelectric Predict what happens when: 1. You increase the A. Number of intensity of UV light 2. You keep the intensity electrons emitted the same and increase increases the wavelength (decreasing the B. Number of frequency) to the blue 3. You keep the intensity electrons emitted the same and increase the wavelength decreases (decreasing the C. No change frequency) into the yellow D. Zero electrons are 4. You keep the yellow light and increase the emitted intensity  Summary of Evidence from the Photoelectric Effect When light shines on a metal surface, the outcome depends on the frequency of the light. If the frequency of the light is _____ the threshold frequency, _______________ from the metal. (This creates a current.) – When the intensity (brightness) of this light is increased, more electrons are emitted. If the frequency of the light is _____ the threshold frequency, _____________________. (There is no current.) – It doesn’t matter how intense (bright) the light is; electrons are never emitted. Threshold Frequency for Na UV V I B G Y O R IR  What is the significance? If light were a wave, then increasing the intensity of the light should increase the energy of the light. With bright enough light of any frequency, electrons should be emitted. This doesn’t happen! How can we explain this? Light is not (just) a wave. It is a particle! Einstein postulated that light must come in packets of energy (or particles or quanta) – called _________.  Light is a particle (and a wave) The energy of a photon is quantized (it can only have certain values): – E = hν (h= 6.626 × 10–34 J s) – (Planck proposed this idea.) – Light comes in packets of energy. Einstein applied the idea of photons to the photoelectric effect. The energy of light depends on the __________, not on the __________.  E/m radiation is a particle - Photoelectric effect Each photon has a definable energy (E = hν) When that photon hits the metal it transfers its energy to an electron. If the photon has enough energy, it can eject one electron. If the photon does not Photoelectric effect simulation have enough energy, no electron is ejected. What is the energy of a photon of frequency 4.0 × 1018 s–1 (in the X-ray part of the spectrum)? Constant: h = 6.626 × 10–34 J s A. 2.6 × 108 J B. 2.6 × 10–15 J C. 1.7 × 10–52 J D. 6.0 × 1051 J Remember: Light is also a wave! What is the wavelength of a photon with an energy of 6.2 × 10–8 J? Constants: c = 3.0 × 108 m/s, h = 6.626 × 10–34 J s A. 3.2 × 10–34 m B. 3.2 × 1018 m C. 3.2 × 10–18 m D. 9.4 × 1025 m A covalent bond between two H atoms (in H2) requires 7.2 × 10–19 J of energy to break, causing the molecule to fall apart. What frequency of light does this correspond to? Constants: c = 3.0 × 108 m/s, h = 6.626 × 10–34 J s A. 1.1 × 1015 s–1 B. 9.2 × 10 –16 s–1 C. 4.8 × 10 –52 s–1 A covalent bond between two H atoms (in H2) requires 7.2 × 10–19 J of energy to break, causing the molecule to fall apart. What is the wavelength? Constants: c = 3.0 × 108 m/s, h = 6.626 × 10–34 J s A. 3.6 × 106 m B. 3.3 × 1023 m C. 2.8 × 10 –7 m What is the wavelength in nm? A covalent bond between two H atoms (in H2) requires 7.2 × 10–19 J of energy to break, causing the molecule to fall apart. What consequences does this have in everyday life?  Electromagnetic Radiation Can be described as either a particle or a wave. – These are models – not reality! It is truly difficult to imagine these ideas – how can one phenomenon be two different things? ________________is important at very small scales. – Matter and energy don’t behave like they do in our macroscopic world. Absorption and Emission Spectra 45 Visible Spectrum Light from the sun (white light) can be separated by a prism (Isaac Newton did this first). Visible light is only a very small part of the full e/m spectrum. 47  Atoms Emit Light: Atomic Emission Spectra Light from one particular element does not contain all the colors of the spectrum – it has only a few wavelengths! 48  Atomic Emission Spectra Prism Continuous White Light Spectrum Hot Gas (Atoms) Emission Spectrum 49  Atoms Can Also Absorb Light: Atomic Absorption Spectrum 50  Absorption and Emission Spectra Prism Continuous White Light Spectrum Hot Gas (Atoms) Emission Spectrum Cold Gas (Atoms) Absorption White Light Spectrum 51 Notice that the wavelengths of light absorbed and emitted by hydrogen are the same! (The wavelengths can be calculated using the Rydberg Equation which requires that integers (whole numbers, n) be used for each line in the spectrum, 52  Can a hydrogen atom absorb or emit every wavelength of light in the visible spectrum? A. Yes B. No C. I don’t know 53  Atomic absorption and emission spectra show light only of _________ wavelengths/energies. The spectrum of an element is the same whether that element is on Earth, in the Sun, or in a galaxy light years away. If you want another explanation, watch t his video of absorption and emission spe ctroscopy 54  Why are the spectra for each element different? Our model of the atom must be able to explain this. Does the Rutherford model of the atom explain atomic absorption/emission? 55 Changing Model of the Atom Rutherford’s Model: – Electrons circling the nucleus like planets around the sun. – This model does not explain ________________. – Also, would have led to the atom imploding! (Charges circling in an electric field would decay.) This model was not sustainable. The next model of the atom was developed by Niels Bohr. 57  Absorption and Emission Spectra Prism Continuous White Light Spectrum Hot Gas (Atoms) Emission Spectrum Cold Gas (Atoms) Absorption White Light Spectrum 58  Bohr Model Niels Bohr 59  Bohr Model Electrons move in ______ around nucleus. These orbits have definite energies and are at definite distances from the nucleus. So - the energies of electrons in atoms are _________. Explained emission and absorption spectra by invoking discrete energy levels - characterized by quantum numbers (n). Photons of electromagnetic energy are emitted or absorbed by atoms as electrons move from one energy level to another. The energy of the photons corresponds to the ________ in energy between the orbits. 60  An _______ moves to Bohr Model higher energy orbit when a ______ is Atomic e– ________. hν Excitatio e– n hν e– A ______ is Atomic e– _______ when Emission an _______ moves to lower energy 61  Problem Bohr’s model (electrons moving in defined orbits around nucleus at known energy levels) only works for _______ – and there are a lot more elements than that! A better way to represent the transitions of electrons upon absorbing or emitting photons is with ________________. 62 Better to use energy diagrams Each energy level has a quantum number The higher the number the higher the energy Energy levels are NOT ORBITS (they represent energy only, NOT the distance from the electron to the nucleus) Electrons transition between energy levels by absorbing or emitting photons with energies equal to the exact difference in energy between the two levels. 63  Better to use energy diagrams Each energy level has a quantum number The higher the number the higher the energy Energy levels are NOT ORBITS (they represent energy only, NOT the energy distance from the electron to the nucleus) Electrons transition between energy levels by absorbing or emitting photons with energies equal to the exact difference in energy between the two levels. 64  Absorption and Emission Which set of A transitions would produce an energy emission spectrum? B energy 65 Absorption and Emission Which color photon is emitted when an electron moves from level 21? A. Green B. Orange C. I don’t know energy Which transition is most likely level 13? energy A. Green B. Yellow 66  Absorption and Emission A energy Which set of transitions would produce an absorption spectrum? B energy 67 Absorption and Emission Which transition is most likely from level 21? A. Green energy B. Yellow Which photon is absorbed when an electron moves from level 13? A. Green B. Orange C. I don’t know energy 68 Absorption and Emission energy energy 69  Absorption and Emission Spectra Prism Continuous White Light Spectrum Hot Gas (Atoms) Emission Spectrum Cold Gas (Atoms) Absorption White Light Spectrum 70 1 eV (electron volt) = 1.6 × 10–19 J How much energy is required to move an electron from n=1 to n=2 levels? A. 13.6 eV B. 3.4 eV C. 10.2 eV D. 1.51 eV E. 1 eV Is the same amount of energy required to move an electron from n=2 to n=3? F. Yes Which of the following transitions for an electron in a hydrogen atom would release the largest amount of energy? A. n = 3 → n = 2 B. n = 4 → n = 2 C. n = 1 → n = 4 D. n = 2 → n = 1 72 Can a hydrogen atom absorb or emit every wavelength of light in the visible spectrum? A. Yes B. No C. I don’t know Why not? 73  How are absorption and emission different from the photoelectric effect? In the photoelectric effect, an atom absorbs a photon resulting in the ejection of an electron. The electron completely leaves the atom (the equivalent of moving to energy level n=∞). This process is called ionization (more on this later). 74  We know that Bohr’s model only works for the hydrogen atom. There are a lot more elements than that! We can represent the energies of the electrons with an energy diagram. But where are the electrons? To answer this we must see that…. 75 Matter is a Wave 76  Matter is a wave (and a particle) de Broglie: all matter has wave properties and, therefore, a wavelength λ. For any piece of matter: v (velocity), not ν λ = h/(mv) (frequency) Macroscopic objects – Atomic-scale objects (such as electrons) – 77 Matter is a Wave: λ = h/(mv What is the wavelength of an electron moving at 2.65 × 106 m s–1? Useful Information: h = 6.626 × 10–34 J s 1 J = 1 kg m2 s–2 Mass of an electron is 9.1 × 10−31 kg 78 Matter is a Wave: λ = h/(mv In the last example, the electron had a wavelength of 2.75 × 10–10 m, which is about the size of an atom. For comparison, the wavelength of a human running very fast is around 10–37 m. Is the wavelength of a human (10–37 m) comparable to the size of a human (1-2 m)? a. Yes b. No 79 What is the evidence that electrons are waves? Diffraction and Diffraction and interference pattern of interference pattern of waves of light waves of electrons 80 What is the evidence that electrons are waves? Interference pattern of X-rays passing through Al foil. Interference pattern of light passing through double slit. Interference pattern of electrons passing through Al foil. 81 What is the evidence that electrons are waves? When electrons are used in the double slit experiment, they show an interference pattern! Watch the Dr. Quantum Double Slit Experiment for the complete explanation. Extra: Is Schrodinger’s cat dead or alive? 82  How does this affect our model of the atom? Our model of the atoms must treat electrons as waves! The highlights of the Quantum Numbers and Orbitals video are reviewed in the next several slides. If you didn’t watch it already, do so tonight to fill in the gaps in your notes! 83  Are electrons circling around the nucleus in orbits? A. Yes B. No C. Sometimes D. In the H atom, but no other atoms E. I don’t know 84  The Bohr model is wrong! It treats electrons as particles only (not waves). We cannot measure accurately both the energy and position (or velocity or momentum) of an electron (or any other small particle). – This is called the Heisenberg Uncertainty Principle. – Since we know the energy of the electron from absorption and emission spectra, we cannot know where it is! Recall Bohr’s model of the atom specified both the position (of the orbit) and the energy of the electron. – This is why we can’t use the Bohr model! 85  Our model of the atom must include: Electrons (and all particles at the atomic-molecular level) have wave-like properties. – (Evidence from interference pattern.) Electrons in an atom can only have certain energies (their energies are quantized). – (Evidence from spectroscopy.) Since we know the energy of the electron, we can’t know its exact position. – (According to the Heisenberg Uncertainty Principle.) 86  Where are the electrons in atoms? To answer this, Erwin Schrödinger applied quantum mechanics (QM). QM treats electrons as waves derived by mathematical descriptions of the energies and probabilities of electrons. – The equation is called a wave equation and the description of an electron is called a wave function  (psi). The probability of finding an electron is 2. We’ll use the results of QM calculations, not the calculations themselves! 87  Where are the electrons in atoms? Even with QM, we can’t say exactly where the electrons are! However, we do know where they are likely to be found. Atomic orbitals: regions of space where electrons with a certain quantized energy have a high probability of being found. 88  What do atomic orbitals “look like”? s p d *The nucleus is located at the origin.* 89  Atomic Orbitals Atomic orbitals: – Each orbital can be described by a set of quantum numbers (n, l, ml) that are derived from quantum mechanical calculations. – A fourth quantum number, ms, describes the electron spin. 90  s orbitals There is only one s orbital in a “set” or subshell. The larger the principle quantum number, the larger the orbital and the higher the energy. 91  p orbitals pz In a set of p orbitals, there are three orbitals. The larger the principle quantum number, the larger the orbital and the higher the energy. If these are the 2p orbitals, what would the 3p orbitals look like in comparison? 92 d orbitals In a set of d orbitals, there are five orbitals. The larger the principle quantum number, the larger the orbital and the higher the energy. If these are the 3d orbitals, what would the 4d orbitals look like in comparison? 93  Orbitals make up the electron cloud. How many orbitals are in the electron cloud of one atom? A. Only the s orbital B. Only one p orbital C. Only the d orbitals D. All of the orbitals (all s, p, d, etc.) overlap within the electron cloud. 94 Video of atomic orbitals 95  Why do we need to know about orbitals? Understanding the idea that electrons can be described by orbitals of different shapes and definite energies allows us to understand the arrangement of the periodic table and predict how elements bond and react. Atomic Orbitals and the Periodic Table 96 Orbitals Have Different Energies 97  Atomic orbitals and the periodic table 1s1 1s2 98  Outer electrons are in 1s22s1 1s22s2 the three 2p orbitals 99  3p 3s 100  3d 4s 101 s p d f 102 103 Orbital Energy Diagrams Aufbau principle: Hund’s Rule: 104 Which is the orbital diagram for C? A 1s 2s 2p B 1s 2s 2p C 1s 2s 2p 105  We often refer to elements by the location of their electrons in the Alkali and outermost orbitals (valence alkaline earth electrons) Non-metals are metals are part part of the “p- of the “s-block” block” 106  Core and Valence Electrons Electron configurations aren’t that interesting, but core and valence electrons are really important! Core electrons are – A closed shell of electrons is very stable – they don’t participate in reactions. – To identify: use the last noble gas (group 18, eg Ne or Xe) and any full d shell (transition metals). Valence electrons are – These are the electrons that determine reactivity! 107 Write the electron configuration for carbon: chlorine: 108 What is the core/valence electron configuration of O? 109  How many core and valence electrons does N have? How many valence electrons does P have? How about As? And Sb? 110  How many core and valence electrons does Si have? How many valence electrons does C have? How about Ge? And Sn? 111 112 *Electron configurations shown here are for valence electrons only.* 113  Electron Configuration Review What is the electron configuration for Br? A. [Ar] 4s2 4p5 B. [Ar] 4s2 3d10 4p5 C. [Kr] 4p5 How many core and valence electrons does Br have? A. 18 core and 17 valence B. 28 core and 7 valence C. 35 core and 0 valence 118 We are making the claim that electrons are in orbitals with quantized energies. What is the evidence for this? Periodic Trends 119  Atomic Radius Which shows two Half the distance atoms interacting between the two through LDFs (not a nuclei covalent bond)? Depends on whether you are A B measuring the covalent interaction or the Van der Waals interaction (Usually the latter) 120  Atomic Radius Down a Group Make a prediction: The atomic radius of Na is ________ than that of Li. A. larger B. smaller C. no different D. Don’t know Why? What did you base your answer on? 121 Atomic Radius Across a Row Make a prediction: The atomic radius of Ne is ________ than that of Li. A. larger B. smaller C. no different D. Don’t know Why? What did you base your answer on?122 Atomic radius _________ across a row!!! 123 How can that be??? 124  What determines the size of the atom? The size of the atom depends on the balance between the: – – 125  Coulomb’s Law Unlike charges attract Like charges repel 126  Coulomb’s Law What do q1 and q2 represent? What about r? What happens to the force (F) if q1 or q2 increases? A) increase B) decrease C) stays same What happens to the force as r increases? A) increase B) decrease C) stays same 127  Coulomb’s law explains both attractions between the protons and electrons and the repulsions between the electrons. 128 The atomic radius represents the state where the _______________ between the electrons and protons are _____ to the _________________ between the electrons. 129 So why does atomic radius decrease across a row? 130  Video of atomic orbitals A filled shell (or subshell) “screens” or “shields” the nucleus. 131  Effective Nuclear Charge Electrons in same Positively or larger shell charged (other valence nucleus electrons) have no effect on charge experienced by + electron of Electron interest. of interest Electrons between (valence nucleus and electron electron of interest ) Every valence electron is (core electrons) attracted by the shield some of the 132 nuclear charge.  Effective Nuclear Charge Core electrons “cancel out” the positive charge from the same number of protons. Each electron in the valence shell feels the effect of the protons that are left. – e.g. Carbon: – __ core and __ valence electrons (and __ protons) – The __ core e– screen the positive charge of __ protons – The valence electrons are only attracted by what remains (__ protons) – The effective nuclear charge is: 133 Effective Nuclear Charge Charge screened by core electrons (# core electrons) Zeff = Z - S Effecti Actual ve Nuclear Nuclea Charge r (# protons) Charge 134  Zeff = Z - S What is the effective nuclear charge of A. +3 Boron? B. +4 Nitrogen? C. +5 Oxygen? D. +6 Fluorine? E. +7 135  Zeff = Z - S Li Be B C N O F Ne # proton s (Z) # core e– (S) ENC (Zeff) 136  Effective Nuclear Charge & Coulomb’s Law What represents effective nuclear charge in Coulomb’s Law? Effective nuclear charge _________ across a row. What happens to the attractive force across a row? A) increases B) decreases C) stays same 137 Atoms with a high effective nuclear charge “hold on to their electrons” tightly. The electrons are more strongly attracted to the nucleus (ES force is stronger). This is the reason why atomic radius decreases across a row. This is really important!!! 138 Which of the following has the smallest atomic radius? A. Ar B. Al C. Ga D. Kr Use Coulomb’s law to explain why. 139  Ionic Radius Ions are atoms in which electrons have been added or removed. 144  If an electron is removed from an atom, what is the charge of the ion? A. Neutral B. Positive C. Negative D. It depends on the element E. I don’t know The outermost electron is removed resulting in a cation. Cations are ___________ charged. 145  Formation of Cations Which has a larger radius? Why? A. Li B. Li+ C. Same D. I don’t know Which electron was removed? 146 147 If an electron is added to an atom, what is the charge of the ion? A. Neutral B. Positive C. Negative D. It depends on the element E. I don’t know An electron is added to the next available (lowest energy) orbital. Anions are ________ charged. 148  Formation of Anions Which has the largest radius? Why? A. F B. F– C. Same D. I don’t know Where was the electron added? 149 150  Which is larger? Na or F + – A. Na+ B. F– C. Same D.I don’t know 151  Isoelectronic Series All have the same ___________________ but ____________________. The attraction between the electrons and protons increases as the charge of the nucleus increases. The electron-electron repulsion is the same for each. 152 Ionization Energy 153  Ionization Energy Ionization Energy: energy required to remove an electron from an atom in the gas phase: Li(g)  Li+(g) + e– (IE = 520 kJ/mol) Why is energy required? 154  Ionization Energy Down a Group Compare Li and Na. Which is easier to remove an electron from? A. Li B. Na C. Same D. Don’t know The ionization energy of Na is __________ than that of Li. E. larger F. smaller 155  Ionization Energy Decreases Down a Group The I.E. of lithium (Li) is ___ kJ/mol. The I.E. of sodium (Na) is ___ kJ/mol. Why? 156  Ionization Energy Across a Row Compare Li and Ne. Which is it easier to remove an electron from? A. Li B. Ne C. Same D. Don’t know The ionization energy of Ne is _________ than that of Li. E. larger F. smaller 157 Ionization Energy Increases Across a Row The I.E. of Lithium (Li) is ___ kJ/mol. The I.E. of Neon (Ne) is _____ kJ/mol. Why? 158 Ionization Energy Trends 159 More evidence for electron shells: periodic trend in IEs 160  Recap: Trends Down a Group Atomic radius IE decreases increases  Li = 520 kJ/mol  Li = 152 pm  Na = 496 kJ/mol  Na = 186 pm  K = 419 kJ/mol  K = 227 pm  Rb = 403 kJ/mol  Rb = 248 pm Why? What happens to the force between each valence electron and the nucleus? Consider: – Charge of electron – Effective nuclear charge – Distance of electron from nucleus (orbitals) 161 Recap: Trends Across a Row Atomic radius decreases. Ionization energy increases. Why? What happens to the force between each valence electron and the nucleus? Consider: – Charge of electron – Effective nuclear charge – Distance of electron from nucleus (orbitals) 162 Periodic trends in atomic radius and ionization energy Smaller atoms have higher ionization energies. The trends in radius and IE are inversely related but caused by the same phenomenon – the effective nuclear charge. 163 Periodic trends in atomic radius and ionization energy 164  Successive Ionizations You can remove more than one electron: First ionization energy – M(g)  M+(g) + e– Second IE – M+(g)  M2+(g) + e– Third IE – M2+(g)  M3+(g) + e– 165  Which ionization energy is largest? A. First IE: Mg (g)  Mg+(g) + e– B. Second IE: Mg+(g)  Mg2+(g) + e– C. Third IE: Mg2+(g)  Mg3+(g) + e– 166  Which ionization energy is largest? A. First IE: Mg(g)  Mg+(g) + e– ___ kJ/mol B. Second IE: Mg+(g)  Mg 2+(g) + e– ____ kJ/mol C. Third IE: Mg2+(g)  Mg3+(g) + e– ____ kJ/mol 167 Note the huge jump for the 3rd ionization energy of Mg. Why? 168  Ionization energies are affected by Effective nuclear charge (larger ENC – larger IE) Size of atom/ion (smaller size – higher IE) The shell that the electron is removed from (IE of core e– >> IE of valence e–) Make sure you can explain WHY! 169 Consider the following successive ionization energies: IE1 1,012 Which element in kJ/mol period three would IE2 1,900 most likely show this kJ/mol trend in ionization IE3 2,910 energies? kJ/mol A. Mg D. P IE4 4,960 B. Al E. S kJ/mol C. Si IE5 6,270 kJ/mol IE6 22,200 kJ/mol 170 Consider the following successive ionization energies: IE1 1,012 kJ/mol IE2 1,900 kJ/mol IE3 2,910 kJ/mol IE4 4,960 kJ/mol IE5 6,270 kJ/mol IE6 22,200 kJ/mol 171 This is the end of Chapter 2! You should be able to apply all the Learning Objectives from Chapter 2 at this point. 172