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Chapter Two: ATOMIC STRUCTURE 2–2 EARLY HISTORY OF CHEMISTRY Egypt Practical chemical arts – very advanced Embalming and preservation of dead Metallurgical expertise Developed use of pigments Mineral and plant infusions Where does the word Chemistry come from? Kham = Egypt, derived into khemeia (Art of Egypt) Khumos = plant juice (Greek) 3 INDIA In 4000 B.C. India was working with copper, gold, paints, pottery, bricks, glass and bronze. In 2500 B.C. working with silver, lead, terra cotta, glazed pottery, gypsum cement 4 INDIA In 1500 B. C. India was working with fermented juice (wine), tanning of leather, understood the origin of matter, the theory of atoms and molecules, atomic theory and resins, 5 Greeks The greatest philosophers Greek hypotheses Concerning the nature of matter Concerning the interactions of matter No experimentation Hindered chemical advancement 6 The Greek Elements Nature of the universe Structure of materials Philosophers (lover of wisdom) Studied the “why” Chemical theory Thales – first recorded Greek philosopher Considered one of the 7 sages of Greece “can a substance be changed from one material into another?” Blue stone – heat became red Cu “Can any substance be changed into another?” “Are all substances different aspects of one basic material?” 7 Democritus He called these small pieces of matter "atomos," the Greek word for indivisible. Democritus, theorized that atoms were specific to the material which they composed. In addition, Democritus believed that the atoms differed in size and shape, were in constant motion in a void, collided with each other; and during these collisions, could rebound or stick together. Although Democritus' theory was remarkable, it was rejected by Aristotle, one of the most influential philosophers of Ancient Greece; and the atomic theory was ignored for nearly 2,000 year 2–8 Four Elements Fire, Air, Water, and Earth Accepted by Aristotle Combination of Properties Fire = hot and dry Air = hot and moist Water = cold and moist Earth = cold and dry 9 Aristotle 10 Four Elements Heavenly bodies did not appear to change Properties must be different Composed of a fifth element Ether (glow) Ether was perfect, eternal, incorruptible Four earthly elements were very different These ideas lasted for 2000 years Why did Aristotle’s philosophy hold sway for so long? Thoughts had an intuitive appeal Tutored the son of Philip of Macedonia Alexander the Great Conquered most of the known world Spread Greek culture and philosophy China to Spain 11 Change Persians conquer Greece Rule was harsh Scientific thought was suppressed Philosophers moved west Pythagoras went to Italy Founded a substantial school Empedocles, Sicily, eminent scholar Why was there a single element? 12 Political Changes Constantinople – sacked by the Crusaders Most of Greek work was lost City was recovered by Greeks City was lost to Turks where it remains Compass discovered Led to great voyages of discovery Europeans no longer felt Greeks knew it all Printing press was invented Unpopular views could not be suppressed simply by not copying the work 13 Europe’s Revival Translated Arab text into Latin Gerbert (French Scholar) later Pope Sylvester II encouraged this work Others followed By 1200 alchemical knowledge was available to European Scholars Albert of Bollstadt (Albert the Great) Studied Aristotle Given credit for discovering Arsenic Roger Bacon 14 Roger Bacon Today he is known for the idea that experimentation and mathematics would advance science Not widely accepted during his life Wrote a universal encyclopedia of knowledge where he described gunpowder Did not discover it Permitted its use to bring down medieval order Earliest symbol of use of technological proficiency Allowed Europe to conquer much of the world over the next 5 centuries 15 Chemistry Late 18th Century Sought an understanding of matter Philosophy changed into Science Before: Air – an element After: Air – a mixture of gases; individual gases were discovered Curiosity, careful observation, experimentation, measurement, and publication 16 Lavoisier and Proust Late 1700’s – Antoine Lavoisier Law of conservation of matter: mass did not change after chemical reactions. 1799 – Joseph Louis Proust Any given compound always contains the same elements in same proportions by mass. Water is always 88.9% O and 11.1% H 1896 – J.J. Thomson Cathode ray studies – stream of particles deflected with magnet or charged plates. Named these particles electrons. Could not determine mass but found charge to mass ratio = 1.76 x 108 coulomb/sec. Franklin and Faraday 1700’s – Benjamin Franklin. Objects can have 2 types of charges: + and –. Like charges repel and opposites attract. Objects can pick up negative charges – static electricity – which can suddenly discharge. 1839 – Michael Faraday Structure of atoms somehow related to electricity. 2–20 Dalton’s Atomic Theory Atomic Theory of Matter The theory that atoms are the fundamental building blocks of matter reemerged in the early nineteenth century, championed by John Dalton. Dalton’s Postulates 1) Each element is composed of extremely small particles called atoms. Dalton’s Postulates 2) All atoms of a given element are identical to one another in mass and other properties, but the atoms of one element are different from the atoms of all other elements. Dalton’s Postulates 3) Atoms of an element are not changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions. Dalton’s Postulates 4) Atoms of more than one element combine to form compounds; a given compound always has the same relative number and kind of atoms. 2–26 Fundamental Chemical Laws Three Important Laws Law of conservation of mass (Lavoisier): Mass is neither created nor destroyed in a chemical reaction. Law of definite proportion (Proust): A given compound always contains exactly the same proportion of elements by mass. 27 Three Important Laws (continued) Law of multiple proportions (Dalton): When two elements form a series of compounds, the ratios of the masses of the second element that combine with 1 gram of the first element can always be reduced to small whole numbers. 28 2–29 Copyright © Houghton Mifflin Company. All rights reserved. Early Experiments to Characterize the Atom J. J. Thomson (1898—1903) Postulated the existence of negatively charged particles, that we now call electrons, using cathode-ray tubes. Determined the charge-to-mass ratio of an electron. The atom must also contain positive particles that balance exactly the negative charge carried by electrons. Copyright © Cengage Learning. All rights reserved 30 The Atom, circa 1900 The prevailing theory was that of the “plum pudding” model, put forward by Thomson. It featured a positive sphere of matter with negative electrons embedded in it. Deflection of Cathode Rays by an Applied Electric Field Robert Millikan (1909) Performed experiments involving charged oil drops. Determined the magnitude of the charge on a single electron. Calculated the mass of the electron (9.11 × 10-31 kg). 33 A Schematic Representation of the Apparatus Millikan Used to Determine the Charge on the Electron Ernest Rutherford (1911) Explained the nuclear atom. The atom has a dense center of positive charge called the nucleus. Electrons travel around the nucleus at a large distance relative to the nucleus. 35 Lead block Uranium Florescent Screen Gold Foil What he got Rutherford's Experiment On a-Particle Bombardment of Metal Foil 2–40 Copyright © Houghton Mifflin Company. All rights reserved. The Modern View of Atomic Structure: An Introduction The Modern View of Atomic Structure The atom contains: electrons protons: found in the nucleus; positive charge equal in magnitude to the electron’s negative charge. neutrons: found in the nucleus; no charge; virtually same mass as a proton. 2–41 The Mass and Charge of the Electron, Proton, and Neutron 2–42 Copyright © Houghton Mifflin Company. All rights reserved. Atomic Number on the Periodic Table Atomic Number 6 Symbol C Name Carbon Atomic Mass 12.01 All atoms of an element have the same number of protons 6 protons Carbon 6 C Carbon 12.01 Isotope Notation 45 Atomic Symbols Show the name of the element, a hyphen, and the mass number in hyphen notation sodium-23 Show the mass number and atomic number in nuclear symbol form mass number 23 Na atomic number 11 Atomic Number, Z All atoms of the same element have the same number of protons in the nucleus, Z 13 Al 26.981 Atomic number Atom symbol AVERAGE Atomic Mass Mass Number, A C atom with 6 protons and 6 neutrons is the mass standard = 12 atomic mass units Mass Number (A) =# protons + # neutrons NOT on the periodic table…(it is the AVERAGE atomic mass on the table) A boron atom can have A = 5 p + 5 n = 10 amu A 10 Z 5 B Isotopes 49 Atoms with the same number of protons but different numbers of neutrons. Show almost identical chemical properties; chemistry of atom is due to its electrons. In nature most elements contain mixtures of isotopes. Different Forms of the Same Element In any element, the # of protons is always constant. Unlike the number of protons, the number of electrons and neutrons can vary within an element without changing the identity of the element. Ex. Carbon (C) ALWAYS has 6 protons, but it can have anywhere from 6-8 neutrons and 2-10 electrons Isotopes An ISOTOPE is a form of an element that has a different number of neutrons than “normal” Carbon has three isotopes Notice how the # of protons does NOT change! Masses of Atoms A scale designed for atoms gives their small atomic masses in atomic mass units (amu) An atom of 12C was assigned an exact mass of 12.00 amu Relative masses of all other atoms was determined by comparing each to the mass of 12C An atom twice as heavy has a mass of 24.00 amu. An atom half as heavy is 6.00 amu. Atomic Mass Listed on the periodic table Gives the mass of “average” atom of each element compared to 12C Average atom based on all the isotopes and their abundance % Atomic mass is not a whole number How to calculate Avg. Atomic Mass To calculate the atomic mass of an element, multiply the mass of each isotope by its natural abundance, expressed as a decimal, and then add the products. Sample Problem 1 Rubidium has two common isotopes, 85-Rb and 87-Rb. If the abundance of 85-Rb is 72.2% and the abundance of 87Rb is 27.8%, what is the average atomic mass of rubidium? Isotope % abundance Fraction of abundance 85-Rb 72.2% 0.722 mass X 85 amu = 61.37 87-Rb 27.8% 0.278 X 87 amu = 24.186 +____________ __ 85.556 amu Atomic Masses 3–57 Elements occur in nature as mixtures of isotopes Carbon = 98.89% 12C 1.11% 13C <0.01% 14C Carbon atomic mass = 12.01 amu Examples 3–58 There are two isotopes of carbon 12C with a mass of 12.00000 amu(98.892%), and 13C with a mass of 13.00335 amu (1.108%) There are two isotopes of nitrogen , one with an atomic mass of 14.0031 amu and one with a mass of 15.0001 amu. What is the percent abundance of each? IONS An atom usually has a neutral charge. That means it has the same number of protons as electrons Remember, a proton has a positive charge and an electron has a negative charge ION – an atom that has lost or gained one or more electrons and has become charged either positively or negatively Ions are atoms with charges The charges can be positive or negative Positive charge is a Cation Negative charge is Anion Positive Ions When an atom LOSES electrons, it becomes more POSITIVE Why? If you are getting rid of negative particles (electrons) but your number of positive particles (protons) are staying the same. In other words, you are subtracting negative numbers Examples What would the charge be if: The neutral form of Gold (Au) lost 4 of its 79 electrons. It now has 79 protons and 75 electrons The neutral form of Mg lost 2 of its 12 electrons. It now has 12 protons and 10 electrons. Negative Ions When an atom GAINS electrons it becomes more NEGATIVE Why? Electrons have a negative charge, so the more you have, the more negative you become Representing Ions Ions are represented by placing a “superscript” charge number next to the atomic symbol. Ex. O-2 = oxygen with a negative 2 charge K+ = potassium with a positive 1 charge N-3 = nitrogen with a negative 3 charge And so on The Nature of Matter and Energy All Waves Have 4 Properties • Amplitude: height from origin to crest. – Determines the brightness or intensity of light. • Wavelength: distance between crests. – Visible light between 400 – 750 nanometers. • Frequency: how fast wave oscillates up and down. – Cycles per second = a hertz (Hz) • Speed of light: 3.00 x 108 m/s λ = c/ν (λ-lambda=wavelength, ν-nu=frequency) Properties of Waves All electromagnetic waves travel at the same speed The speed of light: 300,000 km/s crest trough Properties of Waves Wavelength crest (length/cycle) trough Wavelength (): the length of one complete cycle Properties of Waves Wavelength crest (length/cycle) Amplitude trough Amplitude: 1/2 height between trough and crest Properties of Waves Wavelength crest (length/cycle) Amplitude trough Frequency (): the number of cycles/second Speed = wavelength x frequency c= (length/second) = (length/cycle) x (cycle/second) Hence, =c/ and =c/ Examples 5 Hz? What is the wavelength of light with a frequency 5.89 x 10 What is the frequency of blue light with a wavelength of 484 nm? The Nature of Waves In 1900 Matter and energy were seen as different from each other in fundamental ways Matter was particles Energy could come in waves, with any frequency. Max Planck found that the cooling of hot objects couldn’t be explained by viewing energy as a wave. Energy is Quantized Planck found DE came in chunks with size h DE = hν where n is an integer. and h is Planck’s constant h = 6.626 x 10-34 J s these packets of hν are called quantum Quantum Theory • Planck: proposed that objects absorb and emit energy in restricted amounts. – Called energy pieces quantum = fixed amount. • E = hν – E = energy – h = Planck’s constant (6.6262 x 10-34 J-s) – ν = frequency Energy of a wave Energy is proportional to frequency, and inversely proportional to wavelength E=h = h (c/ ) where h = Planck’s constant In other words, waves with shorter wavelengths (or higher frequency) have higher energy • Calculate the energy of one phonton of yellow light whose wavelength is 589nm • Ans: 3.37x10-19J Photoelectric Effect • Einstein proposed that light consists of quanta of energy that behave like tiny particles he called photons. • When a photon strikes metal, energy is transferred from the photon to an electron. – Electron either absorbs entire photon or none of it. (cannot used part, and cannot collect energy from several photons. If energy is too small, electron doesn’t escape) – Determined by energy of frequency, not number or intensity of photons. The Photoelectric Effect • Electrons are ejected from the surface of a metal when light shines it. – For each metal, a minimum frequency of light is needed to release electrons. Na metal Einstein is next Said electromagnetic radiation is quantized in particles called photons Each photon has energy = h Combine this with E = mc you get the apparent mass of a photon m = h / (c) ν = hc/ 2 The Photoelectric Effect Dual Nature of Radiant Energy • 1923 – Arthur Compton showed a photon could collide with an electron like two tiny balls. • Light is both a particle and a wave. Which is it? Is energy a wave like light, or a particle? Yes Concept is called the Wave Particle duality. What about the other way, is matter a wave? Yes The Wave Nature of Matter The wave nature of light is used to produce this electron micrograph. Louis de Broglie theorized that if light can have material properties, matter should exhibit wave properties. He demonstrated that the relationship between mass and wavelength was h = mv Matter as a wave Using the velocity v instead of the frequency ν we get De Broglie’s equation = h/mv can calculate the wavelength of an object Electromagnetic Radiation Light Made up of electromagnetic radiation Waves of electric and magnetic fields at right angles to each other. Light: Electromagnetic Radiation • Electric and magnetic fields oscillating at right angles to each other. Classification of Electromagnetic Radiation (a) Diffraction Pattern (b) Constructive Interference of Waves (c.) Destructive Interference of Waves Summary Electromagnetic radiation, which at the turn of the 20th century was thought to be pure waveform, was found to possess particulate properties Electrons which were thought to be particles were found to have wavelength associated with them The significance is that matter and energy are not distinct Energy is realy a form of matter, and all matter shows the same type of properties All matter exhibits both particulate and wave properties The Atomic Spectrum of Hydrogen Spectrum The range of frequencies present in light. White light has a continuous spectrum. All the colors are possible. A rainbow. A Beautiful Rainbow What this means Only certain energies are allowed for the hydrogen atom. Can Use only give off certain energies. DE = h = hc / Energy in the atom is quantized Hydrogen spectrum Emission spectrum because these are the colors it gives off or emits Called a line spectrum. There are just a few discrete lines showing 656 nm 434 nm 410 nm 486 nm •Spectrum The Bohr Model Niels Bohr Developed the quantum model of the hydrogen atom. He said the atom was like a solar system The electrons were attracted to the nucleus because of opposite charges. Didn’t fall in to the nucleus because it was moving around The Bohr Ring Atom He didn’t know why but only certain energies were allowed. He called these allowed energies energy levels. Putting energy into the atom moved the electron away from the nucleus From ground state to excited state. When it returns to ground state it gives off light of a certain energy The Bohr Model Niels Bohr adopted Planck’s assumption and explained these phenomena in this way: 1. Electrons in an atom can only occupy certain orbits (corresponding to certain energies). The Bohr Model 2. Electrons in permitted orbits have specific, “allowed” energies; these energies will not be radiated from the atom. 3. Energy is only absorbed or emitted in such a way as to move an electron from one “allowed” energy state to another; the energy is defined by E = h Limitations of the Bohr Model It only works for hydrogen! Classical physics would result in an electron falling into the positively charged nucleus. Bohr simply assumed it would not! Circular motion is not wave-like in nature. Important Ideas from the Bohr Model Points that are incorporated into the current atomic model include the following: Electrons exist only in certain discrete energy levels. 2) Energy is involved in the transition of an electron from one level to another. 1) The Bohr Ring Atom n=4 n=3 n=2 n=1 An excited lithium atom emitting a photon of red light to drop to a lower energy state. An excited H atom returns to a lower energy level. The Bohr Model n is the energy level for each energy level the energy is Z is the nuclear charge, which is +1 for hydrogen. E = -2.178 x 10-18 J (Z2 / n2 ) n = 1 is called the ground state We are worried about the change When the electron moves from one energy level to another. DE = Efinal - Einitial DE = -2.178 x 10-18 J Z2 (1/ nf2 - 1/ ni2) When is it true? Only for hydrogen atoms and other monoelectronic species. Why the negative sign? To increase the energy of the electron you make it further to the nucleus. the maximum energy an electron can have is zero, at an infinite distance. The Bohr Model Doesn’t only work works for hydrogen atoms electrons the don’t move in circles quantization of energy is right, but not because they are circling like planets. The Uncertainty Principle Heisenberg showed that the more precisely the momentum of a particle is known, the less precisely is its position is known: Electron Configurations Electron Configurations The quantum mechanical model of the atom predicts energy levels for electrons; it is concerned with probability, or likelihood, of finding electrons in a certain position. ORBITALS ORBITAL: the regions in an atom where there is a high probability of finding electrons. s is the lowest energy orbital, and p is slightly higher d and f are the next two orbitals. They occupy even higher energy levels and take on more complex shapes than s & p Electron Configurations Regions where electrons are likely to be found are called orbitals. EACH ORBITAL CAN HOLD UP TO 2 ELECTRONS! Electron Configurations Sublevels (l) Principal Energy Level Sublevels Orbitals n=1 1s 2s 2p one (1s) one (2s) three (2p) n=3 3s 3p 3d one (3s) three (3p) five (3d) n=4 4s 4p 4d 4f one (4s) three (4p) five (4d) seven (4f) n=2 Sublevels (l) Sublevel # of orbitals Max # of electrons s 1 2 p 3 6 d 5 10 f 7 14 The Aufbau Principle Electron Configurations 5 4p The way electrons are distributed in an atom is called its electron configuration. The most stable organization is the lowest possible energy, called the ground state. Each component consists of a number denoting the energy level; Electron Configurations 5 4p The way electrons are distributed in an atom is called its electron configuration. The most stable organization is the lowest possible energy, called the ground state. Each component consists of a number denoting the energy level; a letter denoting the type of orbital; Electron Configurations 5 4p The way electrons are distributed in an atom is called its electron configuration. The most stable organization is the lowest possible energy, called the ground state. Each component consists of a number denoting the energy level; a letter denoting the type of orbital; a superscript denoting the number of electrons in those orbitals. Electron Configuration 1 1s group # row # # valence eshell # possibilities are: possibilities are 1-7 s: 1 or 2 7 rows subshell p: 1-6 possibilities are d: 1-10 s, p, d, or f f: 1-14 4 subshells Total e- should equal Atomic # What element has an electron configuration of 1s1? The Orbitals Being Filled for Elements in Various Parts of the Periodic Table Aufbau Principle Assumes that all atoms have the same type of orbitals that the Hydrogen atom does As Protons are added to the nucleus to build up the elements, similarly electrons are added to the hydrogen like orbitals Electron configuration is how the electrons are distributed among the various atomic orbitals in an atom. No more than 2 Electrons in Any Orbital…ever. Wolfgang Pauli, yet another German Nobel Prize winner The next rule is the Pauli Exclusion Principal. The Pauli Exclusion Principle states that an atomic orbital may have up to 2 electrons and then it is full. The spins have to be paired. We usually represent this with an up arrow and a down arrow. Since there is only 1 s orbital per energy level, only 2 electrons fill that orbital. Orbital Diagrams Each box in the diagram represents one orbital. Half-arrows represent the electrons. The direction of the arrow represents the relative spin of the electron. Hund’s Rule “For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized.” This means that, for a set of orbitals in the same sublevel, there must be one electron in each orbital before pairing and the electrons have the same spin, as much as possible. Hund’s Rule Don’t pair up the 2p electrons until all 3 orbitals are half full. Hunds Rule states that when you get to orbitals, you fill them all half way first, and then you start pairing up the electrons. Increasing energy 7s 6s 5s 7p 6p 5p 6d 5d 6f 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s Orbitals available to a Hydrogen atom Increasing energy 7s 6s 5s 4s 3s 2s 1s 7p 6p 5p 4p 6d 5d 4d 3d 3p 2p With more electrons, repulsion changes the energy of the orbitals. 5f 4f Increasing energy 7s 6s 5s 4s 3s 2s 1s 7p 6p 5p 4p 6d 5d 4d 3d 3p 2p He with 2 electrons 5f 4f Details Valence electrons- the electrons in the outermost energy levels (not d). Core electrons- the inner electrons Hund’s Rule- The lowest energy configuration for an atom is the one have the maximum number of unpaired electrons in the orbital. C 1s2 2s2 2p2 Details Valence electrons- the electrons in the outermost energy levels (not d). Core electrons- the inner electrons Hund’s Rule- The lowest energy configuration for an atom is the one have the maximum number of unpaired electrons in the orbital. C 1s 2 2s2 2p2 Shorthand Notation A way of abbreviating long electron configurations Since we are only concerned about the outermost electrons, we can skip to places we know are completely full (noble gases), and then finish the configuration Shorthand Notation Step 1: It’s the Showcase Showdown! Find the closest noble gas to the atom (or ion), WITHOUT GOING OVER the number of electrons in the atom (or ion). Write the noble gas in brackets [ ]. Step 2: Find where to resume by finding the next energy level. Step 3: Resume the configuration until it’s finished. 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 So you start at level 3 on the diagonal rule (all levels start with s) and finish the configuration by adding 7 more electrons to bring the total to 17 [Ne] 3s2 3p5 Practice Shorthand Notation Write the shorthand notation for each of the following atoms: Cl K Ca I Bi The Shorthand Write the symbol of the noble gas before the element Then the rest of the electrons. Aluminum - full configuration 1s 2s 2p 3s 3p Ne is 1s 2s 2p so Al is [Ne] 3s 3p 2 2 6 2 1 2 2 6 2 1 IB sometimes uses another form of notation for this… 2 2 6 2 1 1s 2s 2p 3s 3p becomes… 2e ) Electrons in 1st energy level (shell) 8e Electrons in 2nd energy level (shell) ) 3e Electrons in 3rd energy level (shell) 2.8.3 or 2,8,3 IB will only use this notation for the first 20 elements Exceptions you should know: Copper: Chromium: 1s2 2s2 2p6 3s2 3p64s13d10 or [Ar]4s1 3d10 1s2 2s2 2p6 3s2 3p64s13d5 or [Ar]4s1 3d5 Why do these elements do this? Lower energy arrangement of electrons: this configuration allows for half-filled (less electron repulsion because no pairing in the 3d sublevel) or filled d-sublevel. Electron configurations of Ions N3- [He]2s22p6 Al6+ 1s22s22p3 Cr3+ [Ar]3d3 Cu+ [Ar]3d10 = [Ne] Se2- [Kr] Mg2+ [Ne] Ionization Energy AS IT RELATES TO ELECTRON CONFIGURATION AHL 12.1.1 and 12.1.2 12.1.1 Explain how evidence from first ionization energies across periods accounts for the existence of the main energy levels and sub-levels in an atom. 12.1.2 Explain how successive ionization energy data is related to the electron configuration of an atom. Ionization Energy Definition: The ionization energy of an atom is the minimum amount of energy required to remove a mole of electrons from a mole of gaseous atoms to form a mole of gaseous ions https://www.youtube.com/watch?v=5CBs36jt ZxY First Ionization Energies 2500 First Ionization Energies (KJ/mol) 2000 1500 1000 500 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 Atomic Number First Ionization Energy The Magnitude of Ionization Energies Ionization energy depends on charge of the nucleus. This will be counteracted by the repulsion, or “shielding” of electrons in filled inner orbitals. Each electron in a filled inner shell will cancel one unit of nuclear charge and after these have been subtracted, the remaining nuclear charge is referred to as the effective nuclear charge (ENC) Effective Nuclear Charge Diagram the effective nuclear charge (ENC) for the electrons in magnesium http://www.youtube.com/watch?v=IvSmfgxCSNQ Shielding Effect The shielding (or screening) effect is similar to effective nuclear charge. The core electrons repel the valence electrons to some degree. The more electron shells there are (a new shell for each row in the periodic table), the greater the shielding effect is. Essentially, the core electrons shield the valence electrons from the positive charge of the nucleus. Successive Ionization Numbers First Ionization Q(g) Q+ Second Q+(g) Third (g) + e− Ionization Q2+(g) + e− Ionization Q2+(g) Q3+(g) + e− Note that these are all endothermic changes, because work has to be done to remove a negatively charged electron from the attraction of a positively charged nucleus. http://www.chemguide.co.uk/atoms/properties/moreies.html Graph of Successive Ionization Energies http://legacy.chemgym.net/as_a2/topics/successive_ionisation/graphical.html Successive Ionization Energies 1st Ionization Energy (kJ) This graph provides evidence that the levels can contain different numbers of electrons before they become full. 2500 Ne 2000 Ar 1500 1000 500 Li 0 0 Na 5 10 Atomic Number K 15 20 Level 1 (K shell) 2 (L shell) 3 (M shell) Max # electrons 2 8 8 (or 18) Second ionization energy energy required to remove the second electron. 2 X (g) X (g) e -
