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Atoms, Molecules, and Ions Chemistry Timeline #1 B.C. 400 B.C. Demokritos and Leucippos use the term "atomos” 2000 years of Alchemy 1500's Georg Bauer: systematic metallurgy Paracelsus: medicinal application of minerals 1600's Robert Boyle:The Skeptical Chemist. Quantitative experimentation, identification of elements 1700s' Georg Stahl: Phlogiston Theory Joseph Priestly: Discovery of oxygen Antoine Lavoisier: The role of oxygen in combustion, law of conservation of mass, first modern chemistry textbook Chemistry Timeline #2 1800's Joseph Proust: The law of definite proportion (composition) John Dalton: The Atomic Theory, The law of multiple proportions Joseph Gay-Lussac: Combining volumes of gases, existence of diatomic molecules Amadeo Avogadro: Molar volumes of gases Jons Jakob Berzelius: Relative atomic masses, modern symbols for the elements Dmitri Mendeleyev: The periodic table J.J. Thomson: discovery of the electron Henri Becquerel: Discovery of radioactivity 1900's Robert Millikan: Charge and mass of the electron Ernest Rutherford: Existence of the nucleus, and its relative size Meitner & Fermi: Sustained nuclear fission Ernest Lawrence: The cyclotron and trans-uranium elements Laws • Conservation of Mass • Law of Definite Proportion – – compounds have a constant composition. – They react in specific ratios by mass. • Multiple Proportions- – When two elements form more than one compound, the ratios of the masses of the second element that combine with one gram of the first can be reduced to small whole numbers. Proof • Mercury has two oxides. – One is 96.2 % mercury by mass, the other is 92.6 % mercury by mass. • Show that these compounds follow the law of multiple proportion. • Speculate on the formula of the two oxides. Dalton’s Atomic Theory (1808) All matter is composed of extremely small particles called atoms Atoms of a given element are identical in size, mass, and other properties; atoms of different John Dalton elements differ in size, mass, and other properties Atoms cannot be subdivided, created, or destroyed Atoms of different elements combine in simple whole-number ratios to form chemical compounds In chemical reactions, atoms are combined, separated, or rearranged Modern Atomic Theory Several changes have been made to Dalton’s theory. Dalton said: Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties Modern theory states: Atoms of an element have a characteristic average mass which is unique to that element. Modern Atomic Theory #2 Dalton said: Atoms cannot be subdivided, created, or destroyed Modern theory states: Atoms cannot be subdivided, created, or destroyed in ordinary chemical reactions. However, these changes CAN occur in nuclear reactions! Discovery of the Electron In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle. Cathode ray tubes pass electricity through a gas that is contained at a very low pressure. Thomson’s Atomic Model Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model. Rutherford’s Gold Foil Experiment Alpha particles are helium nuclei Particles were fired at a thin sheet of gold foil Particle hits on the detecting screen (film) are recorded Atomic Particles Particle Charge Mass (kg) Location Electron -1 9.109 x 10-31 Electron cloud Proton +1 1.673 x 10-27 Nucleus 0 1.675 x 10-27 Nucleus Neutron The Atomic Scale Most of the mass of the atom is in the nucleus (protons and neutrons) Electrons are found outside of the nucleus (the electron cloud) Most of the volume of the atom is empty space “q” is a particle called a “quark” About Quarks… Protons and neutrons are NOT fundamental particles. Protons are made of two “up” quarks and one “down” quark. Neutrons are made of one “up” quark and two “down” quarks. Quarks are held together by “gluons” Isotopes Isotopes are atoms of the same element having different masses due to varying numbers of neutrons. Isotope Protons Electrons Neutrons Hydrogen–1 (protium) 1 1 0 Hydrogen-2 (deuterium) 1 1 1 Hydrogen-3 (tritium) 1 1 2 Nucleus Atomic Masses Atomic mass is the average of all the naturally isotopes of that element. Carbon = 12.011 Symbol Composition of the nucleus % in nature Carbon-12 12C 6 protons 6 neutrons 98.89% Carbon-13 13C 6 protons 7 neutrons 1.11% Carbon-14 14C 6 protons 8 neutrons <0.01% Isotope Molecules Two or more atoms of the same or different elements, covalently bonded together. Molecules are discrete structures, and their formulas represent each atom present in the molecule. Benzene, C6H6 Covalent Network Substances Covalent network substances have covalently bonded atoms, but do not have discrete formulas. Why Not?? Graphite Diamond Ions Cation: A positive ion • Mg2+, NH4+ Anion: A negative ion Cl-, SO42- Ionic Bonding: Force of attraction between oppositely charged ions. Ionic compounds form crystals, so their formulas are written empirically (lowest whole number ratio of ions). Periodic Table with Group Names Predicting Ionic Charges Group 1: Lose 1 electron to form 1+ ions H+ Li+ Na+ K+ Predicting Ionic Charges Group 2: Loses 2 electrons to form 2+ ions Be2+ Mg2+ Ca2+ Sr2+ Ba2+ Predicting Ionic Charges B3+ Al3+ Ga3+ Group 13: Loses 3 electrons to form 3+ ions Predicting Ionic Charges Caution! C22- and C4are both called carbide Group 14: Loses 4 electrons or gains 4 electrons Predicting Ionic Charges N3- Nitride P3- Phosphide As3- Arsenide Group 15: Gains 3 electrons to form 3- ions Predicting Ionic Charges O2- Oxide S2- Sulfide Se2- Selenide Group 16: Gains 2 electrons to form 2- ions Predicting Ionic Charges F- Fluoride Br- Bromide Cl- Chloride I- Iodide Group 17: Gains 1 electron to form 1- ions Predicting Ionic Charges Group 18: Stable Noble gases do not form ions! Predicting Ionic Charges Groups 3 - 12: Many transition elements have more than one possible oxidation state. Iron(II) = Fe2+ Iron(III) = Fe3+ Predicting Ionic Charges Groups 3 - 12: Some transition elements have only one possible oxidation state. Zinc = Zn2+ Silver = Ag+ Cadmium = Cd2+ Writing Ionic Compound Formulas Example: Barium nitrate 1. Write the formulas for the cation and anion, including CHARGES! 2. Check to see if charges are balanced. 2+ Ba ( NO3- ) 2 3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced! Writing Ionic Compound Formulas Example: Ammonium sulfate 1. Write the formulas for the cation and anion, including CHARGES! 2. Check to see if charges are balanced. ( NH4+) SO42- 3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. 2 Not balanced! Naming Ionic Compounds • 1. Cation first, then anion • 2. Monatomic cation = name of the element • Ca2+ = calcium ion • 3. Monatomic anion = root + -ide • Cl- = chloride • CaCl2 = calcium chloride Naming Ionic Compounds (continued) Metals with multiple oxidation states some metal forms more than one cation use Roman numeral in name PbCl2 Pb2+ is the lead(II) cation PbCl2 = lead(II) chloride Naming Binary Compounds Compounds between two nonmetals First element in the formula is named first. Second element is named as if it were an anion. Use prefixes Only use mono on second element P2O5 = diphosphorus pentoxide CO2 = carbon dioxide CO = carbon monoxide N2O = dinitrogen monoxide Acids • Substances that produce H+ ions when dissolved in water. • All acids begin with H. • Two types of acids: • Oxyacids • Non-oxyacids Naming acids • If the formula has oxygen in it • write the name of the anion, but change – ate to -ic acid – ite to -ous acid • Watch out for sulfuric and sulfurous • H2CrO4 • HMnO4 • HNO2 Naming acids • • • • • • If the acid doesn’t have oxygen add the prefix hydrochange the suffix -ide to -ic acid HCl H2S HCN Formulas for acids • • • • • • • Hydrofluoric acid Dichromic acid Carbonic acid Hydrophosphoric acid Nitric acid Perchloric acid Phosphorous acid HF H2Cr2O7 H2CO3 H3P HNO3 HClO4 H3PO3 Selenium would commonly form this ion: Se2+ Se+ Se2Sl2S2SeSe36 7 8 9 10 21 22 23 24 25 26 27 28 29 30 11 12 13 14 15 16 0% 17 0% 18 19 0% 20 Se 3- 5 0% Se - 4 0% S2 - 3 Se + 2 Se 2+ 1 0% Sl2 - 0% Se 2- 1. 2. 3. 4. 5. 6. 7. Cesium would commonly form this ion: Ce2+ Cs+ Cs2CCsCs2+ Cm+ 7 8 9 10 21 22 23 24 25 26 27 28 29 30 11 12 13 14 15 16 17 0% 18 19 0% 20 + 6 0% Cm 5 0% Cs 2+ 4 0% Cs - 3 Cs + 2 Ce 2+ 1 0% C- 0% Cs 2- 1. 2. 3. 4. 5. 6. 7. This is the formula for zinc hydroxide: 9 10 21 22 23 24 25 26 27 28 29 30 11 12 13 14 15 16 17 H) 2 18 19 0% Zn 2H 8 0% Zn 2( O 7 OH 2) 2 6 0% Zn ( 5 0% Zn H2 4 0% OH )2 3 Zn O 2 Zn O 1 0% H2 0% Zn ( ZnOH ZnOH2 Zn(OH)2 ZnH2 Zn(OH2)2 Zn2(OH)2 Zn2H H 1. 2. 3. 4. 5. 6. 7. 20 This is the formula for hydrochloric acid: 8 9 10 21 22 23 24 25 26 27 28 29 30 11 12 13 14 15 16 17 18 0% H2 Cl O 7 0% H2 Cl 6 4 5 0% HC lO 4 3 3 0% HC lO 2 0% 2 1 0% HC lO 0% HC lO HCl HClO HClO2 HClO3 HClO4 H2Cl H2ClO HC l 1. 2. 3. 4. 5. 6. 7. 19 20 Iron would commonly form this ion: Fe2+ Fe+ Fe2FeIr2+ Ir+ Fe3+ 6 7 8 9 10 21 22 23 24 25 26 27 28 29 30 11 12 13 0% Fe 3+ 5 0% Ir+ 4 0% Ir2 + 3 0% Fe - 2 0% Fe 2- 1 0% Fe + 0% Fe 2+ 1. 2. 3. 4. 5. 6. 7. 14 15 16 17 18 19 20 This slide contains classified material and cannot be shown to high school students. Please continue as if everything is normal. Which points of Dalton’s theory are not true based on current understanding of the atom? 4 5 6 7 8 9 10 21 22 23 24 25 26 27 28 29 30 11 12 13 m 18 l.. . ica ch em 17 0% 5. In 16 4. At o 15 0% so fd i. . . 0% no ... . .. m m so fa r. .. at te m 14 3. At o 3 2. At o 2 0% 1. Al l 1 0% sc an 1. All matter is composed of extremely small particles called atoms 2. Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties 3. Atoms cannot be subdivided, created, or destroyed 4. Atoms of different elements combine in simple whole-number ratios to form chemical compounds 5. In chemical reactions, atoms are combined, separated, or rearranged 19 20 Quantum Mechanics The Puzzle of the Atom Protons and electrons are attracted to each other because of opposite charges Electrically charged particles moving in a curved path give off energy Despite these facts, atoms don’t collapse Wave-Particle Duality JJ Thomson won the Nobel prize for describing the electron as a particle. His son, George Thomson won the Nobel prize for describing the wave-like nature of the electron. The electron is a particle! The electron is an energy wave! Toupee? The Wave-like Electron The electron propagates through space as an energy wave. To understand the atom, one must understand the behavior of electromagnetic waves. Louis deBroglie Electromagnetic radiation propagates through space as a wave moving at the speed of light. c = C = speed of light, a constant (3.00 x 108 m/s) = frequency, in units of hertz (hz, sec-1) = wavelength, in meters Types of electromagnetic radiation: The energy (E ) of electromagnetic radiation is directly proportional to the frequency () of the radiation. E = h E = Energy, in units of Joules (kg·m2/s2) h = Planck’s constant (6.626 x 10-34 J·s) = frequency, in units of hertz (hz, sec-1) Long Wavelength = Low Frequency = Low ENERGY Short Wavelength = High Frequency = High ENERGY Wavelength Table Relating Frequency, Wavelength and Energy c E h Common re-arrangements: E hc hc E Spectroscopic analysis of the visible spectrum… …produces all of the colors in a continuous spectrum Spectroscopic analysis of the hydrogen spectrum… …produces a “bright line” spectrum Electron transitions involve jumps of definite amounts of energy. This produces bands of light with definite wavelengths. Bohr Model Energy Levels Electron Energy in Hydrogen Eelectron - 2.178 x10 -18 Z J 2 n 2 Z = nuclear charge (atomic number) n = energy level ***Equation works only for atoms or ions with 1 electron (H, He+, Li2+, etc). Calculating Energy Change, E, for Electron Transitions 2 2 Z -18 Z E - 2.178 x 10 J 2 - 2 n n initial final Energy must be absorbed from a photon (+E) to move an electron away from the nucleus Energy (a photon) must be given off (-E) when an electron moves toward the nucleus Quantum Numbers Each electron in an atom has a unique set of 4 quantum numbers which describe it. Principal quantum number (n) Angular momentum quantum number (l) Magnetic quantum number (m) Spin quantum number (s) Pauli Exclusion Principle No two electrons in an atom can have the same four quantum numbers. Wolfgang Pauli Principal Quantum Number Generally symbolized by n, it denotes the shell (energy level) in which the electron is located. Number of electrons that can fit in a shell: 2n2 Angular Momentum Quantum Number The angular momentum quantum number, generally symbolized by l, denotes the orbital (subshell) in which the electron is located. l =3 f Magnetic Quantum Number The magnetic quantum number, generally symbolized by m, denotes the orientation of the electron’s orbital with respect to the three axes in space. Assigning the Numbers The three quantum numbers (n, l, and m) are integers. The principal quantum number (n) cannot be zero. n must be 1, 2, 3, etc. The angular momentum quantum number (l ) can be any integer between 0 and n - 1. For n = 3, l can be either 0, 1, or 2. The magnetic quantum number (ml) can be any integer between -l and +l. For l = 2, m can be either -2, -1, 0, +1, +2. Principle, angular momentum, and magnetic quantum numbers: n, l, and ml Spin Quantum Number Spin quantum number denotes the behavior (direction of spin) of an electron within a magnetic field. Possibilities for electron spin: 1 2 1 2 An orbital is a region within an atom where there is a probability of finding an electron. This is a probability diagram for the s orbital in the first energy level… Orbital shapes are defined as the surface that contains 90% of the total electron probability. Schrodinger Wave Equation d V 8 m dx h 2 2 2 2 E Equation for probability of a single electron being found along a single axis (x-axis) Erwin Schrodinger Heisenberg Uncertainty Principle “One cannot simultaneously determine both the position and momentum of an electron.” You can find out where the electron is, but not where it is going. Werner Heisenberg OR… You can find out where the electron is going, but not where it is! Sizes of s orbitals Orbitals of the same shape (s, for instance) grow larger as n increases… Nodes are regions of low probability within an orbital. Orbitals in outer energy levels DO penetrate into lower energy levels. Penetration #1 This is a probability Distribution for a 3s orbital. What parts of the diagram correspond to “nodes” – regions of zero probability? Which of the orbital types in the 3rd energy level Does not seem to have a “node”? WHY NOT? Penetration #2 The s orbital has a spherical shape centered around the origin of the three axes in space. s orbital shape P orbital shape There are three peanut-shaped p orbitals in each energy level above n = 1, each assigned to its own axis (x, y and z) in space. d orbital shapes Things get a bit more complicated with the five d orbitals that are found in the d sublevels beginning with n = 3. To remember the shapes, think of: “double peanut” …and a “peanut with a donut”! Shape of f orbitals Things get even more complicated with the seven f orbitals that are found in the f sublevels beginning with n = 4. To remember the shapes, think of: Flower Element Lithium Configuration notation 1s22s1 [He]2s1 ____ 1s Beryllium ____ ____ 2p ____ ____ 2s ____ ____ 2p ____ [He]2s2p2 ____ 2s ____ ____ 2p ____ 1s22s2p3 [He]2s2p3 ____ 2s ____ ____ 2p ____ 1s22s2p4 [He]2s2p4 ____ 2s ____ ____ 2p ____ 1s22s2p5 [He]2s2p5 ____ 1s Neon ____ 2s 1s22s2p2 ____ 1s Fluorine ____ [He]2s2p1 ____ 1s Oxygen ____ 2p 1s22s2p1 ____ 1s Nitrogen ____ [He]2s2 ____ 1s Carbon ____ 2s 1s22s2 ____ 1s Boron Noble gas notation Orbital notation ____ 2s ____ ____ 2p ____ 1s22s2p6 [He]2s2p6 ____ 1s ____ 2s ____ ____ 2p ____ Orbital filling table Electron configuration of the elements of the first three series Irregular confirmations of Cr and Cu Chromium steals a 4s electron to half fill its 3d sublevel Copper steals a 4s electron to FILL its 3d sublevel In Bohr’s atomic theory, when an electron moves from one energy level to another energy level more distant from the nucleus. 1. 2. 3. 4. 5. energy is emitted energy is absorbed no change in energy occurs light is emitted none of these 0% 0% 0% 0% 0% itt er en gy is ed em er en gy no 1 2 3 4 5 6 7 8 9 10 21 22 23 24 25 26 27 28 29 30 11 12 13 14 15 16 is c d be r so ab e ng a h in e en y rg 17 s ur c oc h lig 18 itt s ti ed em 19 ne no o h ft e es 20 Which form of electromagnetic radiation has the longest wavelengths? io tio ra d ia 0% in fra ra d n av es w ow icr m 0% xra ys 0% av e s ra y a m 0% s 0% re d gamma rays microwaves radio waves infrared radiation x-rays ga m 1. 2. 3. 4. 5. 1 2 3 4 5 6 7 8 9 10 21 22 23 24 25 26 27 28 29 30 11 12 13 14 15 16 17 18 19 20 How many electrons in an atom can have the quantum numbers n = 3, l = 2? 1 2 3 4 5 6 7 8 9 10 21 22 23 24 25 26 27 28 29 30 11 12 13 14 15 16 17 0% 18 19 0% 6 0% 18 0% 5 0% 10 2 5 10 18 6 2 1. 2. 3. 4. 5. 20 Which of the following combinations of quantum numbers is not allowed? ½ –½ ½ –½ ½ 1 2 3 4 5 6 7 8 9 10 21 22 23 24 25 26 27 28 29 30 11 12 13 14 15 0% 16 0% 17 0% 18 0% 0% 19 0% 42 0½ 0 0 –1 –2 0 43 –2 –½ 1 0 1 3 2 21 –1 ½ 1 3 2 4 4 30 0– ½ s 11 0½ m s l nlm 1. 2. 3. 4. 5. n 20 1. 2. 3. 4. 5. The electron configuration of indium is 1s22s22p63s23p64s23d104p65s24d105p15d10 1s22s22p63s23p64s23d104d104p1 1s23s22p63s23p64s24d104p65s25d105p1 1s22s22p63s23p64s23d104p65s24d105p1 none of these 1 2 3 4 5 6 7 8 9 10 21 22 23 24 25 26 27 28 29 30 11 12 13 14 15 0% 0% 0% 0% 1s 22 s2 2p 63 s2 1s 3p 22 64 s2 s2 2p 3d 63 1. s2 1s .. 3p 23 64 s2 s2 2p 3d 63 1. s2 1s .. 3p 22 6 s2 4s 2p 24 63 d1 s2 ... 3p 64 s2 3d 1. .. no ne of th es e 0% 16 17 18 19 20 Ag has __ electrons in its d orbitals. 1 2 3 4 5 6 7 8 9 10 21 22 23 24 25 26 27 28 29 30 11 12 13 14 15 16 17 18 19 20 Periodicity Determination of Atomic Radius: Half of the distance between nuclei in covalently bonded diatomic molecule "covalent atomic radii" Periodic Trends in Atomic Radius Radius decreases across a period Increased effective nuclear charge due to decreased shielding Radius increases down a group Addition of principal quantum levels Table of Atomic Radii Ionization Energy - the energy required to remove an electron from an atom Increases for successive electrons taken from the same atom Tends to increase across a period Electrons in the same quantum level do not shield as effectively as electrons in inner levels Irregularities at half filled and filled sublevels due to extra repulsion of electrons paired in orbitals, making them easier to remove Tends to decrease down a group Outer electrons are farther from the nucleus Ionization of Magnesium Mg + 738 kJ Mg+ + eMg+ + 1451 kJ Mg2+ + eMg2+ + 7733 kJ Mg3+ + e- Table of 1st Ionization Energies Another Way to Look at Ionization Energy Yet Another Way to Look at Ionization Energ Electron Affinity - the energy change associated with the addition of an electron Affinity tends to increase across a period Affinity tends to decrease as you go down in a period Electrons farther from the nucleus experience less nuclear attraction Some irregularities due to repulsive forces in the relatively small p orbitals Table of Electron Affinities Summary of Periodic Trends