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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! 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 This slide contains classified material and cannot be shown to high school students. Please continue as if everything is normal. 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 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 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: Long Wavelength = Low Frequency = Low ENERGY Short Wavelength = High Frequency = High ENERGY Wavelength Table Toupee? The Wave-like Electron The electron propagates through space on an energy wave. To understand the atom, one must understand the behavior of electromagnetic waves. Louis deBroglie The Great Niels Bohr (1885 - 1962) 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 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! 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) 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 Pauli Exclusion Principle No two electrons in an atom can have the same four quantum numbers. Wolfgang Pauli 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. 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? 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 Orbital filling table 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 ____ 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 energy is emitted 2. energy is absorbed 3. no change in energy occurs 4. light is emitted 5. none of these 1. 0% it. er en 1 2 3 4 5 6 7 8 9 10 11 12 21 22 23 24 25 26 27 28 29 30 31 32 13 14 gy is .. em er en 15 gy is 0% .. o. s ab no 16 ch ge an 17 0% in 0% .. e. i ht g li 18 se itt 0% ... m ne no 19 o h ft e es 20 2 3 4 5 6 7 8 9 10 11 12 21 22 23 24 25 26 27 28 29 30 31 32 13 14 ra d ia ... av es w 17 0% re d in fra 16 io 15 0% xra ys 0% s 0% ra d icr m 1 ga m m a 5. 0% av e 4. ow 3. s 2. gamma rays microwaves radio waves infrared radiation x-rays ra y 1. 18 19 20 1. 2 2. 5 3. 10 4. 18 1 2 3 4 5 6 7 8 9 10 11 12 21 22 23 24 25 26 27 28 29 30 31 32 13 14 15 16 17 0% 18 19 0% 6 0% 18 0% 5 2 0% 10 5. 6 20 1 2 3 4 5 6 7 8 9 10 11 12 21 22 23 24 25 26 27 28 29 30 31 32 13 14 15 0% 16 0% 17 0% 18 0% 0% 19 0% 42 0½ ½ –½ ½ –½ ½ 43 –2 –½ 0 0 –1 –2 0 21 –1 ½ 1 0 1 3 2 30 0– ½ s 11 0½ m s 1 2. 3 3. 2 4. 4 5. 4 1. l nlm n 20 1s22s22p63s23p64s23d104p65s24d105p15d10 2. 1s22s22p63s23p64s23d104d104p1 3. 1s23s22p63s23p64s24d104p65s25d105p1 4. 1s22s22p63s23p64s23d104p65s24d105p1 5. none of these 1. 1 2 3 4 5 6 7 8 9 10 11 12 21 22 23 24 25 26 27 28 29 30 31 32 13 14 15 0% 0% 0% 0% 1s 22 s2 2p 63 s2 3p 1s .. . 22 s2 2p 63 s2 3p 1s .. . 23 s2 2p 63 s2 3p 1s .. . 22 s2 2p 63 s2 3p .. . no ne of th es e 0% 16 17 18 19 20 1 2 3 4 5 6 7 8 9 10 11 12 21 22 23 24 25 26 27 28 29 30 31 32 13 14 15 16 17 18 19 20 Periodicity } Atomic Radius Radius = half the distance between two nuclei of a diatomic molecule. Influenced by three factors. Energy Level › Higher energy level is further away. Charge on nucleus › More charge pulls electrons in closer. Shielding › Layers of electrons shield from nuclear pull. The electron on the outside energy level has to look through all the other energy levels to see the nucleus The electron on the outside energy level has to look through all the other energy levels to see the nucleus. A second electron has the same shielding. As we go down a group Each atom has another energy level, So the atoms get bigger. H Li Na K Rb As you go across a period the radius gets smaller. Same energy level. More nuclear charge. Outermost electrons are closer. Na Mg Al Si P S Cl Ar Table of Atomic Radii Cations form by losing electrons. Cations are smaller that the atom they come from. Metals form cations. Cations of representative elements have noble gas configuration. Anions form by gaining electrons. Anions are bigger that the atom they come from. Nonmetals form anions. Anions of representative elements have noble gas configuration. Rb K Atomic Radius (nm) Na Li Kr Ar H Ne 10 Atomic Number The amount of energy required to completely remove an electron from a gaseous atom. Removing one electron makes a +1 ion. The energy required is called the first ionization energy. The second ionization energy is the energy required to remove the second electron. Always greater than first IE. The third IE is the energy required to remove a third electron. Greater than 1st of 2nd IE. Symbol First H He Li Be B C N O F Ne 1312 2731 520 900 800 1086 1402 1314 1681 2080 Second Third 5247 7297 1757 2430 2352 2857 3391 3375 3963 11810 14840 3569 4619 4577 5301 6045 6276 The greater the nuclear charge the greater IE. Distance from nucleus increases IE Filled and half filled orbitals have lower energy, so achieving them is easier, lower IE. Shielding As you go down a group first IE decreases because The electron is further away. More shielding. All the atoms in the same period have the same energy level. Same shielding. Increasing nuclear charge So IE generally increases from left to right. Exceptions at full and 1/2 fill orbitals. He He has a greater IE than H. same shielding greater nuclear charge First Ionization energy H Atomic number He l l l First Ionization energy l Li has lower IE than H more shielding further away outweighs greater nuclear charge H Li Atomic number He l l First Ionization energy l H Be has higher IE than Li same shielding greater nuclear charge Be Li Atomic number He l l l First Ionization energy l H B has lower IE than Be same shielding greater nuclear charge By removing an electron we make s orbital half filled Be Li B Atomic number First Ionization energy He H Be Li C B Atomic number He First Ionization energy N H C Be Li B Atomic number He First Ionization energy N H C O Be Li Breaks the pattern because removing an electron gets to 1/2 filled p orbital B Atomic number He First Ionization energy N F H C O Be Li B Atomic number Ne He First Ionization energy N F H C O Be Li Ne has a lower IE than He Both are full, Ne has more shielding Greater distance B Atomic number Ne He N F l First Ionization energy l H C O Be Li B l l Na Atomic number Na has a lower IE than Li Both are s1 Na has more shielding Greater distance Atomic number First Ionization energy Full Energy Levels are very low energy. Noble Gases have full orbitals. Atoms behave in ways to achieve noble gas configuration. 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 The tendency for an atom to attract electrons to itself when it is chemically combined with another element. How fair it shares. Big electronegativity means it pulls the electron toward it. Atoms with large negative electron affinity have larger electronegativity. The further down a group the farther the electron is away and the more electrons an atom has. More willing to share. Low electronegativity. Metals are at the left end. They let their electrons go easily Low electronegativity At the right end are the nonmetals. They want more electrons. Try to take them away. High electronegativity. Ionization energy, electronegativity Electron affinity INCREASE Atomic size increases, shielding constant Ionic size increases Another Way to Look at Ionization Energy Yet Another Way to Look at Ionization Energ Summary of Periodic Trends