Evolution of Atomic Theory PDF

Evolution of Atomic Theory 1  Evolution of Atomic Theory Dalton (1803) Thomson (1904) ( + & - charges ) Rutherford (1911) (the nucleus) Bohr (1913) (energy levels) Schrödinger (19...

Evolution of Atomic Theory 1  Evolution of Atomic Theory Dalton (1803) Thomson (1904) ( + & - charges ) Rutherford (1911) (the nucleus) Bohr (1913) (energy levels) Schrödinger (1926) (electron-cloud model) 2 Before Dalton… Law of conservation of mass (Antoine Lavoisier,1789) There is no detectable change in mass during an ordinary chemical reaction. Matter can be neither created nor destroyed. 3 Before Dalton… Law of conservation of mass (Antoine Lavoisier,1789) There is no detectable change in mass during an ordinary chemical reaction. Matter can be neither created nor destroyed. 16 X + 8Y 8 X2Y 4 Before Dalton… Law of definite propotions (Joseph Proust, 1799) Different samples of the same compound always contain its constituent elements in the same proportion by mass (also called Law of Constant Composition). 5 Dalton’s Atomic Theory (1808) 1. Each element is composed of extremely small particles called atoms. 2. All atoms of a given element are identical, having the same size, mass and chemical properties. The atoms of one element are different from the atoms of all other elements. 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. (Law of conservation of mass) 4. Atoms of more than one element combine to form compounds. In any compound, the ratio of the numbers of atoms of any two of the elements present is either an integer or a simple fraction. (Law of definite/multiple proportions) 6 Law of Multiple Proportions If two elements, C and O, form more than one compound, the masses of O that combine with a given mass of C are in the ratio of small whole numbers. (CO 1:1, CO2 1:2) Dalton predicted this law and observed it while developing his atomic theory. When two or more compounds exist from the same elements, they can not have the same relative number of atoms. 7 Discovery of Subatomic Particles In Dalton’s view, the atom was the smallest particle possible. Many discoveries led to the fact that the atom itself was made up of smaller particles. Electrons and cathode rays Radioactivity Nucleus, protons, and neutrons 8 Forces between electrically charged objects repulsion attraction Effect of a magnetic field on charged particles: 9 The Electron (Cathode Rays) Streams of negatively charged particles were found to emanate from cathode tubes, causing fluorescence. 10 The Electron (Cathode Rays) J.J. Thomson, measured mass/charge of e- as 1.76 × 108 coulombs/gram (C/g). (1906 Nobel Prize in Physics) 11 The Electron (Cathode Rays) https://www.youtube.com/watch?v=GzMh4q-2HjM 12 The Electron (Millikan Oil-Drop Experiment) ▪ Once the charge/mass ratio of the electron was known, determination of either the charge or the mass of an electron would yield the other. ▪ Robert Millikan determined the charge on the electron in 1909. (1923 Nobel Prize in Physics) 13 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. 14 Radioactivity Radioactivity is the spontaneous emission of high-energy radiation by an atom. Its discovery showed that the atom had more subatomic particles and energy associated with it. Three types of radiation were discovered by Ernest Rutherford: – α particles (positively charged) – β particles (negatively charged, like electrons) – γ rays (uncharged) 15 Discovery of the Nucleus Ernest Rutherford shot α particles at a thin sheet of gold foil and observed the pattern of scatter of the particles. (1908 Nobel Prize in Chemistry) 16 Discovery of the Nucleus 17 Rutherford’s Model Since some particles were deflected at large angles, Thomson’s model could not be correct. Rutherford postulated a very small, dense nucleus with the electrons around the outside of the atom. Most of the volume is empty space. Atoms are very small; 1 – 5 Å or 100 – 500 pm. Other subatomic particles (protons and neutrons) were discovered. Mass of a proton=1.67x10-24 g The Neutron J. Chadwick (1932) Mass of a neutron=1.67x10-24 g 1 amu = 1.67 × 10–24 g 18 Subatomic Particles Protons (+1) and electrons (–1) have a charge; neutrons are neutral. Protons and neutrons have essentially the same mass (relative mass 1). The mass of an electron is so small we ignore it (relative mass 0). Protons and neutrons are found in the nucleus; electrons travel around the nucleus. 19 Atoms of an Element Elements are represented by a one or two letter symbol. This is the symbol for carbon. All atoms of the same element have the same number of protons, which is called the atomic number, Z. It is written as a subscript BEFORE the symbol. The mass number is the total number of protons and neutrons in the nucleus of an atom. It is written as a superscript BEFORE the symbol. 20 Isotopes & Ions Isotopes are atoms of the same element with different masses. Isotopes have different numbers of neutrons, but the same number of protons. Ions are atoms that have the same atomic and mass number but different number of electrons + charged ions: cations, - charged ions: anions 21 Atomic Weight Because in the real world we use large amounts of atoms and molecules, we use average masses in calculations. An average mass is found using all isotopes of an element weighted by their relative abundances. This is the element’s atomic weight. That is, Atomic Weight = Ʃ [(isotope mass) × (fractional natural abundance)] Note: the sum is for ALL isotopes of an element. 22 Bohr’s atomic model Bohr suggested that only certain stable electron orbits around the nucleus were allowed. The electron moves in a circular orbit at a fixed distances from the nucleus. 23 Wave Mechanical Model Niels Bohr: Exact atomic orbitals Louis de Broglie proposed that electrons act as both waves and particles. (Wave Particle Duality) Heisenberg uncertainty principle: It is impossible to know simultaneously both the momentum (mass times velocity) and the position of a particle with certainty. Thus, it is not appropriate to imagine the electron circling the nucleus in well- defined orbits. Erwin Schrodinger proposes electron probability through wave function equations. 24 Schrödinger Equation Where are you most likely to find e-? Instead of orbits, electrons reside in orbitals or 3D spaces around the nucleus Electron cloud around nucleus with fuzzy edges, no solid boundry. HΨ (R,r)= EΨ(R,r) H: Hamiltonian operator with Kinetic and Potential energy terms Ψ: Wavefunction with dynamical information about the system E: Energy of the system P,+ e,- 25  Atomic Orbital: Wavefunction of the electron Principle energy level (n) designates how close an e- is to the nucleus. Subshell the shape of the probability of e- location (orbital). – s-1 orbital – p-3 orbitals – d-5 orbitals – f- 7 orbitals 26 Atomic Orbitals Ψ2 , probability density: probability of finding an electron at a particular point in space. s orbital p orbital d orbital 27 Aufbau (Building-up) Principle Electrons fill orbitals from lowest to highest enery level. The order of occupation of orbitals: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 5d 4f 6p... 28 Aufbau (Building-up) Principle Pauli Exclusion Principle: An orbital can hold a maximum of two electrons spinning in opposite directions. 29 Aufbau (Building-up) Principle Hund’s rule: To minimize repulsion and maintain low energy, electrons fill orbitals singly. When all orbitals are occupied by at least one e-, then electrons will pair up. ☑ ☒ 30 Aufbau (Building-up) Principle Hund’s rule: In its ground state, an atom adopts a configuration with the greatest number of unpaired electrons. ☑ ☒ 31 Electronic Configurations Electron configuration is used to determine location of e-. Identifies number of electrons in ground state in each subshell at each energy level for that atom. 32 Electronic Configurations 33 Electronic Configurations 2. Shorthand notation for electron configuration Core electrons (inner electrons) vs valence electrons (outermost electrons) Closed-shell atom 34 Electron Configurations of Ions Cations: The electrons are lost from the highest energy level (n value). 3Li+ is 1s2 (losing a 2s electron). 26Fe2+ is 1s22s22p63s23p63d6 (losing two 4s electrons). Anions: The electron configurations are filled to ns2np6; e.g., 9F– is 1s22s22p6 (gaining one electron in 2p). 35 Periodic Table The periodic table is a systematic organization of the elements. 2. Elements are arranged in order of atomic number. 36 Reading the Periodic Table Boxes on the periodic table list the atomic number ABOVE the symbol. The atomic weight of an element is listed below the symbol on the periodic table. 37 Organization of the Periodic Table The rows on the periodic table are called periods. Columns are called groups. Elements in the same group have similar chemical properties. 38 Periodicity When one looks at the chemical properties of elements, one notices a repeating pattern of properties and reactivity. 39 Periodic Table 2. 40 Periodic Table 2. 41 Metals Metals tend to lose electrons to attain noble gas electronic configuration and tend to form cations. Some properties of metals include shiny luster. conducting heat and electricity. Malleable, ductile. Tend to react with nonmetals to form ionic substances. 42 Nonmetals Nonmetals tend to gain electrons to attain noble gas electronic configuration and tend to form anions. Some properties of nonmetals include Vary greatly in appearance. Generally poor conductors of heat and electricity. Brittle in the solid stage. Not lustrous Melting point lower than metals. 43 Metalloids Elements on the steplike line are metalloids. These are Boron (B), Silicon (Si), Germanium (Ge), As (Arsenic), Antimony (Sb), and Tellurium (Te). Their properties are sometimes like metals and sometimes like nonmetals. Silicon example: Looks like metal, but rather brittle and a poor conductor of heat and electricity. 44 Chemical Formulas The subscript to the right of the symbol of an element tells the number of atoms of that element in one molecule of the compound. Molecular compounds are composed of molecules and almost always contain only nonmetals. 45 Diatomic Molecules These seven elements occur naturally as molecules containing two atoms: – Hydrogen – Nitrogen – Oxygen – Fluorine – Chlorine – Bromine – Iodine 46 Types of Formulas Empirical formulas give the lowest whole-number ratio of atoms of each element in a compound. Molecular formulas give the exact number of atoms of each element in a compound. If we know the molecular formula of a compound, we can determine its empirical formula. The converse is not true without more information! 47 Picturing Molecules Structural formulas show the order in which atoms are attached. They do NOT depict the three- dimensional shape of molecules. Perspective drawings, ball-and-stick models, and space-filling models show the three-dimensional order of the atoms in a compound. 48 Ions When an atom of a group of atoms loses or gains electrons, it becomes an ion. Cations are formed when at least one electron is lost. Monatomic cations are formed by metals. Anions are formed when at least one electron is gained. Monatomic anions are formed by nonmetals, except the noble gases. 49 Common Cations 50 Common Anions 51 Polyatomic Ions Sometimes a group of atoms will gain or lose electrons. These are polyatomic ions. A polyatomic cation: – Ammonium ≡ NH4+ A polyatomic anion: – Sulfate ≡ SO42– 52 Ionic Compounds Ionic compounds (such as NaCl) are generally formed between metals and nonmetals. Electrons are transferred from the metal to the nonmetal. The oppositely charged ions attract each other. Only empirical formulas are written. 53 Writing Formulas Because compounds are electrically neutral, one can determine the formula of a compound this way: – The charge on the cation becomes the subscript on the anion. – The charge on the anion becomes the subscript on the cation. – If these subscripts are not in the lowest whole- number ratio, divide them by the greatest common factor. 54 Sample Exercise 2.1 Atomic Size The diameter of a U.S. dime is 17.9 mm, and the diameter of a silver atom is 2.88 Å. How many silver atoms could be arranged side by side across the diameter of a dime? 55 Sample Exercise 2.2 Determining the Number of Subatomic Particles in Atoms How many protons, neutrons, and electrons are in an atom of (a) 197Au, (b) strontium-90? 56 Sample Exercise 2.3 Writing Symbols for Atoms Magnesium has three isotopes with mass numbers 24, 25, and 26. (a) Write the complete chemical symbol (superscript and subscript) for each. (b) How many neutrons are in an atom of each isotope? 57 Sample Exercise 2.4 Calculating the Atomic Weight of an Element from Isotopic Abundances Naturally occurring chlorine is 75.78% 35Cl (atomic mass 34.969 amu) and 24.22% 37Cl (atomic mass 36.966 amu). Calculate the atomic weight of chlorine. 58 Sample Exercise 2.5 Using the Periodic Table Which two of these elements would you expect to show the greatest similarity in chemical and physical properties: B, Ca, F, He, Mg, P? Solution Elements in the same group of the periodic table are most likely to exhibit similar properties. We therefore expect Ca and Mg to be most alike because they are in the same group (2A, the alkaline earth metals). 59 Sample Exercise 2.6 Relating Empirical and Molecular Formulas Write the empirical formulas for (a) glucose, a substance also known as either blood sugar or dextrose—molecular formula C6H12O6; (b) nitrous oxide, a substance used as an anesthetic and commonly called laughing gas—molecular formula N2O. Sample Exercise 2.7 Writing Chemical Symbols for Ions Give the chemical symbol, including superscript indicating mass number, for (a) the ion with 22 protons, 26 neutrons, and 19 electrons; and (b) the ion of sulfur that has 16 neutrons and 18 electrons. 61 Sample Exercise 2.8 Predicting Ionic Charge Predict the charge expected for the most stable ion of barium and the most stable ion of oxygen. Solution We will assume that barium and oxygen form ions that have the same number of electrons as the nearest noble-gas atom. From the periodic table, we see that barium has atomic number 56. The nearest noble gas is xenon, atomic number 54. Barium can attain a stable arrangement of 54 electrons by losing two electrons, forming the Ba2+ cation. Oxygen has atomic number 8. The nearest noble gas is neon, atomic number 10. Oxygen can attain this stable electron arrangement by gaining two electrons, forming the O2− anion. 62 Sample Exercise 2.9 Identifying Ionic and Molecular Compounds Which of these compounds would you expect to be ionic: N2O, Na2O, CaCl2, SF4? Sample Exercise 2.10 Using Ionic Charge to Write Empirical Formulas for Ionic Compounds Write the empirical formula of the compound formed by (a) Al3+ and Cl– ions, (b) Al3+ and O2– ions, (c) Mg2+ and NO3– ions.  Chemical Nomenclature The system of naming compounds is called chemical nomenclature. We will learn how to name: 1) Ionic compounds 2) Acids 3) Binary Molecular Compounds 4) Simple Organic Compounds – Alkanes – Alcohols  Inorganic Nomenclature Write the name of the cation. If the cation can have more than one possible charge, write the charge as a Roman numeral in parentheses. If it is a polyatomic cation, it will end in -ium. If the anion is an element, change its ending to -ide; if the anion is a polyatomic ion, simply write the name of the polyatomic ion. Patterns in Oxyanion Nomenclature When there are two oxyanions involving the same element – the one with fewer oxygens ends in -ite. – the one with more oxygens ends in -ate. NO2− : nitrite; NO3− : nitrate SO32− : sulfite; SO42− : sulfate Patterns in Oxyanion Nomenclature Central atoms on the second row have a bond to, at most, three oxygens; those on the third row take up to four. Charges increase as you go from right to left. Patterns in Oxyanion Nomenclature The one with the second fewest oxygens ends in -ite: ClO2– is chlorite. The one with the second most oxygens ends in -ate: ClO3– is chlorate. The one with the fewest oxygens has the prefix hypo- and ends in -ite: ClO– is hypochlorite. The one with the most oxygens has the prefix per- and ends in -ate: ClO4– is perchlorate.  Acid Nomenclature If the anion in the acid ends in -ide, change the ending to -ic acid and add the prefix hydro-. – HCl: hydrochloric acid – HBr: hydrobromic acid – HI: hydroiodic acid If the anion ends in -ite, change the ending to -ous acid. – HClO: hypochlorous acid – HClO2: chlorous acid If the anion ends in -ate, change the ending to -ic acid. – HClO3: chloric acid – HClO4: perchloric acid  Nomenclature of Binary Molecular Compounds The name of the element farther to the left in the periodic table (closer to the metals) or lower in the same group is usually written first. A prefix is used to denote the number of atoms of each element in the compound (mono- is not used on the first element listed, however). Nomenclature of Binary Compounds The ending on the second element is changed to -ide. – CO2: carbon dioxide – CCl4: carbon tetrachloride If the prefix ends with a or o and the name of the element begins with a vowel, the two successive vowels are often elided into one. – N2O5: dinitrogen pentoxide – CO: carbon monoxide