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Chapter-3 Atoms: The Building Blocks of Matter I. Early Theories of Matter A. Greek Philosophers 1. Democritus and Leucippus I. Early Theories of Matter “Atomos” indivisible Democritus, 400 BC “Everything is made up of a few simple parts called atomos.” Atomos means “uncuttable or indivisible” in Greek He envisioned atomos as small, solid particles of many different sizes and shapes *His ideas were rejected later by Aristotle I. Early Theories of Matter A. Greek Philosophers 1. Democritus and Leucippus 2. Aristotle I. Early Theories of Matter Matter IS infinitely divisible. He was so influential and this is what people believed for the next 2000 years! He believed that all matter was “continuous” Why? Aristotle invented logic: Terms: "A", "B", "rabbit", "Peter" Propositions: "A" is "B" or "A" is not "B“ "Peter is a Rabbit" Syllogisms: "A is B. B is C. Therefore, A is C." Two kinds of Syllogisms: "All rabbits are brown. Peter is a rabbit. Therefore, Peter is brown. (Induction) "Flopsy, Mopsy, Cottontail and Peter are brown. Flopsy, Mopsy, Cottontail and Peter are rabbits. Therefore, all rabbits are brown." (Deduction – what scientists do) B. Alchemists Contributions of Alchemists • Invention of Lab Equipment • Important Chemical Procedures • Discovery of Elements • Descriptive Chemistry C. Lavoisier & Proust Lavoisier – Law of Conservation of Mass C. Lavoisier & Proust 2. Proust – Law of Definite Proportions Law of Definite Proportions Compounds always use specific proportions of their elements, regardless of what is available. Water will always contain: 11.2% H & 88.8% O by mass Hydrogen peroxide contains: 5.9% H & 94.1% O by mass. Different ratios mean different properties! Percent Mass part × 100 % Mass = whole mass of element × 100 % Mass = mass of compound Practice Problem Two samples of compounds made of only H and O have the following compositions: Compound A: 15.0 g Hydrogen 120.0 g Oxygen Compound B: Are they same ? 2.0 g Hydrogen 34.0 g Oxygen 1 × 16.00 = 1 × 12.00 2 × 16.00 = 1 × 12.00 Atomic Theory – John Dalton 1808 3. Dalton – Law of Multiple Proportions Law of Multiple Proportions When two different compounds are made of the same set of elements, different masses of one element combine with the same mass of the other in small, whole number ratios. HUH? Example: Carbon dioxide CO2 has exactly twice as much oxygen as carbon monoxide CO. Law of Multiple Proportions Example Law of Multiple Proportions Dalton’s Atomic Theory 1) Matter is made of tiny particles called atoms 2) All atoms of an element are: identical in size, mass and other properties different from those of the other elements 3) Atoms cannot be subdivided, created or destroyed 4) Atoms of one element can combine with atoms of other elements to form compounds. A given compound always has the same relative number and types of atoms 5) In chemical reactions, atoms are combined, separated or rearranged Modern Theory: • Atoms of an element have a characteristic average mass which is unique to that element, but elements CAN have different isotopes! • Atoms cannot be subdivided, created, or destroyed in ordinary chemical reactions. But, changes CAN occur in nuclear reactions Dalton actually did have evidence from experiments: Mass is conserved in chemical reactions Elements combine to make compounds in predictable mass ratios: Law of Definite Proportions Law of Multiple Proportions Summary of Law’s of 1790’s The transformation of a substance(s) into one or more new substances is called “chemical reaction” Law of conservation of mass: mass is neither created nor destroyed during ordinary chemical reactions or physical changes. Law of definite proportions: A pure compound always contains the same elements in the same proportions by mass. Law of multiple proportions: When elements combine to form more than a single compound, the ratios of the masses of the combining elements can be expressed by a ratio of small whole numbers. Dalton’s Model/Chart Today we can even see atoms! A penny contains 29,000,000,000,000,000,000,000 atoms III. Dalton’s model vs. Modern Model How has it changed? IV. Finding Subatomic Particles A. Radioactivity - the phenomenon of rays being produced spontaneously by unstable atomic nuclei. B. X-rays & radioactivity 1. Wilhelm Roentgen - 1895 a. shot cathode rays at metal and radiation was given off that would darken photographic film. b. the rays could pass through flesh, but not bone to expose photo plate. c. rays were unknown, so they were given the name "x-rays“. 2. Becquerel - obtained a sample of uranium ore (pitchblende) to study phosphorescence a. He noticed uranium also darken photographic paper b. Radioactive Decay c. Transmutation - change of one element into another 3. Marie Curie- Noticed polonium and radium also gave off radiation: a. radiation of radium is two million times greater than uranium. b. it conducted electricity in air "ionized air“. c. ionizing radiation has sufficient energy to change atom and molecules into ions. d. non-ionizing radiation does not, such as radio waves. Non-ionizing vs. Ionizing Radiation C. Discovery of Electrons 1. William Crookes (1879) discovers cathode rays (streams of electrons) 2. JJ Thomson (1897) – Cathode Ray Tube Experiment Cathode ray tube (CRT) to deduce the presence of a negatively charged particles. Discovery of electrons using CRT Electron’s charge-to-mass ratio The beam of negative particles bends downward The magnet is turned around and the beam bends in opposite direction atoms were not indivisible particles 2. JJ Thomson (1897) – Plum Pudding Model An electron is much smaller than a hydrogen atom? Dalton was wrong: atoms can be divided! 2. JJ Thomson (1897) All elements should contain electrons All atoms are electrically neutral (so there must be positive particles present) Mass of electrons are so small other particles? CRT experiment Electrons discovered Plum-pudding model Mass/charge ratio Mass of electron (1/2000th of a H atom) 3. Robert A. Millikan (1909) – Oil Drop Experiment Oil drop experiment findings: Measured the charge of an electron (1.602× 10-19 C) Calculated the mass of an electron (9.109 × 10-31 kg) 4. Rutherford, Marsden & Geiger (1911) 4. Ernest Rutherford – Gold foil Experiment Conclusions– Nucleus is: small dense positively charged atoms are mostly empty space 4. Ernest Rutherford – Gold foil Experiment --A question arose: Why don’t the electrons fall into the nucleus ? 4. Ernest Rutherford – Other Findings Characterized radioactivity as 3 types of radiation: Alpha rays (): positively charged particles Beta rays (): negatively charged particles Gamma rays (): no charge Nuclear Symbols: Neutron Proton Electron Alpha Beta Gamma Positron High energy Lower energy High energy Beta particles and gamma rays are usually emitted together! 5. Eugen Goldstein’s Canal Ray Experiment (1886) Goldstein used a perforated anode in the cathode ray tube and found that there were these positively charged rays (canal rays or anode rays), which consisted on positively charged particles called as protons. 6. James Chadwick – Discovery of Neutron (1932) He discovered that most nuclei also contain another neutral particle: neutron. It is slightly more massive than proton but has no charge. D. Bohr’s Model (to be continued…) -Niels Bohr proposed a model in which orbiting electrons don't lose energy. A change to the Laws of Physics! Five Major Atomic Models 1. Dalton’s Atomic Model 1808 2. Plum Pudding Model 1909 3. Rutherford’s Nuclear Model 1910 4. Bohr’s Solar System Model 1913 5. Shrodinger’s Wave Mechanical Model 1927 E. Subatomic Particles of Atom Nuclear particles are held together by Nuclear Forces! Location e- cloud Nucleus Nucleus Nucleons: particles in the nucleus - Proton (p+) neutron (n0) E. Subatomic Particles of Atom Thomson electrons Chadwick neutrons Rutherford protons E. Subatomic Particles of Atom Of course – scientists have found even more particles – here’s a brief introduction! Did you know that if a nucleus were the size of the dot in the exclamation point at the end of this sentence, its mass would be approximately as much as that of 70 automobiles! The Modern View of the Atomic Structure Atoms: Consist of tiny, dense, positively charged nucleus surrounded by cloud of (-) electrons Nucleus: Contains positively charged protons and neutral neutrons Atomic number(Z): number of protons (determines the element) Mass number: sum of protons and neutrons (determines the isotope) Isotopes: atoms of an element that differ in the number of neutrons Atomic Mass: Average mass of all isotopes based on the abundance (expressed in amu) 23 sodiu 23 Na 11 Atomic Vocabulary Mass number = p+ + n0 (in the nucleus) Element symbol A E Z Atomic number = # p + (= # e– ) Elements are put in order of increasing atomic number on the periodic table, identifies an element. Atomic Vocabulary Carbon 6 C 12.011 name Atomic number Element symbol Average atomic mass (in amu) F. Isotopes Atoms of one element that have different mass numbers (different # of neutrons) There are 3 naturally occurring isotopes of hydrogen: Isotope Name #p Protium #n % of all H 1 0 99.985 Deuterium 1 1 0.015 Tritium 1 2 < 0.001 Periodic Table by Number of Stable Isotopes Neutral Isotopes Arsenic 18 75 Phosphorous 8 8 33 16 15 18 75 31 Example: An atom of carbon with 7 neutrons: 13C 6 An atom of lead with 125 neutrons: 207Pb 82 Relative Atomic Masses It is convenient to use relative atomic mass because the atoms are too small. (Oxygen atom weighs 2.657 ×10-23g). The atomic mass of any nuclide is determined by comparing it with carbon-12 atom. One atomic mass unit (amu) is exactly 1/12 of the mass of one C-12 atom. (1.660540 ×10-27 kg) Average Atomic Masses of Elements Different isotopes will have different masses. Most elements consist of mixtures of isotopes. Average Atomic Mass is the weighted average of the element’s average atomic masses of the naturally occurring isotopes of an element There are three naturally occurring isotopes of uranium: 234U 235U 238U 92 92 92 How to Calculate Average Atomic Masses? Mass # 39 19 K 40 19 K 41 19 K Atomic # Abundance of isotopes? 93.2581 % 0.0117 % 6.7302 % Atomic Masses of isotopes? 38.963707 amu 39.963998 amu 40.961826 amu How to Calculate Average Atomic Masses? % Symbol abundance 39 K amu decimal product 93.2581 38.963707 × 0.932581 = 36.3368 0.0117 39.963998 × 0.000117 = 0.00468 6.7302 40.961826 × 0.067302 = 2.7568 19 40 K 19 41 K 19 + 39.0983 Avogadro’s Number It is defined as the number of particles in one mole of any pure substance. This number is found to be 6.022×1023. The MOLE One mole of something contains 6.022×1023 particles of that substance. Example: One mole of Hydrogen atoms contains exactly 6.022×1023 atoms of Hydrogen. Example: 1 mole of H2O has 6.022×1023 water molecules. Example: 1 mole of NaCl has 6.022×1023 formula units of NaCl. Molar Mass Mass of one mole of a substance. By definition, a molar mass is the mass of 1 mol of a substance (i.e., g/mol). Molar mass can be calculated by expressing the mass in amu or in grams. Molar mass of C = 12 g/mol Mole Relationships One mole of atoms, ions, or molecules contains Avogadro’s number of those particles. One mole of molecules or formula units contains Avogadro’s number times the number of atoms or ions of each element in the compound. Mole is a Bridge to Mass, Particles, Volume Moles provide a bridge from the molecular scale to the real-world scale. Moleville Molar Mass Railroad Mass Junction Avagodro’s Number Particles Conversions