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Download Lecture 1 – Matter, Atomic Structure
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Chemistry Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Chemistry: A Science for the 21st Century • Health and Medicine • Sanitation systems • Surgery with anesthesia • Vaccines and antibiotics • Gene therapy •Energy and the Environment • Fossil fuels • Solar energy • Nuclear energy 2 Chemistry: A Science for the 21st Century • Materials and Technology • Polymers, ceramics, liquid crystals • Room-temperature superconductors? • Molecular computing? • Food and Agriculture • Genetically modified crops • “Natural” pesticides • Specialized fertilizers 3 The Study of Chemistry Macroscopic Microscopic 4 The scientific method is a systematic approach to research. A hypothesis is a tentative explanation for a set of observations. tested modified 5 A law is a concise statement of a relationship between phenomena that is always the same under the same conditions. Force = mass x acceleration A theory is a unifying principle that explains a body of facts and/or those laws that are based on them. Atomic Theory 6 Chemistry is the study of matter and the changes it undergoes. Matter is anything that occupies space and has mass. A substance is a form of matter that has a definite composition and distinct properties. liquid nitrogen gold ingots silicon crystals 7 A mixture is a combination of two or more substances in which the substances retain their distinct identities. 1. Homogenous mixture – composition of the mixture is the same throughout soft drink, milk, solder 2. Heterogeneous mixture – composition is not uniform throughout cement, iron filings in sand 8 Physical means can be used to separate a mixture into its pure components. magnet distillation 9 An element is a substance that cannot be separated into simpler substances by chemical means. •114 elements have been identified • 82 elements occur naturally on Earth gold, aluminum, lead, oxygen, carbon, sulfur • 32 elements have been created by scientists technetium, americium, seaborgium 10 Replace with Table 1.1 from 7e page 6 11 A compound is a substance composed of atoms of two or more elements chemically united in fixed proportions. Compounds can only be separated into their pure components (elements) by chemical means. lithium fluoride quartz dry ice – carbon dioxide 12 Classification of Matter Replace with Table 1.5 from 7e page 7 13 A Comparison: The Three States of Matter 14 The Three States of Matter: Effect of a Hot Poker on a Block of Ice 15 Types of Changes A physical change does not alter the composition or identity of a substance. sugar dissolving ice melting in water A chemical change alters the composition or identity of the substance(s) involved. hydrogen burns in air to form water 16 Extensive and Intensive Properties An extensive property of a material depends upon how much matter is being considered. • mass • length • volume An intensive property of a material does not depend upon how much matter is being considered. • density • temperature • color 17 Matter - anything that occupies space and has mass mass – measure of the quantity of matter SI unit of mass is the kilogram (kg) 1 kg = 1000 g = 1 x 103 g weight – force that gravity exerts on an object weight = c x mass A 1 kg bar will weigh on earth, c = 1.0 1 kg on earth on moon, c ~ 0.1 0.1 kg on moon 18 Atoms, Molecules, and Ions Dalton’s Atomic Theory (1808) 1. Elements are 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. Compounds are composed of atoms of more than one element. 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. 4. A chemical reaction involves only the separation, combination, or rearrangement of atoms; it does not result in their creation or destruction. 20 Dalton’s Atomic Theory Law of Multiple Proportions 21 16 X + 8Y 8 X2Y Law of Conservation of Mass 22 Cathode Ray Tube J.J. Thomson, measured mass/charge of e(1906 Nobel Prize in Physics) 23 Cathode Ray Tube 24 Millikan’s Experiment Measured mass of e(1923 Nobel Prize in Physics) e- charge = -1.60 x 10-19 C Thomson’s charge/mass of e- = -1.76 x 108 C/g e- mass = 9.10 x 10-28 g 25 Types of Radioactivity (uranium compound) 26 Thomson’s Model 27 Rutherford and the Nuclear Atom • Ernest Rutherford directed Hans Geiger and Ernst Marsden’s experiment in 1910. – - particle scattering from thin Au foils – Gave us the basic picture of the atom’s structure. 28 Rutherford’s Experiment (1908 Nobel Prize in Chemistry) particle velocity ~ 1.4 x 107 m/s (~5% speed of light) 1. atoms positive charge is concentrated in the nucleus 2. proton (p) has opposite (+) charge of electron (-) 3. mass of p is 1840 x mass of e- (1.67 x 10-24 g) 29 Rutherford and the Nuclear Atom • In 1912 Rutherford decoded the -particle scattering information. – Explanation involved a nuclear atom with electrons surrounding the nucleus . 30 Rutherford and the Nuclear Atom Rutherford’s major conclusions from the -particle scattering experiment 1. The atom is mostly empty space. 2. It contains a very small, dense center called the nucleus. 3. Nearly all of the atom’s mass is in the nucleus. 4. The nuclear diameter is 1/10,000 to 1/100,000 times less than atom’s radius. 31 Rutherford and the Nuclear Atom • Because the atom’s mass is contained in such a small volume: – The nuclear density is ~1015g/mL. – This is equivalent to ~3.72 x 109 tons/in3. – Density inside the nucleus is almost the same as a neutron star’s density. 32 Rutherford’s Model of the Atom atomic radius ~ 100 pm = 1 x 10-10 m nuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m “If the atom is the Houston Astrodome, then the nucleus is a marble on the 50-yard line.” 33 Chadwick’s Experiment (1932) (1935 Noble Prize in Physics) H atoms: 1 p; He atoms: 2 p mass He/mass H should = 2 measured mass He/mass H = 4 + 9Be 1n + 12C + energy neutron (n) is neutral (charge = 0) n mass ~ p mass = 1.67 x 10-24 g 34 mass p ≈ mass n ≈ 1840 x mass e35 Atomic Number, Mass Number, and Isotopes Atomic number (Z) = number of protons in nucleus Mass number (A) = number of protons + number of neutrons = atomic number (Z) + number of neutrons Isotopes are atoms of the same element (X) with different numbers of neutrons in their nuclei Mass Number A ZX Atomic Number 1 1H 235 92 2 1H U Element Symbol (D) 238 92 3 1H U (T) 36 The Isotopes of Hydrogen 37 A molecule is an aggregate of two or more atoms in a definite arrangement held together by chemical forces. H2 H2O NH3 CH4 A diatomic molecule contains only two atoms: H2, N2, O2, Br2, HCl, CO diatomic elements A polyatomic molecule contains more than two atoms: O3, H2O, NH3, CH4 38 An ion is an atom, or group of atoms, that has a net positive or negative charge. cation – ion with a positive charge If a neutral atom loses one or more electrons it becomes a cation. Na 11 protons 11 electrons Na+ 11 protons 10 electrons anion – ion with a negative charge If a neutral atom gains one or more electrons it becomes an anion. Cl 17 protons 17 electrons Cl- 17 protons 18 electrons 39 A monatomic ion contains only one atom: Na+, Cl-, Ca2+, O2-, Al3+, N3- A polyatomic ion contains more than one atom: OH-, CN-, NH4+, NO3- 40 2.1 Give the number of protons, neutrons, and electrons in each of the following species: (a) (b) (c) (d) carbon-14 2.1 Strategy Recall that the superscript denotes the mass number (A) and the subscript denotes the atomic number (Z). Mass number is always greater than atomic number. (The only exception is H, where the mass number is equal to the atomic number.) In a case where no subscript is shown, as in parts (c) and (d), the atomic number can be deduced from the element symbol or name. To determine the number of electrons, remember that because atoms are electrically neutral, the number of electrons is equal to the number of protons. 2.1 Solution (a) The atomic number is 11, so there are 11 protons. The mass number is 20, so the number of neutrons is 20 − 11 = 9. The number of electrons is the same as the number of protons; that is, 11. (b) The atomic number is the same as that in (a), or 11. The mass number is 22, so the number of neutrons is 22 − 11 = 11. The number of electrons is 11. Note that the species in (a) and (b) are chemically similar isotopes of sodium. 2.1 (c) The atomic number of O (oxygen) is 8, so there are 8 protons. The mass number is 17, so there are 17 − 8 = 9 neutrons. There are 8 electrons. (d) Carbon-14 can also be represented as 14C. The atomic number of carbon is 6, so there are 14 − 6 = 8 neutrons. The number of electrons is 6.