Atoms and Atomic Theory Chapter 2
2.1 Early Chemical Discoveries and the Atomic Theory
Law of Conservation of Mass The total mass of substances present after a chemical reaction is the same as the total mass of substances before the reaction. In short, Mass of Reactants = Mass of Products.
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5 Law of Constant Composition Joseph Proust (1754–1826) Also known as the law of definite proportions. All samples of compounds have the same composition- the same proportions by mass of the constituent elements.
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Dalton’s Atomic Theory 1.Each element is composed of small particles called atoms. Atoms are neither created nor destroyed in chemical reactions. 2.All atoms of a given element are identical and differ from all other elements 3.Compounds are formed when atoms of more than one element combine in simple numerical ratios.
8 When two elements form more than one compound (if the mass of one element is kept constant) the ratios of the masses of the other element are in small whole numbers
In forming carbon monoxide, 1.0 g of carbon combines with 1.33 g of oxygen. In forming carbon dioxide, 1.0 g of carbon combines with 2.66 g of oxygen. If two elements form more than a single compound, the masses of one element combined with a fixed mass of the second are in the ratio of small whole numbers. Law of Multiple Proportions John Dalton 1803
10 Multiple Proportions Example Consider water (H 2 O) and hydrogen peroxide (H 2 O 2 ) The mass ratio for water is 2:16 or simplified 1:8 The mass ratio for hydrogen peroxide is 2:32 or 1:16 The mass ratio of oxygen in both compounds is 2:1
2.2 Electrons and Other Discoveries in Atomic Physics Review +/- Attract -/- and +/+ Repel All matter is made of charged particles + > - Positive + < - Negative + = - Neutral
Cathode ray tube
14 The Electron Streams of negatively charged particles were found to come from cathode tubes. J. J. Thomson is credited with their discovery (1897). Thomson measured the charge/mass ratio of the electron to be 1.76 10 8 coulombs/g.
15 The Atom, circa 1900: “Plum pudding” model, put forward by Thompson. Positive sphere of matter with negative electrons imbedded in it.
16 Millikan Oil Drop Experiment Robert Millikan showed ionized oil drops can be balanced against the pull of gravity by an electric field. 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.
17 Millikan Oil Drop Experiment Robert Millikan determined the charge on the electron in He determined the charge of an electron to be 1.60 X C
Millikin’s experiment He then used Thomsons charge-to- mass ratio to calculate the mass of an electron 9.10 x g 18
19 Radioactivity: The spontaneous emission of radiation by an atom. First observed in 1896 by Henri Becquerel while he was studying the properties of a uranium compound. Marie and Pierre Curie further experimented and discovered radium and poloniium All three shared the Nobel prize in 1903.
20 Radioactivity Two types of radiation were discovered by Ernest Rutherford: – particles – particles – rays (Paul Villard)
2.3 The Nuclear Atom
22 Discovery of the Nucleus Ernest Rutherford shot particles at a thin sheet of gold foil and observed the pattern of scatter of the particles.
23 The Nuclear Atom Since some particles were deflected at large angles, Thompson’s model could not be correct.
24 The Nuclear Atom Rutherford postulated a very small, positive, dense nucleus with the electrons around the outside of the atom. Most of the volume of the atom is empty space.
The nuclear atom – illustrated by the helium atom Rutherford protons 1919 James Chadwick neutrons 1932
26 Subatomic Particles Protons and electrons are the only particles that have a charge. Protons and neutrons have essentially the same mass. The mass of an electron is so small we ignore it.
2.4 Chemical Elements Atomic Number: Number of protons (Z) – Also equal to number of electrons Mass Number: Total protons and neutrons (A) Number of Neutrons = A-Z Atomic Mass Unit (u): 1/12 the mass of a carbon-12 atom – Mass of a proton/neutron is ~1u – Mass of an electron is ~ 1/2000u
2-4 Chemical Elements To represent a particular atom we use symbolism: A= mass numberZ = atomic number E= element
29 Symbols of Elements Elements are symbolized by one or two letters.
30 Atomic Number All atoms of the same element have the same number of protons: The atomic number (Z)
31 Atomic Mass The mass of an atom in atomic mass units (amu) is the total number of protons and neutrons in the atom.
Isotopic Notation Ne Ne Ne Isotopes: Same atomic number but different mass number Percent Natural Abundance: On Earth, 90.51% Ne % Ne % Ne-22
33 How many protons, neutrons, and electrons are in (a)an atom of 197 Au (b) an atom of strontium-90?
Ions Atoms that have gained or lost electrons and carry a charge
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2.5 Atomic Mass Why is carbon’s mass in the PT not ? – Must take into account all naturally occurring isotopes. – Find average atomic mass Atomic Mass = Fractional Abundance of Isotope 1 X Mass of Isotope 1 ( ) + Fractional Abundance of Isotope 2 X Mass of Isotope 2 ( ) +
37 Naturally occurring chlorine is 75.78% 35 Cl, which has an atomic mass of amu, and 24.22% 37 Cl, which has an atomic mass of amu. Calculate the average atomic mass (that is, the atomic weight) of chlorine. Average atomic mass = (0.7578)( amu) + (0.2422)( amu) = amu amu = amu
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2.6 Introduction to the Periodic Table
Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 2Slide 41 of 27 The Periodic table Alkali MetalsAlkaline EarthsTransition MetalsHalogensNoble Gases Lanthanides and Actinides Main Group
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2.7 The Concept of the Mole and the Avogadro Constant Mole: the amount of a substance that contains the same number of entities (atom or molecules) as there are in exactly 12g of carbon-12 N A = x units/mole
Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 2Slide 44 of 27 Molar Mass The molar mass, M, is the mass of one mole of a substance. (Get from Atomic Mass on PT) Lithium average atomic mass: 6.941u Lithium molar mass: grams/mole
2.8 Using the Mole Concept in Calculations Conversion factor 1 mole S = x S atoms 1 mole S = grams S
Copyright 2011 Pearson Canada Inc
Copyright 2011 Pearson Canada Inc
Copyright 2011 Pearson Canada Inc