Chapter 2: Atoms, Molecules and Ions. Dalton s Atomic Theory ( ) Postulates

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1 Chapter 2: Atoms, Molecules and Ions The topics in this chapter should be review from a previous course. It is expected that you are able to review and master this material quickly and somewhat independently. From this Chapter you should: understand and be able to discuss/describe the atomic theory of matter. understand and be able to discuss/describe the history of the discovery of subatomic particles. understand and be able to discuss/describe the structure of the atom, atomic number and mass number. understand atomic mass, isotopes, mass spectra and be able to calculate average atomic mass. become more familiar with the Periodic table: Families (columns) and Periods. understand compound formulas: molecular, empirical, structural. understand the difference between molecular and ionic compounds. become proficient in inorganic nomenclature: names and formulas of compounds (LAB). be able to name and write formulas for a few types of simple organic compounds. 1 Dalton s Atomic Theory ( ) Postulates 1. Each element is composed of extremely small, indivisible particles called atoms. 2. All atoms of a given element are identical to one another in mass and other properties, but 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. 4. Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms. Which of these postulates is consistent with the Law of Conservation of Mass? Which of these postulates is consistent with the Law of Definite Proportions? Which of these postulates have since been modified? 2

2 History of Modern Atomic Theory In the late 19th and early 20 th century, a series of experimentation allowed scientists to establish a model of the atom that remains today the foundation of modern atomic theory. The Discovery of the Electron-J.J. Thomson Cathode ray tubes (CRT) - mid 1800 s - electron beam Thompson concluded that cathode rays are streams of negatively charged particles based upon their behavior. He also measured the charge/mass ratio of the electron to be C/g in The exact mass of the electron was still unknown. (C is the symbol for coulomb, the SI unit for electric charge.) 3 History of Modern Atomic Theory Millikan Oil Drop Experiment (1909) Once the charge/mass ratio of the electron was known, determination of either the charge or the mass of an electron would enable you to determine the other. Robert Millikan (University of Chicago) determined the charge on the electron in Electron Charge: 1.602x10-19 C Electron Mass =? (calculate it) 4

3 History of Modern Atomic Theory Radioactivity-further evidence that the atom is divisible! The spontaneous emission of radiation by an atom. First observed by Henri Becquerel (1896). Also studied by Marie and Pierre Curie. Three types of radiation were identified by Ernest Rutherford: α particles (He nucleus, + charge and heavy) β particles (electron, charge and light) γ rays (high energy electromagnetic radiation: no charge, no mass) 5 Discovery of the Nucleus In 1910, Ernest Rutherford directed a beam of α particles at a thin sheet of gold foil and observed the pattern of scatter of the particles. 6

4 Rutherford is credited with the discovery of the nuclear atom. Based on the alpha particle scatter pattern, Rutherford postulated in 1911 that the atom contains a very small, dense nucleus with the electrons surrounding the nucleus. Most of the volume of the atom is empty space. Protons were later discovered by Rutherford in (They were first detected by Eugen Goldstein in 1886, emitted in the opposite direction compared to electrons from a CRT.) Neutrons were finally discovered by James Chadwick in Model of the Atom Electrons ( ), e Protons (+), p + Neutrons (0), n 0 Neutral atoms: number of protons=number of electrons Summary of Subatomic Particles 8

5 Mass Spectrometry: The Discovery of Isotopes! A mass spectrometer measures the relative mass and abundance of an element s isotopes. Mass Spectrum of Cl The two isotopes of Cl are clearly defined. Determine their approximate percent abundance. 9 Some Examples of Isotopic Abundances and Masses Determined Using Mass Spectrometry 10

6 Atomic Number, Mass Number and Isotopic Notation Atomic Number, Z: The number of protons in the nucleus of an element. Each specific element has a unique atomic number. Mass number, A: The number of protons + neutrons in the nucleus of an element. Mass number is NOT unique. Isotopes Atoms of the same element with different masses. Isotopes have different numbers of neutrons. 11 6C 12 6C 13 6C 14 6C Isotopic notation : Z A X carbon -12 : 12 6 C What is the isotopic notation for Uranium-235? 11 Atomic Mass Mass values on the periodic table are relative to carbon-12. The mass values are given as amu (atomic mass units). 1 amu = 1/12 the mass of an atom of carbon-12 (defined). We convert to absolute units (g, kg, lbs. etc.) by using the conversions: 1 amu = 1.661x10-24 g 1 g = x10 23 amu (look familiar?) The atomic masses listed in the periodic table are average masses. These are weighted averages based upon the naturally occurring isotopic abundances. Average Atomic Mass = 12

7 Calculating Atomic Masses Textbook Problem 2.35: Naturally occurring magnesium has the following isotopic abundances. (a)what is the average atomic mass of magnesium? (b)sketch the mass spectrum of magnesium. 13 Mendeleev is given credit for first proposal of the Periodic Table, published in Groups or families; these are the columns of the periodic table and contain elements with similar properties. Rows; these are called periods. The Periodic Table 14

8 Navigating the Periodic Table Make sure you can locate and identify the various classifications of the elements. Alkali metals: Group 1A Alkaline earth metals: Group 2A Halogens: Group 7A Nobel gases : Group 8A The 7 diatomic elements MUST be memorized! 15 Representing Molecular Compounds Molecule: smallest identifiable unit of a pure COVALENT (not ionic) compound. Molecules are primarily formed from the nonmetal elements. Molecular compounds can be gases, liquids or solids at room temperature and pressure. There are various ways to represent molecules structurally: Perspective drawings show the three-dimensional arrangement of atoms. Structural formulas show the order in which atoms are bonded. Condensed formulas (condensed structural) group atoms that are bonded together. Example: Ethylene Glycol 16

9 Types of Chemical Formulas Molecular formulas give the actual number of atoms of each element in one molecule of a compound. (Used only for covalent compounds.) Ethylene Glycol Example: Different compounds can have the same molecular formula with a different arrangement of the atoms. Empirical formulas give the lowest whole-number ratio of atoms of each element in a compound. Ethylene Glycol Example: Note: For organic compounds, the standard order of the elemental symbols is C, followed by H, then the remaining symbols in alphabetical order. Compounds with different molecular formulas can have the same empirical formula. Empirical formulas are always used for ionic compounds. Do you know why? 17 Molecular Compounds-More Examples Let s write the structural, condensed, molecular and empirical formulas for acetic acid. Let s compare ethanol (b.p. = 78.5 C) with dimethyl ether (b.p. = 23.6 C). 18

10 Ionic Compounds: Cations + Anions CATIONS are formed by the loss of ELECTRONS Cations are formed when an atom LOSES electrons (oxidation). lost e (to anion) Cations are (+) charged ions. 19 Common monatomic CATIONS formed from the group 1A, 2A and 3A METALS The monatomic cations are all metals. They lose valence shell electrons = column #. Group 1A : lose how many e? Group 2A: lose how many e? Metals of group 3A: lose how many e? 20

11 Ionic Compounds: Cations + Anions ANIONS are formed by the GAIN of ELECTRONS Anions are formed when an atom GAINS electrons (reduction). gained e (from cation) Anions are ( ) charged ions. 21 Common monatomic ANIONS are formed from the NONMETAL elements The monatomic anions are all nonmetals. They gain electrons to fill their outer valence shell (octet rule). Group 4A nonmetal: gain how many e? Group 5A nonmetals: gain how many e? Group 6A nonmetals: gain how many e? Group 7A: gain how many e? H when bonded to a metal: gains how many e? 22

12 Some Common Ions Cd 2+ Formulas of ionic compounds are empirical formulas, they reflect the smallest whole number ratio of cations to anions that results in a net charge of zero. Examples: 23 Coulomb s Law for Charged Particles: IONIC Attractions! The force of attraction between two charged particles is given by Coulomb s Law: (Electrostatic force) Force = k (n+ e)(n e) k is a constant d 2 n + and n - are the magnitude of the (+) and ( ) charges e is the charge of an electron d is the distance between the atoms 24

13 IONIC Forces of Attraction The electrostatic forces that hold cations and anions together in an extended array (a lattice) are STRONG! As a result of these strong attractive forces, ALL ionic compounds are solids at room temperature. (Ionic Compounds have high melting points.) Solid ionic compounds DO NOT conduct electricity. In the molten state ionic compounds DO conduct electricity. Water solutions containing dissolved ionic compounds DO conduct electricity. 25 Polyatomic Ions: Polyatomic ions consist of more than one atom. Polyatomic ions are usually groups of NONMETAL elements covalently bonded together. This group as a whole has a net charge (+ or ). 26

14 Chemical Nomenclature: Names and Formulas Inorganic Compounds: The naming of ionic and molecular inorganic compounds will be covered during a LAB period. This should be review for you! You are expected to be proficient in inorganic nomenclature. This is essential for success in some future topics covered in general chemistry! Some simple Organic (carbon based) Compounds: Alkanes: contain only C and H, all bonds are single Alcohols: contain the OH functional* group, covalently bonded to a carbon atom Carboxylic acids: contain the COOH functional group *Functional Group: An atom or group of atoms that imparts characteristic properties to an organic compound. 27 Questions and Problems Complete the following problems from the textbook: 2.1 A charged particle is caused to move between two electrically charged plates, as shown below. a) Why does the path of the charged particle bend? b) What is the sign of the electrical charge on the particle? c) As the charge on the plates is increased, would you expect the bending to increase, decrease, or stay the same? d) As the mass of the particle is increased while the speed of the particles remains the same, would you expect the bending to increase, decrease, or stay the same? 28

15 Questions and Problems 2.4 Does the following drawing represent a neutral atom or an ion? Write its complete chemical symbol including mass number, atomic number, 2.5 Which of the following diagrams is most likely to represent an ionic compound, and which a molecular one? Explain your choice. 29 Questions and Problems 2.6 Write the chemical formula for the following compound. Is the compound ionic or molecular? Name the compound a) Which two of the following are isotopes of the same element: b) What is the identity of the element whose isotopes you have selected? 30

16 2.49 Fill in the gaps in the following table: Questions and Problems 2.53 Using the periodic table to guide you, predict the chemical formula and name of the compound formed by the following elements: (a) Ga and F, (b) Li and H, (c) Al and I, (d) K and S a) What is a hydrocarbon? Questions and Problems b) Butane is the alkane with a chain of four carbon atoms. Write a structural formula for this compound, and determine its molecular and empirical formulas a) What is a functional group? b) What functional group characterizes an alcohol? c) With reference to Exercise 2.75, write a structural formula for 1-butanol, the alcohol derived from butane, by making a substitution on one of the end carbon atoms. 32

17 Questions and Problems 2.83 A cube of gold that is 1.00 cm on a side has a mass of 19.3 g. A single gold atom has a mass of amu. (a) How many gold atoms are in the cube? (b) From the information given, estimate the diameter in Å of a single gold atom. (c) What assumptions did you make in arriving at your answer for part (b)? 33 Questions and Problems 2.92 There are two different isotopes of bromine atoms. Under normal conditions, elemental bromine consists of two atoms (a diatomic molecule) and the mass of a Br 2 molecule is the sum of the masses of the two atoms in the molecule. The mass spectrum of Br 2 consists of three peaks: Mass (amu) Relative Peak Size (a) What is the origin of each peak (of what isotopes does each consist)? (b) What is the mass of each isotope? (c) Determine the average molecular mass of a Br 2 molecule. (d) Determine the average atomic mass of a bromine atom. (e) Calculate the abundances of the two isotopes. 34