Chapter #2 Atoms, Molecules and Ions

Transcription

1 Chapter #2 Atoms, Molecules and Ions 2.1 The Early History of Chemistry 2.2 Fundamental Chemical laws 2.3 Dalton s Atomic Theory 2.4 Cannizzaro s Interpretation 2.5 Early experiments to Characterize the Atom 2.6 The Modern View of Atomic Structure: An Introduction 2.7 Molecules and Ions 2.8 An Introduction to the Periodic Table 2.9 Naming Simple Compounds 1 Chapter 2 Topics 1: The Observations That Led to an Atomic View of Matter 2: The Observations That Led to the Nuclear Atom Model 3: Dalton s Atomic Theory and Today s Version 4: Molecules and Ions 5. Elements: A First Look at the Periodic Table 6: Compounds: Introduction to Bonding 7: Compounds: Formulas and Names 2 1

2 Law of Conservation of Mass The total mass of substances does not change during a chemical reaction. 3 Law of Definite Proportions No matter what its source, a particular chemical compound is composed of the same elements in the same parts (fractions) by mass. 4 2

3 Law of Definite Proportions WATER H 2 O No matter what the source water is ALWAYS 2 parts hydrogen to 1 part oxygen 5 Law of Definite Proportions Chemical analysis of a 9.07 g sample of calcium phosphate shows that it contains 3.52 g of Ca. How much Ca could be obtained from a kg sample? kg kg kg kg 6 3

4 Law of Definite Proportions Chemical analysis of a 9.07 g sample of calcium phosphate shows that it contains 3.52 g of Ca. How much Ca could be obtained from a kg sample? Mass fraction of Ca = 3.52 g Ca = * 100% = 38.8% 9.07 g sample (i.e., ANY sample of Ca 3 (PO 4 ) 2 is 38.8% Ca by mass ) Mass of Ca in kg of sample = kg sample 38.8 kg Ca 100 kg sample = kg Ca 7 Law of Multiple Proportions In a nutshell, two (or more) compounds can contain different relative amounts of the same elements: If elements A and B react to form two compounds, the different masses of B that combine with a fixed mass of A can be expressed as a ratio of small whole numbers. (Evidence of the existence tiny individual particles.) 8 4

5 Law of Multiple Proportions Mass of Oxygen that Combines with 1.00 g of Carbon Compound #1 Compound # g O per g C 2.66 g O per g C mass of O in compound #2 = 2.66 g = 2 mass of O in compound # g 1 EXACT 2:1 RATIO

6 Atomic Theory Protons, Neutrons and Electrons Theory and Discovery 11 Daltons Atomic Theory 1. All matter consists of tiny particles called atoms. 2. Atoms of an element are identical in mass and other properties and are different from atoms of any other element. 3. Compounds result from the chemical combination of a specific ratio of atoms of different elements. 4. Chemical reactions involve reorganization of the atoms changes in the way they are bonded. Atoms of one element do not change and cannot be converted into atoms of another element during chemical reactions. John Dalton 12 6

7 Thomson and Cathode Rays Aim: To study the structure of the atom. Investigate the electrical discharges of atoms. Procedure: Use a partially evacuated tube to apply high voltage to a screen coated with a chemical compound. Measure the deflection. 13 Thomson and Cathode Rays Thomson used partially evacuated glass tubes to discover the existence of negatively charged particles called electrons. Developed the plum pudding model. Plum Pudding Model 14 7

8 Millikan and Oil Drops Aim: To study the electronic properties of the atom. Investigate the mass of an electron. Procedure: Use x-rays to produce charges on oil drops. Measure the magnitude of the electron charge. 15 With Thomson s cathode ray experiment, determined the mass of an electron, 9.11x10-31 kg. Millikan and Oil Drops 16 8

9 Rutherford Experiment Aim: To study the internal structure of the atom. Investigate the mass distribution in the atom. Procedure: Use a radioactive source to bombard a thin piece of gold foil. Collect the radiation. 18 9

10 Original and Revised Theories Original Theory: Plum pudding model Revised Theory: Nucleus (dense positive charge) at the center of the atom. Large amount of space between nucleus and electron. 19 Moving electron cloud surrounding nucleus. Almost all the mass is in the nucleus! 20 10

11 Notes: mass of e - tiny relative to p +, n. p +, n have same mass (almost). e -, p + have same charge, opposite sign. 21 Atomic Definitions I: Symbols, Isotopes, Numbers A Z X The Nuclear Symbol of the Atom X = Atomic symbol of the element Z = The atomic number: the number of protons in the nucleus (All atoms of the same element have the same # of protons.) A = The mass number (protons plus neutrons: A = Z + N) N = The number of neutrons in the nucleus N = A - Z Isotopes = atoms of an element with the same number of protons, but different number of neutrons in the nucleus 22 11

12 Figure 2.14 Isotopes of sodium 23 What is the nuclear particle with no charge? 1. Electron 2. Proton 3. Neutron 4. Positron 5. Alpha 24 12

13 Atomic Nucleus contains protons with positive charge contains neutrons with no charge chemistry of atom is determined by its electronic structure atoms with same number of protons have identical chemical properties isotopes are atoms with same number of protons but different number of neutrons 25 Which of the following statements are true? I. The number of protons is the same for all neutral atoms of an element. II. The number of electrons is the same for all neutral atoms of an element. III. The number of neutrons is the same for all neutral atoms of an element. 1. I, II, and III 2. I and II only 3. II and III only 4. I and III only 5. None 26 13

14 Available on the course website under Exam Info and Lecture Notes 27 A neutral atom of rhenium(re)-185 contains: p +, 75 n, 110 e p +, 111 n, 75 e p +, 110 n, 75 e p +, 110 n, 74 e

15 Neutral ATOMS If neutral, then # e - = # p + = atomic number. Remember: # n = A - # p + Numbers of each particle: 51 Cr = p + (24), e - (24 ), n ( 27) 239 Pu = p + (94), e - (94), n (145) 15 N = p + (7), e - (7), n (8) 56 Fe = p + (26), e - (26), n (30) 235 U = p + (92), e - (92), n (143) 29 Modern Reassessment of the Atomic Theory 1. All matter is composed of atoms. Although atoms are composed of smaller particles (electrons, protons, and neutrons), the atom is the smallest body that retains the unique identity of the element. 2. Atoms of one element cannot be converted into atoms of another element in a chemical reaction. Elements can only be converted into other elements in nuclear reactions in which protons are changed. 3. All atoms of an element have the same number of protons and electrons, which determine the chemical behavior of the element. Isotopes of an element differ in the number of neutrons, and thus in mass number, but not in chemical behavior (much). A sample of the element is treated as though its atoms have an average mass. 4. Compounds are formed by the chemical combination of two or more elements in specific ratios, as originally stated by Dalton

16 What separates isotopes of the same element? 1. Different number of protons 2. Different number of electrons 3. Different number of neutrons 4. Different overall charge 5. None of the above 31 Definitions Element - The simplest type of substance with unique physical and chemical properties. An element consists of only one type of atom. It cannot be broken down into any simpler substances by physical or chemical means. Compound - A substance composed of two or more elements that are chemically combined. Pure Substances - Their compositions are fixed! Elements and compounds are examples of pure substances. Mixture - Is a group of two or more elements and/or compounds that are physically intermingled. Molecule - A structure that consists of two or more atoms that are chemically bound together and thus behave as an independent unit

17 Definitions Chemical Formula The symbols of for the elements are used to indicate the types of atoms present, and the subscripts are used to indicate the relative numbers of atoms present. Structural Formula A formula in which the bonds are shown along with the elemental symbols and order of atom arrangement. 33 See figures in the textbook

18 More on amu in Ch Review: Based on the Law of Definite Proportions, what H:O ratio do you expect to find in water (H 2 O)? 1. 1:1 2. 2:1 3. 1:2 4. 2:

19 Forming Bonds Sharing and Transfer of Electrons 37 Chemical Compounds and Bonds Chemical Bonds The forces that hold the atoms of elements together in the compound. Covalent Bonds Electrons are shared between atoms of different elements to form covalent compounds. Ionic Bonds- Attractive forces between two oppositely charged ions to form ionic compounds

20 Chemical Compounds and Bonds Cation An atom that has lost electron(s) to form a + ion. Common for metal elements. Anion An atom that has gained electron(s), to form a - ion. Common for nonmetal elements. Monatomic (single atom) ions form binary ionic compounds. Polyatomic ions have many atoms per ion and an overall charge. 39 A: No interaction, atoms too far apart. B: Atoms move closer, electron clouds distort. C: Covalent bond forms. D: The protons from each nucleus share the overlapping electron cloud

21 Forming an Ionic Bond Step 1: Electron transfer from sodium atom (neutral sodium atom becomes sodium + ion) 41 Copyright Houghton Mifflin Company. All rights reserved. Forming an Ionic Bond Step 2: Electron added to chlorine atom (neutral chlorine atom becomes chloride - ion) 42 Copyright Houghton Mifflin Company. All rights reserved. 21

22 Arrangement of sodium ions and chloride ions 43 Figure 2.18: Sodium metal reacts with chlorine gas 2 Na(s) + Cl2(g) 2NaCl(s) 44 22

23 How is an ion formed? 1. By either adding or subtracting protons from the atom. 2. By either adding or subtracting electrons from the atom. 3. By either adding or subtracting neutrons from the atom. 4. All of the above are true. 5. Two of the above are true. 45 PERIODIC TABLE MENDELEEV TABLE A tabular arrangement of the elements based upon their chemical properties most elements are metals (left side of line) and tend to lose electrons nonmetals (right side) tend to gain electrons vertical columns are groups - similar properties such as: alkali metals, halogens, noble gases, etc. horizontal rows are periods 46 23

24 Figure 2.20: The periodic table 47 Groups in the Periodic Table Main Group Elements (Vertical Groups) Group 1A - Alkali Metals Group 2A - Alkaline Earth Metals Group 3A - Boron Family Group 4A - Carbon Family Group 5A - Nitrogen Family Group 6A - Oxygen Family Group 7A - Halogens Group 8A - Noble Gases Other Groups (Vertical and Horizontal Groups) Groups Transition Metals Period 6 - Lanthanides (Rare Earth Elements) Period 7 - Actinides (Radiocative & Artificial) 48 24

25 POSITION of ELEMENTS on the PERIODIC TABLE Tells you about: Properties Reactivity with other elements Most probable oxidation state(s) Composition of compounds 49 WRAP TABLE AROUND! nonmetals tend to gain electrons metals tend to lose electrons 50 25

26 What monatomic ion would chlorine (Cl) form? 1. Cl 2-2. Cl - 3. Cl + 4. Cl Predicting the Ion that an Element will Form Problem: What monatomic ions will each of the elements form? (a) barium(z=56) (b) sulfur(z=16) (c) titanium(z =22) (d) fluorine(z=9) Plan: We use the z value to find the element in the periodic table and which is the nearest noble gas. Elements that lie after a noble gas will lose electrons and those before a noble gas will gain electrons. Solution: (a) Ba 2+, barium - an alkaline earth element (Group 2A) - will lose two electrons - attains the same number of electrons as the xenon (Xe) (b) S 2-, sulfur - in the oxygen family (Group 6A) - expected to gain two electrons - attains the same number of electrons as argon (Ar) (c) Ti 4+, titanium - in Group 4B - expected to lose 4 electrons - attains the same number of electrons as argon (Ar). (d) F -, fluorine - a halogen (Group 7A) - expected to gain one electron 52 - attains the same number of electrons as neon (Ne). 26

27 +1 H Li K Rb Cs Fr +2 Be Na Mg Ca Sr Ba Ra Sc Y La Ac The Periodic Table of the Elements Ti Most Probable Oxidation States (more on these in Ch 4) Hf Rf V Zr Nb Mo Tc Ta Cr Mn Fe Co Ni Cu W Re Ru Os Ir Pt Db Sg Bh Hs Mt Ds He B Al Rh Pd Ag Cd Au In Tl C Si Sn Pb N P Zn Ga Ge As Hg Sb Bi O S Se Te Po F Cl Br I At Ne Ar Kr Xe Rn Ce Th Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Pa U Np Pu Am Cm Bk Cf Es FmMd No Lr 53 Naming Simple Compounds BINARY: comprised of two elements IONIC: cation + anion (metal + nonmetal; two types of metals: Type I and Type II) COVALENT: between two nonmetals POLYATOMIC: several elements 54 27

28 55 (Type II: metals that have >1 common oxidation state) 56 28

29 Rules for Naming Binary Ionic Compounds Always list cation (+) name first, anion (-) name second. Monatomic anion s name = first part of element name + -ide If the ion is not listed in the Type II table, assume it s Type I. 57 Which of the following compounds is not properly named? 1. FeCl 3 : iron(iii) chloride 2. NaBr : sodium bromide 3. Ca 3 N 2 : calcium nitride 4. PbO : lead(ii) oxide 5. MgH 2 : magnesium hydride 6. None of the above 58 29

30 59 Which of the following compounds is not properly named? 1. Fe 2 (SO 4 ) 3 : iron(iii) sulfate 2. Ba(OH) 2 : barium hydroxide 3. CaCrO 4 : calcium chromate 4. Cr(NO 2 ) 3 : chromium(iii) nitrite 5. KClO 3 : potassium chloride 6. RbHCO 3 : rubidium hydrogen carbonate 60 30

31 Naming Binary Covalent Compounds The element further to the left in the periodic table goes first, named as if it were the cation. Second element is named as if it were the anion. Use prefixes to say how many atoms of each element are present, EXCEPT never use mono- on first element. EXAMPLES: N 2 O dinitrogen monoxide NO nitrogen monoxide N 2 O 4 dinitrogen tetroxide 61 Which of the following compounds is not properly named? 1. CCl 4 : carbon tetrachloride 2. CO 2 : carbon dioxide 3. H 2 S : dihydrogen sulfide 4. NF 3 : nitrogen trifluoride 5. SF 6 : sulfur hexafluoride 6. None of the above 62 31

32 HYDRATES Compounds containing WATER molecules MgSO 4 7H 2 O magnesium sulfate heptahydrate CaSO 4 2H 2 O calcium sulfate dihydrate Ba(OH) 2 8H 2 O barium hydroxide octahydrate CuSO 4 5H 2 O copper(ii) sulfate pentahydrate Na 2 CO 3 10H 2 O sodium carbonate decahydrate 63 Many common anions, when combined with H + as the cation(s), make a very reactive compound called an acid, which dissociates in water to give the two separate ions both dissolved in water. Examples: Acids H 3 PO 4 (aq) H + (aq) + H 2 PO 4- (aq) HCl (aq) H + (aq) + Cl - (aq) H 2 SO 4 (aq) 2 H + (aq) + SO 4 2- (aq) phosphoric acid hydrochloric acid sulfuric acid Litmus paper indicates if it s an acidic solution ph = quantitative measure of H + concentration in water 64 32

33 Rules for naming acids Acids = Molecules which dissociate when dissolved in water to give H + (aq) Binary acids Oxoacids Keep any prefixes (hypoor per-) 65 Which of the following compounds is not properly named? 1. HNO 3 : nitric acid 2. HCl : hydrochloric acid 3. H 2 SO 4 : sulfuric acid 4. HClO 4 : chloric acid 5. H 3 PO 4 : phosphoric acid 6. None of the above 66 33