Chapter 2 notes
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This document provides an overview of atomic structure and chemical bonding. It begins with a brief history of atomic theory from Dalton to Rutherford, including Thomson's discovery of the electron. Key topics covered include atomic number, mass number, isotopes, ions, molecules, chemical formulas and nomenclature. The periodic table is introduced as a way to organize elements based on atomic structure. Chemical formulas are used to represent ionic compounds, molecular compounds and acids/bases. Hydrates are also defined as compounds that contain bonded water molecules.
Chapter 2 notes
- 1. Atoms, Molecules and Ions Chapter 2 2.1 The Atomic Theories 2.2 The structure of the atom 2.3 Atomic number, mass number and isotopes 2.4 The periodic table 2.5 Molecules and ions 2.6 Chemical formula 2.7 Naming compounds
- 2. 2 Dalton’s Atomic Theory 1. Elements are composed of extremely small particles called atoms. Atoms of the same element all have the same size, mass and chemical properties. The atoms of one element are different from the atoms of all other element. 2. Compounds are composed of atoms of two or more elements. 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. 3. A chemical reaction involves only the separation, combination, or rearrangement of atoms; it does not result in their creation or destruction.
- 3. 3 Law of Definite Proportions -Different samples of the same compound always contains its elements in a definite proportion by mass. Dalton’s Atomic Theory H2O
- 4. 4 Dalton’s Atomic Theory Law of Multiple Proportions -In different compounds of the same elements, the various masses of one element that combine with a fixed mass of another element are related by small whole-number ratios.
- 5. 5 8 X2Y16 X 8 Y+ Law of Conservation of Mass - Matter is neither created nor destroyed Dalton’s Atomic Theory
- 6. 6 The Modern View of Atomic Structure Atom- the basic unit of an element that can enter into chemical combination (extremely small and indivisible) Three subatomic particles - electrons , protons, and neutrons.
- 7. 7 Thomson Cathode Ray Tube experiment - Thomson together with Millikan concluded that the mass of an e- is exceedingly small (e- mass = 9.10 x 10-28 g). - The cathode ray consist of negatively charged particles found in all matter
- 9. Three types of rays produced by decay of radioactive substances such as “Uranium”.. (i) Alpha (α ) rays..positively charged particles (α ) particles.. deflected by positively charged plate (ii) Beta (β) rays..electrons..deflected by negatively charged plate (iii) Gamma (γ) rays..high-energy rays..no charge and are not affected by an external field.
- 10. 10 Thomson’s Model – a spherical atom composed of diffuse, positively charge matter, in which e- embedded like “raisin in a plum pudding”.
- 11. Rutherford’s gold foil α-scattering experiment Atoms of gold foil
- 12. 12 atomic radius ~ 100 pm = 1 x 10-10 m nuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m Rutherford’s Model of the Atom 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)
- 13. 13 Chadwick’s Experiment (1932) H atoms - 1 p; He atoms - 2 p mass He/mass H should = 2 measured mass He/mass H = 4 neutron (n) is neutral (charge = 0) n mass ~ p mass = 1.67 x 10-24 g 2n0n
- 14. 14 mass p ≈ mass n ≈ 1840 x mass e-
- 15. 15 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) that have the same atomic number but different mass numbers XA Z H1 1 H (D)2 1 H (T)3 1 U235 92 U238 92 Mass Number Atomic Number Element Symbol Atomic number, Mass number and Isotopes
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- 17. 17 The Isotopes of Hydrogen
- 18. 18 6 protons, 8 (14 - 6) neutrons, 6 electrons 6 protons, 5 (11 - 6) neutrons, 6 electrons How many protons, neutrons, and electrons are in C 14 6 ? How many protons, neutrons, and electrons are in C 11 6 ?
- 19. 19 Naturally occurring carbon consists of three isotopes, 12 C, 13 C, and 14 C. State the number of protons, neutrons, and electrons in each of the following. 12 C 13 C 14 C 6 6 6 Proton 6 6 6 Neutron 6 7 8 Electron 6 6 6
- 20. 20 In naturally occurring magnesium, there are three isotopes. 24 Mg 25 Mg 26 Mg 12 12 12 Isotopes of Mg
- 21. 21 The Modern Periodic Table Period Group AlkaliMetal NobleGas Halogen AlkaliEarthMetal
- 22. 22 A molecule is an aggregate of two or more atoms in a definite arrangement held together by chemical forces H2 H2O NH3 CH4 H2, N2, O2, Br2, HCl, CO A polyatomic molecule contains more than two atoms O3, H2O, NH3, CH4 diatomic elements A diatomic molecule contains only two atoms
- 23. 23 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. anion – ion with a negative charge If a neutral atom gains one or more electrons it becomes an anion. Na 11 protons 11 electrons Na+ 11 protons 10 electrons Cl 17 protons 17 electrons Cl- 17 protons 18 electrons
- 24. 24 A monatomic ion contains only one atom A polyatomic ion contains more than one atom Na+ , Cl- , Ca2+ , O2- , Al3+ , N3- OH- , CN- , NH4 + , CO3 2− ,HCO3 − , SO4 2− , PO4 3− ,NO3 - ,NO2 −
- 25. 25 The names of common polyatomic anions • end in ate. NO3 − nitrate PO4 3− phosphate • with one oxygen less end in ite. NO2 − nitrite PO3 3− phosphite • with hydrogen attached use the prefix hydrogen (or bi). HCO3 − hydrogen carbonate (bicarbonate) HSO3 − hydrogen sulfite (bisulfite)
- 26. 26 Common Ions Shown on the Periodic Table
- 27. 27 13 protons, 10 (13 – 3) electrons 34 protons, 36 (34 + 2) electrons How many protons and electrons are in ?Al27 13 3+ How many protons and electrons are in ?Se78 34 2-
- 29. 29 A molecular formula shows the exact number of atoms of each element in the smallest unit of a substance An empirical formula shows the simplest whole-number ratio of the atoms in a substance H2OH2O molecular empirical C6H12O6 CH2O O3 O N2H4 NH2
- 30. 30 ionic compounds consist of a combination of cations and an anions • The formula is usually the same as the empirical formula • The sum of the charges on the cation(s) and anion(s) in each formula unit must equal zero The ionic compound NaCl
- 31. 31 The most reactive metals (green) and the most reactive nonmetals (blue) combine to form ionic compounds.
- 32. 32 Formula of Ionic Compounds Al2O3 2 x +3 = +6 3 x -2 = -6 Al3+ O2- CaBr2 1 x +2 = +2 2 x -1 = -2 Ca2+ Br- Na2CO3 1 x +2 = +2 1 x -2 = -2 Na+ CO3 2-
- 33. 33 Chemical Nomenclature • Ionic Compounds – Most are binary compounds, some are ternary compounds – Often a metal + nonmetal – Anion (nonmetal), add “ide” to element name BaCl2 barium chloride K2O potassium oxide Mg(OH)2 magnesium hydroxide KNO3 potassium nitrate
- 34. 34 • Transition metal ionic compounds – indicate charge on metal with Roman numerals FeCl2 2 Cl- -2 so Fe is +2 iron(II) chloride FeCl3 3 Cl- -3 so Fe is +3 iron(III) chloride Cr2S3 3 S-2 -6 so Cr is +3 (6/2) chromium(III) sulfide
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- 39. 39 • Molecular compounds − Nonmetals or nonmetals + metalloids − Common names − H2O, NH3, CH4, − Element furthest to the left in a period and closest to the bottom of a group on periodic table is placed first in formula − If more than one compound can be formed from the same elements, use prefixes to indicate number of each kind of atom − Last element name ends in ide
- 40. 40 HI hydrogen iodide NF3 nitrogen trifluoride SO2 sulfur dioxide N2Cl4 dinitrogen tetrachloride NO2 nitrogen dioxide N2O dinitrogen monoxide Molecular Compounds
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- 42. 42 An acid can be defined as a substance that yields hydrogen ions (H+ ) when dissolved in water. For example: HCl gas and HCl in water •Pure substance, hydrogen chloride •Dissolved in water (H3O+ and Cl− ), hydrochloric acid
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- 44. 44 An oxoacid is an acid that contains hydrogen, oxygen, and another element (the central element). H3PO4 phosphoric acid HNO3 nitric acid H2SO4 sulfuric acid HNO2 nitrous acid H2SO3 sulfurous acid H2CO3 carbonic acid
- 45. 45 Naming Oxoacids and Oxoanions
- 46. 46 The rules for naming oxoanions, anions of oxoacids, are as follows: 1. When all the H ions are removed from the “-ic” acid, the anion’s name ends with “-ate.” 2. When all the H ions are removed from the “-ous” acid, the anion’s name ends with “-ite.” 3. The names of anions in which one or more but not all the hydrogen ions have been removed must indicate the number of H ions present. For example: – H2PO4 - dihydrogen phosphate – HPO4 2- hydrogen phosphate – PO4 3- phosphate
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- 48. 48 A base can be defined as a substance that yields hydroxide ions (OH- ) when dissolved in water. NaOH sodium hydroxide KOH potassium hydroxide Ba(OH)2 barium hydroxide
- 49. 49 Hydrates are compounds that have a specific number of water molecules attached to them. BaCl2•2H2O LiCl•H2O MgSO4•7H2O Sr(NO3)2 •4H2O barium chloride dihydrate lithium chloride monohydrate magnesium sulfate heptahydrate strontium nitrate tetrahydrate CuSO4•5H2O CuSO4
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