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Chemistry 101 : Chap. 2

Chemistry 101 : Chap. 2. Atoms, Molecules and Ions. Atomic Theory of Matter (2) The discovery of Atomic Structure (3) The Modern View of Atomic Structure (4) Atomic Weight (5) Periodic Table (6) Molecules and Molecular Compounds (7) Ions and Ionic Compounds (8) Naming Inorganic Compounds.

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Chemistry 101 : Chap. 2

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  1. Chemistry 101 : Chap. 2 Atoms, Molecules and Ions • Atomic Theory of Matter • (2) The discovery of Atomic Structure • (3) The Modern View of Atomic Structure • (4) Atomic Weight • (5) Periodic Table • (6) Molecules and Molecular Compounds • (7) Ions and Ionic Compounds • (8) Naming Inorganic Compounds

  2. The Atomic Theory of Matter The history of development of atomic theory of matter begins in ancient Greece. However, modern atomic theory has it’s origin in a burst of scientific discovery between 1870 and 1930. Democritus (460 ~ 370 BC) Democritus proposed atomic theory of matter. He and other Greek philosophers believed that material world must be made up of hard and tiny indivisible particles that they called atomos, which are in constant motion.

  3. The Atomic Theory of Matter Aristotle (384 ~ 322 BC) Aristotle proposed 4 element theory of matter. Fire dry hot Air Earth wet cold Water The school of thought laid out by Socrates, Plato and Aristotle dominated the western philosophy for 2000 years and the atomic theory of matter was completely buried.

  4. The Atomic Theory of Matter John Dalton (1766 ~ 1844) • Dalton’s Atomic Theory • Each element is composed of atoms • All atoms of a given element are identical, • but they are different from the atoms of • all other elements • (3) Atoms are neither created nor destroyed • in chemical reactions. • (4) Compounds are formed from chemical • combination of two or more atoms. Dalton proposed that all matter is made up of atoms and stated that elements are the simplest form of matter.

  5. The Atomic Theory of Matter • What can Dalton’s theory explain? (1) Law of constant composition  In a given compound, the relative numbers and kinds of atoms are constant. [postulate 4] (2) Law of conservation of mass  The total masses of material present before and after a chemical reaction are identical [postulate 3] (3) Law of multiple proportions •  If elements A & B combine to form more than one compound, the • masses of B which can combine with a given mass of A are in • the ratios of small whole numbers 12g C + 16g O  CO or 12g C + 32g O  CO2 16g : 32g = 1:2

  6. The Discovery of Atomic Structure After Dalton’s atomic theory, not much of progress had been made and no one had direct evidence for the existence of atom. Then, things started to change in late 1800s…  William Crooks (1832 ~ 1919): Cathode-ray tube (CRT) [1879] A high voltage between two electrodes in a partially evacuated tube generates electrical discharge (cathode ray)

  7. The Discovery of Atomic Structure • J. J. Thomson (1856 ~ 1940) : Discovery of electron [1897] • Rays are the same regardless of the • identity of the cathode material • (2) Conduct quantitative analysis of the • effect of electric and magnetic field •  determine the charge to mass ratio He discovered that cathode rays are negatively charged particles, which he originally called ``corpuscles’’ . He won a Nobel prize in physics [1906]. charge/mass = 1.76  108 C/g

  8. The Discovery of Atomic Structure • Robert Millikan (1868 ~ 1953) : Determine the charge of electron [1907] Millikan’s oil-drop experiment Measured charge = 1.60 10-19 C Electron mass = charge/[charge/mass] = 9.10  10-28 g The machine on the right hand side is the original apparatus Millikan used to perform his oil-drop experiment. He won a Nobel prize in physics [1923].

  9. The Discovery of Atomic Structure • Ernest Rutherford (1871~1937): Discovery of nucleus [1911] Rutherford’s -particle [4He2+] scattering experiment He directed his graduate student Hans Geiger and undergraduate student Ernest Marsden to carry out -paticle experiment. He won a Nobel prize in chemistry [1908].

  10. The Discovery of Atomic Structure Radioactivity: Generation of  - particles   - ray: particles with +2 charge •  - ray: particles with 1 charge   - ray: high energy radiation with no charge

  11. The Discovery of Atomic Structure • From the scattering experiment…. (1) Most -particles simply pass through the gold foil. (2) Small amount of scattering was observed at large angles. • Rutherford postulated that.. (1) Most of the total volume of an atom is empty space. • (2) Most of the mass of an atom and all of its positive • charge reside in a very small region, • called nucleus. Rutherford also found the existence of protons inside of nucleus [1919]. Another particle in nucleus, neutron, was found by James Chadwick in 1932.

  12. Rutherford's Model: + Early Models of an Atom  Rutherford’s model • J. J. Thomson’s model “plum-pudding model” Electrons are negatively charged, but atoms as a whole are neutral.

  13. Modern View of Atomic Structure The list of subatomic particles has grown considerably since the discovery of electrons, but only the electron, proton and neutron have a bearing on chemical behavior. A convenient unit (non-SI) to describe the dimensions of atoms and molecules is Angstrom (Å). 1 Å = 1 10-10 m = 100 pm

  14. Modern View of Atomic Structure Properties of subatomic particles Particle Charge (C) Mass (g) Mass (amu) Proton +1.60  10-19 (+1) 1.6727  10-24 1.0073 Neutron 0 ( 0) 1.6750  10-24 1.0087 Electron 1.60  10-19 (1) 9.1097  10-28 5.486  10-4 Every atom has an equal number of protons and electrons so that it has no electrical charge Atomic Mass Unit (amu) 1 amu = 1/12 of the mass of carbon (12C) atom = 1.66054  10-24 (g)

  15. Modern View of Atomic Structure The characteristics of each atom are determined by the numbers of proton, neutron and electrons. Hydrogen: 1 proton Helium: 2 protons 2 neutrons Lithium: 3 protons 4 neutrons Beryllium: 4 protons 5 neutrons • Atomic Number: The number of protons in the nucleus of an atom. • Mass Number: The total number of protons plus neutrons in the atom • Isotopes : Atoms with identical atomic numbers but different mass numbers such as C-14 and C-12.

  16. Modern View of Atomic Structure Same information : An element is defined by the number of protons

  17. Atomic Weight Atomic Mass Unit (amu) = 1.66054  10-24 g 12C = 12 amu (exact), 1H = 1.0078 amu, 16O = 15.9949 amu • Average Atomic Masses :Weighted average of all the isotopes of an element found in nature. Example : Naturally occurring carbon is composed of 98.93% 12C and 1.07 % 13C. What is the average mass of carbon? (0.9893)(12 amu) + (0.0107)(13.00335) = 12.01 amu This is the mass of carbon atom shown in the periodic table fractional abundance of C-12 mass of C-12 mass of C-13

  18. Atomic Weight Example: Boron has two naturally occurring isotopes: 10B (10.01 amu) and 11B (11.01 amu). If the average atomic weight of Boron is 10.81, what are the fractional abundances of the two isotopes?

  19. Periodic Table If the elements are arranged in order of increasing atomic number, their chemical properties are found to show a repeating, or periodic, pattern. period group Elements having similar properties are placed in vertical columns

  20. Periodic Table Halogen Alkaline earth metal Transition metals Alkali metal rare gas = H2, N2, O2, F2, Cl2, Br2, I2

  21. Molecules and Molecular Compounds Chemical Compounds Molecular Ionic • Molecular compounds are composed of more than • one type of atom • H2O, NH3, CH3OH, O2 (2) Most molecular substances contain only non-metallic atoms O2, H2O, H2O2, CO, CO2, CH4

  22. Molecules and Molecular Compounds  Chemical Formulars • Molecular Formulas : Indicate the actual numbers and types • of atoms in a molecule Ex. C2H4O2 (2) Empirical Formulas : Indicate the relative number of atoms of each type in a molecule Ex. CH2O (3) Structural Formulas : H O H – C – C – O – H H

  23. Molecules and Molecular Compounds Picturing Molecular Compounds (Ex. Methane) Structural Formula Perspective drawing Space-filling model Ball-and-stick model

  24. Ions and Ionic Compounds • Ion : Atoms can readily gain or loose electrons and become ions. Anion: An ion with a negative charge Cation: An ion with a positive charge Cl Na+

  25. Noble Gases Alkaline Earth Metals Alkali Metals Halogens Ions and Ionic Compounds Which elements form cations and which form anions? Metalstend to form Cations Nonmetalstend to form Anions VIII A I A II A III A IV A VA VI A VIIA

  26. Ions and Ionic Compounds How many electrons each element can gain or loose? Each element tends to have the same number of electrons as noble gases (rare gases).

  27. Ions and Ionic Compounds • Example: Determine the number of electrons, protons and neutrons in each of the following ions No. of Protons No. of Neutrons No. of Electrons 16O2- 40Ca2+ 58Fe3+ 80Br 

  28. Ions and Ionic Compounds • Ionic Compounds :Cations(metals) andanions (non-metal) combine to form ionic compounds NaCl Alternating positive and negative charges

  29. Ions and Ionic Compounds Ionic compounds : • Ionic compounds are generally combination of metals • and nonmetals NOTE: Molecular compounds are generally composed of nonmetals only (H2O , CH3OH , CH3CH2Cl , …) (2) Ionic compounds are represented by empirical formulas  use simplest whole-number ratio of cations and anions NOTE: There is no discrete (or isolated) molecule of NaCl (3) Ionic compounds are always neutral. Therefore, the total positive charge equals the total negative charge Mg2+ and N3- form Mg3N2 : 3(+2) + 2(3) = 0

  30. Ions and Ionic Compounds • Example : Find the empirical formula for the ionic compound made of given cation and anion Na, O => Al, O => Ca, O =>

  31. Naming Ions and Ionic Compounds Names of ionic compounds consist of the cation name followed by the anion name CaCl2 = calcium + chloride  calcium chloride • Names of Positive Ions (cations) : (1) Cations formed from metal atoms have the same name as the metal. Na+ sodium ion, Zn+  zinc ion, Al3+  aluminum ion NOTE: Ions formed from a single atom are called monatomic ions

  32. Naming Ions and Ionic Compounds (2) If a metal can form different cations, the positive charge is indicated by a Roman numerical in parenthesis following the name of the metal Fe2+ iron (II) ion Cu+  copper (I) ion Fe3+  iron (III) ion Cu2+  copper (II) ion These ions are usuallytransition metals NOTE: Metals that form only one cation group 1A  Na+, K+, Rb+ group 2A  Mg2+, Ca2+, Sr2+, Ba2+ and Al3+ (group 3A), Ag+ (group 1B), Zn2+ (group 2B)

  33. Naming Ions and Ionic Compounds (3) Cations formed from nonmetal atoms have names that end in -ium NH4+ ammonium ion H3O+  hydronium ion NOTE: These ions are examples of polyatomic ions

  34. Naming Ions and Ionic Compounds • Names of Negative Ions (anion) : (1) The names of monatomic anions are formed by replacing the ending of the name of the element with –ide. H- hydrogen  hydride ion, O2- oxygen  oxide ion, NOTE: polyatomic anions with common names ending with –ide OH- hydroxide ion, CN-  cyanide ion (2) Polyatomic anions containing oxygen (oxyanions) • ending with –ate : reserved for the most common oxyanion • NO3- nitrate ion, SO42-  sulfate ion

  35. Naming Ions and Ionic Compounds b) ending with –ite : used for oxyanion with the same charge, but one fewer O atom than those ending with –ate. NO2- nitrite ion, SO32-  sulfite ion c) If a series of oxyanions extends to more than two members, use prefix per- (one more) or hypo- (one fewer) ClO4- perchlorate ion (one more than –ate) ClO3-  chlorate ion ClO2-  chlorite ion ClO- hypochlorite ion (one fewer than -ite)

  36. Naming Ions and Ionic Compounds NOTE: Oxyanions with the maximum number of oxygens (i) Charges increase from right to left. (ii) Second row elements (C, N) have maximum 3 oxygen atoms and third row elements (P, S, Cl) have maximum 4 oxygen atoms (row # + 1). (iii) All names end with –ate except for ClO4-

  37. Naming Ions and Ionic Compounds (3) Anions derived by adding H+ to an oxyanion are named by adding as a prefix the word hydrogen or dihydrogen. CO32- : carbonate ion  HCO3- : hydrogen carbonate ion PO43- : phosphate ion  H2PO4- : dihydrogen phosphate ion Halogen (7A)

  38. Names of Binary Molecular Compounds • The name of the element farther to the left in the periodic table • appear first. (NOTE: Oxygen is always written last except when • combined with fluorine.) (2) If both elements are in the same group, the one having the higher atomic number is named first (3) The name of the second element is given an –ide ending (4) Greek prefixes are used to indicate the number of atoms of each element (1  mono-, 2di-, 3 tri-, 4 tetra-, 5  penta-, 6  hexa- ) Cl2O : dichloro monoxide NF3 : nitrogen trifluoride P4S10 : tetraphosphorous decasulfide N2O4 : dinitrogen tetroxide

  39. Naming Compounds : Examples • Before you try to name a compound : (1) Is the compound ionic or molecular? (2) For ionic compounds, find the name of each ion. For molecular compounds, find the number of each atom. BF3 : NiO : KMnO4 : SO

  40. Naming Compounds : Examples • Write down the chemical formulas for the following compounds (1) Sodium Nitride, Q:Is this ionic or molecular? Q: Is anion monatomic or polyatomic ion? (2) Diphosphorus pentoxide,

  41. Naming Compounds : Examples (1) NaClO : (2) Fe2(CO3)3 : (3) SF6 : (4) aluminium hydroxide : (5) ammonium sulfate : (6) NaH2PO4 :

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