Chapter 2


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Chapter 2

  1. 1. Chapter 2: Atoms, Ions, and Molecules Outline2.1: Atoms/Atomic Theory 2.4: Molecules/Ions2.2: Components of Atom 2.5: Ionic Compounds2.3: Periodic Table 2.6: Names of Compounds
  2. 2. Building Blocks of ChemistryAtoms: electrons, protons, and neutronsMolecules: fundamental unit of compounds;Identified by formulas and namesIons: a charged species, either positive or negative,as in ionic compounds
  3. 3. 2.1: Atoms and the Atomic TheoryDalton’s model of the atom (1808):I. Element is composed of tiny particles calledatoms. All atoms of given element have sameproperties. Atoms of different elements havedifferent properties
  4. 4. 2.1: Atoms and the Atomic TheoryII. In an ordinary chemical reaction, atoms movefrom one substance to another, but are not createdor destroyed or converted into an atom of anotherelement (law of conservation of matter).III. Compounds are formed when atoms of two ormore elements combine. In a given compound, theproportions of the atoms are definite and constant(law of constant proportions).
  5. 5. 2.2: Components of the AtomI. Electrons: First evidence observed in studies of conduction of electricity through gases at low pressures. a. J.J. Thomson (1897): cathode rays produced consisted of a stream of negatively charged particles. b. Electrons are common to all atoms, carry a unit negative charge of –1, and have a very small mass, roughly 1/2000 that of the lightest atom (Plum Pudding Model).
  6. 6. RadioactivityRoentgen discovered X-rays in 1895, while studying the glow produced in certain substances by cathode rays. Noted glow on piece of paper some distance from tube; remained glowing when taken into another room.
  7. 7. Radioactivity (Cont’d) Fluoresence: chemicals which continue to glow after being exposed to radiation (light). Becquerel studied fluoresence by wrapping photographic paper in black paper, placing a few crystals of material on the paper and exposing it to strong sunlight. Ordinary light would not pass through the black paper, however, X-rays would.
  8. 8. Radioactivity (Cont’d)Looking at Uranium compounds, He discovered that they fogged the paper even when not exposed to sunlight. His graduate student, Marie Sklodowska (Curie), called it radioactivity and they won the Nobel Prize in 1903. She won in again in 1911.
  9. 9. 3 Types of RadioactivityDiscovered by Ernest Rutherford. I. Alpha (α): beam of positively charged particles. Mass 4 times that of a hydrogen atom and a charge twice the magnitude of, but opposite in sign to, an electron II. Beta (β): negatively charged particles, electrons. III. Gamma (γ): not deflected by magnetic field
  10. 10. Electrons (Cont’d)II. Millikan’s Oil Drop Experiment: Thomson able to measure mass to charge ratio for an electron but could not measure the mass or charge. a. Millikan discovered the charge in 1909. Mass measured using the rate at which the oil drops fall under the influence of gravity. b. Mass of electron determined to be 9.1 x 10 - 28 g.
  11. 11. 2.2: Components of the AtomI. Protons and Neutrons: Ernest Rutherford (1911) a. Bombarded a piece of thin gold foil with α- particles (helium minus its electrons) b. Most of the particles passed through the foil with no change in direction; some, however, were reflected back at acute angles, suggesting the mass of the atom was concentrated in the center and was positively charged
  12. 12. Components of the NucleusI. Proton: mass nearly equal to that of a single hydrogen atom. Carries a charge of (+1), equal in magnitude to that of an electron (-1).II. Neutron: uncharged particle with a mass slightly greater than that of a proton.Because protons and neutrons are much heavier than electrons, >99.9% of the mass of the atom is in the nucleus, even though the volume of the nucleus is much smaller than that of the atom.
  13. 13. Atomic NumberAll the atoms of a particular element have the same number of protons. This number is a basic property of an element and is called its atomic number and is given the symbol Z. Z = number of protonsIn a neutral atom, the number of protons in the nucleus is exactly equal to the number of electrons outside the nucleus
  14. 14. ExamplesH atom: 1 proton, 1 electron; Z = 1C atom: 6 protons, 6 electrons; Z = 6Fe atom: 26 protons, 26 electrons; Z = 26
  15. 15. Mass NumbersThe mass number of an element, A, is found by adding up the number of protons and neutrons in the nucleus: A = # of protons + # of neutronsAll atoms of a given element have the same number of protons, hence same atomic number. They may differ, however, from one another in mass and therefore, mass number
  16. 16. IsotopesWhile the number of protons in an atom is fixed, the number of neutrons is not (elements are distinguished by the number of protons in the nucleus).Atoms that contain the same number of protons, but a different number of neutrons are called isotopes.
  17. 17. Examples of IsotopesZ = 1: Hydrogen, A = 1 (no neutrons) Deuterium, A = 2 (1 neutron) Tritium, A = 3 (2 neutrons)Z = 92: Uranium 235, A = 235 (143 neutrons) Uranium 238, A = 238 (146 neutrons)
  18. 18. Nuclear SymbolComposition of the nucleus is shown by its nuclear symbol, where A is the mass number, Z is the atomic number and X is the element symbol. A X Z
  19. 19. Introduction to the Periodic TableAn element is a substance all of whose atoms have the same number of protons and thus the same atomic number. The periodic table is a listing of all known elements and their atomic numbers. a. Horizontal rows are called periods. b. Vertical columns are called families (groups). i. Designated by numbers 1-18 (IUPAC, 1985)
  20. 20. Periodic TableElements falling in groups 1, 2, 13, 14, 15, 16, 17, and 18 (formerly groups 3A-8A) are referred to as main-group elements.Elements in groups 3-12 are called transition metals.Elements in group 1 are called the alkali metals.Elements in group 2 are called the alkaline earth metals.Elements in group 7 are called the halogens.Elements in group 8 are called the noble gases.
  21. 21. Trends in the Periodic TableElements in the same group have similar chemical properties, i.e. Li, Na, and K all react vigorously with water to produce hydrogen gas. He, Ne, and Ar do not react with any other substances.The periodic table is an arrangement of elements, in order of increasing atomic number, and in horizontal rows of such a length, that elements with similar chemical properties fall directly beneath one another in vertical groups (Mendeleev).
  22. 22. Metals and NonmetalsThe diagonal line that starts from boron separates metals and nonmetals; elements along this line are called metalloids (B, Si, Ge, As, Sb, Te).Metals: conductive of heat and electricity, malleable, ductile, shiny, solids.Non-metals: not always solids, poor conductors of heat/electricity
  23. 23. Molecules and IonsMolecules and Ions result from the combination of atoms from two or more elements. a. A molecule is uncharged and is usually composed of nonmetallic elements. The atoms are held together by covalent bonds (shared electrons). b. Represented by molecular formula, the number of atoms of each element indicated by a subscript written after the elemental symbol. e.g. H2O CO2 CH4 NH3 CH3OH
  24. 24. IonsIf an atom gains or loses electrons, charged particles called ions are formed. a. Metals lose electrons to form positively charged ions called cations. e.g. Na0  Na+ + 1 electron b. Nonmetals typically gain electrons to form negatively charged particles called anions. e.g. Cl0 + 1 electron  Cl-1
  25. 25. Ions (Cont’d)If an ion is derived from a single atom, they are said to be monatomic.Many of the most important ions in chemistry are polyatomic, that is, containing more than 1 atom. e.g. OH-1 NH4+1A polyatomic ion is essentially a charged molecule.
  26. 26. Ionic CompoundsIonic compounds are compounds in which the ions are held together by an ionic bond (transfer of electrons).The molecular formula for an ionic compound is illustrated the same way as for molecules. e.g. NaCl CaCl2KBrNote that the metal is always shown first.
  27. 27. Cations and Anions With Noble Gas StructureThe charges of ions formed by main-group elements can be predicted in a straight-forward manner.Atoms that are close to a noble gas (Group 18) in the periodic table form ions that contain the same number of electrons as the closest noble gas atom.The unreactivity of noble gases suggest a stable electronic configuration, which other atoms would like to achieve.
  28. 28. Trends in the Formation of Ions of the Main Group ElementsGroup No. of Electrons in Atom Charge Examples 1 1 more than noble gas +1 Na, K 2 2 more than noble gas +2 Mg, Ca 16 2 less than noble gas -2 O, S 17 1 less than noble gas -1 F, Cl
  29. 29. Polyatomic IonsThere are only 2 common polyatomic cations: NH4+, Hg2+All other cations considered here will be derived from individual metal atoms.Most of the polyatomic anions contain 1 or more oxygen atoms. These species are referred to as oxoanions
  30. 30. Names of CompoundsA compound can be identified by one of 2 ways: chemical formula or chemical nameNomenclature of Ions: Monoatomic cations take the name of the metal from which they are derived. e.g. Na+ sodium K+ potassiumComplications arise from metals which can adopt more than one ion (oxidation state)
  31. 31. Metals With Multiple Oxidation StatesIn order to distinguish between these different ions, a roman numeral is used to indicate the charge e.g. Fe2+ iron(II) Fe3+ iron(III)An older system uses the endings –ic for the ion of higher charge and –ous for the ion of lower charge. e.g. Fe2+ ferrous Fe3+ ferric
  32. 32. Monoatomic AnionsMonoatomic anions are named by adding the suffix -ide to the stem of the name of the nonmetal from which they are derived. e.g. N3- nitride O2- oxide H- hydride S2- sulfide F- fluoride Se2- selenide Cl- chloride Te2- telluride Br- bromide I- iodide
  33. 33. Multiple Polyatomic IonsWhen a nonmetal forms two oxoanions, the suffix -ate is used for the anion with the larger number of oxygen atoms. The suffix –ite is used for the anion containing the fewest number of oxygens.When a nonmetal forms more than 2 oxoanions, the prefixes per- (largest number of oxygen atoms) and hypo- (fewest oxygen atoms) are used.
  34. 34. Ionic CompoundsThe name of an ionic compound has two parts: the first names the cation and the second names the anion (same order in which the ions appear).In naming compounds containing transition metals, the charge is indicated by a Roman numeral: Cr(NO3)3 chromium(II) nitrate SnCl2 tin(II) chloride
  35. 35. Binary Molecular CompoundsWhen two nonmetals combine, the product is often a binary molecular compound (covalent compound).There is no simple way to deduce the formulas of covalent compounds; however, there is a systematic way of naming them
  36. 36. Nomenclature of Covalent CompoundsI. The first word gives the name of the element that appears first in the formula; a greek prefix (2 = di, 3 = tri, 4 = tetra, etc) is used to show how many atoms of that element are presentII. The second word consists of the appropriate greek prefix indicating how many atoms of the second element are present; the stem of the name of the second element and the ending –ide.
  37. 37. ExamplesN2O5NO2PCl3Cl2O7
  38. 38. Common Covalent CompoundsH2O waterH2O2 hydrogen peroxideNH3 ammoniaC2H2 acetylenePH3 phosphineNO nitric oxideCH4 methane
  39. 39. AcidsSome covalent compounds containing H atoms ionize in water to form H+ ions.Such compounds are called acids. Bases ionize to form OH-.An example of an acid is hydrogen chloride (HCl). In water, HCl ionizes to H+ + Cl-. In solution, it is referred to as hydrochloric acid.
  40. 40. Acids (Cont’d)Most acids contain oxygen in addition to hydrogen. Such acids are called oxoacids. Two of the most common are sulfuric (H2SO4) and nitric (HNO3).The names of the oxoacids are derived from the names of the corresponding oxoanions. The –ate ending of the anion is replaced by –ic in the acid and the –ite is replaced by –ous.