1 / 56

CHAPTER 2

CHAPTER 2. Atoms, Molecules, and Ions. Law of Mass Conservation. Mass is neither created nor destroyed in chemical reactions. Also known as Law of Conservation of Matter. 243.7 g. 243.7 g. 159.7 g. 84 g. 111.7 g. 132 g. Law of Definite Proportions.

judith-vega
Download Presentation

CHAPTER 2

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. CHAPTER 2 • Atoms, Molecules, and Ions

  2. Law of Mass Conservation • Mass is neither created nor destroyed in chemical reactions. • Also known as Law of Conservation of Matter 243.7 g 243.7 g 159.7 g 84 g 111.7 g 132 g

  3. Law of Definite Proportions • Different samples of a pure chemical substance always contain the same proportion of elements by mass. • Elements combine in specific proportions to form compounds. • Different samples of a pure chemical substance always have identical chemical properties.

  4. Dalton’s Atomic Theory • An element is composed of extremely small, indivisible particles called atoms. • All atoms of a given element have identical properties that differ from those of other elements. • Atoms cannot be created, destroyed, or transformed into atoms of another element. • Compounds are formed when atoms of different elements combine with one another in small whole-number ratios. • The relative numbers and kinds of atoms are constant in a given compound.

  5. Law of Multiple Proportions • Elements can combine in different ways to form different substances, whose mass ratios are in whole number multiples of each other. • Elements combine in whole number ratios to form compounds.

  6. Structure of Atoms • Atoms composed of protons, neutrons and electrons • Molecules are composed of atoms • Elements are composed of atoms – these may or may not be molecules.

  7. Fundamental Particles • Three fundamental particles make up atoms. Particle Mass Charge Electron (e–) 0.00054858 –1 Proton (p, p+) 1.0073 +1 Neutron (n, n0) 1.0087 0 Know the relative size, charge, and nuclear location of each particle.

  8. Discovery and Properties of Electrons • Humphrey Davy (early 1800’s) passed electricity through molten salts and separated the elements (he discovered most of the IA and IIA metals). • Compounds are held together by electrical forces. • Michael Faraday found that the amount of metal formed during electrolysis is proportional to current passed through the solution. • Metals have distinct oxidation states (positive charges) associated with each element.

  9. Discovery and Properties of Electrons • Cathode ray tubes have been known since the 1800’s (same as still found in most TV or computer monitors). • The glow is caused by “rays” emitted from cathode (– end) to anode (+ end). • Cathode rays must be negatively charged!

  10. Discovery and Properties of Electrons • J.J. Thomson added two adjustable voltage electrodes into the experiment • This allowed the charge to mass ratio of electrons to be measured. • e–/m = –1.75881  108 coulomb/g of e– • Thomson named cathode rays electrons.

  11. Discovery and Properties of Electrons • Robert A. Millikan used an oil drop experiment to determined the charge and mass of the electron. 1 electron charge = –1.60219  10-19 coulomb mass of 1 electron = 9.11  10-28 g FYI. Millikan was the first American to win a Nobel Prize – which was awarded for this work on electrons.

  12. Canal Rays and Protons • Goldstein (1886) postulated that if there were cathode rays, that there would also be streams of positively charged particles going to opposite direction in these cathode rays. • These “Canal Rays” would be positively charged. • Goldstein postulated the existence of a “proton”. • Canal rays have variable masses unlike the particles in cathode rays. • Canal rays are composed of ionized gas molecules and their mass depends upon the mol.wt. of the gas used in the tube.

  13. Rutherford and the Atom • Ernest Rutherford (1910) developed the basic picture of atom by following Geiger & Marsden’s experiment on -particle scattering from thin Au foils

  14. Rutherford and the Atom Rutherford concluded that: • the atom is mostly empty space • the atom has a very small, dense center called the nucleus • nearly all the atom’s mass is found in the nucleus • nuclear diameter is 1/10,000 to 1/100,000 times less than the atom’s radius • nuclear density is 1015g/mL (this is equivalent to 3.72 x 109 tons/in3 or equal to the density of a neutron star or black hole)

  15. Atomic Number • Atomic number = number of protons in the nucleus • The number of protons determines the identity of the element. • The atomic number is sometimes designated as the letter Z. This number is seldom written with the chemical symbol for the element (it is redundant). The atomic number is equal to the number of electrons in a neutral atom of that element.

  16. Neutrons • James Chadwick analyzed evidence from -particle scattering off Be and realized a large neutral subatomic particle was being knocked out of the nuclei. • This led to the proof of the massive neutral particles in the nucleus called “neutrons”

  17. Mass Number and Isotopes • H.G. J. Moseley recognized that atomic number is the defining difference between elements. • He discovered that it was possible to calculate the atomic number of an element based upon its atomic spectrum. • This validated Mendeleev’s periodic law and justified position of K and Ar • His discovery is still used to identify new elements by their atomic spectra. • FYI. Moseley was killed in World War I at Gallipoli.

  18. Mass Number and Isotopes • The mass number of an element (specifically the isotope of the element) is given the symbol A. • A is sum of protons + neutrons Specific isotopes of an element (E) can be shown as A = mass number Z = atomic number N = neutrons in that isotope

  19. Mass Number & Isotopes • When representing an element, N is very rarely (almost never) shown. • Isotopes can be written as More often the isotopes are written as

  20. H D T Mass Number and Isotopes • Hydrogen is treated differently With the exception of H, no difference is seen in the relative reactivity of the isotopes of an element. Plants do differentiate isotopes and this is used in SIRA (stable isotope ratio analysis) for identification of both plants and animals.

  21. Atomic Weights • This is the lower number on periodic chart • This is weighted average of the masses of the constituent isotopes naturally occurring.

  22. The Atomic Weight Scale and Atomic Weights • IUPAC defined the mass of 12C as 12 amu exactly. • 1 amu = (1/12) mass of 12C • By definition, each 1 amu = 1 g • Remember, new terminology • 1 amu = 1 Dalton or 1 Da • The atomic weight of any element is defined as the weighted average of its naturally occurring isotopes (as found on earth).

  23. The Atomic Weight Scale and Atomic Weights • Example: Naturally occurring Cu consists of 2 isotopes. It is 69.1% 63Cu with a mass of 62.9 amu, and 30.9% 65Cu, which has a mass of 64.9 amu. Calculate the at. wt. of Cu to one decimal place.

  24. The Atomic Weight Scale and Atomic Weights • Example: Naturally occurring Cr consists of four isotopes. It is 4.31% 50Cr, mass = 49.946 amu; 83.76% 52Cr, mass = 51.941 amu; 9.55% 53Cr, mass = 52.941 amu; and 2.38% 54Cr, mass = 53.939 amu. Calculate the at. wt. of chromium. You may choose to do this as a single calculator entry.

  25. The Atomic Weight Scale and Atomic Weights • Example: The atomic number of boron is 10.811 amu. The masses of the two naturally occurring isotopes 10B and 11B, are 10.013 and 11.009 amu, respectively. Calculate the fraction and percentage of each isotope. • Remember 1 = x + (1-x)

  26. Mixtures Pure Substances Heterogeneous Homogeneous ChemicalCompounds Elements Compounds and Mixtures Matter

  27. Compounds and Mixtures • Compounds • substances composed of two or more elements in a definite ratio by mass • can be decomposed into the constituent elements (usually by chemical reaction)

  28. Compounds and Mixtures • Mixtures • composed of two or more substances in variable ratios • can be separated into components by simple physical or chemical methods homogeneous mixtures– cannot be separated by physical inspection heterogeneous mixtures– usually can be separated by simple physical means

  29. Chemical Bonding • Chemical bonding is the attractive forces that hold atoms together. • Covalent bonds result from the sharing of two (usually) or more electrons between 2 atoms. • Ionic bonds are electrostatic forces describing the attraction between positive and negative ions.

  30. Chemical Formulas • This shows the ratio of the atoms of the elements present in the molecule or compound. • Elements • He, Au, Na – monatomic elements • O2, H2, Cl2– diatomic molecules • O3, S4, P8– polyatomic molecules • Allotropes • O2 and O3, S4 and S8, Cgraphite and Cdiamond • Compounds • H2O, C12H22O11 - compounds

  31. Molecular Representations • A great deal of information can be conveyed with a chemical formula. • Representations of ethanol (ethyl alcohol) C2H6O CH3CH2OH Chemical formula Condensed formula Structural formula Ball and stick model Space filling model

  32. Formulas for Compounds • Compound Contains • HCl 1 H atom and 1 Cl atom • H2O 2 H atoms and 1 O atom • NH3 3 H atoms and 1 N atom • C3H8 3 C atoms and 8 H atoms

  33. Ions & Ionic Compounds • Ions are atoms or groups of atoms that are charged. • Polyatomic ions are treated in a special way in formulas. • There are two types of ions. • Cations– positive ions • Cations have one or more electrons fewer than the neutral atom (or group of atoms) should have. • Anions– negative ions • Anions one or more electrons more than the neutral atom (or group of atoms) should have.

  34. Ionic or Covalent? • Is the compound ionic or covalent? • Generally covalent bonds consist of nonmetals (usually) sharing electrons. • Ionic bonds exist between metals and nonmetals (usually) – the metal gave up its electron to the nonmetal. • Polyatomic ions have covalent bonds but also carry a charge.

  35. . .. .. .. . . . _ _ _ : : : : : : : : : : : : Na Cl Cl Cl Cl Cl Cl .. .. .. .. .. .. . Mg . Ionic Bonding (or NaCl) + Na+ + (or MgCl2) + Mg 2+ +

  36. Formulas for Compounds Containing Polyatomic Ions • Compound Contains • NH4NO3 2 N atoms, 4 H atoms, and 3 O atoms • Ca(OH)2 1 Ca atom, 2 O atoms, and 2 H atom • Al2(SO4)3 2 Al atoms, 3 S atoms, and 12 O atoms • (NH4)3PO4 3 N atoms, 12 H atoms, 1 P atom, and 4 O atoms • MgSO4·7 H2O 1 Mg atom, 1 S atom, 11 O atoms (4 + 7 O atoms) and 14 H atoms. • The “·” in this formula means that 7 water molecules are associated with the MgSO4.

  37. Ion Names and Formulas • Anions • Cl–chloride HCl hydrochloric acid • Br–bromide HBr hydrobromic acid • F– fluoride HF hydrofluoric acid • I–iodide HI hydroiodic acid • OH–hydroxide (strong baseapolyatomicanion) • O–2oxide • S–2sulfide You are expected to have these MEMORIZED! Know both the name and the charge of the ion.

  38. Ion Names and Formulas Polyatomic Anions • SO4–2sulfate H2SO4sulfuric acid • CO3–2 carbonate H2CO3carbonic acid • Some ions also have common or archaic names that must also be memorized. • HCO3–hydrogen carbonate or bicarbonate • PO4–3 phosphate or tribasic phosphate • HPO4–2hydrogen phosphate or dibasic phosphate • H2PO4–dihydrogen phosphate or monobasic phosphate

  39. Ion Names and Formulas Polyatomic Anions • CH3CO2– acetate CH3COOH acetic acid • (vinegar) • MnO4–permanganate • CrO4 –2chromate • Cr2O7 –2dichromate

  40. Ion Names and Formulas • The Halogen Oxides • All follow the same nomenclature pattern ClO4–perchlorate(perbromate……) ClO3–chlorate (bromate, iodate) ClO2–chlorite (bromite, iodite) ClO–hypochlorite (hypobromite, ……) The Corresponding Acids ClO4–perchlorateHClO4 perchloric acid ClO3–chlorate HClO3chloric acid ClO2– chlorite HClO2 chlorous acid ClO–hypochloriteHClO hypochlorous acid

  41. Ion Names and Formulas • Cations • NH4+ammonium • Most of the metal cations are named exactly the same as the element. • Na+sodium • Ca+2calcium • Al+3aluminum • You are expected to know these by memory.

  42. General Rules for Chemical Formulas • These are stable ions (they do not decompose). If the ion is monatomic (one element), the number of ions is indicated by a subscript If the ion is polyatomic: If there is only a single ion present in the compound, no subscript is necessary (one ion is understood) If there are multiple ions present within the compound, parentheses must be used with the appropriate subscript to indicate the number of ions present.

  43. Ionic Compounds • Ionic compounds have no overall charge. • This means that the total charge of the cations and anions must cancel out (equal zero). • NaCl sodium chloride (+1 and –1 = 0) • KOH potassium hydroxide (+1 and –1 = 0) • CaSO4 calcium sulfate (+2 and –2 = 0) • Al(OH)3 aluminum hydroxide (+3 and –3 = 0)

  44. Naming Compounds • By convention, chemical formulas are written with the most positive element listed first. • Binary ionic compounds. For metal ions you only need to indicate the charge on the ion (usually necessary for transition and post-transition metals). The number of that ion present is assumed. • Binary covalent compounds. For nonmetals, a prefix is only used for the cation (first element listed) if more than one atom is present. Since nonmetals may combine in numerous ratios, the anion portion of the name always is given with a numerical prefix.

  45. Ion Names and Formulas • Metals with “Latin” Symbols • Many metal ions also have archaic names. • You are required to know both names. • Iron, Fe • Fe+2iron(II) or ferrous ion • Fe+3iron(III) or ferric ion • Copper, Cu • Cu+copper(I) or cuprous • Cu+2copper(II) or cupric • Tin, Sn • Sn+2tin(II) orstannous • Sn+4tin(IV) orstannic

  46. Nomenclature of Inorganic Compounds • This is where the memorizing the ions helps! • binary compounds– only two elements(or ions) • metal + nonmetal = ionic compound • nonmetal + nonmetal = covalent compound • Chemical formulas of inorganic compounds are written with the more positive portion of the compound shown first and the negative portion shown last. • Common nomenclature is to ‘name” the less electronegative element (or ion) first (the one with the positive charge) and the more electronegative element (or ion) second.

  47. Nomenclature of Inorganic Compounds LiBr lithiumbromide MgCl2 magnesiumchloride lithiumsulfide Li2S Al2O3 aluminumoxide Na3P sodiumphosphide Mg3N2 magnesiumnitride

  48. Nomenclature of Inorganic Compounds • Transition metals typically have more than one common oxidation state (charge). • IUPAC - Roman numerals to indicate metal’s oxidation state • older system - add “ic” to element’s Latin root for higher oxidation state • older system - add “ous” to element’s Latin root for lower oxidation state • The older nomenclature is commonly encountered (biology, medicine, etc.) so you need to know both.

  49. Nomenclature of Inorganic Compounds IUPAC Archaic FeBr2 iron(II) bromide ferrous bromide FeBr3 iron(III) bromide ferric bromide SnO tin(II) oxide stannous oxide SnO2 tin(IV) oxide stannic oxide TiCl2 titanium(II) chloride titanous chloride TiCl3 titanium(III) chloride titanic chloride (none) TiCl4 titanium(IV) chloride

  50. Nomenclature of Inorganic Compounds • The binary compounds of hydrogen and a nonmetal are named as hydrogen (stem)ide. • The aqueous solutions of these compounds are acidic and are named as hydro(stem)ic acid. NameAcidic Solution HF hydrogen fluoride hydrofluoric acid HCl hydrogen chloride hydrochloric acid HBr hydrogen bromide hydrobromic acid H2S hydrogen sulfide hydrosulfuric acid

More Related