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CHAPTER 2

CHAPTER 2. ATOMS, MOLECULES AND IONS. Contents. The Atomic Theory of Matter The Discovery of Atomic Structure The Modern View of Atomic Structure Atomic weights The Periodic Table Chemical Substances: Formulae and Names Ions and Ionic Compounds Naming Inorganic Compounds

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CHAPTER 2

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  1. CHAPTER 2 ATOMS, MOLECULES AND IONS

  2. Contents • The Atomic Theory of Matter • The Discovery of Atomic Structure • The Modern View of Atomic Structure • Atomic weights • The Periodic Table • Chemical Substances: Formulae and Names • Ions and Ionic Compounds • Naming Inorganic Compounds • Chemical Reactions

  3. ATOMIC THEORY OF MATTER (DALTON'S THEORY) • Law of definite proportions led to theory that all matter made up of atoms. • Atoms- basic building blocks and don't change when react with other atoms. • Element- describes matter composed of only one type of atom. • Compound- combination of atoms in specific proportions. • Chemical reaction- atoms exchange partners producing other compounds.

  4. LAW OF MULTIPLE PROPORTIONS • Some elements can form more than one compound when they react together (C & O: CO and CO2; N & O: N2O, NO, NO2, etc.). Dalton’s law predicted that the mass proportions should be proportional. Experiment confirmed this leading to this law. • Law of multiple proportions: when two elements form more than one compound, the ratio of the masses in one compound divided by the ratio of these masses in the other compound gives a ratio of small whole numbers. E.g. There are three binary compounds that form between barium and nitrogen. There was 4.9021 g , 9.8050 g and 14.7060 g of Ba per 1.0000 g N in the three compounds. Show that these compounds obey the law of multiple proportions.

  5. The Discovery of Atomic Structure • Thompson - Cathode rays. • Milliken - Oil drops. • Rutherford - backscattering -particles. • Radioactivity, the spontaneous emission of radiation from an atom led to the discovery of -, -, and -rays.

  6. The Modern View of Atomic Structure • Building Blocks • protons have a + charge, mp = 1.67 x 10-27kg • neutrons are neutral, mn = mp • electrons have a  charge, me = mp/1835 • Neutral atom: # electrons = # protons • Mass of atom is found by adding the mass of protons and neutrons • Protons identify the element (# protons called the atomic number, Z). • Isotopes have varying numbers of neutrons,

  7. Atomic and Molecular Weights Atomic mass scale: • A relative mass scale with the mass of the isotope is defined as exactly 12 amu (daltons) and used to determine the relative mass of all elements. • 1 amu = 1.6605x1024 g • Mass of other atoms reported relative to this. Average Atomic Masses • Reported mass of carbon and other elements on periodic table do not correspond to the expected value since they are weighted averages from all of the isotopes. where f1 = the fractional abundance of isotope 1 and AM = the atomic mass. E.g. Determine the atomic mass of boron if the masses of the two isotopes are 10.013 amu and 11.009 amu and the fractional abundances are 0.1978 and 0.8022, respectively. E.g. 2 Using the periodic table, determine the fractional abundance of the 35Cl and 37Cl isotopes if their masses are 34.969 and 36.966, respectively.

  8. Periodic Table • Elements in a column have similar reactivity. • Blue elements are semi-metals (metalloids). • Elements to left of blue area are metals. • Elements to right of blue area are non-metals. • 1A = Alkali metals • 2A = Alkaline earth metals • 6A =Chalcogens • 7A = Halogens • 8 A = Noble gases

  9. BONDING • An electron from each atom strongly interacts to form a bond. • Bonding can be either: • Ionic: Occurs between metal atoms or between metal and nonmetal atoms. • Covalent: occurs between nonmetal atoms and forms molecules. • Covalent: sharing of electrons occurs. • Ionic: electron(s) is (are) transferred from one atom to the other to produce • a positively charged substance (cation) and • a negatively charged substance (anion). • Ions bound by electrostatic attractions. • The formula of the ionic compound can be determined from the charge on the cation and the anion (even with a polyatomic ion): • Mg2+ and Cl form an ionic compound with the formula: MgCl2. • Fe3+ and O2 form an ionic compound with the formula: Fe2O3. • Fe2+ and form a compound with the formula: Fe(NO3)2

  10. NAMING INORGANIC COMPOUNDS IUPAC and Common Usage Methods of naming used for both organic and inorganic compounds. Organic compounds = predominantly carbon containing compounds. Inorganic compounds = all others Binary ionic compounds: compounds having a cation (a metal) and an anion (one of the main group anions). Cations: • Accepted method: Roman numeral in parentheses to indicate the charge. Not necessary when only one ionic charge possible. See Table 2.2. • Common method: Remove ending and add either -ous or -ic to • Latin form of the element used instead on some: Stannous, Stannic; Ferrous, Ferric; Cuprous, Cupric. Anions: • For monoatomic anions add -ide to the stem. Fluor-, Ox-, Nitr-, Sulf-, etc.

  11. NOMENCLATURE 2 Binary molecular compounds: • More cationlike appears first; -ide ending is placed on the anionlike substance. • The more cationlike element appears to the left of or below the other element in the periodic table • With hydrogen as one of the two, hydrogen is first; place an -ide on the other element. • If a compound contains a group VI or VII element, an –ide ending is added to it. Numerical prefixes are used especially with the element listed second. Mono- not used with the first element. • When oxygen with fluorine, oxygen first in name: E.g. Oxygen difluoride = OF2.

  12. Compounds with Polyatomic ions • Name these the same way that ionic binary compounds are named replacing the name of the anion (in a few cases cation) with the name of the polyatomic ion. • Oxy anions (anions containing at least 1 oxygen atom): difficult because there are often several possible anions. • Add -ate to the stem (carbonate, sulfate, nitrate, etc); but • Many oxyanions can have other proportions of oxygen e.g. , These can be named by adding: • per- x -ate for most oxygen • -ite for a small number of oxygen • hypo- x -ite gives least • E.g. oxychlorides. • If ion contains H, write: “hydrogen” + name of ion without hydrogen, e.g. hydrogen sulfate. • A prefix of mono- or di- to indicate the number of hydrogens may be needed. E.g. dihydrogen phosphate- .

  13. Oxyacids • Binary compounds which are acids: Prefix hydro- added to the anion part; change the -ide ending to -ic. • Hydrogen chloride becomes hydrochloric acid. • The ending -ite is changed to -ous and -ate to -ic.

  14. CHEMICAL EQUATIONS • Chemical reaction indicates the reactants and products. • Reactants on left; Products on right • 2H2(g) + O2(g)  2H2O(l); • 4Fe(s) + 3O2(g)  2Fe2O3(s) • NaCl(aq) + AgNO3(aq) AgCl(s) + NaNO3(aq) • Total number of each atom should be the same on each side of the reaction. Law of Conservation of mass should guide us. • Steps to Balancing an Equation (simple): • Write an unbalanced equation with all reactants and products: • Balance by changing co-efficients • Reduce coefficients E.g. the combustion of methanol produces water and carbon dioxide: CH3OH(l) + O2(g)  CO2(g) + H2O(l) E.g. Balance: • P4 + N2O(g)  P4O6(s) + N2(g) • P2O5 + H2O  H3PO4

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