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AP Chemistry Chapter 2

AP Chemistry Chapter 2. Section 2.1 Timeline for Chemistry Prior to 1000 BC-natural ores-ornaments and weapons, embalming fluids 400 BC- Greeks : fire,earth,water and air. Democritus and Leucippos : matter made of small indivisible parts (atomos)

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AP Chemistry Chapter 2

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  1. AP Chemistry Chapter 2 Section 2.1 Timeline for Chemistry Prior to 1000 BC-natural ores-ornaments and weapons, embalming fluids 400 BC-Greeks: fire,earth,water and air

  2. Democritus and Leucippos: matter made of small indivisible parts (atomos) Next 2000 years-Alchemy: pseudoscience: metal into gold. Made advances in experimental techniques and equipment 1500’s-Bauer: metallurgy-extraction of metals from ores. Paracelsus: medicinal applications of minerals 1600’s-Boyle: first chemist with quantitative data (pressure and volume) Greek system left and entered era of elements

  3. 1700-Stahl: phlogiston substance released during burning -Oxygen disovered by Priestley or dephlogistonated air. -CO2, N2, H2 and O2 discovered -Lavoisier-combustion with oxygen (not phlogiston). Created law of conservation of mass. Created balance to 0.005g. • 1800-Proust-law of definite proportions: compounds contain same proportions of elements by mass.

  4. Section 2.2 Fundamental Chemical Laws • Law of multiple proportions: when two elements form a series of compounds, the ratio of the masses of the second element that combine with 1 gram of the first element can always be reduced to small whole numbers. Example: The following data were collected for several compounds of nitrogen and 1 g of oxygen. Compound A 1.750 g Compound B 0.8750 g Compound C 0.4375 g A/B = 1.750/0.875=2/1; B/C=0.875/0.4375=2/1; A/C=1.750/0.4375=4/1

  5. Section 2.3 • Dalton’s Atomic Theory • 1808 • 1. each element is made atoms • 2. Atoms are the same for one type and different for another type • 3. Compounds formed when atoms combine with the same relative numbers and types of atoms • 4. atoms do not change in reactions, they just get reorganized. • Dalton created a table of atomic masses (it was incorrect) but it lead the way to organizing the elements.

  6. Gay-Lussac 1809-1811 • Experiment of gases showed that 1 volume of Oxygen gas reacted with 2 volumes of hydrogen gas to form 2 volumes of gaseous water. • Avogadro’s Hypothesis explained that the volume of gas is determined by number of molecules and not the size. • This then showed that water is H2O and not OH as Dalton showed. • This hypothesis was not accepted for 50 years…

  7. Section 2.4 JJ Thomson • Electron discovered using the cathode ray tube during experiments from 1898-1903 • Plum pudding model for atom: positive sphere with negative charges embedded throughout

  8. Robert Millikan • 1909 Millikan charged oil drops in an experiment that allowed him to calculate the charge to mass ratio of an electron. The mass of an electron is 9.11 x 10-31 kg.

  9. Henri Becquerel • 1896 Becquerel discovered that images could be made on photographic plates by the element uranium. He called that radioactivity or spontaneous emission of radiation. • Further studies revealed the radiation as gamma rays, beta particles and alpha particles.

  10. Ernest Rutherford • 1911 Gold Foil experiment done by Rutherford showed that a stream of alpha particles was unexpectedly deflected. • Plum pudding model for the atom was abandoned and a new model developed with a positive center (nucleus) and negative charges moving about the center.

  11. Section 2.5Modern View of the Atomic Structure • Nucleus contains positively charged protons (of which determines the type of atom) and neutrally charged particles called neutrons. (each have about an equal mass to the other) • Electrons reside outside the atom at certain energy levels. • The nucleus is very small compared to the rest of the atom and yet it is very dense containing almost all the atom’s mass.

  12. The arrangement of the electrons gives each element different chemical properties • Isotopes are atoms with the same number of protons but different number of neutrons (different mass) • Atomic symbol is X. Atomic number is the letter Z and mass number is A. The atomic symbol is written as follows: AZX • Example: Write the symbol for Fluorine

  13. Section 2.6 Molecules and Ions • Chemical Bonds: force that holds atoms together • Covalent bonds: share electrons • Ionic bonds: give up or take on electrons Chemical Formulas: CO2 Structural Formulas: bonds shown by lines H-O-H Space filling models Ball and stick models

  14. Ions • An atom or group of atoms that has a net positive or negative charge. • Example: NaCl (table salt) • Na + Cl NaCl • Ionic solid or salt can consist of simple ions or polyatomic ions

  15. Section 2.7 Periodic Table • Periodic table: chart that shows all the known elements organized by increasing atomic number from left to right. • Atomic number= number of protons and also the number of electrons in a neutral atom. • Most elements are metals. Conductors of heat and electricity, malleable, ductile and lustrous. Tend to lose electrons and form positive ions

  16. Nonmetals appear in the upper right-hand corner of the table. They tend to gain electrons and form negative ions when bonding to a metal. • Nonmetals bond with other nonmetals and form covalent bonds.

  17. Groups or families: similar chemical properties and same number of valence electrons • Alkali metals: Group 1: form 1+ charges, very reactive to form ionic bonds with nonmetals. • Alkaline earth metals: Group 2: form 2+ charges when reacting with nonmetals. • Halogens: Group 7A or 17: form salts with 1- charges with metals • Noble Gases: Group 8A or 18:little reactivity, have full octet in valence shell.

  18. Periods: horizontal rows that show trends in atoms. Indicate number of energy levels in an atom. • Symbols for elements consist of 1, 2 or 3 letter designations: • Carbon C • Calcium Ca • Unnilpentium Unp 105

  19. Section 2.8 Naming Single Compounds • Binary ionic Compounds Type I: compounds with two elements • Cation always named first, anion second • Cation keeps its name • Anion changes ending of name to –ide • Example NaCl • KI • CaS • CsBr • MgO • AlCl3

  20. Binary Ionic Compounds Type II • If metal has more than one type of cation use roman numerals • The ion with the higher charge ends in –ic, and the one with the lower charge has the ending –ous • Example FeCl3 and FeCl2 • CuCl • HgO

  21. Ionic compounds with polyatomic ions • Consider the compound NH4NO3. Memorize the list on page 63 of common polyatomic ions. • The list shows anions or oxyanions and the one with the smaller number of oxygen atoms ends in –ite and larger gets -ate. • Example SO32- and SO42-

  22. Binary Compounds (Type III; covalent-contain two nonmetals) • The first element in the formula is named first, using the full element name • The second element is named with -ide • Prefixes denotes the numbers of atoms present • Mono is never used for first element • Example: N2O PCl5 • Water and ammonia are referred to by their common name, not their systematic name.

  23. Prefixes used to indicate number in chemical names • Prefix number • Mono- 1 • Di- 2 • Tri- 3 • Tetra- 4 • Penta- 5 • Hexa- 6 • Hepta- 7 • Octa- 8 • Nona- 9 • Deca- 10

  24. Formulas from names • Given the name Calcium Hydroxide write the formula • Ca(OH)2 • Example: iron(II)oxide

  25. Acids: rules for naming • Rule for naming acids depends on whether or not the anion contains oxygen. • Anion no oxygen: prefix Hydro- and the suffix –ic • Example: HCl hydrochloric acid • Example: H2S • Anion with oxygen: • -ate switch to –ic • Example: contains sulfate (SO4) H2SO4 • -ite switch to –ous • Example: contains sulfite(SO3) H2SO3

  26. Acids with Chlorine • HClO4 Perchlorate perchloric acid • HClO3 chlorate • HClO2 chlorite chlorous acid • HClO hypochlorite

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