Atoms, Molecules and Ions

1. Atoms, Molecules and Ions Lecture 2 郭修伯

2. Atoms, Molecules and Ions • 2.1 The Early History of Chemistry • 2.2 Fundamental Chemical Laws • 2.3 Dalton’s Atomic Theory • 2.4 Cannizzaro’s Interpretation • 2.5 Early Experiments to Characterize the Atom • 2.6 The Modern View of Atomic Structure • 2.7 Molecules and Ions • 2.8 An Introduction to the Periodic Table • 2.9 Naming Simple Compounds

3. Early history of chemistry • ~ 400 B.C., the Greeks: • all matter was composed of 4 fundamental substances: fire, earth, water and air • Fundamentals of modern chemistry: • laid in the 16th century • development of systematic metallurgy (extraction of metals from ores) • Irish scientist Robert Boyle (1627-1691) • first “chemist” who performed truly quantitative experiments: the relationship between P and V of gases (the book “the Sceptical Chemist” in 1661)

4. Early history of chemistry • Combustion evokes intense interests in 17 & 18 centuries • German chemist Georg Stahl (1660-1734) : a substance burning in a closed container eventually stopped burning because the air in the container became saturated with “phlogiston”. • English scientist Joseph Priestley (1733-1804): discovered oxygen gas (originally called “dephlogisticated air”)

5. Priestley Medal Source: Roald Hoffman, Cornell University

6. Fundamental chemical laws • A French chemist Antoine Lavoisier (1743-1794) • named “oxygen” • carefully weighted the reactant and products of various reactions • “mass is neither created nor destroyed”  LAW of CONSERVATION of MASS • published “Elementary Treatise on Chemistry” in 1789, the first modern chemistry textbook

7. Fundamental chemical laws • An English schoolteacher John Dalton (1766-1844) • Published “A New System of Chemical Philosophy” in 1808 • Atomic theory: • Each element is made up of tiny particles called atoms • The atoms of a given element are identical; the atoms of different elements are different in some fundamental way(s) • Chemical compounds are formed when atoms combine with each other. A given compound always has the same relative numbers and types of atoms • Chemical reactions involve reorganization of the atoms – changes in the way they are bound together. The atoms themselves are not changed in a chemical reaction.

8. Fundamental chemical laws • Determining absolute formulas for compounds • French chemist Joseph Gay-Lussac (1778-1850) • Italian chemist Amedeo Avogadro (1776-1856) • Avogadro’s hypothesis: at the same T and P, equal V of different gases contain the same number of particles. • No general agreement exists concerning the formula for elements • Chaos reigned in the first half of the nineteenth century • Mess till the leadership of the Italian chemist Stanislao Cannizzaro (1826-1910)

9. Cannizzaro’s Interpretation • Italian chemist Stanislao Cannizzaro: • Guided by two beliefs: • Compounds contained whole numbers of atoms as Dalton postulated • Avogadro’s hypothesis was correct-equal volumes of gases under the same conditions contain the same number of molecules • Relative atomic mass to hydrogen (arbitrary assigned to be 2) was determined

10. Early Experiments to Characterize the Atom • English physicist J.J. Thomson (1856-1940) studied electrical discharges in partially evacuated tube cathode-ray tubes during the period from 1898 to 1903

11. Cathode-ray tubes • Thomson postulated that the ray was a stream of negatively charged particles, now called electrons. • Deflection of the beam in a magnetic field

12. Thomson’s Plum pudding model • An atom consisted of a diffuse cloud of positive charge with the negative electrons embedded randomly in it.

13. Radioactivity • 3 types of radioactive emission (early 20th century): • gamma (γ) rays, beta (β) particles, and alpha (α) particles. • A γray is high-energy “light”; a β particle is a high-speed electron; and an α particle has a 2+ charge (i.e., a charge twice that of the electron and with the opposite sign). The mass of an α particle is 7300 times that of the electron.

14. Nuclear atom • In 1911 Ernest Rutherford carried out an experiment to test Thomson’s plum pudding model.

15. Nuclear atom • Most of the αparticles passed straight through, many of the particles were deflected at large angles Thomson’s plum pudding model  X Actual result  proposed model

16. The simplest view of the atom is that it consists of a tiny nucleus with a diameter of about 10-13 cm and electrons that move about the nucleus at an average distance of about 10-8 cm away from it The Modern View of Atomic Structure: An Introduction

18. “If all atoms are composed of these same components, why do different atoms have different chemical properties?” • The number and the arrangement of the electrons • The electrons constitute most of the atomic volume • The number of electrons possessed by a given atom greatly affects its ability to interact with other atoms

19. Isotopes • Atoms with the same number of protons but different numbers of neutrons Sodium-23 is the only naturally occurring form of sodium. Sodium-24 does not occur naturally but can be made artificially.

20. A • where the atomic number Z (number of protons) is written as a subscript and the mass number A (the total number of protons and neutrons) is written as a superscript. (The particular atom represented here is called “sodium-23.” It has 11 electrons, 11 protons, and 12 neutrons.) Z

21. Molecules and Ions • Chemical bonds • Atoms have electrons and these electrons participate in the bonding of one atom to another • The forces that hold atoms together in compounds • Covalent bonds: bonds formed by sharing electrons • Molecule • Collection of atoms • The simplest method to present a molecule is the chemical formula • CO2: each molecule contains 1 atom of carbon and 2 atoms of oxygen

22. Structural formula • Solid lines: in the plane of the paper • Dashed lines: behind the plane of the paper • Wedges: in front of the plane of the page

23. Space-filling model shows both the relative sizes of the atoms in the molecule and their spatial relationships Ball-and-stick models Methane

24. Ion • An atom or group of atoms that has a net positive or negative charge cation

25. Because anions and cations have opposite charges, they attract each other. This force of attraction between oppositely charged ions is called ionic bonding. anion

26. Ionic solid (also named salt) • A solid consisting of oppositely charged ions • can consist of simple ions, as in sodium chloride (NaCl), or of polyatomic (many-atom) ions as ammonium nitrate (NH4NO3)

27. Periodic Table

28. Periodic Table • Shows all the known elements • the letter is the symbol for that element • the number is the atomic number (number of protons) for that element • Horizontal rows in the table are called periods: • first period: contains H and He • second period: contains elements Li through Ne

29. Periodic Table • Most of the elements are metals • Shown in the center of the periodic table • Efficient conduction of heat and electricity • Malleability (can be hammered into thin sheets) • Ductility (can be pulled into wire) • (often) a lustrous appearance • Relatively few nonmetals appear in the upper right-hand corner of the table • tend to gain electrons to form anions in the reactions with metals • often bond to each other by forming covalent bonds

30. Periodic Table • Elements in the same vertical columns (called groups or families) have similar chemical properties • Ex: alkali metals (Group 1A) – Li, Na, K, Rb, Cs and Fr - are very active elements that readily form ions with a 1+ charge when they react with nonmetals • Ex: alkali earth metals (Group 2A) – Be, Mg, Ca, Sr, Ba and Ra - form ions with a 2+ charge when they react with nonmetals • Ex: halogens (Group 7A) – F, Cl, Br, I and At - form diatomic molecules. React with metals to form salts containing ions with a 1- charge. • Ex: Noble gases (Group 8A) – He, Ne, Ar, Kr, Xe and Rn - form monatomic gases and have little chemical reactivity

31. Naming Simple Compounds: Binary compounds • Ionic compound (Type I) • The metal involved forms only a single type of cation • The cation is always named first and the anion second • A cation takes its name from the name of the element • Ex: Na+ is called sodium • An anion is named by taking the first part of the element name and adding – ide. • Ex: Cl- ion is called chloride.

33. Ionic compound (Type II) • Most* commonly occurs for compounds containing transition metals, which often form more than one cation • Ex: The compound FeCl2 contains Fe2+ ions, and the compound FeCl3 contains Fe3+ ions. • Naming system (I): • iron (II) chloride (FeCl2) and iron (III) chloride (FeCl3) • Roman numeral indicates the charge of the cation • Naming system (II): • Ferrous chloride (FeCl2) and ferric chloride (FeCl3) • The ion with the higher charge has a name ending in –ic, and the one with the lower charge has a name ending in –ous • *common transition metals that form only one ion are zinc (Zn2+) and silver (Ag+)

35. Naming binary ionic compounds

36. Example 2.2 • Give the systematic name of each of the following compounds. a. CoBr2 b. CaCl2 c. Al2O3 d. CrCl3 • Solution

37. Polyatomic ions are assigned special names that must be memorized to name the compounds containing them. • Ex: the compound ammonium nitrate (NH4NO3) contains the polyatomic ions NH4+ and NO3- • Oxyanions: several series of anions contain an atom of a given element and different numbers of oxygen atoms • When there are two members in such a series, the name of the one with the smaller number of oxygen atoms ends in –ite, and the name of the one with the larger number ends in –ate • Ex:sulfite (SO32-) and sulfate (SO42-). • When more than two oxyanions make up a series, hypo- ( less than) and per-(more than) are used as prefixed to name the members of the series with the fewest and the most oxygen atoms, respectively.

38. Ionic compounds with Polyatomic Ions

39. Binary covalent compound (Type III) • Containing two nonmetals (i.e., do not contain ions) • Naming rules: • The first element in the formula is named first, using the full element name. • The second element is named as if it were an anion. • Prefixes are used to denote the numbers of atoms present • The prefix mono- is never used for naming the first element. • Ex: CO is carbon monoxide, not monocarbon monoxide

40. avoid awkward pronunciations, we often drop the final o or a of the prefix when the element begins with a vowel. For example, N2O4 is called dinitrogen tetroxide, not dinitrogen tetraoxide; and CO is called carbon monoxide, not carbon monooxide.

41. Naming compounds

42. Naming compounds

43. Example 2.3 • Give the systematic name of each of the following compounds. a. Na2SO4 b. KH2PO4 c. Fe(NO3)3 d. Mn(OH)2e. Na2SO3 f. Na2CO3 g. NaHCO3 h. CsClO4i. NaOCl j. Na2SeO4 k.KBrO3 • Solution

45. Example 2.4 • Given the following systematic names, write the formula for each compound. a. ammonium sulfate b. vanadium (V) fluoride c. dioxygen difluoride d. rubidium peroxide e. gallium oxide

47. Naming Acids • Acid: • When dissolved in water, certain molecules produce a solution containing free H+ ions (protons). • The rules for naming acids depend on whether the anion contains oxygen • If the anion does not contain oxygen, the acid is named with the prefix hydro- and the suffix –ic. • Ex: HCl called hydrochloric acid, HCN called hydrocyanic acid and H2S called hydrosulfuric acid • When the anion contains oxygen, the acid name is formed from the root name of the anion with a suffix of –ic or –ous.

48. Anion containing oxygen • If the anion name ends in –ate, the acid name ends with –ic (or sometimes –ric). • Ex: H2SO4 contains the sulfate anion (SO42-) and is called sulfuric acid; H3PO4 contains the phosphate anion (PO43-) and is called phosphoric acid; and HC2H3O2 contains the acetate ion (C2H3O2-) and is called acetic acid • If the anion has an –ite ending, the acid name ends with –ous. • Ex: H2SO3, which contains sulfite (SO3 2-), is named sulfurous acid; and HNO2, which contains nitrite (NO2-), is named nitrous acid.