1. Chapter 2 Atoms, Molecules, and Ions

2. LAW OF CONSERVATION OF MASS Antoine Lavoisier (1743-1794) Carefully measured and provided a quantitative interpretation of the chemical reaction associated with combustion. Law: During a chemical change, the total mass remains constant.

3. LAW OF DEFINITE PROPORTIONS Joseph Proust (1754-1826) Law: Different samples of a pure compound always contain the same proportion of elements by mass. Also called Law of Constant Composition

4. LAW OF MULTIPLE PROPORTIONS John Dalton (1766-1844) Law: When two or more elements form more than one compound, the ratio of the masses of one element in these compounds for a fixed mass (i.e. 1 gram) of the other element is a small whole number.

5. ATOMIC THEORY OF MATTER (1808) John Dalton (1766-1844) Elements (matter) are composed of atoms The atoms of a given element are identical. Each element is characterized by the mass of its atoms. Compounds are formed when atoms of different elements combine with each other (Multiple Proportions).

6. ATOMIC THEORY OF MATTER (2) A given compound is a chemical combination of the same atoms in the same relative numbers. (Definite Proportions). A chemical reaction is the rearrangement of atoms leading to new compounds. Atoms are neither destroyed nor created in a chemical reaction (Conservation of Mass).

7. AVOGADRO’S HYPOTHESIS Amadeo Avogadro (1776-1856) Equal volumes of two gases at the same pressure and temperature contain the same number of particles (atoms or molecules).

8. ATOMIC STRUCTURE Subatomic Particles (Table 2.1) ELECTRON (cathode ray) –1899 J.J. Thomson: electron charge = -1 and e/m = - 1.76E+8 C/g –1909 R. Millikan: electron mass = 9.11E-31 kg or about 1/1836 of proton or neutron mass PROTON1911Rutherford –Positive charge, +1Relative mass = 1 amu NEUTRON1932Chadwick –No charge, 0Relative mass = 1 amu

9. Table 2.1 The Mass and Charge of the Electron, Proton, and Neutron

10. MODELS OF THE ATOM Thomson (Plum Pudding) Model: positive mass with electrons embedded in it Rutherford Model (1911): positive charge in small volume with (diameter = 1E-15 m) electrons occupying mostly empty space (d = 1E-10 m) around the nucleus Bohr Atom - Chapter 7 Quantum Mechanical Atom - Chapter 7

11. Figure 2.13 a & b (a) Expected Results of the Metal Foil Experiment if Thomson's Model Were Correct (b) Actual Results

12. ATOMIC STRUCTURE Atomic Symbol –Shorthand notation for element –One or two letters on Periodic Table Atomic Structure –Atomic Number (Z) = #protons, uniquely defines an atom –Mass Number (A) = #protons + #neutrons –If atom is neutral, Z = #electrons

13. NUCLIDE SYMBOL Atomic symbol, E; symbol in the middle of each element box on the Periodic Table. Z (left subscript); number on the top of each element box on the Periodic Table. A (left superscript) If species is an ion (has a charge), add + or - charge (right superscript) A Z E ch 23 11 Na has 11 p +, 11 e -, 12 n o

14. IONS A charged species with unequal numbers of protons and electrons. If # protons > # electrons, the ion has a net positive charge and is called a cation If # protons < # electrons, the ion has a net negative charge and is called a anion An ion may consist of an atom or a group of atoms 23 11 Na + has 11 p +, 10 e -, 12 n o

15. MOLECULES Molecules or compounds form when atoms are connected by chemical bonds in which electrons act as the “glue” between atoms. –If electrons are shared between two atoms, the bond is a covalent bond. I.e., the bond between two non-metal atoms. –If electrons are transferred to produce ions, the bond is ionic. Ions are charged particles which form via the gain (anion, commonly formed from nonmetal elements) or loss (cation, commonly formed from metal elements) of electrons. Oppositely charged ions attract and form an ionic bond. Type of bond between a metal and a non-metal atom. Polyatomic ions are charged groups of atoms; they can also form ionic bonds.

16. ISOTOPE Atoms which have the same Z (same # p + )but a different A (different # n 0 ) Most elements have isotopes that occur in nature in precise proportions (fractional abundances, %). A few elements have no naturally occurring isotopes.

17. Figure 2.15 Two Isotopes of Sodium

18. PERIODIC TABLE An arrangement of elements according to increasing atomic number which shows the periodic or regularly repeating nature of elemental properties. –Rows = periods –Columns = groups or families; note similarity of properties –Metals NonmetalsSemimetals –Main group (A)Transition Metals

19. Figure 2.21 The Periodic Table

20. NOMENCLATURE or NAMING COMPOUNDS Binary Ionic Compounds –Metal atoms tend to lose electrons and form cations. –Nonmetal atoms tend to gain electrons and form anions. –Use Periodic Table to determine charges and number of each ion in the compound. Note that the ionic compound must be neutral overall. –Name cation first as element and anion second with “ide” ending. Binary Ionic Compound Type I (T2.3) –Some atoms form only a single type of ion (Binary Ionic Cmp Type I)

21. Figure 2.22 Common Cations and Anions

22. NOMENCLATURE (con’t) Binary Ionic Compound Type II (T2.4, Fig 2.22) –Some transition metal elements form more than one common ion. Designate charge with Roman numeral (II) Polyatomic Ions (Table 2.5) –Memorize –Oxoanion = nonmetal + oxygen

23. Table 2.6 Prefixes Used to Indicate Number in Chemical Names

24. NOMENCLATURE (con’t) Binary Covalent Compound Type III (Table 2.2) –Compounds formed from two nonmetals in which electrons are shared in chemical bond. –Name more “cation-like” first, then the more “anion- like) second with “ide” ending. Hydrogen is almost always named first. –Indicate number of each using prefix as needed. (T2.6) –Note historic names (water, ammonia)

25. Table 2.6 Prefixes Used to Indicate Number in Chemical Names