Chapter 2 Atoms, Molecules & Ions 2.1 The Atomic Theory of Matter Dalton’s Theory 1807 1) Elements are composed of small particles called atoms. 2) Atoms of an element are identical; atoms of different elements are different. 3) Atoms cannot be created or destroyed during chemical reactions. 4) Compounds are formed when atoms of different elements combine. Law of Conservation of mass: during a chemical reaction, mass is conserved. Total mass of the reactants = Total mass of the products Law of Multiple proportions: If 2 elements A and B combine to form more than 1 compound, the mass of B which combines with a mass of A is a ratio of small whole numbers. E.g. H2 O and H2 O2 2.2 The Discovery of the Atomic Structure Electrons 1. Results of Thomson's experiments (1890's) on the behavior of cathode rays in electric & magnetic fields: • cathode rays consist of negatively charged particles called electrons • charge to mass ratio of e- = 1.76 x 108 C/g 2. Results of Millikan’s oil drop experiment: • charge of e- = 1.60 x 10-19 C; mass of e- = 9.11 x 10-28 g Radioactivity: spontaneous emission of radiation 1. Becquerel discovered radioactivity in a Uranium pitchblende sample (1896) 2. Madame Curie discovered polonium and radium (early 1900’s) 3. Results of Rutherford's experiments on radiation (1910-1920) • An atom has a nucleus - a small, dense, positively charged region. • Electrons are located outside of the nucleus. Most of the atom is empty space. • Three types of radiation exist: alpha: high mass particles with +2 charge (helium nuclei) beta: low mass particles with -1 charge (electrons) gamma: neutral, high energy radiation similar to x-rays 2.3 The Modern View of Atomic Structure Nuclear Model: 1) Atoms consist of protons, electrons & neutrons (3 subatomic particles). 2) Protons & neutrons are located in the nucleus; it contains most of the mass. Size analogy: If an atom were the size of a football stadium, nucleus would be the size of a marble. 3) Electrons move rapidly in region outside of the nucleus. •

Masses of atoms are so small that we use the atomic mass unit (amu) to scale up the numbers. 1 amu = 1.66054 x 10-24 g

Electron: -1 charge, mass = 9.11 x 10-28 g = 5.486 x 10-4 amu, Thomson (1897) proton: +1 charge, mass = 1.672 x 10-24 g = 1.0073 amu, Rutherford (1919) neutron: no charge, mass = 1.675 x 10-24 g = 1.0087 amu, Chadwick (1932)

Isotopes • The number of protons in an atom defines what an element is. This is the atomic number, Z. (top # on periodic table) Isotopes are atoms of an element that have a different number of neutrons. Isotopes of an element have the same atomic number, but a different mass number. Mass number = A = number of nucleons (protons & neutrons) in nucleus nuclide: atom of a specific isotope • •

A Isotope symbol for element X Z X For neutral atoms: # protons = # electrons

E.g. 3 H isotopes: normal H: 11 H stable, 1 p, 0 n

A = p + n; Z = p

deuterium: 21 H stable, 1 p, 1 n

tritium: 31 H radioactive, 1 p , 2 n

2.4 The Periodic Table -1st table 1869 Features of modern periodic table: 1) Elements arranged in order of increasing atomic number. 2) Horizontal Rows in periodic table are called periods. 7 periods exist 3) Vertical Columns are groups or families; elements have similar properties. Group names: Group 1A: alkali metals Group 2A: alkaline earth metals Group 7A: halogens Group 8A: noble gases 4) representative elements: A Group; transition elements: B Group 5) Metals are located to the left of the stair-step line. Nonmetals are located to the right of the stair-step. Elements located at the stair-step are intermediate in character - semiconductor or metalloid elements: B, Si, Ge, As, Sb, Te, At Physical state of elements at 25 °C & 1 atm: Gases: O2, N2, H2, F2, Cl 2, and Noble gases Liquids: Br2, mercury Solids: everything else 2.5 Molecules and Molecular Compounds molecule: 2 or more atoms bonded together; discrete entities. •

Many elements exist as diatomic molecules: H2, N2, O2, F2, I2, Cl 2, Br2

Molecular compounds consist of nonmetal elements. Molecular formulas give the actual numbers and types of atoms in a molecule. E.g. CH4, H2O2, C2H4, C6 H12 O6 Empirical formulas give the smallest whole number ratio of atoms in a molecule. E.g. CH4, HO, CH2, CH2 O 2.6 Ions and Ionic Compounds Many chemical reactions involve transfer of electrons between atoms: Metal atoms tend to lose electrons & form + charged cations. Nonmetal atoms tend to gain electrons & form - charged anions. Generally, atoms gain or lose enough electrons to have same number of electrons as nearest noble gas: Group 1A 2A 3A 5A 6A 7A Charge of ion 1+ 2+ 3+ 321-

Ionic compound: consists of metals and nonmetals (or polyatomic ions); ionic compound is a long 3-D array of cations & anions; not individual molecules. Ionic formulas: the number of electrons lost & gained must be equal, so + and charge cancel out. Rules for writing ionic formula: 1) Write down formulas of ions 2) Combine the smallest # of ions to give the charge sum equal to 0; if the charges are not equal, find the lowest common multiple E.g. Predict the formula for the compound formed from the following elements: Ca & O: Ca 2+O2- → CaO Mg & N Mg2+N3- → Mg3 N2 Al & Cl Al 3+Cl - → AlCl 3 Ions have a different # of electrons & protons Ex.

23 11

Na+ : 11p; 12 n; 10 e-

35 17

Cl - : 17 p; 18 n, 18 e-

NOMENCLATURE- IONIC COMPOUNDS A. Naming Cations: 1. Fixed charge metals: Cations have same name as the metal element. Groups 1A, 2A, Al, Ag, and Zn are fixed charge metals – cations that have 1 specific charge. E.g. Ag+ silver ion Zn2+ zinc ion Al 3+ Aluminum ion + 2+ Li lithium ion Ca calcium ion 2. Variable charge metals: If the metal can form more than 1 cation, the charge is indicated by a Roman numeral in parenthesis after the metal name. Most of the transition metals are variable charge metals. E.g. Common metals which exist in more than one positive state: Fe 2+ iron(II) Au+ gold(I) Cu+ copper(I) Hg 22 + mercury(I) Fe 3+ iron(III) Au3+ gold(III) Cu2+ copper(II) Hg2+ mercury(II) 3. Polyatomic Cations: consist of nonmetals H3O+ hydronium NH +4 ammonium B. Naming Anions 1. monoatomic anions: change ending to -ide E.g. oxygen → oxide sulfur → sulfide

hydrogen → hydride

2. Polyatomic anions: most end in -ate or -ite; usually contain O (oxy) Know polyatomic anions on handout. a. Rule for naming oxy series anions: per-....-ate 1 more O than -ate -ate -ite 1 less O than -ate hypo-...-ite 2 less O than -ate b. If H+ is added to a polyatomic ion, write hydrogen (or bi-) in front of name. HCO 3− hydrogen carbonate or bicarbonate H2PO −4

dihydrogen phosphate

NOMENCLATURE RULES I. IONIC COMPOUNDS contain cations & anions 1) Name metal cation. 2) Include Roman numeral in parenthesis ONLY IF metal has variable charge. Fixed charge metals: Group 1A, 2A, Ag, Zn, and Al; others are variable. 3) Name anion. E.g. MgBr2 magnesium bromide PbS lead(II) sulfide 2+ 3barium nitride Ba N → Ba3N2 iron(III) sulfite Fe 3+SO 32 − → Fe 2 (SO3)3 Ca(ClO2)2 calcium chlorite Cr2(CO 3)3 chromium(III) carbonate II. Binary Molecular compounds: contain 2 nonmetals 1) Name 1st element & use a prefix (table 2.6) to indicate the number of atoms. Note that mono- is never used for the first element. 3) Name 2nd element & include prefix for number of atoms (see table 2.6). 4) Change ending of 2nd element to –ide. E.g. N2O5 dinitrogen pentoxide ICl 3 iodine trichloride tetraphosphorus hexasulfide P4S6 dibromine heptaoxide Br2O7 III. Acid: substance that yields H + ions in aqueous solution A. Binary Acids (Nonoxy acids): contain H & 1 nonmetal 1) Name hydrogen as hydro-. 2) Name nonmetal & change ending to -ic acid. E.g. HCl (aq) hydrochloric acid H2S(aq) hydrosulfuric acid + 3hydrophosphoric acid H P → H3P(aq) hydroselenic acid H+Se2- → H2Se(aq) B. Ternary Acids (Oxy Acids): contain H & polyatomic ion 1) no hydro prefix 2) Change ending of polyatomic ion: -ate → -ic acid -ite → -ous acid E.g. HNO3(aq) nitric acid H2SO4(aq) sulfuric acid H2SO3(aq) sulfurous acid H3PO4(aq) phosphoric acid − + chlorous acid H ClO 2 → HClO2 perbromic acid H+BrO −4 → HBrO 4

ELEMENTS st

Know names of 1 36 elements (H – Kr) & Ag silver I iodine

Au gold Sn tin

Ba barium Hg mercury

Pb lead U uranium

Pt platinum Cd cadmium

Know these Polyatomic ions Cations: Ammonium

NH +4

Hydronium

H3O+

Anions: Acetate

C2H3O −2

Carbonate

CO 32 −

Oxalate

C2O 24 −

Cyanide

CN-

Thiocyanate

Bicarbonate

HCO 3−

SCN-

Perchlorate

ClO −4

Hydroxide

OH-

Chlorite

ClO −2

Chlorate

ClO 3−

Hypochlorite

ClO-

Nitrate

NO 3−

Nitrite

NO −2

Sulfate

SO 24 −

Sulfite

SO 32 −

Phosphate

PO 34−

Phosphite

PO 33 −

Permanganate

MnO −4

Chromate

CrO 24 −

Dichromate

Cr2O 72 −

*Note: Names of ions on the right can be derived from the lefthand group. The oxy-anions for bromine and iodine can be named in a manner analogous to that listed for chlorine.

NOMENCLATURE Worksheet A. Write the correct formula: Fe3+S2- →

1. iron (III) sulfide

Fe2S3

2. silver dichromate Ag+Cr2O 72 − → Ag2Cr2O7 3. sodium phosphide

Na+P3- →

4. cobalt (III) nitrite

Co3+NO −2 → Co(NO2)3

5. tin(IV) perchlorate

Sn+4ClO −4

6. diphosphorus pentasulfide 7. calcium phosphite

Na3P

→ Sn(ClO4)4

P2 S5

Ca2+PO 33 − → Ca3(PO3)2

8. magnesium permanganate Mg2+MnO −4 → Mg(MnO4)2 9. chlorous acid

H+ClO −2 → HClO2(aq)

10. hydrosulfuric acid

H+S2- →

H2S(aq)

B. Write the correct name: 1. S 4O8

tetrasulfur octoxide

2. AlH3

aluminum hydride

3. Cr(SCN)3

chromium(III) thiocyanate

4. PbO2

lead(IV) oxide

5. HBr(aq)

hydrobromic acid

6. Zn(HSO4)2 zinc bisulfate (or zinc hydrogen sulfate) 7. MnC 2O4

manganese(II) oxalate

8. NH4C2H3O2

ammonium acetate

9. H2CO3(aq)

carbonic acid

10. Fe(BrO)2

iron(II) hypobromite