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
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
• 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
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)
• 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
Mass number = A = number of nucleons (protons & neutrons) in nucleus
nuclide: atom of a specific isotope
• Isotope symbol for element X Z X
A = p + n; Z = p
• For neutral atoms: # protons = # electrons
E.g. 3 H isotopes: normal H: 1 H
1 deuterium: 2 H
1 tritium: 3 H
stable, 1 p, 0 n stable, 1 p, 1 n 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+ 3- 2- 1-
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
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
Li lithium 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 2 + mercury(I)
Fe 3+ iron(III) Au3+ gold(III) Cu2+ copper(II) Hg2+ mercury(II)
3. Polyatomic Cations: consist of nonmetals
H3O+ hydronium NH + 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
-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
4 dihydrogen phosphate
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
barium nitride Ba N → Ba3N2
iron(III) sulfite Fe 3+SO 3 − → 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
hydrophosphoric 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 − → HBrO 4
Know names of 1 36 elements (H – Kr) &
Ag silver Au gold Ba barium Pb lead Pt platinum
I iodine Sn tin Hg mercury U uranium Cd cadmium
Know these Polyatomic ions
Ammonium NH +
Acetate C2H3O −
Carbonate CO 3 −
Oxalate C2O 2 −
Thiocyanate SCN- Perchlorate ClO −
Hydroxide OH- Chlorite ClO −
ClO 3 Hypochlorite ClO-
Nitrate NO 3 Nitrite NO −
Sulfate SO 2 −
4 Sulfite SO 3 −
Phosphate PO 3 −
4 Phosphite PO 3 −
Permanganate MnO −
Chromate CrO 2 −
Dichromate Cr2O 7 −
*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.
A. Write the correct formula:
1. iron (III) sulfide Fe3+S2- → Fe2S3
2. silver dichromate Ag+Cr2O 7 − → Ag2Cr2O7
3. sodium phosphide Na+P3- → Na3P
4. cobalt (III) nitrite Co3+NO − → Co(NO2)3
5. tin(IV) perchlorate Sn+4ClO −
4 → Sn(ClO4)4
6. diphosphorus pentasulfide P2 S5
7. calcium phosphite Ca2+PO 3 − → Ca3(PO3)2
8. magnesium permanganate Mg2+MnO − → Mg(MnO4)2
9. chlorous acid H+ClO − → 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