Download Chapter 2. Atoms, Molecules and Ions

GENERAL CHEMISTRY I
CHEM-1030
INSTRUCTOR’S LECTURE NOTES
CHANG, CHEMISTRY CHAPTER 2
Atomic Theories:
Democritus’ Atomic Theory, 5th Century BC:
Dalton’s Atomic Theory, 1808:
1.
Elements are composed of extremely small, eternal spheres called atoms.
2.
All atoms of an element are identical with the same mass and properties.
3.
Atoms of different elements have different weights and properties.
4.
Compounds composed of atoms chemically combined, in small whole no. ratios.
5.
Chemical reactions involve the rearrangement of atoms.
Law of Definite Proportions
Law of Multiple Proportions
Law of Conservation of Mass
Atomic Structure:
The three “subatomic particles”
Location
Mass (g)
Mass (amu)
Charge (coulomb)
Charge (charge unit)
Proton
nucleus
1.67262 x 10-24
~1
+ 1.6022 x 10-19
+1
Neutron
nucleus
1.67493 x 10-24
~1
0
0
Electron
“electron cloud”
9.10939 x 10-28
~0
-1.6022 x 10-19
-1
Discovery of the Electron: Cathode ray tube, define electric charge, electric current,
battery, high voltage source, anode, cathode, (evacuated) cathode ray tube (ancestor of
TV), how to make rays visible, effect on cathode ray path of magnet and electric charge.
Millikan’s Oil Drop Experiment showing the charge of the electron and its uniformity.
Discovery of Radioactivity:
Henri Becquerel’s experiment, nature of alpha, beta and gamma radiation.
Discovery of the Nucleus: Plum Pudding atomic model. Ernest Rutherford’s experiment
showing small size, hardness and positive charge of the nucleus. Modern view of the
atom with most mass concentrated in the nucleus and the size due to orbiting electrons.
Discovery of the Proton: Magnitude of + charge is identical to that of the negative
electron. However, mass of the proton is 1820 times that of the electron.
Discovery of the Neutron: No charge, so harder to detect. Existence deduced from the
relative masses of H and He nuclei. Refer to table on text page 43. Proton and neutrons
have similar but not exactly the same mass.
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Atomic Nuclei: One to ~100 protons. Number of protons determines identity of atom.
Protons are mutually repulsive so neutrons needed to allow cohesion. Only “allowed”
combinations of protons and neutrons have survived to the present. Z >83. Unstable
combinations decay. Atomic nuclei (or elements) represented by atomic symbols.
Atomic number, Z in lower left (redundant). Mass number A, in upper left. A is the sum
of Z and no. of neutrons. Z varies without altering atomic identity. Isotopes Examples of
H, Cl and Sn isotopes. Number of electrons = no. of protons in electrically neutral atom.
The Periodic Table: Robert Mosely’s Discovery of Z for each element. Arranged by Z,
properties of the elements vary periodically. (Periodic table always available for tests.)
Define Period, Family or Group, Metal, Nonmetal, Metalloid, Alkali Metals, Alkaline
Earth Metals, Halogens, Noble (once Inert) Gases, Gaseous, Liquid and Solid Elements.
Molecules: Define Monatomic, Diatomic and Polyatomic Molecules. Monatomic and
Diatomic Gases. Polyatomic Molecules H2O and CO2.
Ions: Definition, How do they come about. Monatomic and Diatomic. Define anion and
cation. Examples: Cl-1, O-2, K+1, Ca+2, etc.
Chemical Formulas: Define Molecular Formula. Examples; O2, O3, (allotropes) CO,
CO2, C6H12O6, etc.
Molecular Models: Represent shapes of molecules. (More later)
Empirical Formulas: Empirical means experimentally determined. An empirical
formula tells what atoms are present and in what simplest whole number ratio. It does
not give the number of each kind of atom in the molecule. Some empirical and molecular
formulas are the same. e.g. H2O, HCl, CO. Benzene (C6H6) and acetylene (C2H2) both
have a 1:1 ratio of C to H, so both have the empirical formula CH. The empirical
formula of glucose (C6H12O6) is CH2O. Many chemical analyses require a determination
of the empirical formula and the molecular mass to get the molecular formula. Exs.
Ionic compounds: Driving force behind the formation of ionic compounds. Sum of the
ionic charges must be zero. For sodium fluoride, Na+1 and F-1 gives NaF. Calcium
oxide Ca+2 and O-2 gives CaO. Magnesium phosphide, Mg+2 and P-3 give Mg3P2. Etc.
Nomenclature: Division of chemical compounds into organic and inorganic. Inorganic
compounds may be ionic or molecular (covalent). Acids, bases and hydrates are
considered separately.
Nomenclature of Ionic Compounds: Consist of cations and anions. Many are binary; a
few are ternary. Must learn formulas and names of each type to be able to name
compounds. Some can be predicted from periodic table.
Most metals and a few other elements form cations. Cations to learn;
Group IA cations of Li, Na, K, Rb, Cs.
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Group IIA cations of Be, Mg, Ca, Sr, Ba, Ra
Cation of Al
H+1 and NH4+1
Common and Stock names of transition metal cations Cd, Cr, Co, Cu, Fe, Pb, Mn, Hg,
Ag, Sn and Zn
Anions to Learn: Common monatomic anions (and a few others) take -ide suffix. These
can be predicted from periodic table.
Anions of F, Cl, Br, I, H.
Anions of O, S, (Se and Te)
Anions of N and P
Anion of C
Other –ide anions are hydroxide, peroxide, superoxide and cyanide.
Oxyanions to learn: sulfate & sulfite, nitrate & nitrite, hypochlorite, chlorite, chlorate &
perchlorate, carbonate, thiocyanate, permanganate, chromate, dichromate,.
Common and Stock names of hydrogen-containing anions.
Examples of Common Ionic compounds:
sodium chloride, sodium fluoride, potassium iodide
calcium oxide, magnesium oxide, magnesium hydroxide
aluminum oxide, aluminum sulfide, aluminum hydride
ferrous sulfate, ferric carbonate, cupric sulfate, cuprous sulfite
ammonium chloride, ammonium phosphate, sodium phosphate
calcium carbonate, lithium carbonate
potassium nitrate, potassium nitrite, barium nitrate
sodium bicarbonate, strontium hydrogen sulfite
calcium hypochlorite, magnesium perchlorate
sodium dihydrogen phosphate, potassium monohydrogen phosphate
Molecular Compounds: Binary covalent compounds between two nonmetals by sharing,
not transfer, of electrons. Put the name of the more metallic element first. If more than
one compound between two nonmetals can exist, use Greek prefixes to denote.
Hydrogen chloride, hydrogen bromide, hydrogen fluoride
Carbon monoxide, carbon dioxide
Sulfur dioxide, sulfur trioxide
The six nitrogen oxides denoted by NOx
Phosphorus trichloride and pentachloride
Trivial names of water, methane, ammonia, phosphine
Acids and Bases: Definitions, importance of, must learn names and formulas
Meaning of hydrogen ion or proton; actual existence of H+1 ion.
The mineral acids may be strong or weak, binary or oxyacids
Hydro acids, most of which are binary:
HF, HCl, HBr, HI have pure compound and (aqueous) acid names
Also HCN, H2S
Anion names
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Oxyacids and their anion names: nitric, nitrous, sulfuric, sulfurous, phosphoric, carbonic
Also perchloric, chloric, chlorous, hypochlorous
List of Must-know acids:
Bases: Hydroxide Ion. Relation to hydrogen ion, neutralization
Define strong, weak, soluble and insoluble bases.
The four strong, soluble bases are NaOH, KOH, Mg(OH) 2 and Ca(OH) 2
Importance of the base ammonia; cf liquid, gaseous and aqueous ammonia
Hydrates: Define, Mechanism of formation
Ex: copper(II) sulfate pentahydrate, sodium carbonate decahydrate, magnesium sulfate
heptahydrate, calcium sulfate hemihydrate and dihydrate