Chem 1411 Chapt. 2
Atom- Smallest particle of an element. Basic chemical building block of
matter.
Dalton’s Atomic Theory:
1. Each element is composed of atoms.
2. Atoms of an element are identical. Atoms of different elements are
different.*
3. In chemical reactions:
Atoms of one element are not changed into atoms of a different element.*
Atoms can neither be created nor destroyed*.
4. Compounds are formed when atoms of more than one type of element
combine.
Law of Multiple Proportions- If two substances are made of the same
types of elements, but the elements are in different proportions, then the two
substances are different. Example: NO, NO2, N2O, H2O, H2O2
Law of Constant Composition- All samples of a given substance must
have the same composition. Example, each water molecules has 2 hydrogen
atoms for every oxygen atom.
Law of Conservation of matter- Matter can neither be created nor
destroyed.
Radioactivity- Spontaneous emission of radiation, atomic and subatomic
particles.
X-Rays- High energy light(photons).
Gamma Rays(γ)- Light with energy higher than x-rays
Alpha Rays(α)- A stream of helium nuclei(24He+2)
Beta Rays(β)- A stream of electrons emitted from the nucleus of an atom.
Atomic structure
J.J. Thompson, while studying cathode rays (electrical discharges in
evacuated tubes), calculated the ratio of an electron’s charge to its mass
(-1.76 x 108 C/g).
Robert Millikan, with his oil-drop experiment, calculated the charge on an
electron (-1.6 x 10-19C)
The mass of an electron is (-1.6 x 10-19C)/(-1.76 x 108C/g ) = 9.1 x 10-28g
Ernest Rutherford discovered the proton by bombarding thin gold foil with
alpha particles in 1919.
James Chadwick discovered the neutron in 1932.
Structure of the Atom
Proton(p)
Symbol
1
+1 H,
Charge
+1
Mass
p
neutron(n)
e-, (-10e)
0
-1
1 a.m.u.
Location
electron(e)
nucleus
1
n
0
0a.m.u.
electron shells
1 a.m.u.
nucleus
Atomic Mass Unit(a.m.u.)- 1 a.m.u. is equal to 1.66 x 10-24g
Atomic Number- Number of protons in an atom’s nucleus (Z).
Mass Number- The sum of the number of protons and neutrons in an atom’s nucleus.
Number of Neutron- Mass number minus atomic number.
Example for neutral atoms:
23
+
11 Na
31
15 P
#proton is 11; #electron is 10; #neutron is 23 -11 = 12
15
15
31 -15 = 16
Isotopes- atoms with the same number of protons, but different number of neutrons.
Hydrogen isotopes: 11H(hydrogen, 99.99%);
1
2
H(deuterium,0.015%);
1
3
H(tritium)
37
Chlorine isotopes: 1735Cl (75%) Cl-35
Cl-37
17 Cl (25%)
Atomic Mass or Atomic Weight- is the weighted average of the isotope masses
For Chlorine: (35 x 0.75) + (37 x 0.25) = 35.45 a.m.u.
Elemental Arrangement in Periodic Table
Periods (rows)- horizontal; There are 7 periods.
Period 1 normally ranges from H to He; Period 2 ranges from Li to Ne, etc..
Groups (columns, also known as families)- vertical
With the exception of the transition and the inner transition metals, there are 8 groups,
called the representative groups.
Group IA has the alkali metals
Group IIA has the alkaline earth metals
Group IIIA
Group IV
Group VA has the pnicogens
Group VIA has the chalcogens
Group VIIA has the halogens
Group VIIIA has the noble gases.
General Classification of Elements
Metals (electropositive)- Element to the left of the metalloids
Non-metals (electronegative)- Elements to the right of the metalloids
Metalloids (semi-metallic)-Consists of Boron(B), Silicon(Si), Germanium(Ge),
Arsenic(As), Antimony(Sb), Tellerium(Te) and Astatine(At).
Symbols
Letter symbols represent the elements. Letter is always capitalized. For symbols that have
2 letters, the second is lowercase.
Types of Formulas
Molecular Formula- Provides the exact number ratios of atoms in a compound.
Ex. C6H6, CO2, C3H8O, C2H6, HCl
Empirical Formula- Provides the reduced whole number ratio of the atoms of that
compound. Ex. CH, CO2, C3H8O, CH3, HCl
Structural Formula- Provides exact number ratio of atoms, connectivity, and orientation
in space. (water)H2O => H-O-H
(acetylene)C2H2 => H-C≡C-H
(formic acid) H2CO
(ammonia) NH3
(phosphorus pentachloride) PCl5
(sulfur hexafluoride) SF6
Molecule- At least two atoms chemically bonded.
Ex. O2, H2, HBr, C3H8O
Compound- At least two different types of elements chemically combined.
NaCl, HBr, H2SO4
Types of MoleculesDiatomic- O2, H2, N2, Br2, F2, I2, Cl2, CO, NaCl,
Polyatomic- P4, S8, H2O, C2H5OH
Types of CompoundsIonic- Consists of metals and non-metals (Or in general cations and
anions).
NaCl, MgCl2, K2S, Na2SO4
Molecular (covalent)- Consists of non-metals only.
HCl, N2O4, C3H6O, C6H12O6
Note- All compounds can be molecules; not all molecules can be
compounds.
Ions- Are chemical species that have a net charge.
Monatomic- cations: K+, Na+, Mg+2, Al+3
Anions: Cl-, O2-, BrThe monatomic ions like to take charges that will give them the electronic
stability of their nearest noble gas, when they combine with other atoms.
Example Na atom loses 1 electron to have 10 electrons like Ne and becomes
Na+ ion. Oxygen gains 2 electrons to have 10 electrons like Ne and become
O2- ion.
Polyatomic Ions-consist of two or more atoms and a net charge.
Cations: NH4+, H3O+ (ammonium NH4+ , hydronium H3O+ )
Anions: NO3-, NO2-(nitrate, nitrite)
SO42-, SO32-(sulfate, sulfite)
PO43-, PO33-(phosphate, phosphite)
CO32-(carbonate)
ClO4-(perchlorate), (BrO4-, IO4- ), etc...
ClO3-(chlorate), (BrO3-, IO3- ), etc…
ClO2-(chlorite), (BrO2-, IO2- ), etc..
ClO-(hypochlorite)
OH-(hydroxide)
C2H3O2-(acetate)
CN-(cyanide)
CrO4-2 (chromate), Cr2O72- (dichromate)
Writing Formulas for Ionic Compounds
The metal is named first. Then the non-metal ion. The absolute value of the
cation charge becomes the subscript for the anion; the absolute value of the
anion charge becomes the subscript for the cation. The subscripts must be
reduced to their lowest set of prime numbers
Ex. Ba2+ and Cl- form BaCl2
Mg2+ and NO3- form Mg(NO3)2
Fe2+ and PO43- form Fe3(PO4)2
Ca2+ and O-2 form Ca2O2 => CaO !!!!!!
Mg2+ and N-3 form Mg3N2
2Na+ + SO4-2 = Na2SO4
2Fe3+ + 3SO4-2 = Fe2(SO4)3
NOMENCLATURE
Naming Inorganic Compounds
Ionic Compounds- Consist of metal and non-metal. For metals that have
only one charge state, name its compound by naming the metal first, then the
anion. Example:
Na+ + Cl- = NaCl Sodium chloride
2Na+ + SO42- = Na2SO4 Sodium sulfate
For metals that can have multiple charge states ( Or valence state), one must
use roman numerals to indicate the charge state of the metal. Example:
Fe2+ + 2Cl- = FeCl2 Iron(II) chloride OR Ferrous chloride
2Fe3+ + 3SO42- = Fe2(SO4)3 Iron (III) sulfate OR Ferric sulfate
Cu2S Copper(I) sulfide OR Cuprous sulfide
CuS Copper(II) sulfide OR Cupric sulfide
What about Tin (Sn, stannum, Sn4+, Sn2+)?
Molecular (covalent) Binary Compounds- Consist of 2 non-metal
elements.
In naming the covalent compounds, the more electropositive element’s name
comes first; then the more electronegative element follows with the suffix ‘ide’. Also, one uses prefixes (mono, di, tri, tetra, penta, hexa, hepta, octa,
nona, deca) to indicate the number of atoms of each element. If there is only
one atom of the more electropositive element, do not use the prefix ‘mono’.
Example:
CO carbon monoxide; CO2 carbon dioxide
NO2 nitrogen dioxide; N2O dinitrogen monoxide;
N2O5 dinitrogen pentoxide
P4O10 tetraphosphorus decoxide
Naming Acids
Hydroacids- Consist of hydrogen and a non-metal other than oxygen.
Hydrogen is the cation; the anion name ends as “-ide” for example chloride,
sulfide, cyanide.
You can name as molecular binary compounds or as acids. To name as an
acid, use prefix hydro, followed by the modified anion name by replacing
the “ide” with “ic”, followed by the word “acid”.
Example: HCl named as a molecular binary compound is hydrogen
chloride.
HCl named as an acid is : hydrochloric acid.
For the acids containing Sulfur, add uric.
H2S named as a molecular binary compound is :Dihydrogen monosulfide.
H2S named as an acid is: hydrosulfuric acid.
HCN (derived from the cyanide ion CN-) is named hydrocyanic acid.
As a molecular compound?
Oxyacids- Consist of hydrogen, oxygen and other non-metal(s).
One can name them as molecular compounds or as acids. In naming them as
acids drop the hydro prefix; if the anion suffix is “ate” replace the with “ic”.
If the anion suffix is “ite” replace it with “ous”.
For the acids containing Sulfur, add ‘ ur’.
Example: H2SO4 named as a molecular compound is: Dihydrogen sulfate.
H2SO4 named as an acid is: Sulfuric acid
H2SO3 named as a molecular compound is: Dihydrogen sulfite.
H2SO3 named as an acid is: Sulfurous acid.
What about nitrate and nitrite oxoanions?
HClO4 (derived from the perchlorate ion ClO4-) is named Perchloric acid
HClO3 (derived from the chlorate ion ClO3-) is named Chloric acid
HClO2 (derived from the chlorite ion ClO2-) is named Chlorous acid
HClO (derived from the hypochlorite ion ClO-) is named Hypochlorous
acid
An old way of naming the ions of polyprotic acids that are missing one
proton (H+) is by putting the prefix ‘bi’ in front of the name of the anion. Ex:
HS- is called bisulfide (also known as the hydrogen sulfide ion).
HSO4- is called bisulfate ( also known as hydrogen sulfate ion).
HSO3- is called bisulfite ( also known as hydrogen sulfite ion).
HCO3- is called bicarbonate ( also known as hydrogen carbonate ion).
H2PO4- is called biphosphate ( also known as dihydrogen phosphate ion).
H2PO3- is called biphosphite ( also known as dihydrogen phosphite ion)
Organic Compounds- Compounds made mostly of carbon and hydrogen.
Simple(normal) Alkanes- CnH2n+2