Download Brown, Le May, and Bursten: Chapter 2

CHAPTER 2: ATOMS, MOLECULES AND IONS
 The Atomic Theory of Matter
 The Discovery of Atomic Structure
 The Modern View of Atomic Structure
 The Periodic Table
 Ions and Ionic Compounds
 Naming Inorganic Compounds
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CONSERVATION OF MASS
Law of conservation of mass: mass is neither created or
destroyed during a reaction.
 The atoms form new bonds and thus are present after reaction only
bound to some other atoms.
E.g. 2H2(g) + O2(g)  2H2O(l); 2 g of H2 plus 16 g of O2 produce how
many grams of water?
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ATOMIC THEORY OF MATTER (DALTON'S THEORY)
Law of definite proportions led to theory that all matter made up of
atoms.
1. Atoms- basic building blocks and don't change when react with
other atoms.
2. Element- describes matter composed of only one type of atom.
3. Compound- combination of atoms in specific proportions.
4. Chemical reaction- atoms exchange partners producing other
compounds.
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LAW OF MULTIPLE PROPORTIONS
Some elements can form more than one compound when they
react together (C & O: CO and CO2; N & O: N2O, NO, NO2, etc.).
Dalton’s law predicted that the mass proportions should be
proportional. Experiment confirmed this leading to this law.
Law of multiple proportions: when two elements form more
than one compound, the ratio of the masses in one compound
divided by the ratio of these masses in the other compound
gives a ratio of small whole numbers.
E.g. There are three binary compounds that form between barium
and nitrogen. There was 4.9021 g , 9.8050 g and 14.7060 g of Ba
per 1.0000 g N in the three compounds. Show that these
compounds obey the law of multiple proportions.
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The Discovery of Atomic Structure
1. Thompson - Cathode rays.
2. Milliken - Oil drops.
3. Rutherford - backscattering -particles.
4. Radioactivity, the spontaneous emission of radiation from
an atom led to the discovery of -, -, and -rays.
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The Modern View of Atomic Structure
1. Building Blocks
a) protons have a + charge, mp = 1.67 x 1027kg
b) neutrons are neutral,
mn = mp
c) electrons have a  charge, me = mp/1835
2. Neutral atom: # electrons = # protons
3. Mass of atom is found by adding the mass of protons and
neutrons
4. Protons identify the element (# protons called the atomic number,
Z).
5. Isotopes have varying numbers of neutrons, 11H , 12 H , 13 H
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The Periodic Table
1A
H
Li
Na
K
Rb
Cs
Fr




2A
Be
Mg
Ca
Sr
Ba
Ra
8A
3A 4A 5A 6A 7A He
B C N O F Ne
Al Si P S Cl Ar
Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I
Xe
La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
Ac
Elements in a column have similar reactivity.
Grayed elements are semi-metals (metalloids).
Elements to left of gray area are metals.
Elements to right of gray area are non-metals.
Special names:
Group
Name
Elements
1A
Alkali metals
Li, Ma, K, Rb,
Cs, Fr
2A
Alkaline earth
Be, Mg, Ca,
metals
Sr, Ba, Ra
6A
Chalcogens
O. S, Se, Te,
Po
7A
Halogens
F, Cl, Br, I, At
8A
Noble gases
He, Ne, Ar, Kr,
Xe, Rn
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BONDING
An electron from each atom strongly interacts to
form a bond.
 Bonding can be either:
a) Ionic: Occurs between metal atoms or
between metal and nonmetal atoms.
b) Covalent: occurs between nonmetal atoms
and forms molecules.
 Covalent: sharing of electrons occurs.
 Ionic: electron(s) is (are) transferred from one
atom to the other to produce
a) a positively charged substance (cation) and
b) a negatively charged substance (anion).
c) Ions bound by electrostatic attractions.
d) The formula of the ionic compound can be
determined from the charge on the cation
and the anion (even with a polyatomic ion):
i) Mg2+ and Cl form an ionic compound
with the formula: MgCl2.
ii) Fe3+ and O2 form an ionic compound
with the formula: Fe2O3.
iii) Fe2+ and NO 3 form a compound with
the formula: Fe(NO3)2
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ACIDS AND BASES
 Acid = a substance that donates an H+ when it
dissolves in water.
HA 2 H  (aq )  A  (aq ) where A is the
rest of the compound.
H O
a) HCl: HCl 2 H  (aq )  Cl  (aq )
b) H3PO4:
H O
H3PO 4 (l) 2 H  (aq )  H 2 PO 4 (aq )
H O
H 2 PO 4 (aq ) 2 H  (aq )  HPO 24  (aq )
H O
HPO 24  (aq ) 2 H  (aq )  PO34 (aq )
 Base = a substance that donates OH when it
dissolves in water.
a) NaOH(s):
H O
NaOH (s) 2 Na  (aq )  OH  (aq )
H O
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NAMING INORGANIC COMPOUNDS
IUPAC and Common Usage Methods of naming
used for both organic and inorganic compounds.
a) Organic compounds = predominantly
carbon containing compounds.
b) Inorganic compounds = all others
 Binary ionic compounds: compounds having
a cation (a metal) and an anion (one of the
main group anions).
a) Cations:
i) Accepted method: Roman numeral in
parentheses to indicate the charge. Not
necessary when only one ionic charge
possible. See Table 2.2.
ii) Common method: Remove ending and
add either -ous or -ic to
iii) Latin form of the element used instead
on some: Stannous, Stannic; Ferrous,
Ferric; Cuprous, Cupric.
b) Anions:
i) For monoatomic anions add -ide to the
stem. Fluor-, Ox-, Nitr-, Sulf-, etc.
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NOMENCLATURE 2
Binary molecular compounds:
 More cationlike appears first; -ide ending is
placed on the anionlike substance.
The more cationlike element appears to the
left of or below the other element in the
periodic table
 With hydrogen as one of the two, hydrogen is
first; place an -ide on the other element.
 If a compound contains a group VI or VII
element, an –ide ending is added to it.
Numerical prefixes are used especially with
the element listed second. Mono- not used
with the first element.
a) When oxygen with fluorine, oxygen first in
name: E.g. Oxygen difluoride = OF2.
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Compounds with Polyatomic ions
 Name these the same way that ionic binary
compounds are named replacing the name of
the anion (in a few cases cation) with the
name of the polyatomic ion.
 Oxy anions (anions containing at least 1
oxygen atom): difficult because there are often
several possible anions.
1) Add -ate to the stem (carbonate, sulfate,
nitrate, etc); but
2) Many oxyanions can have other proportions
of oxygen e.g. SO 32  , SO 24  . These can be
named by adding:
a) per- x -ate for most oxygen
b) -ite for a small number of oxygen
c) hypo- x -ite gives least
E.g. oxychlorides.
 If ion contains H, write: “hydrogen” + name of
ion without hydrogen, e.g. hydrogen sulfate.
 A prefix of mono- or di- to indicate the number
of hydrogens may be needed. E.g.
dihydrogen phosphate- H 2 PO 4 .
Oxyacids:
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 Binary compounds which are acids: Prefix
hydro- added to the anion part; change the ide ending to -ic.
H O
HCl(g ) 2 HCl(aq) . Hydrogen chloride
becomes hydrochloric acid.
 The ending -ite is changed to -ous and -ate to
-ic.
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