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Unit 2
Atoms, Molecules, &
Ions
Law of Constant Composition (Definite Proportion): All
samples of a compound have the same composition, the same
proportion by mass of the constituent elements.
H2O by mass is always
11.19 % H
88.81 % O
A 0.1 g sample of Mg, when combined with O2 (g) yields
0.166 g MgO. A second sample, 0.144 g of Mg is also
combined with O2 (g).
What is the mass of MgO produced?
Law of Multiple Proportions: When 2 elements form a
series of compounds, the ratio of the masses of the
second element that combine with 1.0 g of the first
element can always be reduced to small whole numbers.
The following data were collected for several compounds of
nitrogen combining with 1 g oxygen. Determine the formulas.
Compound A = 1.750 g N
Compound B = 0.8750 g N
Compound C = 0.4375 g N
Dalton’s Atomic Theory




Each element is made up of tiny particles called atoms.
The atoms of a given element are identical; atoms of
different elements are different in some fundamental
way.
Chemical compounds are formed when atoms of different
elements combine with each other. They combine in
simple whole number ratio, always with the same relative
number and types of atoms.
Chemical reactions involve reorganization of the atoms,
changes in the way they are bonded together. The atoms
themselves are indestructible.
History
1803 John Dalton discovered that elements are made of atoms. He
thought that atoms were solid, like a marble.
Early 1800’s Michael Faraday discovered that a cathode ray is
negatively charged particles
1874 George Stoney coined the term “electron”
1875 Crooks discovered the electron has a mass of 9.1 x 10-28 g and
have a negative (-) charge.
1897 JJ Thomson used a cathode ray tube to discover negative ray
consisting of electrons with a charge to mass
ratio c/m = -1.76 x 108 c/g
1895 Roentgen discovered when using a cathode ray tube, that
there was fluorescence in the room. He found X rays.
1909 Robert Millikan worked with an oil drop to determine the
magnitude of e- charge to be – 1.6022 x 19-19 coulomb.
1896 Henri Bacquerel (1852 – 1908)
discovered radioactivity
a alpha particle is a helium nucleus
b beta particle is an electron
g gamma ray is pure energy
4He
2
e-
1911 Earnest Rutherford (1871 – 1937) discovered that
there is a mass in the center of the atom with a
positive charge containing p+, the e- are around the
nucleus, with the atom mostly empty space.
1932 Chadwick discovered the neutron which is also found
in the nucleus.
Subatomic particles
Particle
amu
mass (g)
Atomic
charge
Electrical
charge
(coulomb)
Proton (p+)
1.0073
1.673 x 10-24
+1
+1.602 x 10-19
Electron (e-)
.0005486
9.109 x 10-28
-1
-1.602 x 10-19
Neutron (n)
1.0087
1.675 x 10-24
0
0
Nucleus has a density of 1 x 1013 to 1 x 1014 g/cm3
There are 4 forces that hold an atom together:
1. gravitational – attraction between 2 bodies with mass
2. electromagnetic – attractive and repulsive forces due to
charged or magnetic objects
3. strong nuclear force – keep protons which all have positive
charges from flying apart
4. weak nuclear force – responsible for (b) beta decay in
radioactivity
Periodic Table
Atomic mass
12.0111
C
Element symbol
6
Atomic number
Elements with the atomic mass in ( ) means that very
little of it exists at any given time. Scientists have
not been able to attain an good average atomic mass.
Nuclear Notation
Atomic Mass
Atomic number
A
Z
X
Chemical symbol
Atomic mass = p+ + N
Atomic number = p+
Isotopes = Atoms of the same element (so the same
atomic number) but with a different mass because
they have a different number of neutrons.
Kinds of Decay
Alpha decay a
A
Z
X 
A-4
N+
Z-2
4
He
N = new element
Beta decay b
X 
Z
A
N
+
e
Z+1
A
N  p+ + e-
Atomic Mass is the average of the number of
protons and Neutrons found in the nucleus
Average
fraction of abundance ) ( Mass of
=
S
(
isotope
Atomic Mass
of each isotope
each
)
In naturally occurring Argon, 99.600% of the atoms are
40Ar with a mass of 39.9624 amu (u), 0.337% 36Ar with
18
18
38
a mass of 35.96755 u, and 0.063% 18Ar with a mass of
37.96272 u. Calculate the Average mass of Argon.
Classification of elements as metals,
nonmetals, and metalloids.
Alkali Earth metals +2 ion
Alkali metals +1 ion
Rare Earth Metals
Noble gases
Halogens -1 ion
Chalcogens -2 ion
Transition metals
Periodic table
with atomic
symbols,
atomic numbers,
and partial
electron
configurations.
Diatomic molecules
There are 7 elements that do not exist as 1
atom alone. As an element, they are always
as 2 atoms bonded together as a molecule.
Location of the
diatomic elements
Molecular Compound – molecule with more than 1
type of atom
Molecular Formula – shows the actual # of atoms
Empirical Formula – shows the relative # of atoms in
the smallest possible whole # ratio
Structural Formula – shows how the atoms are
assembled into the molecule
Chemical Bonds –
the forces holding atoms together

Ionic bonds – transfer of e- resulting in ions
attracted to each other
Ions – atom or group of atoms which have gained or lost
e- (s), therefore carrying an electrical charge.
cation (+) positive charged ion
anion (-) negative charged ion

Covalent bonds – sharing of e-
Polar covalent bond – an uneven sharing of e- resulting
in a molecule which appears to have opposites
charges on different parts of the molecule.
Nonpolar covalent bond – an even sharing of e-
Naming compounds


Inorganic binary compounds type I
• one atom (+) cations maintain atom name
• one atom (-) anion changes name (-ide)
• (+) cation is followed by the (-) anion
Inorganic type II – some anions form multiple charges
• Use roman numerals to indicate charge
example: CuCl – copper (I) chloride,
CuCl2 – copper (II) chloride
• Or/ higher charge name ends with
(-ic)
while lower charge name ends with (-ous)
example: FeCl3 ferric chloride
FeCl2 ferrous chloride
Element
Latin naming
Copper
Cu+2 Cuprous, Cu+3 Cupric
Iron
Fe+2 Ferrous, Fe+3 Ferric
Mercury
Hg2+2 Mercurous, Hg+2 Mercuric
Lead
Pb+2 Plumbous, Pb+4 Plumbic
Tin
Sn+2 Stannous, Sn+4 Stannic
Chromium
Cr+2 Chromous, Cr+3 Chromic
Manganese
Mn+2 Manganous, Mn+3 Manganic
Cobalt
Co+2 Cobaltous, Co+3 Cobaltic
Polyatomic – must memorize the names!

Oxy anions: elements with oxygen acting as
polyatomic
• Less than small number of oxygen (hypo-)
• Small number of oxygen (-ite)
• Large number of oxygen (-ate)
• More than large number of oxygen (per-)
Example: hypochlorite
ClO-1
chlorite
ClO2-1
chlorate
ClO3-1
perchlorate
ClO4-1
Naming Acids – when dissolved in H2O
they produce free protons (H+)
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If anion does not contain oxygen
•
Use prefix (hydro-) with suffix (-ic)
example HS hydrosulfuric acid
If anion does contain oxygen
•
•
•
If anion ends in (-ate) use root name with (-ic)
H2SO4 sulfuric acid
If anion ends in (-ite) use root name with (-ous)
H2SO3 sulfurous acid
When multiple oxyanion use
Example HClO4
HClO3
HClO2
HClO
perchlorate
chlorate
chlorite
hypochlorite
perchloric acid
chloric acid
chlorous acid
hypochlorous acid
Binary Covalent type III – 2 nonmetals
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1st element in formula named first with full element
name
2nd element named as if anion
Prefixes used to denote number of atoms present
Mono is never used for naming 1st element
mono - 1
hexa - 6
di
- 2
hepta - 7
tri
- 3
octa - 8
tetra - 4
nona - 9
penta - 5
dec
- 10
Naming Organic Molecules


Name root name according to the number
of carbons in the longest continuous
chain.
Name the suffix according to the
bonding, carbon to carbon.
• Single bond C – C
• Double bond C = C
• Triple bond C C

end with -ane
end with -ene
end with -yne
Name substitutions on the main chain
before the root name, giving the lowest
possible number.
Hydrocarbon
Substitute Naming
Group Names
Methyl -CH3
Ethyl -CH2CH3
Propyl -CH2CH2CH3
Butyl
-CH2CH2CH2CH3
Halides
Chloro -Cl
Bromo -Br
Iodo -I
Fluoro -F
Prefixes used when
more than one group
of the same kind is
attached
ditritetrapentahexa-
two substitutes
three
four
five
six
Naming organic molecules containing
functional groups
Drop the e, add
Functional Group
ending
-OH
Group Name
Name
Alcohol
- ol
Carboxylic Acid
Ketone
- oic Acid
- one
Functional Group
Group Name
Aldehyde
Name ending
- al
Drop –ane, add
Double bond
Triple bond
- ene
- yne
Isomers: Molecules with the same chemical
formula, but different arrangement of atoms.
With different arrangement of atoms, the
properties of the chemicals are different.
Small molecules only have one arrangement possible.
CH4 Methane
C2H6 Ethane
C5H12 has 3 isomers