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Chemistry 101 : Chap. 2
Atoms, Molecules and Ions
(1) Atomic Theory of Matter
(2) The discovery of Atomic Structure
(3) The Modern View of Atomic Structure
(4) Atomic Weight
(5) Periodic Table
(6) Molecules and Molecular Compounds
(7) Ions and Ionic Compounds
(8) Naming Inorganic Compounds
The Atomic Theory of Matter
The history of development of atomic theory of
matter begins in ancient Greece. However, modern
atomic theory has it’s origin in a burst of scientific
discovery between 1870 and 1930.
 Democritus (460 ~ 370 BC)
Democritus proposed atomic theory of matter.
He and other Greek philosophers believed that
material world must be made up of hard and tiny
indivisible particles that they called atomos,
which are in constant motion.
The Atomic Theory of Matter
 Aristotle (384 ~ 322 BC)
Aristotle proposed 4 element theory
of matter.
Fire
dry
hot
Air
Earth
wet
cold
Water
The school of thought laid out by Socrates, Plato and Aristotle
dominated the western philosophy for 2000 years and the atomic
theory of matter was completely buried.
The Atomic Theory of Matter
 John Dalton (1766 ~ 1844)
Dalton’s Atomic Theory
Dalton proposed that all matter is
made up of atoms and stated that
elements are the simplest form of
matter.
(1) Each element is composed of atoms
(2) All atoms of a given element are identical,
but they are different from the atoms of
all other elements
(3) Atoms are neither created nor destroyed
in chemical reactions.
(4) Compounds are formed from chemical
combination of two or more atoms.
The Atomic Theory of Matter
 What can Dalton’s theory explain?
(1) Law of constant composition
 In a given compound, the relative numbers and kinds of atoms
are constant. [postulate 4]
(2) Law of conservation of mass
 The total masses of material present before and after a chemical
reaction are identical [postulate 3]
(3) Law of multiple proportions
 If elements A & B combine to form more than one compound, the
masses of B which can combine with a given mass of A are in
the ratios of small whole numbers
12g C + 16g O  CO or 12g C + 32g O  CO2
16g : 32g = 1:2
The Discovery of Atomic Structure
After Dalton’s atomic theory, not much of progress had been made and
no one had direct evidence for the existence of atom. Then, things started
to change in late 1800s…
 William Crooks (1832 ~ 1919): Cathode-ray tube (CRT) [1879]
A high voltage between two electrodes in a partially
evacuated tube generates electrical discharge
(cathode ray)
The Discovery of Atomic Structure
 J. J. Thomson (1856 ~ 1940) :
Discovery of electron [1897]
He discovered that cathode rays
are negatively charged particles, which
he originally called ``corpuscles’’ .
He won a Nobel prize in physics [1906].
(1) Rays are the same regardless of the
identity of the cathode material
(2) Conduct quantitative analysis of the
effect of electric and magnetic field
 determine the charge to mass ratio
charge/mass = 1.76  108 C/g
The Discovery of Atomic Structure
 Robert Millikan (1868 ~ 1953) :
Determine the charge of
electron [1907]
The machine on the right hand side is
the original apparatus Millikan used
to perform his oil-drop experiment.
He won a Nobel prize in physics [1923].
Millikan’s oil-drop experiment
Measured charge = 1.60 10-19 C
Electron mass = charge/[charge/mass]
= 9.10  10-28 g
The Discovery of Atomic Structure
 Ernest Rutherford (1871~1937):
Discovery of nucleus [1911]
He directed his graduate student
Hans Geiger and undergraduate student
Ernest Marsden to carry out -paticle
experiment. He won a Nobel prize in
chemistry [1908].
Rutherford’s -particle [4He2+]
scattering experiment
The Discovery of Atomic Structure
 Radioactivity: Generation of  - particles
  - ray: particles with +2 charge
  - ray: particles with 1 charge
  - ray: high energy radiation with no charge
The Discovery of Atomic Structure
 From the scattering experiment….
(1) Most -particles simply pass through the gold foil.
(2) Small amount of scattering was observed at large
angles.
 Rutherford postulated that..
(1) Most of the total volume of an atom is empty space.
(2) Most of the mass of an atom and all of its positive
charge reside in a very small region,
called nucleus.
Rutherford also found the existence of protons inside of nucleus [1919].
Another particle in nucleus, neutron, was found by James Chadwick
in 1932.
Early Models of an Atom
 J. J. Thomson’s model
“plum-pudding model”
 Rutherford’s model
Rutherford's Model:
+
Electrons are negatively charged, but atoms as a whole are neutral.
Modern View of Atomic Structure
The list of subatomic particles has grown considerably since the
discovery of electrons, but only the electron, proton and neutron have
a bearing on chemical behavior.
A convenient unit (non-SI) to describe the dimensions of
atoms and molecules is Angstrom (Å).
1 Å = 1 10-10 m = 100 pm
Modern View of Atomic Structure
 Properties of subatomic particles
Particle
Proton
Neutron
Electron
Charge (C)
Mass (g)
+1.60  10-19 (+1)
Mass (amu)
1.6727  10-24
1.0073
( 0)
1.6750  10-24
1.0087
1.60  10-19 (1)
9.1097  10-28
5.486  10-4
0
Every atom has an equal number of protons and electrons so that it has
no electrical charge
 Atomic Mass Unit (amu)
1 amu = 1/12 of the mass of carbon (12C) atom
= 1.66054  10-24 (g)
Modern View of Atomic Structure
The characteristics of each atom are determined by the numbers of
proton, neutron and electrons.
Hydrogen:
1 proton
Helium:
2 protons
2 neutrons
Lithium:
3 protons
4 neutrons
Beryllium:
4 protons
5 neutrons
 Atomic Number: The number of protons in the nucleus of an atom.
 Mass Number: The total number of protons plus neutrons in the atom
 Isotopes : Atoms with identical atomic numbers but different mass
numbers such as C-14 and C-12.
Modern View of Atomic Structure
Same information : An element is defined by
the number of protons
Atomic Weight
Atomic Mass Unit (amu) = 1.66054  10-24 g
12C
= 12 amu (exact), 1H = 1.0078 amu,
16O
= 15.9949 amu
 Average Atomic Masses : Weighted average of all the isotopes of
an element found in nature.
Example : Naturally occurring carbon is composed of 98.93% 12C
and 1.07 % 13C. What is the average mass of carbon?
(0.9893)(12 amu) + (0.0107)(13.00335) = 12.01 amu
fractional
abundance of
C-12
mass of C-12
mass of C-13
This is the mass
of carbon atom
shown in the
periodic table
Atomic Weight
Example: Boron has two naturally occurring isotopes: 10B (10.01 amu)
and 11B (11.01 amu). If the average atomic weight of Boron
is 10.81, what are the fractional abundances of the two isotopes?
Periodic Table
If the elements are arranged in order of increasing atomic number,
their chemical properties are found to show a repeating, or periodic,
pattern.
period
group
Elements having similar properties are placed in vertical columns
Periodic Table
Halogen
Alkaline earth
metal
Alkali
metal
Transition metals
rare
gas
= H2, N2, O2, F2, Cl2, Br2, I2
Molecules and Molecular
Compounds
Chemical Compounds
Molecular
Ionic
(1) Molecular compounds are composed of more than
one type of atom
H2O, NH3, CH3OH, O2
(2) Most molecular substances contain only non-metallic
atoms
O2, H2O, H2O2, CO, CO2, CH4
Molecules and Molecular
Compounds
 Chemical Formulars
(1) Molecular Formulas : Indicate the actual numbers and types
of atoms in a molecule
Ex. C2H4O2
(2) Empirical Formulas : Indicate the relative number of atoms of
each type in a molecule
Ex. CH2O
(3) Structural Formulas :
H
O
H–C–C–O–H
H
Molecules and Molecular
Compounds
 Picturing Molecular Compounds (Ex. Methane)
Structural Formula
Perspective drawing
Space-filling model
Ball-and-stick model
Ions and Ionic Compounds
 Ion : Atoms can readily gain or loose electrons and
become ions.
Cation: An ion with a positive charge
Na+
Anion: An ion with a negative charge
Cl
Ions and Ionic Compounds
 Which elements form cations and which form anions?
Metals tend to form Cations
Nonmetals tend to form Anions
VIII A
I A II A
Alkaline Earth Metals
Alkali Metals
III A IV A VA VI A VIIA
Halogens
Noble
Gases
Ions and Ionic Compounds
 How many electrons each element can gain or loose?
Each element tends to have the same number of
electrons as noble gases (rare gases).
Ions and Ionic Compounds
 Example: Determine the number of electrons, protons and
neutrons in each of the following ions
No. of Protons
16O240Ca2+
58Fe3+
80Br

No. of Neutrons No. of Electrons
Ions and Ionic Compounds
 Ionic Compounds : Cations (metals) and anions (non-metal)
combine to form ionic compounds
NaCl
Alternating positive and negative
charges
Ions and Ionic Compounds
 Ionic compounds :
(1) Ionic compounds are generally combination of metals
and nonmetals
NOTE: Molecular compounds are generally composed of
nonmetals only (H2O , CH3OH , CH3CH2Cl , …)
(2) Ionic compounds are represented by empirical formulas
 use simplest whole-number ratio of cations and anions
NOTE: There is no discrete (or isolated) molecule of NaCl
(3) Ionic compounds are always neutral. Therefore, the total
positive charge equals the total negative charge
Mg2+ and N3- form Mg3N2 : 3(+2) + 2(3) = 0
Ions and Ionic Compounds
 Example : Find the empirical formula for the ionic compound
made of given cation and anion
Na, O =>
Al, O =>
Ca, O =>
Naming Ions and Ionic Compounds
Names of ionic compounds consist of the cation name
followed by the anion name
CaCl2 = calcium + chloride  calcium chloride
 Names of Positive Ions (cations) :
(1) Cations formed from metal atoms have the same name
as the metal.
Na+  sodium ion, Zn+  zinc ion, Al3+  aluminum ion
NOTE: Ions formed from a single atom are called
monatomic ions
Naming Ions and Ionic Compounds
(2) If a metal can form different cations, the positive charge
is indicated by a Roman numerical in parenthesis following
the name of the metal
Fe2+  iron (II) ion
Fe3+  iron (III) ion
Cu+  copper (I) ion
Cu2+  copper (II) ion
These ions are usually transition metals
NOTE: Metals that form only one cation
group 1A  Na+, K+, Rb+
group 2A  Mg2+, Ca2+, Sr2+, Ba2+
and Al3+ (group 3A), Ag+ (group 1B), Zn2+ (group 2B)
Naming Ions and Ionic Compounds
(3) Cations formed from nonmetal atoms have names that
end in -ium
NH4+  ammonium ion
H3O+  hydronium ion
NOTE: These ions are examples of polyatomic ions
Naming Ions and Ionic Compounds
 Names of Negative Ions (anion) :
(1) The names of monatomic anions are formed by replacing
the ending of the name of the element with –ide.
H- hydrogen  hydride ion, O2- oxygen  oxide ion,
NOTE: polyatomic anions with common names ending with –ide
OH-  hydroxide ion, CN-  cyanide ion
(2) Polyatomic anions containing oxygen (oxyanions)
a) ending with –ate : reserved for the most common oxyanion
NO3-  nitrate ion, SO42-  sulfate ion
Naming Ions and Ionic Compounds
b) ending with –ite : used for oxyanion with the same charge,
but one fewer O atom than those ending with –ate.
NO2-  nitrite ion,
SO32-  sulfite ion
c) If a series of oxyanions extends to more than two members,
use prefix per- (one more) or hypo- (one fewer)
ClO4ClO3ClO2ClO-
 perchlorate ion (one more than –ate)
 chlorate ion
 chlorite ion
 hypochlorite ion (one fewer than -ite)
Naming Ions and Ionic Compounds
NOTE: Oxyanions with the maximum number of oxygens
(i) Charges increase from right to left.
(ii) Second row elements (C, N) have maximum 3 oxygen atoms
and third row elements (P, S, Cl) have maximum 4 oxygen
atoms (row # + 1).
(iii) All names end with –ate except for ClO4-
Naming Ions and Ionic Compounds
(3) Anions derived by adding H+ to an oxyanion are named by
adding as a prefix the word hydrogen or dihydrogen.
CO32- : carbonate ion  HCO3- : hydrogen carbonate ion
PO43- : phosphate ion  H2PO4- : dihydrogen phosphate ion
Halogen (7A)
Names of Binary Molecular
Compounds
(1) The name of the element farther to the left in the periodic table
appear first. (NOTE: Oxygen is always written last except when
combined with fluorine.)
(2) If both elements are in the same group, the one having the higher
atomic number is named first
(3) The name of the second element is given an –ide ending
(4) Greek prefixes are used to indicate the number of atoms of each element
(1  mono-, 2 di-, 3 tri-, 4 tetra-, 5  penta-, 6  hexa- )
Cl2O : dichloro monoxide
NF3 : nitrogen trifluoride
N2O4 : dinitrogen tetroxide P4S10 : tetraphosphorous decasulfide
Naming Compounds : Examples
 Before you try to name a compound :
(1) Is the compound ionic or molecular?
(2) For ionic compounds, find the name of each ion.
For molecular compounds, find the number of each atom.
BF3 :
NiO :
KMnO4 :
SO
Naming Compounds : Examples
 Write down the chemical formulas for the following compounds
(1) Sodium Nitride, Q: Is this ionic or molecular?
Q: Is anion monatomic or polyatomic ion?
(2) Diphosphorus pentoxide,
Naming Compounds : Examples
(1) NaClO :
(2) Fe2(CO3)3 :
(3) SF6 :
(4) aluminium hydroxide :
(5) ammonium sulfate :
(6) NaH2PO4 :