Chapter 2
Atoms, Molecules and Ions
The Early History of Chemistry
Before 16th Century
 Greeks were the first to attempt to explain why
chemical changes occur.
 Alchemy: Attempts to change cheap metals into
gold. They invented the idea of atoms, that matter
is not continuous. They discovered several
elements and learned to prepare mineral acids.
The Early History of Chemistry
16th Century
 German develop the systematic metallurgy
(extraction of metal from ores)
 Swiss develop the medicinal application of
minerals
17th Century
 Robert Boyle: First chemist to perform
quantitative experiments
Fundamental Chemical Laws
Law of Conservation of Mass (Antoine
Lavoisier, 18th Century) – Mass is neither
created nor destroyed
Law of Definite Proportion (Joseph Proust,
19th Century) – A given compound always
contains exactly the same proportion of
elements by mass. This principle of constant
composition of compounds is a law of definite
proportion. Carbon tetra chloride is always 1
atom carbon per 4 atoms chlorine
Fundamental Chemical Laws
Law of Multiple Proportions (John Dalton,
19th Century) – When two elements form a
series of compounds, the ratio of the masses of
the second element that combine with 1g of the
first element can always be reduced to small
whole numbers. The ratio of the masses of
oxygen in H2O and H2O2 will be a small whole
number (“2”).
Dalton’s Atomic Theory (1808)
1. Each element is made up of tiny particles called atoms
2. The atoms of a given element are identical; the atoms
of different elements are different in some fundamental
way or ways
3. Chemical compounds are formed when atoms of
different elements combine with each other. A given
compound always has the same relative numbers and
types of atoms
4. Chemical reactions involve reorganization of the atoms
– changes in the way they are bound together. The
atoms themselves are not changed in a chemical
reaction.
Avogadro’s Hypothesis (1811)


At the same temperature and pressure,
equal volumes of different gases contain
the same number of particles
5 liters of oxygen and 5 liters of nitrogen
contain the same number of particle
Thomson Atomic Model (1903)



An atom consists of a diffuse cloud of positive
charge with the negative electrons embedded
randomly in it. This model is often called plum
(or raisin) pudding model.
Observed cathode ray (produced at the
negative electrode and repelled by the negative
pole of an applied electric field.
Cathode ray was a stream of negatively
charged particles now called electrons
Deflection of Cathode Rays by an Applied Electric Field
The Plum Pudding Model of the Atom
Rutherford Atomic Model (1911)
An atom with a dense center of positive
charge (the nucleus) with electrons moving
around the nucleus at a distance that is large
relative to the nuclear radius. Nucleus is
very small compared with the overall size of
the atom. Nucleus is extremely dense,
accounts for almost all of the atom’s mass.
Radioactivity
Spontaneous emission of radiation
 Gamma () rays: high energy light
 Beta () particles: high speed electron
 Alpha () particles (He2+): 2+ charge,
charge twice that of electron and with
opposite sign. The mass of an  -particle
is 7300 times that of the electron
Rutherford’s Experiment on -particle Bombardment of Metal Foil
Expected and Actual Results of Rutherford’s Experiment
The Modern View of Atomic Structure
The atom contains:
 Electrons: move around the nucleus (mass:
9.11 X 10-31 kg, Charge 1-)
 Protons: found in the nucleus, they have a
positive charge equal in magnitude to the
electron’s negative charge (mass: 1.67 X 1027 kg, charge 1+)
 Neutrons: found in the nucleus, virtually
same mass as a proton but no charge. (mass:
1.67 X 10-27 kg, charge: 0)
Nuclear Atom Viewed in Cross Section
The Chemists’ Shorthand Atomic Symbols
Mass number 
Atomic number
39
K
 19
 Element Symbol
Atomic number (Z): number of protons, gives
the symbol of the element (X)
Mass number (A): Total number of protons
and neutrons
Elemental form = Zero net charge
Therefore, # electrons = # of protons
Isotopes
Atoms with the same number of protons but
different number of neutrons. In nature most
elements contain mixtures of isotopes
23
11Na
: 11 protons, 11 electrons, and 12 neutrons
24 Na : 11 protons, 11 electrons, and 13 neutrons
11
Two Isotopes of Sodium
Molecules and Ions




Chemical Bonds: The forces that hold atoms
together in compounds. H2O, NO, CO2
Covalent bonds: Covalent bonds result from
atoms sharing electrons. Cl2
Ionic bonds: Force of attraction between
oppositely charged ions.
Molecule: A collection of covalently-bonded
atoms. H2, O2
Ions

Ions: An ion is an atom or group of atoms
that has a net positive charge or negative
charge particle (an unequal number of
protons and electrons) is obtained by
removing or adding electrons. Na+, Cl-

Cation: A positive ion (Na+, Mg2+, NH4+)
Anion: A negative ion (Cl-, SO42-)

Formulas


Chemical Formula: In which the symbols for
the elements are used to indicate the types of
atoms present and subscripts are used to
indicate the relative numbers of atoms.
CO2 indicates each molecule contains 1 atom of
carbon and 2 atoms of oxygen.
Structural Formula: In which the individual
bonds are shown by lines. It may or may not
indicates the actual shape of the molecules.
O=C=O
Periodic Table






Periodic table is organized based on the
properties that elements have in common with
one another.
Groups: Elements in the same vertical columns
are in the same group have similar chemical
properties.
Group 1A: Alkali metals: Li, Na, K, Rb, Cs, Fr
Group 2A: Alkaline earth metals: Be, Mg, Ca, Sr, Ba,
Ra
Group 7A: Halogens: F, Cl, Br, I, At (astatine)
Group 8A: Noble gases: He, Ne, Ar, Kr, Xe, Rn (radon)
The Periodic Table
Periodic Table






Periods: The horizontal rows of elements in the
periodic table are called periods.
First period: horizontal row one contains H and He
Second period: row two contains Li through Ne
Letters in the boxes are the symbols for the
elements
Abbreviations are based on the current element
names or the original names.
The number above each symbol is the atomic
number (number of protons)
Periodic Table



Most of the elements are metals in the periodic
table.
Metals: Conduction of heat and electricity,
malleability, ductility, lustrous, form positive
ions
Nonmetals: appear in the upper right hand
corner of the periodic table except hydrogen.
Nonmetals lack the physical properties that
characterize the metal, gain electrons in
chemical reaction and form negative ions, form
covalent bond to each other.
Naming Compounds



Binary Compounds: Compounds composed of
two elements
Binary Ionic Compounds (Type 1): contains a
positive ion (cation) always written first in the
formula and a negative ion (anion)
Rules:
1. The cation is always named first and the anion
second
2. A monatomic (meaning one atom) cation takes its
name from the name of the element
3. A monatomic anion is named by taking the root of
the element name and adding –ide
Binary Ionic Compounds
Compound
Ions Present
Name
NaCl
Na+, Cl-
Sodium chloride
KI
K+, I-
Potassium iodide
CaS
Ca2+, S2-
Calcium sulfide
Li3N
Li+, N3-
Lithium nitride
CsBr
Cs+, Br-
Cesium bromide
MgO
Mg2+, O2-
Magnesium oxide
Naming compounds





Binary Ionic Compounds (Type II): Metals that
form more than one type of positive ion. Fe2+ and
Fe3+
Transition metals form several positive oxidation
states
Charge on the metal ion must be specified
Roman numeral indicates the charge of the
cation. Iron (II) chloride and iron (III) chloride
The ion with the higher charge has a name ending
in –ic and the one with the lower charge has a
name ending in –ous; ferrous chloride and ferric
chloride
Naming Compounds
Ionic Compounds with Polyatomic Ions:
 Need to know the names of the polyatomic ions
(Table 2.5).
NH4+  ammonium, SO42-  sulfate
 Na2SO4 Sodium sulfate
 KH2PO4 Potassium dihydrogen phosphate
 Fe(NO3)3 Iron(III) nitrate
 CsClO4 Cesium perchlorate
 NaOCl sodium hypochlorite
 Al2(Cr2O7)3 Aluminum dichromate
 Sr(CN)2 Strontium cyanide
Naming compounds
Binary Covalent Compounds (Type III):
Formed between two nonmetals
Rules:
 The first element in the formula is named
first, using the full element’s name
 Second element is name as if it were an anion
 Use prefixes to denote the number of atoms
present
 Never use mono – prefix for naming the first
element
CO ==> carbon monoxide, not
monocarbon monoxide
P2O5 ==> diphosphorus pentoxide
S2Cl4 ==> disulfur tetrachloride
NO2 ==> nitrogen dioxide
N2O5 ==> dinitrogen penoxide
Common Cations and Anions
Formulas from Names
Name
Chemical Formula
Diphosphorus pentasulfide
P2S5 (two non metals)
Cesium peroxide
Cs2O2 (Cs1+, O22-)
Aluminum fluoride
AlF3 (Al3+, F-1)
Vanadium (v) fluoride
VF5 (V5+, F-1)
Dioxygen difluoride
O2F2 (two non metals)
Gallium oxide
Ga2O3 (Ga3+, O2-)
Ammonium dichromate
(NH4)2Cr2O7 (NH4+, Cr2O72-)
Cupric phosphate
Cu3(PO4)2 (Cu2+, PO43-)
Flowchart for Naming Binary Compounds
Acids



When dissolved in water produce a solution
containing free H+ ions (protons)
An acid is a molecule with one or more H+
ions attached to an anion
If the anion does not contain oxygen, the acid
is named with the prefix hydro – and the
suffix –ic
HCl Hydrochloric acid
HCN Hydrocyanic acid
Acids

If the anion contains oxygen, the acidic name
is formed from the root name of the anion with
the suffix of –ic or –ous depending on the
anion
HNO3 Nitric acid (Nitrate anion)
H2SO4 Sulfuric acid (Sulfate anion)
H3PO4 Phosphoric acid (Phosphate anion)
HC2H3O2 Acetic acid (Acetate anion)
H2SO3 Sulfurous acid (sulfite anion)
HNO2 Nitrous acid (nitrite anion)
Flowchart for Naming Acids
Summary











Laws
Dalton’s atomic Theory
Avogadros’ Hypothesis
Various Atomic Models
Radio Activity (, ,  rays)
Atomic Symbol (atomic #, mass#, isotopes)
Chemical Bonds (covalent bonds, Ionic bonds)
Periodic Table
Naming various types of compounds
Formulas
Acids