Chemistry 101 : Chap. 2

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Chemistry 101 Chap. 2
Atoms, Molecules and Ions

• 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

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The Atomic Theory of Matter
The history of development of atomic theory of
matter begins in ancient Greece. However, modern
atomic theory has its 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.
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The Atomic Theory of Matter
? Aristotle (384 322 BC)
Aristotle proposed 4 element theory of matter.
Fire
dry
hot
Earth
Air
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.
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The Atomic Theory of Matter
? John Dalton (1766 1844)

• Daltons Atomic Theory
• Each element is composed of atoms
• 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.

Dalton proposed that all matter is made up of
atoms and stated that elements are the simplest
form of matter.
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The Atomic Theory of Matter

• What can Daltons 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 12
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The Discovery of Atomic Structure
After Daltons 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)
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The Discovery of Atomic Structure

• J. J. Thomson (1856 1940)
• Discovery of electron 1897

• 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

He discovered that cathode rays are negatively
charged particles, which he originally called
corpuscles . He won a Nobel prize in physics
1906.
charge/mass 1.76 ? 108 C/g
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The Discovery of Atomic Structure

• Robert Millikan (1868 1953)
• Determine the charge of
• electron 1907

Millikans oil-drop experiment
Measured charge 1.60 ?10-19 C Electron mass
charge/charge/mass
9.10 ? 10-28 g
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.
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The Discovery of Atomic Structure

• Ernest Rutherford (18711937)
• Discovery of nucleus 1911

Rutherfords ?-particle 4He2 scattering
experiment
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.
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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

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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.
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Early Models of an Atom

• J. J. Thomsons model
• plum-pudding model

? Rutherfords model
Electrons are negatively charged, but atoms as a
whole are neutral.
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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
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Modern View of Atomic Structure
? Properties of subatomic particles
Particle Charge (C)
Mass (g) Mass (amu) Proton
1.60 ? 10-19 (1) 1.6727 ? 10-24
1.0073 Neutron 0 (
0) 1.6750 ? 10-24
1.0087 Electron ?1.60 ? 10-19 (?1)
9.1097 ? 10-28 5.486 ? 10-4
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)
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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.

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Modern View of Atomic Structure
Same information An element is defined by
the number of protons
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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
This is the mass of carbon atom shown in
the periodic table
fractional abundance of C-12
mass of C-12 mass of C-13
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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?
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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
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Periodic Table
Halogen
Alkaline earth metal
Transition metals
Alkali metal
rare gas
H2, N2, O2, F2, Cl2, Br2, I2
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Molecules and Molecular Compounds
Chemical Compounds
Molecular
Ionic

• 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
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Molecules and Molecular Compounds
? Chemical Formulars

• 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
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Molecules and Molecular Compounds
? Picturing Molecular Compounds (Ex. Methane)
Structural Formula
Perspective drawing
Space-filling model
Ball-and-stick model
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Ions and Ionic Compounds

• Ion Atoms can readily gain or loose electrons
  and
• become ions.

Anion An ion with a negative charge
Cation An ion with a positive charge
Cl?
Na
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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
III A
IV A
VA
VI A
VIIA
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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).
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Ions and Ionic Compounds

• Example Determine the number of electrons,
  protons and
• neutrons in each of the
  following ions

No. of Protons No. of Neutrons No. of
Electrons
16O2- 40Ca2 58Fe3 80Br ?
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Ions and Ionic Compounds

• Ionic Compounds Cations (metals) and anions
  (non-metal)
• combine to
  form ionic compounds

NaCl
Alternating positive and negative charges
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Ions and Ionic Compounds
? Ionic compounds

• 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
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Ions and Ionic Compounds

• Example Find the empirical formula for the
  ionic compound
• made of given cation and
  anion

Na, O gt
Al, O gt
Ca, O gt
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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
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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 Cu ? copper
(I) ion Fe3 ? iron (III) 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)
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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
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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)

• ending with ate reserved for the most common
  oxyanion
• NO3- ? nitrate ion, SO42- ? sulfate
  ion

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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)
ClO4- ? perchlorate ion (one more than
ate) ClO3- ? chlorate ion
ClO2- ? chlorite ion ClO- ?
hypochlorite ion (one fewer than -ite)
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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-
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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)
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Names of Binary Molecular Compounds

• 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
P4S10 tetraphosphorous decasulfide
N2O4 dinitrogen tetroxide
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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
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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,
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Naming Compounds Examples
(1) NaClO
(2) Fe2(CO3)3
(3) SF6
(4) aluminium hydroxide
(5) ammonium sulfate
(6) NaH2PO4