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Chapter 2
Atoms, Molecules, and Ions
Chapter 2: Topics
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Early history of chemistry
Fundamental chemical laws
Dalton’s atomic theory
Early experiments to characterize the
atom
The modern view of atomic structure
Molecules and ions
An introduction to the Periodic Table
Naming simple compounds
2.1 The early history of chemistry
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Greeks
Democritus and others - atomos
Alchemy
1660 - Robert Boyle- experimental
definition of element.
• Lavoisier- Father of modern chemistry.
Greeks
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Matter is composed of fire, earth, water and air
The Greek philosopher Democritus (460 B.C. –
370 B.C.) was among the first to suggest the
existence of atoms (from the Greek word
“atomos”)
 He believed that atoms were indivisible and
indestructible
 His ideas did agree with later scientific
theory, but did not explain chemical
behavior, and was not based on the scientific
method – but just philosophy
Alchemy
• Turning Cheep metals into gold
• Alchemists discovered several
elements and prepared mineral acids
17th Century
• Robert Boyle: First “chemist” to perform
quantitative experiments
• He published his first book: “The Skeptical
Chemist” in 1661.
• He talked about elements
18th Century
• George Stahl: Phlogiston flows
out of a burning material.
• Joseph Priestley: Discovers
oxygen gas, “dephlogisticated
air, i.e., low in phlogistone”
2.2 Fundamentals chemical Laws
• Law of Conservation of Mass
• Law of Definite Proportion
• Law of Multiple
Law of Conservation of Mass
 It was discovered by Antoine Lavoisier
 It was the basis for development of
chemistry in the 19th century
 Mass is neither created nor destroyed
 Combustion involves oxygen, not
phlogiston
Law of Definite Proportion
(Proust’s Law)
A given compound always contains
exactly the same proportion of
elements by mass.
Water is composed of 11.1% H and
88.9% O (w/w)
Law of Multiple Proportions
When two elements form a series of
compounds, the ratios of the masses of
the second element that combine with 1
gram of the first element can always be
reduced to small whole numbers.
The ratio of the masses of oxygen that
combine with 1g of H in H2O and H2O2
will be a small whole number (“2”).
Example
• Water, H2O has 8 g of oxygen per 1g of
hydrogen.
• Hydrogen peroxide, H2O2, has 16 g of
oxygen per 1g of hydrogen.
• 16/8 = 2/1
• Small whole number ratios.
• This fact could be explained in terms of
atoms
2.3 Dalton’s Atomic Theory (1808)
 Elements are made up of small particles
called atoms
 Atoms of each element are identical.
Atoms of different elements are
different.
 Compounds are formed when atoms
combine. Each compound has a always
same type and relative number of atoms
 Chemical reactions are rearrangement
of atoms but atoms are never changed
into atoms of other element. , or
created or destroyed.
Gay-Lussac hypothesis (1809)
• Provided basics to determining absolute formulas of
compounds
• Gay-Lussac- under the same conditions of
temperature and pressure, compounds always react
in whole number ratios by volume.
– 2volumes of H react with one volume of O to form
2volumes of gaseous water and
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
• 5 liters of nitrogen
• Same number of particles!
•If Avogadro's hypothesis is correct,
Gay-Lussac’s can be interpreted as
follows:
• 2 molecules of H react with 1 molecule
of O
2 molecules of H2O
2.4 Early experiments to characterize the atom
 Based on Dalton, Gay-Lussac,
Avogadro and others, work started
to identify the nature of the atom
 What is an atom made of? How
do atoms of various elements
differ?
The electron
 J. J. Thomson - postulated the existence of
electrons using cathode ray tubes.
 Ernest Rutherford - explained the nuclear
atom, containing a dense nucleus with
electrons traveling around the nucleus at a
large distance.
Thomson’s Experiment
Voltage source
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When high voltage is applied to the tube a ray
emanates from the cathode is called cathode ray.
Thomson’s Experiment
Voltage source
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Thomson’s Experiment
Voltage source
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Passing an electric current makes a beam
appear to move from the negative to the
positive end.
Thomson’s Experiment
Voltage source
• By adding an electric field
Thomson’s Experiment
Voltage source
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 By adding an electric field, he found that the
moving particles were negatively charged
Results of Thomson Experiment
• Electrons are produced from electrodes
made from various types of metals, all
atoms must contain electrons.
• Since atoms are electrically neutral,
they must contain positively charged
particles.
• Thomson determined charge-to-mass
ratio of an electron:
• e/m = -1.76X108C/g
Thomson’s Model
• Atom consisted of a
diffuse cloud of positive
charge with negative
electrons embedded
randomly
• Atom was like plum
pudding.
• Thomson believed that
the electrons were like
plums embedded in a
positively charged
“pudding,” thus it was
called the “plum pudding”
model.
Millikan’s Experiment
Atomizer
Oil droplets
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Telescope
Oil
Millikan’s Experiment
X-rays
X-rays give some electrons a charge.
Millikan’s Experiment
From the mass of the drop and the charge on
the plates, the mass of an electron is calculated
Radioactivity
• Certain elements produce high energy
radiation
• Discovered by accident and was a result of
spontaneous emission by uranium
• Bequerel (1896) found that a piece of mineral
containing uranium could produce an image
on a photographic plate in the absence of
light.
• Three types of radiation were known:
– alpha- helium nucleus (+2 charge, 7300 times that
of the electron)
– beta- high speed electron
– gamma- high energy light
The nuclear atom
Rutherford’s Experiment
• Aimed at testing Thomson’s plum pudding model
• Used uranium to produce alpha particles.
• Alpha particles are directed at gold foil through
hole in lead block.
• Since the mass is evenly distributed in gold
atoms alpha particles should go straight through.
• Used gold foil because it could be made atoms
thin.
Lead
block
Uranium
Florescent
Screen
Gold Foil
What he expected
Because
Because, he thought the mass was
evenly distributed in the atom.
What he got
How he explained it
• Atom is mostly empty
• Small dense,
positive particle at
center.
• Alpha particles
are deflected by
it if they get close
enough.
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Proof for nuclear atom
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Nuclear atom model
• According to Rutherford:
The atom consists of a dense center of
positive charge (Nucleus) with
electrons moving around it at distance
that is large relative to the nuclear
radius
2.5 The modern view of an atomic structure:
An introduction
• The atom is mostly
empty space.
• Two regions
• Nucleus- protons and
neutrons.
• It is characterized by
small size and high
density
• Electron cloud- region
where you might find an
electron.
• The chemistry of atom
• Results mainly from
electrons
A cross section of nuclear atom
Mass and charge of nuclear
particles
Particle
Electron
Proton
Neutron
Mass (Kg)
9.11X10-31
1.67X10-27
1.67X10-27
Charge
-1
+1
None
Why atoms of different elements
have different properties?
• Atoms of different elements have
different number of protons and
electrons.
• Number and arrangement of electrons
around nucleus differ from one element
to another.
Sub-atomic Particles
• Z - atomic number = number of protons
determines type of atom.
• A - mass number = number of protons +
neutrons.
• Number of protons = number of
electrons if atom is neutral.
Symbols
Mass number
Atomic number
Na-23
A
X
Z
23
Na
11
Isotopes
 Atoms of the same element (same atomic
number) with different mass numbers
 Atoms with the same number of protons,
but different numbers of neutrons.
Isotopes of chlorine
35Cl
37Cl
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17
chlorine - 35
Cl-35
chlorine – 37
Cl-37
Two isotopes of sodium
• Isotopes show almost identical chemical properties. Why?
• They possess same number of electrons
2.6 Molecules and ions
Introduction to chemical bonding
• The forces that hold atoms together are
called chemical bonding
• Covalent bonding - sharing electrons.
• Collection of atoms by covalent bonding lead
to molecules
• Molecules can be represented by formulas
• Chemical formula- Symbol relates number
and type of atoms in a molecule.
• Diatomic molecule: two atoms of same
element are connected by a covalent bond.
Molecular formula and structural formula
• Molecular formulas
– give the actual numbers and types of atoms
in a molecule.
– Examples: H2O, CO2, CO, CH4, H2O2, O2,
O3, and C2H4.
• Structural formula: bonds are shown as lines
H
H C
H
H
C H
H
Representing Structure in Molecules
Accurately represents the angles
at which molecules are attached.
Ions
• Atoms or groups of atoms with a charge.
• Cations- positive ions - get by losing
electrons(s).
• Anions- negative ions - get by gaining
electron(s).
• Ionic bonding- Force of attraction
between oppositely charged ions.
• Ionic solids are called salts.
Formation of Cations
Formation of Anions
Examples of ions
• Cation: A positive ion
Mg2+, NH4+
• Anion: A negative ion
Cl-, SO42-
Polyatomic ions
2.7 Introduction to the Periodic Table
• Elements are classified by:
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properties
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atomic number
• Groups (vertical)
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1A = alkali metals
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2A = alkaline earth metals
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7A = halogens
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8A = noble gases
• Periods (horizontal)
Periodic Table
Groups /Families
Periods
Metals
• Conductors
• Lose electrons
• Malleable and ductile
Nonmetals
• Brittle
• Gain electrons
• Covalent bonds
Semi-metals or Metalloids
Alkali Metals
Alkaline Earth Metals
Halogens
Transition metals
Noble Gases
Inner Transition Metals
Lanthanides
Lanthanides
Select an element
(
= Internet link )
2.8 Naming Simple Compounds
• Binary compounds are composed of
two electrons
• Both ionic and covalent compounds
will be considered
Binary Ionic Compounds(Type I)
1. Cation first, then anion
2. Monatomic cation takes its name from
the name of the element
Ca2+ = calcium ion
3. Monatomic anion takes its name from
the root of the element + -ide
Cl- = chloride
CaCl2 = calcium chloride
Name the following compounds
CsF
Calcium fluoride
AlCl3
Aluminum chloride
LiH
Lithium hydride
Calcium hydroxide
Ca(OH)2
Some Common Cations
Some Common Anions
Binary Ionic Compounds (Type II):
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metal forms more than one cation
use Roman numeral in name
PbCl2
Pb2+ is cation
PbCl2 = lead (II) chloride
PbCl4 = lead (IV) chloride
Example
• FeCl3
Iron(III) chloride
Ferric chloride
• FeCl2
Iron(II) chloride
Ferrous chloride
Example
• Write the name of the compound Fe2O3
Oxidation state of Fe = +3
Iron(III) oxide
• Group 1A, Group 2A and Al3+ do not
take Roman numerals
• Silver, Ag forms more than one
oxidation state but Roman numerals
are not used.
Common Cations and Anions
Ionic compounds with polyatomic ions
• Polyatomic ions are assigned special names that
need to be memorized
• Oxyanions: anions that contain an atom of a given
element and different numbers of oxygen atoms
• Polyatomic anions (with many atoms) containing
oxygen end in -ate or -ite. (The one with more
oxygen is called -ate.)
• Examples: NO3- is nitrate, NO2- is nitrite.
• (Exceptions: hydroxide (OH-), cyanide (CN-),
peroxide (O22-).)
 Examples:
NaNO3
Sodium nitrate
K2SO4
Potassium sulfate
Al(HCO3)3
Aluminum bicarbonate
or aluminum hydrogen
carbonate
Common Polyatomic Ions
Binary Covalent compounds (Type III):
Compounds between two nonmetals
Although they do not contain ions, they are
named similarly to binary ionic
compounds
- First element in the formula is named first
using full element name
- Second element is named as if it were an
anion.
- Use prefixes to denote number of atoms
- Never use mono- for naming first element
P2O5 = diphosphorus pentoxide
Acids
• Substances that produce H+ ions when
dissolved in water.
• All acids begin with H.
• Two types of acids:
– Oxyacids
– Non-oxyacids
Naming acids
• If the formula has oxygen in it
• write the name of the anion, but change
– ate to -ic acid
– ite to -ous acid
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Watch out for sulfuric and sulfurous
HNO3
HNO2
HC2H3O2
Naming acids
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If the acid doesn’t have oxygen
add the prefix hydrochange the suffix -ide to -ic acid
HCl
H2S
HCN