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Transcript
CH 2: Atoms,
Molecules, and Ions
Chapter Outline
 History


of chemistry (2.1 – 2.4)
Chemical laws – start with slide 8 for content, 3-7 FYI slides
Path to the atom
 Modern
atomic structure (2.5)
 Molecules vs. Ions (2.6)
 Naming molecular and ionic compounds
(2.8)
 Introduction to the periodic table (2.7)
History of Chemistry
Greek Philosophers: 5th Century BCE
(BCE = before the common era - replaces BC)
 The
Greek philosophers were the first to
reflect on the nature of matter.

They proposed that all matter is made out of
first 4 elements -- earth, air , water,
fire. Aristotle added a fifth element, plasma
(also called ether).
Greek Philosophers
 Democritus
had an alternate view of
matter. He proposed that matter
was made up of tiny particles called
atoms.

His "theory" was not well accepted at the time.
Alchemy
Alchemy: ~600-1600's CE


(CE = common era, replaces AD)
Alchemy developed at about the same time in
China, India, and Greece. It spread into
Europe in the 8th century.
Alchemy
Alchemists had two pursuits
1.
2.
Search for a means to convert “base” metals
into gold
Search for the elixir of life
•
Substance that would lead to immortality
Alchemy
Advances from Alchemy

Many new substances where identified
• Plaster of Paris, nitric acid….


New lab techniques and equipment
developed
New medicines identified
Modern Chemistry, ~1600 on

First chemists/physicists to use scientific method









Boyle - elements
Lavoisier – law of conservation of matter
Proust - law of definite proportions
Dalton – law of multiple proportions, atomic theory
Avogadro - hypothesis
Thomson – charge to mass ratio for an electron
Millikan – charge on the electron
Bequerel and the Curies - radioactivity
Rutherford – nuclear atom
Modern Chemistry
Robert Boyle: ~1660
 Proposed a substance to be an element
unless it can be broken down into simpler
substances.

Proposed one of the gas laws – CH 5
Lavoisier: ~1760
Law of Conservation of Matter

Matter is neither created nor destroyed in a
chemical reaction.
Proust: late 1700s
Law of Definite Composition/Proportions

A given compound always contains the same
proportion of elements by mass.
John Dalton: ~1800
Law of Multiple Proportions

When two elements form more than one
compound, the ratios of the masses of the
second element that combines with one gram
of the first element can always be reduced to
small whole numbers.
Law of Multiple Proportions
Example
 Consider
two 100.0 g samples 2 different
compounds containing only C and H.
Compound A: 27.2 g of C
 Compound B: 42.9 g of C.
Show how this data illustrates the law of
multiple proportions.
 Dalton
also proposed the first table of
atomic masses

Most masses later need revision
 Dalton
Theory
is best known for proposing Atomic
Dalton’s Atomic Theory
Elements are made up of tiny particles
called atoms.
1.
•
Atoms are indivisible and indestructible
Atoms of a given element are identical
2.

Atoms of different elements differ in some
fundamental way(s)
Dalton’s Atomic Theory
Compounds form when atoms of
different elements combine with each
other.
3.

A given compound always has the same
relative number and types of elements.
Dalton’s Atomic Theory
Chemical reactions occur when atoms
change how they are bound to each
other.
4.

Individual atoms are not changed, just
rearranged
Avogadro: 1811
Avogadro's Hypothesis

At the same temperature and pressure equal
volumes of gases contain the same number of
particles.
• Based on Guy-Lussac’s data
• See page 41/42
From Dalton to Atomic Structure
 Dalton’s
atomic theory lead to much
research on the nature of the atom.
 This research showed the atom to made
up of smaller particles.
J.J. Thomson
 Thomson
measured the deflection of a
cathode ray beam in electrical and
magnetic fields of known strengths.
Applied electrical field
Cathode ray
+
(+)
(-)
Metal
electrode
Cathode ray tube experiment, pg 43
Metal
electrode


Thomson found the cathode rays were
attracted by the positive charge and repelled
by negative
These findings clearly indicated that the rays
consisted of negatively charged particles.
• Today we know these particles as electrons.
 Thomson
measured the deflection of the
beam in a magnetic field and more! From
his data he determined the charge:mass
ratio for an electron
e = - 1.76 x 108 C/g
m
e = charge on the electron in coulombs
M = mass of an electron in grams
Cathode ray tube experiment
 Thomson
also found that the cathode ray
particles were identical regardless of
source.

Concluded all elements contain these
negative particles (electrons)
J.J. Thomson
Thomson:



identified cathode ray beams as a stream of
negatively charged particles
calculated the charge to mass ratio for these
negatively charged particles
proposed the existence of positively charged
particles
• To balance the negative charge of the electrons
Millikan ~1909
 Millikan’s
oil drop experiment allowed him
to determine the charge on an electron

This charge can be plugged into Thomson’s
formula and the mass of the electron
calculated
• Mass electron = 9.11 x 10-31 kg
• Page 44
Radioactivity
Becquerel, Marie and Pierre Curie: ~1896
 Henri Becquerel - observed the natural emission
of energy/rays by uranium.

Marie and Pierre Curie studied “Becquerel's
rays”.


The Curies’ findings suggested that matter was
composed of smaller particles than atoms.
The Curies coined the term radioactivity to describe
the rays emitted.
Radioactivity
 Three



types of radioactivity were identified:
gamma rays - very high energy light
beta particles - high energy electrons
alpha particles - He+2 particles
• 2 protons and 2 neutrons
Plum Pudding Model of the Atom
 In
the early 1900’s the accepted model of
the atom was called the plum pudding
model of the atom

Electrons (tiny and negatively charged) were
pictured to be dispersed in a ‘cloud’ of positive
charge.
• Proposed by JJ Thomson and Lord Kelvin in 1904
Rutherford and the Nuclear Atom
 In
1911 an experiment conducted in
Ernest Rutherford’s lab showed the “plum
pudding” model to be incorrect.


Experiment was conducted by Geiger and
Marsden and the findings interpreted by
Rutherford.
See page 45
Rutherford’s Atom
 First


to propose a nuclear atom.
An atom has a dense positive center
containing all of positive charge and most of
the mass of the atom – the nucleus
Electrons are scattered in the empty space
around the nucleus
• Electrons occupy a volume that is huge as
compared to the size of the nucleus.
A New Model of the Atom
Expected based on
Plum pudding model
Rutherford’s model
Based on ”his” results
Modern Atomic Structure
 Rutherford
continued to study the atom
and the positive matter of the atom.

1919, + particle named the proton
 ~1932
James Chadwick proposed the
existence of a third subatomic particle, the
neutron.
Subatomic Particles
Subatomic
Particle
Charge
Mass, amu
Location in atom
Electron
(e-)
-1
0 amu
Outside of nucleus
Proton (p)
+1
~1 amu
Nucleus
Neutron (n)
0
~1 amu
Nucleus
Mass of Subatomic Particles
 Protons
and neutrons have ~ the same
mass (in the range of 10 -27 kg).

Mass of each and of individual atoms is often
expressed in amu rather than grams
• Atomic mass unit (amu) – 1/12 the mass of a
carbon-12 atom
Mass of Subatomic Particles
 The
mass of the electron (10-31 kg) is tiny as
compared to that of the proton and
neutron (10-27 kg) .

Therefore, the electron’s mass is considered
to be ~0 amu when calculating the mass of an
atom.
Subatomic Particles and the
Elements
 Each
element has a unique number of
protons.

Number of protons defines the element.
Atomic # = # protons
6
C
Subatomic Particles
 Since
atoms are neutral, for every proton
there is a/n _________.
 When
atoms interact to form compounds,
it is their ___________ that interact.
Terms
 Mass
number = sum of the # of protons
and the # neutrons in the nucleus of an
atom

FOR MOST ELEMENTS THE MASS
NUMBER IF NOT ON THE PERIODIC
TABLE.
• You will be given enough information to determine
mass number or number of neutrons.
Terms
 Isotopes
= atoms of a given element that
differ in mass number


Isotopes have the same number of
_____________.
Isotopes differ in the number of _______.
Isotopes
 Writing

atomic symbols for isotopes
pg 46
Mass #
11
5
Atomic #
B
Symbol for
element
FAQ - Isotopes
 When
is mass number found on the
periodic table?
 What’s
the atomic mass? Is it the same
as the mass number?
84
(209)
Po
6
12.0107
C
Molecules and Ions (2.6)
 Atoms
of different elements combine to
form compounds

Atoms in compounds are held together by
chemical bonds.
• Bonds involve interactions of the bonding atoms’
________
Bonding
There are two types of bonds:
1.
Covalent bonds – bonding atoms share
electrons
•
•
•
Atoms are always nonmetal atoms
Covalently bonded atoms form molecules
Ways to represent molecules


Chemical formula; H2O
Structural formula
O
H
H
Bonding
2.
Ionic bonds – attractive force among
oppositely charged ions
•
•
Bond formed between metal cations and
nonmetal anions
No molecules involved
Ions - Terms
 Ion

– charged atom or group of atoms
Formed when atoms gain or lose electrons
 Cation

Formed when an atom _______ electrons
 Anion

– positively charged ion
– negatively charged ion
Formed when an atom ______ electrons
Ions
 Describing
ion formation

Cation example:

Anion example:
Naming Binary Compounds
compounds – compound composed
of 2 elements
 Binary

NaCl

CO

CO2
Types of Binary Compounds
 Type

Metal forms only one ion
 Type


II binary ionic compounds
Metal forms more than one ion
Use roman numerals to indicate the charge
on the ion
 Type

I binary ionic compounds
III binary covalent compounds
Compound between 2 nonmetals
Types I Binary Compounds
 Compound
between a metal and a
nonmetal

Metal forms only one ion
 Name


the cation and then the anion.
Name of the cation is the name of the element
Name of the anion is the name of the
nonmetal with the ending changed to “ide”
Monoatomic cations to know
Group # Charge on ion examples
IA
+1
Na1+ sodium (ion)
K1+ potassium (ion)
IIA
+2
Mg2+ magnesium (ion)
IIIA
metals
+3
Al3+
aluminum (ion)
Monoatomic anions to know
Group # Charge on ion examples
VA
-3
N3P3-
nitride (ion)
phosphide (ion)
VIA
-2
VIIA
-1
O2S2F1Cl1Br1I1-
oxide (ion)
sulfide
fluoride (ion)
chloride (ion)
bromide (ion)
iodide (ion)
Practice
 Name
 chemical formula
 Chemical
formula  name