Transcript Chapter 4 Atomic Structure

Chapter 5
“Atomic Structure”
Chemistry
Olympic High School
Mr. Daniel
Credits: Stephen L. Cotton
Charles Page High School
Section 5.1 Defining the Atom

OBJECTIVES:
 Describe the history of the development of
ideas about atoms.
Explain
Dalton’s atomic theory.
Describe
the size of the atom.
Section 5.1 Defining the Atom
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” –
meaning “unable to be cut”)
Antoine Laurent
Lavoisier (1743-1794)
The Father of Modern
Chemistry: discovered
oxygen & hydrogen,
developed modern
thermodynamics,
invented the first periodic
table
Dalton’s Atomic Theory (experiment based)
1) All elements are composed of tiny
indivisible particles called atoms
John Dalton
(1766 – 1844)
2) Atoms of the same element are
identical. Atoms of any one element
are different from those of any other
element.
3) Atoms of different elements combine in simple wholenumber ratios to form chemical compounds
4) In chemical reactions, atoms are combined, separated,
or rearranged – but never changed into atoms of
another element.
Section 5.2
Structure of the Nuclear Atom

OBJECTIVES:
Identify three types of subatomic particles and
their properties.
 Describe the structure of atoms, according to
the Rutherford Model.

Section 5.2
Structure of the Nuclear Atom

One change to Dalton’s atomic theory is that atoms
are divisible into subatomic particles:
 Electrons, protons, and neutrons are examples of
these fundamental particles
 There are many other types of particles, but we
will study these three
In 1897, J.J. Thomson used a cathode ray tube to deduce
the presence of a negatively charged particle: the electron
JJ Thomsonl
The electron’s charge-to-mass ratio = 1.76  108 C/g
Cathode Ray Tubes
A Standard Television
Tube (First common in
the 1950’s)
A CRT from a Radar
Scope (WW II)
Standard Computer Monitor
(often called a “CRT”)
Cathode ray tubes pass electricity through a gas
that is contained at a very low pressure.
simulation
Robert Millikan
Mass of the Electron
Mass of the
electron is
9.11 x 10-28 g
The oil drop apparatus
1916 – Robert Millikan determined the mass of the
electron: 1/1840 the mass of a hydrogen atom;
has one unit of negative charge
Conclusions from the Study of the Electron:
a) Cathode rays have identical properties
regardless of the element used to produce
them. All elements must contain identically
charged electrons.
b) Atoms are neutral, so there must also be
positive particles in the atom to balance the
negative charge of the electrons
c) Electrons have very little mass, therefore atoms
must contain other particles that account for
most of the mass
Thomson’s Atomic Model
J. J. Thomson
Thomson believed that the electrons were like plums
embedded in a positively charged “pudding,” thus it was
called the “plum pudding” model.
Other Particle Discoveries:
 Eugen Goldstein first observed evidence of
what is now called the “proton” in 1886 particles with a positive charge, and a mass of
1840 times that of an electron. It’s existance
was later confirmed by Ernest Rutherford in
1919.
 1932 – James Chadwick confirmed the
existence of the “neutron” – a particle with no
charge, but a mass nearly equal to a proton
The problem:
In the following pictures, there is a target hidden by a
cloud. To figure out the shape of the target, we shot
some beams into the cloud and recorded where the
beams came out. Can you figure out the shape of the
target?
Target #1
The Answer:
Target #2
The Answer
Ernest Rutherford’s
Gold Foil Experiment - 1911
 Alpha particles are helium nuclei - The alpha
particles were fired at a thin sheet of gold foil
 Particle that hit on the detecting screen (film) are
recorded
Rutherford’s Findings
 Most of the particles passed right through
 A few particles were deflected
 VERY FEW were greatly deflected
“Like howitzer shells bouncing off
of tissue paper!”
Conclusions:
a) The nucleus is small
b) The nucleus is dense
c) The nucleus is positively
charged
The Rutherford Atomic Model
Based on his experimental evidence, Rutherford’s
Nuclear Model was developed and stated:
The atom is mostly empty space
 All the positive charge, and almost
all the mass is concentrated in a
small area in the center. He called
this a “nucleus”
 The electrons are distributed around the nucleus,
and occupy most of the volume
 Because of the exceptionally high mass of the
nucleus, it must contain particles in addition to
protons (neutrons were discovered later)

Ernest Rutherford video
Subatomic Particles
Crash Course Chemistry - History of Chemical Concepts
Particle
Electron
Charge
Mass (g)
Location
-1
9.11 x 10-28
Electron
cloud
+1
1.67 x 10-24
Nucleus
(e-)
Proton
(p+)
Neutron
(no)
(1 atomic mass unit)
0
1.67 x 10-24
(1 atomic mass unit)
Nucleus
Sizing up the Atom
 Elements can be subdivided into smaller and
smaller pieces until there are only single atoms; The
smallest particles of an element that still have the
properties of that element.
100,000,000 copper atoms in a single file, would
be approximately 1 cm long
Despite their incredibly small size, individual
atoms have recently become observable with
scanning tunneling microscopes
The
Beginning:
Atomic
Kanji Xe on Ni
“original
child”
Image:
atoms
on
a Cu
surface
Forming
aFe
Image:
Fe
atoms
on
a Cu surface
Image:
Fe
atoms
on
a of
Cu
Carbon
Monoxide
Man
An
Artistic
View
Nisurface
Section 5.3
Distinguishing Among Atoms

OBJECTIVES:
 Explain what makes isotopes different from
each other.
Calculate the number of protons, neutrons
and electrons in an atom using atomic
number and mass
Calculate
the average atomic mass and
atomic number.
Atomic Number
Atoms are composed of protons, neutrons, and electrons
Element
# of
# of
Atomic #
 How then are atoms of one element different from
protons electrons
(Z)
another element?
Elements
Carbonare different
6 because they
6 contain different
6
numbers of PROTONS
Phosphorus
15
15
15
Atomic number (Z) : the number of protons in the nucleus of
each atom of that element.
Iodine
53
53
53
therefore:
Gold  An atom
79is always neutral,
79
79
# protons in an atom = # electrons
Mass Number
Mass number is the number of protons and
neutrons in the nucleus of an atom:
Mass # = p+ + n0
p+
n0
e- Mass #
8
10
8
18
Arsenic - 75
33
42
33
75
Phosphorus - 31
15
16
15
31
Nuclide
Oxygen - 18
Complete Symbols

Complete Symbols contain the symbol of the
element, the mass number and the atomic
number.
Examples:
Superscript →
Subscript →
Mass
number
Atomic
number
12
6
X
14
6
235
C
C
U
92
Symbols continued…

We can also put the mass number after the name
of the element:
 carbon-12
 carbon-14
 uranium-235
Symbols

Find each of these:
a) number of protons
b) number of neutrons
c) number of electrons
d) Atomic number
e) Mass Number
80
Br
35
a) 35
b) 45
c) 35
d) 35
e) 80
Symbols

If an element has an atomic number of 34
and a mass number of 78, what is the:
a) number of protons
a) 34
b) number of neutrons
b) 44
c) number of electrons
c) 34
78
d) complete symbol
d)
Se
34
Symbols

If an element has 91 protons and 140
neutrons what is the
a) Atomic number
a) 91
b) Mass number
b) 231
c) number of electrons
c) 91 231
d) complete symbol
d)
Pa
91
Symbols

If an element has 78 electrons and 117
neutrons what is the
a) Atomic number
a) 78
b) Mass number
b) 195
c) number of protons
c) 78 195
d) complete symbol
d)
Pt
78
Isotopes
Dalton was wrong about all elements of the same
type being identical:

Atoms of the same element can have different
numbers of neutrons.
Thus, they have different mass numbers.
These are called isotopes.
Isotopes are atoms of the same element having
different masses, due to varying numbers of
neutrons.
Isotope
Protons Electrons
Neutrons
Hydrogen–1
(protium)
1
1
0
Hydrogen-2
(deuterium)
1
1
1
1
1
2
Hydrogen-3
(tritium)
Nucleus
Isotopes
Elements occur
in nature as
mixtures of
isotopes.
Atomic Mass


How heavy is an atom of oxygen?
 It depends, because there are different kinds of
oxygen atoms.
We generally refer to the average atomic mass.
Average atomic mass is based on the abundance
(percentage) of each isotope of an element as it is
found in nature.
 It is the number (red) that we find on the periodic
table
Measuring Atomic Mass

Instead of grams, the mass unit we use for atoms
is the Atomic Mass Unit (amu or µ)

It is defined as one-twelfth the mass of a carbon-12
atom.
Carbon-12 is chosen because of its isotope purity.
Each isotope has its own atomic mass, thus we
determine the average mass from the percent
abundance of each isotope.
Atomic Masses
Atomic mass is the average of all the naturally
occurring isotopes of that element.
Isotope
Symbol
Carbon-12
12C
Carbon-13
13C
Carbon-14
14C
Composition of
the nucleus
6 protons
6 neutrons
6 protons
7 neutrons
6 protons
8 neutrons
Carbon = 12.0107 µ
% in nature
98.89%
1.11%
<0.01%
To calculate the average atomic mass:


Multiply the mass number
of each isotope by its
abundance (expressed as
a decimal), then add the
results.
The mass of the isotope
is expressed in atomic
mass units (amu or µ)
Isotope
Carbon-12
Carbon-13
Carbon-14
Symbol
% in nature
12C
98.89%
13C
14C
1.11%
<0.01%
(12.00 µ) (.9889) + (13.00 µ) (.0111) = 12.01 µ
Atomic Mass Calculation
Element X has two natural isotopes. The isotope
with a mass of 10.012 amu has a relative
abundance of 19.91%. The isotope with a mass
of 11.009 amu has a relative abundance of
80.09%. Calculate the atomic mass of this
element.
(10.012 µ) (.1991) + (11.009 µ) (.8009) = 10.81 µ
Boron
Section 5.4
Organizing the Elements

OBJECTIVES:
 Explain how elements are organized in a
periodic table.
Compare
Identify
early and modern periodic tables.
three broad classes of elements.
Distinguish
different areas of the
periodic table.
Section 5.4
Organizing the Elements

A few elements, such as gold and copper, have
been known for thousands of years.

Yet, only about 13 (of 90) had been identified by
the year 1700.

As more were discovered, chemists realized
they needed a way to organize the elements.


Section 5.4
Organizing the Elements
Chemists used the properties of elements
to sort them into groups.
In 1829 J. W. Dobereiner arranged
elements into triads – groups of three
elements with similar properties
(40 + 137) ÷ 2 = 88
Ca 40
Sr 88
Ba 137
Li 7
Na 23
K 39
Cl 35
Br 80
I 127
John Newlands (1837-1898)


Law of Octaves: noted that after interval of
eight elements, similar physical/chemical
properties reappeared.
Newlands was the first to formulate the concept
of periodicity (repeating patterns) in the
properties of the chemical elements.
Lothar Meyer
Mendeleev’s Periodic Table
By
the mid-1800s, about 70
elements were known to exist
Mendeleev – A Russian
chemist arranged the elements in
order of increasing atomic mass
 It was the beginning of the modern
“Periodic Table”
 Dmitri
His Completed
His Original
Table…
Work…
Co and Ni;
Ar and K;
But,
there
were
problems:
some
discovered,
the elements
When
Mendeleev
left blanks
for elements
he
elementsthat
did not
fit his
with
their
generally
matched
predictions
predicted
existed
such
as groups
Germanium Te and I
A better arrangement…



In 1913, Henry Moseley – British physicist, arranged
elements according to increasing atomic number
His basic arrangement is still used today
The symbol, atomic number & mass are the basic
items included in the periodic table
Squares in the Periodic Table

The periodic table displays the symbols and
names of the elements, along with atomic
number and average atomic mass
Black symbol = solid
Red symbol = gas @
Blue symbol = liquid
All @ 25º C
There are other possible
arrangements. How about a:
Spiral Periodic Table
The Periodic Table:
Your “best friend” is an arrangement of
elements in which they are separated into
groups based on a set of repeating properties.
The periodic
table allows you
to easily
compare the
properties of
one element to
another
The Periodic Law

Vertical column = group (or family)
Horizontal rows Similar
= periods
physical & chemical
When elements are arranged
in order of
Repeated properties
in each row
properties
increasing
atomic
number,
there is
a periodic
Identified
by number
& letter
There are 7 periods
repetition of their physical and chemical
properties.
Areas of the periodic table
Three classes of elements are:
1) Metals 2) Nonmetals 3) Metalloids
Metals
1) Metals: electrical conductors, have luster,
are ductile and malleable
(in pink below)
Nonmetals: generally brittle and dull, are poor
conductors of heat and electricity (in blue below)
Some nonmetals…

are gases (O, N, Cl);

are brittle solids (S, C);

one is dark red liquid (Br)
Notice the heavy, stair-step line…
3) Metalloids: border the line

Their properties are intermediate between
metals and nonmetals (in green below)
Elements in the 1A-7A groups are
called the Representative Elements
1A

2A
8A
Have very predictable properties and
patterns of behavior
3A 4A 5A 6A 7A
Groups of elements
Group IA – Alkali Metals


Are the most reactive metals
Form a “base” when reacting with water
Alkali
Metals
Groups of Elements
Group 2A – Alkaline Earth Metals

Are the second most reactive metals

Also form bases with water, but do not dissolve well; hence “earth metals”
Alkaline
Earth
Metals
Groups of elements
Group 7A – Halogens
The most reactive non-metal group
 Means “salt-forming”

The Halogen Group
Groups of elements
Group 8A- Noble Gases
1)


Previously called “inert gases” because they rarely take part in a
reaction
Noble gases have a completely full electron arrangement
Noble
Gases
The “B” groups are called
the Transition Elements
La
Ac
Lanthanide Series
Actinide Series
The “Inner Transition Metals”
actually belong here
Crash Course Chemistry - The Periodic Table