Atoms and Elements
Chapter 4
Tro, 2nd ed.
ELEMENTS AND ATOMS
All known substances on Earth and probably the
universe are formed by combinations of more than
100 elements.
An element is a fundamental or elementary substance
that cannot be broken down into simpler
substances by chemical means.
Each element has a number.
Beginning with hydrogen as 1, the
elements are numbered in order of
increasing complexity.
ELEMENTS AND ATOMS
Most substances can be decomposed into two or more
simpler substances.
Water can be decomposed into hydrogen and oxygen.
Table salt can be decomposed into sodium and chlorine.
An element cannot be decomposed into a simpler
substance.
An atom is the smallest particle of an element that can
exist.
An atom is the smallest unit of an element that can
enter into a chemical reaction.
An atom is very small
The
of an
atom isof
Thisdiameter
is 1 to 5 ten
billionths
0.1
to 0.5 nm.
a meter.
If the diameter of this dot is 1
Even smaller particles than atoms
mm, then 10 million hydrogen
exist. These are called subatomic
atoms would form a line across
particles.
the dot.
Dalton’s Atomic Theory
1.
Elements are composed of minute
indivisible particles called atoms.
2. Atoms of the same element are alike in
mass and size. Atoms of different elements
have different masses and sizes. (Is this
still correct?)
3. Chemical
compounds
are formed
Modern
research
has demonstrated
thatby the
Atoms
under
special
circumstances
canelements
union
of
two
or
atoms
of
different
atoms are composed of subatomic
be decomposed.
in whole number ratios.
particles.
WHAT’S IN AN ATOM?
DISCOVERING SUBATOMIC
PARTICLES!
In 1897 Sir Joseph Thompson
demonstrated that cathode rays:
- travel in straight lines
- are negative in charge
- are deflected by electric and magnetic fields
- produce sharp shadows
- are capable of moving a small paddle wheel
Subatomic Particles:
Electron
This was the discovery of the
fundamental unit of charge
– the electron.
Subatomic Particles: Protons
Eugen Goldstein, a German physicist, first
observed protons in 1886:
Thompson determined the proton’s
characteristics.
Thompson showed that atoms contained both
positive and negative charges.
This disproved the Dalton model of the atom
which held that atoms were indivisible.
But the mass of the atom could not be accounted
for by the mass of protons inside it. There had
to be something else.
Subatomic Particles: Neutron
James Chadwick discovered the neutron in
1932.
Its actual mass is slightly greater
than the mass of a proton.
4.1
0.00055
1.0073
1.0087
Memorize name, symbol, relative charge, relative mass, and
LOCATION (in nucleus or in orbitals outside of nucleus) of these
three subatomic particles. See table 4.1 on page 90/96 in your
book.
The Nuclear Atom
Radioactivity was discovered by
Becquerel in 1896.
Radioactive elements spontaneously emit
alpha (a) particles, beta (b) particles and
gamma (g) rays from their nuclei.
By 1907, Rutherford found that a particles
emitted by certain radioactive elements
were helium nuclei, consisting of 2
protons and 2 neutrons.
The Rutherford Experiment
In 1911, Rutherford performed experiments that
shot a stream of a particles at a thin piece of
gold foil.
Most of the a particles passed through the foil with little or no
deflection.
He found that a few were deflected at large angles and some
a particles even bounced back.
Rutherford’s alpha particle scattering experiment.
5.5
The Rutherford Experiment
An electron with a mass of 0.00055 amu
could not have deflected an a particle
with a mass of ~4 amu.
Rutherford knew that like charges repel.
Rutherford concluded that each gold atom
contained a positively charged mass that
occupied a tiny volume. He called this mass
the nucleus.
The Rutherford Experiment
If a positive a particle approached close
enough to the positive mass it was deflected.
Most of the a particles passed through the gold
foil. This led Rutherford to conclude that a gold
atom was mostly empty space.
Because a particles have relatively high masses,
the extent of the deflections led Rutherford to
conclude that the nucleus was very heavy and
dense.
Deflection
Scattering
Deflection and scattering of a particles by positive gold nuclei.
5.5
General Arrangement of Subatomic
Particles
Rutherford’s experiment showed that an
atom had a dense, positively charged
nucleus.
Chadwick’s work in 1932 demonstrated the atom
contains neutrons.
Rutherford also noted that light, negatively
charged electrons were present in an atom
and offset the positive nuclear charge.
General Arrangement of Subatomic
Particles
Rutherford put forward a model of the atom in which
a dense, positively charged nucleus is located at
the atom’s center.
The negative electrons surround the nucleus.
The nucleus contains protons and neutrons
5.6
ATOMS AND THEIR SUBATOMIC
PARTICLES
The atomic number of an element is
equal to the number of protons in
the nucleus of that element.
The atomic number of an atom determines which
element the atom is.
Every atom with an atomic number of 1 is a
hydrogen atom.
Every carbon atom contains 6 protons in its
nucleus.
atomic
number
Every atom with an
atomic number of 1
is a hydrogen atom.
1 proton in the
nucleus
atomic
number
Every atom with an
atomic number of 6
is a carbon atom.
6 protons in the
nucleus
Ions: atoms that have lost or
gained electrons
Positive ions were explained by assuming
that a neutral atom loses electrons.
Negative ions were explained by
assuming that extra electrons can be
added to atoms.
When one or more electrons are lost
from an atom, a cation is formed.
5.4
When one or more electrons are added
to a neutral atom, an anion is formed.
5.4
Atomic Structures of Ions
Ion
+
-
p
e
-1
17
18
+1
19
18
-2
16
18
+2
38
36
Cl
K
S
Sr
Isotopes of the Elements
Atoms of the same element can have
different masses.
They always have the same number of protons, but
they can have different numbers of neutrons in
their nuclei.
The difference in the number of neutrons accounts
for the difference in mass.
These are isotopes of the same element.
Isotopic Notation
Isotopic Notation for Carbon-12
6 protons + 6 neutrons
12
C
6
6 protons
Isotopic Notation for Carbon-14
6 protons + 8 neutrons
14
C
6
6 protons
Hydrogen has three isotopes
1 proton
1 proton
1 proton
0 neutrons
1 neutron
2 neutrons
Examples of Isotopes
Element
Hydrogen
Hydrogen-2
Hydrogen-3
Uranium-235
Uranium-238
Chlorine-35
Chlorine-37
Protons
1
1
1
92
92
17
17
Electrons
1
1
1
92
92
17
17
Neutrons
0
1
2
143
146
18
20
Symbol
1
1
2
1
3
1
H
H
H
235
92
238
92
U
U
35
17
Cl
37
17
Cl
Atomic Mass
The mass of a single atom is too small to
measure on a balance.
Using a mass spectrometer, the mass of
the hydrogen atom was determined.
The mass of one hydrogen atom was
determined to be 1.673 x 10-24 g.
A Modern Mass Spectrometer
Positive ions
formed from
sample.
Electrical field
From the intensity and positions
at slits
A mass
of the lines Deflection
on the mass
of
accelerates
spectrogram
spectrogram,
the different
ions
positive
ions. positive
is recorded.
isotopes and
their at
relative
occurs
amounts can
be determined.
magnetic
field.
5.8
A typical reading from a mass spectrometer. The two
principal isotopes of copper are shown with the
abundance (%) given.
5.9
ATOMIC MASS UNITS AND
RELATIVE ATOMIC MASS
Single atomic masses are too small to weigh on a balance. To overcome
this problem a system of relative atomic masses using “atomic mass units”
was devised to express the masses of elements using simple numbers.
The standard to which the masses of all other atoms are compared to was
chosen to be the most abundant isotope of carbon, carbon-12.
A mass of exactly 12 atomic mass units (amu) was assigned to the carbon12 atom. An amu is defined as exactly equal to 1/12th mass of a carbon-12
atom. 1 amu = 1.6606 x 10-24 g
Isotopes of the same element have different masses.
The listed atomic mass of an element is the average relative mass of the
isotopes of that element compared to the mass of carbon-12.
ATOMIC MASS UNITS AND
RELATIVE ATOMIC MASS
Calculating relative atomic mass:
Convert percent abundance to fraction
by dividing by 100
Multiply fraction abundance times each
isotopes mass
Sum up the results
Rel Atomic Mass = f1*m1 + f2*m2 + …
See example on next slide
To calculate the atomic mass multiply the
atomic mass of each isotope by its
fractional abundance and add the results.
Isotope
Isotopic
mass (amu)
Abundance
(%)
63
29
Cu
62.9998
69.09
65
29
Cu
64.9278
30.91
Average
atomic mass
(amu)
63.55
(62.9998 amu) 0.6909 = 43.48 amu
(64.9278 amu) 0.3091 = 20.07 amu
63.55 amu
Relationship between Mass
Number and Atomic Mass
Atomic mass found on the Periodic Table is RELATIVE
atomic mass, the weighted average of all the isotopes.
Each isotope of an element has a unique WHOLE number
called the mass number, the sum of protons and
neutrons. Mass number is NOT found on most Periodic
Tables.
******************************************************************
ATOMIC MASS IS NOT THE SAME AS MASS NUMBER
ATOMIC MASS IS NOT THE SAME AS MASS NUMBER
ATOMIC MASS IS NOT THE SAME AS MASS NUMBER
******************************************************************
REMEMBER: ATOMIC MASS IS AN AVERAGE OF THE
DIFFERENT ISOTOPES’ MASSES. It will have decimal
points!
Relationship Between Mass
Number and Atomic Number
The mass number minus the atomic number equals
the number of neutrons in the nucleus.
mass
number
atomic
number
109
47
Ag
atomic
mass number number
109
47
=
=
number of
neutrons
62
Atomic Mass, Mass Number &
Abundance of Isotopes
EXAMPLE: Chlorine has two isotopes, chlorine-35
and chlorine-37. Write down the atomic number,
mass number and number of neutrons for each.
Then look up the atomic mass. Which isotope is in
greater abundance?
Cl-35: Z = 17, A = 35, 35 - 17 = 18 neutrons
Cl-37: Z = 17, A = 37, 37 - 17 = 20 neutrons
Atomic mass is 35.453 amu.
Closer to 35, therefore chlorine-35 is more abundant.
Elements are not distributed equally
by nature.
Oxygen is the most abundant element in
the human body (65%).
Oxygen is the most abundant element in
the crust of the earth (49.2%).
In the universe, the most abundant
element is hydrogen (91%) and the
second most abundant element is
helium (8.75%).
Distribution of the
common elements
in nature.
3.2
Sources of Element Names
GreekColor
Iodine: from the Greek iodes meaning
violet.
LatinFluorine: from the Latin fluere meaning to
Property flow. The fluorine containing ore
fluorospar is low melting.
German- Bismuth: from the German
Color
weisse mass which means white mass.
Location Germanium: discovered in 1866 by a
German chemist.
Famous- Einsteinium: named for Albert Einstein.
Scientists
These symbols have carried over from the earlier names of the
A number
Mostofsymbols
symbols
appear
withtothe
have
same
no letter
connection
as the with
element.
the element.
elements
(usuallystart
Latin).
Element Names and
Symbols
Element names and symbols are in
Periodic Table inside front cover.
Symbols: Learn first 36 on Periodic
Table, plus Ag, Au, Pt, Hg, Sn, Pb, I,
U, Ba, and Rn.
Also memorize that some elements
exist as diatomic molecules : H2, O2,
N2, F2, Cl2, Br2, and I2
THE PERIODIC TABLE OF THE
ELEMENTS
The periodic table was designed by
Dimitri Mendelev in 1869.
In the table each element’s symbol is
placed inside of a box.
Above the symbol of the element is its
atomic number.
7
N
The elements are arranged in order of
increasing atomic number.
Elements with similar chemical
properties are organized in columns
called families or groups .
He
Ne
Ar
Kr
Xe
Rn
These elements are known
as the noble gases. They are
nonreactive.
THE PERIODIC TABLE OF
THE ELEMENTS
18 columns - called groups
7 rows - called periods
Representative elements are in groups
1,2, 3A-8A (13-18)
Transition metals are in groups 3-12
Groups of elements have similar
chemical properties
THE PERIODIC TABLE OF
THE ELEMENTS
SPECIAL GROUP NAMES:
1A (1) = alkali metals
2A (2) = alkaline earth metals
8A (18) = noble gases (six)
7A (17) = halogens (four)
NOTE:
H really belongs to its own group,
which is why it’s shown by itself in my
periodic table! It has 1 electron in outer
shell, like the 1A elements, but it’s not a
metal, and reacts more like group 7A.
The Periodic Table of Elements
Nonmetals: upper right
Metals: lower left
Metalloids touch stairs, except for Al.
• Metals are solid at room temperature.
– Mercury is an exception. At room temperature it
is a liquid.
• Metals are good conductors of heat and electricity.
Most elements
are metals
physical
properties
of metals
• Metals are malleable (they can be rolled
or
hammered into sheets).
• Metals have high luster (they are shiny).
• Metals are ductile (they can be drawn into wires).
• Most metals have a high melting point.
Most elements
are metals
• Metals have high densities
Chemical Properties of Metals
Metals have little tendency to combine with
each other to form compounds.
Many metals readily combine with nonmetals to form
ionic compounds.
They can combine with oxygen.
sulfur.
chlorine.
In nature, minerals are formed by combinations
of the more reactive metals with other
elements.
Chemical Properties of Metals
A few of the less reactive metals such as copper,
silver and gold are found in the free state.
Metals can mix with each other to form alloys.
Brass is a mixture of copper and zinc.
Bronze is a mixture of copper and tin.
Steel is a mixture of carbon and iron.
Physical Properties of Nonmetals
Everything metals are, nonmetals are
NOT!
Lack luster (they are dull)
Have relatively low melting points
Have low densities.
Poor conductors of heat and electricity
At room temperature, carbon, phosphorous,
sulfur, selenium, and iodine are solids.
Physical State at Room Temperature
phosphorous
carbon
Solid
selenium
sulfur
NONMETALS
iodine
Physical State at Room Temperature
liquid
NONMETALS
bromine
Physical State at Room Temperature
nitrogen,
oxygen
gas
fluorine,
chlorine
NONMETALS
helium, neon, argon, krypton, xenon, radon
The Metalloids
Metalloids have
properties that are
intermediate between
metals and nonmetals
1.
boron
2.
silicon
3.
germanium
4.
arsenic
5.
antimony
6.
tellurium
polonium
7.
= Alkali Metals
= Halogens
= Alkali Earth Metals
= Lanthanides
= Noble Gases
= Actinides
= Transition Metals
Important Groups - Hydrogen
Nonmetallic
Colorless, diatomic gas – exists as H2
Very low melting point & density
Reacts with nonmetals to form molecular
compounds
HCl is acidic gas
H2O is a liquid
Reacts with metals to form hydrides
metal hydrides react with water to form H2
HX dissolves in water to form acids
Important Groups – IA, Alkali Metals
Hydrogen usually placed here,
though it doesn’t belong
Soft, low melting points,low
density
Flame tests  Li = red, Na =
yellow, K = violet
Very reactive, never find
uncombined in nature
Tend to form water soluble
compounds with colorless
solutions
React with water to form basic
(alkaline) solutions and H2
2 Na + 2 H2O  2 NaOH + H2
releases a lot of heat
65
lithium
sodium
potassium
rubidium
cesium
Important Groups – IIA, Alkali Earth
Metals
Harder, higher melting, and denser than
alkali metals
Flame tests  Ca = red, Sr = red, Ba =
yellow-green
Reactive, but less than corresponding
alkali metal
Form stable, insoluble oxides from
which they are normally extracted
Oxides are basic = alkaline earth
Reactivity with water to form H2,  Be =
none; Mg = steam; Ca, Sr, Ba = cold
water
beryllium
magnesium
calcium
strontium
barium
Important Groups – VIIA, Halogens
Nonmetals
F2 & Cl2 gases; Br2 liquid; I2 solid
All diatomic
Very reactive
Cl2, Br2 react slowly with water
Br2 + H2O  HBr + HOBr
React with metals to form ionic
compounds
HX all acids
HF weak < HCl < HBr < HI
fluorine
chlorine
bromine
iodine
Important Groups – VIIIA, Noble Gases
All gases at room
temperature
Very low melting and
boiling points
Very unreactive,
practically inert
Very hard to remove
electron from or give an
electron to
Review
What is the atomic number of boron, B?
What is the atomic mass of silicon, Si?
How many protons does a chlorine atom have?
How many electrons does a neutral neon atom have?
Will an atom with 6 protons, 6 neutrons and 6 electrons
be electrically neutral?
Will an atom with 27 protons, 32 neutrons and 27
electrons be electrically neutral?
Will a Na atom with 10 electrons be electrically neutral?
What is the mass number of fluorine?