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The Atom
Read p. 47-50
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1807 John Dalton
Atomic Theory
• All matter is made of atoms
• Atoms are indestructible and
indivisible
• Atoms combine in small whole
numbers to form compounds
Now
• Atoms are divisible
• Valence electron shell changes in
chemical interactions
Atomic Structure
• So atoms can’t be broken into
smaller pieces, BUT they are made
of smaller pieces!
• Those smaller pieces all on their
own don’t act like the element the
atom is from.
• Atom: the smallest piece of an
element that keeps the chemical
properties of that element.
Subatomic Particles
Particle
Proton
Relative
mass
1
Neutron
1
Electron
1/1840
Relative
charge
+
positive
0
neutral
negative
Location
nucleus
nucleus
shells
around
nucleus
Two areas of an Atom
• Nucleus: made up of:
–Protons (+ positive), mass number = 1
–Neutrons (0 neutral), mass number = 1
–Makes up most of the mass of the atom
+
0
0
+
• Electrons (- negative)
–Around the nucleus
–Tiny, mass number = 1/1840
–Lots of space between the nucleus and
the electrons.
-
Perspective
• If the nucleus were 1 meter across,
the electrons would be ten
kilometers away.
• Most of an atom is empty space!
Atomic Number
• Atomic number (Z): number of
protons
8
• Defines an element
O
16.00
Trends in Atomic Numbers
• Trend: Atomic numbers increase as
you move to the right and as you
move down the Periodic Table of
the Elements.
Atoms are Neutral
• Electrons were discovered first
• We knew there must be something
to balance the charge and give the
atom mass.
• The number of electrons = number
of protons
Holding it all together
• Nuclear forces hold protons and
neutrons together
• The positive force of the nucleus
holds the electrons near the nucleus.
–Positive and negative attract.
Protons
Neutrons
Mass
number
Electrons
Element
-
+
0
0
+
-
-
+
+
0
-
0
0
+
-
-
0
+
0
+
-
-
+
0
0
+
+
0
-
-
-
+
0
-
0
0
-
+
+
0
+
0
0
-
+
+
+
0
-
-
+
0
+
-
+
0
0
0
+
+
0
0
+
0
+
0
0
0
-
+
+
-
+
-
-
-
-
-
+
0
0
+
-
+
0
0
0
+
+
0
0
+
0
+
0
0
0
-
+
+
+
-
+
-
-
-
-
-
0
+
-
+
0
0
0
+
+
0
+
0
0
-
+
+
+
0
-
-
Protons
Neutrons
Mass
number
Electrons Element
2
2
4
2
Helium
3
3
6
3
Lithium
5
5
10
5
Boron
7
7
14
7
Nitrogen
10
10
20
10
Neon
11
11
22
11
Sodium
8
8
16
8
Oxygen
Which element…?
• Has 20 protons?
Which element…?
• Has 15 protons?
Which element…?
• Has 17 protons?
Which element…?
• Has 53 protons?
Which element…?
• Has 79 protons?
Which element…?
• Has 40 electrons?
Which element…?
• Has 77 electrons?
Which element…?
• Has 50 electrons?
Which element…?
• Has 118 electrons?
Mass Number
• Mass number (A): protons +
neutrons
• For light elements, protons = neutrons
• Heavier atoms require higher
proportions of neutrons to overcome
the repulsion between larger numbers
of protons.
• Go back to pictures, ID mass
number
Neutral and Charged
• Neutral atom: electrons =
protons
• Ion: an atom that has gained or
lost electrons
–Gained = - negative = anion
–Lost = + positive = cation
Examples
• H+
–Hydrogen that has lost one electron
–Cation
• Br–Bromine that has gained one electron
–Anion
• Fe3+
–Iron that has lost three electrons
–Cation
-
+
0
0
+
-
-
+
0
0
+
-
-
+
+
0
-
0
0
+
-
-
-
+
+
0
-
0
0
+
-
-
-
+
0
+
-
-
+
0
0
0
+
+
0
-
-
-
-
+
0
+
-
-
+
0
0
0
+
+
0
-
-
-
-
-
-
+
0
-
0
-
+
0
0
+
-
+
0
0
-
+
+
+
0
-
-
• Potassium (K) that has lost one
electron
–K+
–cation
• Chlorine (Cl) that has gained one
electron
–Cl–Anion
• Vanadium (V) that has lost three
electrons
–V3+
–Cation
• Oxygen that has gained two
electrons
–O2–anion
•
•
•
•
•
Li+
FCu2+
Al3+
S2-
•
•
•
•
•
Lost one, cation
Gained one, anion
Lost two, cation
Lost three, cation
Gained two, anion
Isotopes
Average atomic mass
-
-
6
+
C
0
+
0
+
12.01
-
+
+
0
0
-
0
+
0
-
-
• But why is the atomic mass
different from the mass number?
–Carbon atomic mass = 12.011 amu
–Mass number = 12 amu
Isotopes
• Atoms of the same element can
have different numbers of neutrons
• But the number of protons does not
change, so it stays the same
element, with the same properties
• Isotopes of the same element have
different masses
• Isotopes are atoms of the same
element with different masses (and
different numbers of neutrons)
35
• Identified by mass number
– different mass numbers
– slightly different physical properties
17
Cl
37
17
Cl
Naming Isotopes
• Mass number (A): the total number
of protons and neutrons in the
nucleus of an isotope
• Hyphen Notation:
–Element-Mass number
–Hydrogen-3
• Nuclear Symbol
mass _ number
atomic_ number
A
Z
X
Symbol
Nuclear symbols
• Helium-4: 4 is the mass number
–The total number of protons plus
neutrons
• This helium isotope can also be
written: 4
He
2
–4 is the mass number, 2 is the number
of protons (from the periodic table)
Isotope Notation
Mass
number
Atomic
number
• Al-27
• Aluminum-27
27
13
Al
-
+
0
0
+
-
-
0
0
+
-
-
+
0
0
+
-
-
+
0
+
-
-
0
+
0
+
-
-
+
0
0
+
+
0
-
-
0
+
0
+
-
-
+
0
0
+
+
-
-
-
+
0
-
0
0
-
+
+
0
+
0
0
-
+
+
+
0
-
-
-
+
0
-
0
0
-
+
+
0
+
0
0
-
+
+
+
-
• Nuclide: any isotope of any element
• Write the nuclear symbol for
Carbon-13 13
C
6
• Write the nuclear symbol for the
element with 11 protons and 12
neutrons 23
Na
11
• Both are chlorine, both have the
same number of protons (17) and
electrons (17)
• They differ in their number of 35Cl
17
neutrons, and thus, their mass
37
number
17 Cl
• Isotopes have the same chemical
properties (because they have the
same number of electrons)
• They have slightly different physical
properties because they have
different masses.
• Write the hyphen notation and the
nuclear notation
-
+
0
+
-
-
0
+
0
+
-
-
+
0
0
+
+
0
-
-
0
+
0
+
-
-
+
0
0
+
+
-
-
-
+
0
-
0
0
-
+
+
0
+
0
0
-
+
+
+
0
-
-
-
+
0
-
0
0
-
+
+
0
+
0
0
-
+
+
+
-
Practice naming Isotopes
• The notation Sn-117 represents an
isotope of the element
____________.
• Its mass number is ________.
• Since its atomic number is
__________, the atom contains
–_________ protons,
–_________ electrons and
–_________ neutrons.
• Cesium-133 is an isotope whose
mass number is _________.
• Its atomic number is ________.
• The atom contains _______
protons,
• ________ electrons and
• ________ neutrons.
• An isotope contains 27 protons and
39 neutrons. The isotope is an
atom of the element
____________.
• Its atomic number is _________.
• This atom has _______ electrons.
• Its mass number is __________.
Give the element, and number
of subatomic particles
Atomic
number
47
83
36
76
Mass
number
107
210
84
190
Element
Protons
Electrons Neutrons
Give the element, protons,
electrons and neutrons
Atomic
number
26
88
7
78
Mass
number
55
226
14
195
Element
Protons
Electrons Neutrons
Relative Atomic Mass
Relative atomic mass
• Natural abundance: proportions
of different isotopes of an element
• Used to determine relative atomic
mass
Mixtures of Isotopes
• In nature, elements occur as a
mixture of isotopes.
• Average atomic mass = weighted
average of isotope masses
• On the periodic table, the average
atomic mass is below the element
name.
Example
• For chlorine, 75% of atoms have a mass
of 35, and 25% have a mass of 37.
(75  35)  (25  37)
 35.5
100
• Relative atomic mass of chlorine is 35.5
• So, how do we know the natural
abundance of the isotopes of each
element?
The Mass Spectrometer
• Mass spectrometer: an instrument
that separates particles by mass,
and records the relative proportions
of each size particle.
• Use: to determine the natural
abundances of the isotopes of a
particular element and the atomic
mass of an element.
Mixtures of Isotopes
• In nature, elements occur as a
mixture of isotopes.
• Average atomic mass = weighted
average of isotope masses
• On the periodic table, the average
atomic mass is below the element
name.
• Natural abundances: the
proportion of each isotope of an
element found in nature.
• Atomic mass is calculated using
the natural abundance of each
isotope.
• Mass spectrometer: an
instrument that separates particles
by mass, and records the relative
proportions of each size particle.
Thin
tube
A
D
C
F
B
E
• A: vaporizes (gas) substance
• B: particles converted from neutral
atoms to positive ions by
bombarding (hitting repeatedly)
with electrons (ionize)
• C: particles accelerated by an
electrical field
• D: fast-moving ions enter a
magnetic field and are deflected
–smaller mass: deflected more
–greater mass: deflected less
• E: detector records ions of a
particular mass
• F: vacuum prevents collisions with
gas molecules.
• Strength of magnetic field (D) is varied
to bring particles of different masses
to the detector.
• The abundance of each ion in the
sample can be determined because of
the different masses of each kind of
particle.
• Mass spectrum: the relative
abundance of each mass particle
Calculating Atomic Mass
(percent A x mass number A) + (percent B x mass number B)
100
Ex: Natural chlorine contains 75% 35Cl and 25%
37Cl.
(75 x 35) + (25 x 37)
100
atomic mass = 35.5
relative number of
atoms
Mass spectrum of copper
100
80
69.09
60
30.91
40
20
0
63
65
mass number
relative number of
atoms
Mass spectrum of neon
100
91
80
60
40
20
0.3
0
20
21
mass number
9
22
relative proportion
The mass spectrum of magnesium
120
100
80
60
40
20
0
100
24
12.8
14.4
25
m/z
26
Most abundant assigned value of 100, others
given a value in proportion to this.
100
Mg  100 
 78.6%
127.2
12.8
25
Mg  100 
 10.0%
127.2
14.4
26
Mg  100 
 11.3%
127.2
24
relative atomic mass
= (24 x 0.786) + (25 x 0.100) + (26 x 0.113) = 24.3
Uses of Isotopes
Radioactive Isotopes
• Natural element is exposed to
neutrons in a nuclear reactor
• Atoms gain neutrons
• “radioisotopes”
• Give off gamma rays (high energy
electromagnetic rays or photons)
Naturally occurring
radioactive isotopes
• Carbon-14
• Over time, carbon-14 decays (loses
neutrons) and turns into carbon-12.
• The rate of decay can be used to
calculate the age of very old things.
–Carbon dating
–Ancient trees, dinosaur bones
“Tracers”
• In medicine
• To measure the rate of activity in an
organ, ex: thyroid
• Patient drinks radioactive iodine,
which is processed by the thyroid
• Rate of decrease of radioactive
isotope is measured
Radiation
• Intense radioactivity, produce
gamma rays
• Ex: cobalt-60
• Used in chemotherapy to kill cancer
cells