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AP Chemistry
Chapter 3
The Structure of the Atom
1
Beginning of the Atomic
Model
2
Democritus was the early (around
400BC) Greek philosopher who is
credited with the concept of the atom
(atomos) –which means invisible
3
Dalton (around 1800AD) is an English
school teacher who proposed the law of
conservation of mass, the law of
definite proportions, and the law of
multiple proportions.
His many experiments
with gases proved
these laws are true, if
atoms exist.
Dalton is also known as the Father of
the (Modern) Atomic Theory
4
Dalton’s atomic theory:
1.All matter is composed of very small
particles called atoms
2.Atoms of a given element are
identical in size, mass, and other
properties; atoms of different
elements differ in these properties.
5
3.Atoms cannot be subdivided, created,
or destroyed
4.Atoms of different elements combine
in simple whole-number ratios to form
chemical compounds.
5.In chemical reactions, atoms are
combined, separated, or rearranged.
6
Two aspects of Dalton’s atomic theory
proven to be incorrect:
a.We now know atoms are divisible.
b. Atoms of the same element can have
different masses.
7
J. J. Thomson is the man credited with
the discovery of the electrons in the
late 1800’s, using cathode ray tubes
8
Millikan calculated the mass of the
electron (very, very small)
9
Knowledge of electrons led to two
inferences about atomic structure:
1.Because atoms are electrically
neutral, they must contain positive
charge to balance the negative
electrons.
2. Because electrons have so little
mass, atoms must contain other
particles to account for most of their
mass
10
Nucleus of the atom—discovered by
Lord Ernest Rutherford
Gold foil experiment—actually done
by Hans Geiger and Ernest Marsden
11
Observations:
a.Majority of the alpha (α) particles
penetrated foil undeflected.
b.About 1 in 20,000 were slightly
deflected
c.About 1 in 20,000 were deflected
back to emitter
12
13
Conclusions:
1. Mass of the atom and the positive
charge are concentrated in small
regions called nucleus
2.Most of the atom is empty
3.Magnitude of charge on the nucleus
is different for different atoms
14
4. Number of electrons outside the
nucleus = number of units of nuclear
charge (to account for the fact that the
atom is electrically neutral)
Atoms are electrically neutral because
they contain equal numbers of
protons and electrons
15
A couple years later Rutherford
presented evidence for a neutral
particle which was also in the nucleus
and contained a similar mass to that
of a proton – called a neutron
16
Mass of one proton = mass of neutron
= mass of 1837 electrons
Thus the total mass of an atom is
basically the sum of the protons and
neutrons, called the atomic mass or
mass number, abbreviated A
17
Atomic number—the number of protons
in the nucleus of the atom.
--number of protons identifies the
element and is equal to the number of
electrons (of a neutral atom)
--symbol is Z
18
Isotopes are atoms of the same
element that have different masses
because they have different numbers
of neutrons but they still have similar
chemical properties
19
Isotopes of Carbon
Mass Number of
Isotope
Number of Protons
Number of
Neutrons
8
6
2
9
6
3
10
6
4
11
6
5
12
6
6
13
6
7
14
6
8
15
6
9
16
6
10
17
6
11
18
6
12
19
6
13
20
6
14
20
Isotopes of Carbon
Mass Number of Carbon
Isotopes
Name of
Isotopes
8
carbon-8
9
carbon-9
10
carbon-10
11
carbon-11
12
carbon-12
13
carbon-13
14
carbon-14
15
carbon-15
16
carbon-16
17
carbon-17
18
carbon-18
19
carbon-19
20
carbon-20
21
Nuclide is the general term for any
isotope of any element
22
Nuclide is the general term for any
isotope of any element
Atomic Mass Unit (amu) is exactly 1/12
the mass of a carbon-12 atom
23
Average atomic mass is the weight
average of the atomic masses of the
naturally occurring isotopes of an
element.
Ave. Atomic mass = %abundace(mass
of isotope 1) + %abundance(mass of
isotope 2) +…..
24
Example 1:
Element Sciencium has two isotopes.
Sciencium-301 has an abundance of
59.5%, and Sciencium-304 is the
other. What is the average atomic
mass?
301 amu x .595 = 179 amu
304 amu x .405 = 123 amu
302 amu
Example 2:
Element Pepsium has an average
atomic mass of 335. Two isotopes of
Pepsium exist. If Pepsium-327 is
30.5% abundant, then what is the
second isotope?
327 amu x 0.305 = 99.7 amu
? amu x 0.695 =
? amu
335 amu
1 – 0.305 = 0.695
Example 2:
Element Pepsium has an average
atomic mass of 335. Two isotopes of
Pepsium exist. If Pepsium-327 is
30.5% abundant, then what is the
second isotope?
327 amu x 0.305 = 99.7 amu
? amu x 0.695 =
? amu
335 amu
Step 1: Find the missing weighted mass
Example 2:
Element Pepsium has an average
atomic mass of 335. Two isotopes of
Pepsium exist. If Pepsium-327 is
30.5% abundant, then what is the
second isotope?
327 amu x 0.305 = 99.7 amu
? amu x 0.695 = 235 amu
335 amu
335 – 99.7 = 235.3
Example 2:
Element Pepsium has an average
atomic mass of 335. Two isotopes of
Pepsium exist. If Pepsium-327 is
30.5% abundant, then what is the
second isotope?
327 amu x 0.305 = 99.7 amu
? amu x 0.695 = 235 amu
335 amu
Step 2: Find the missing mass
Example 2:
Element Pepsium has an average
atomic mass of 335. Two isotopes of
Pepsium exist. If Pepsium-327 is
30.5% abundant, then what is the
second isotope?
327 amu x 0.305 = 99.7 amu
338 amu x 0.695 = 235 amu
335 amu
235 ÷ 0.695 = 338
Mass Spectrometry
How we know isotopes.
A mass spectrometer has three parts:
1. ionizer
2. magnetic field
3. detector
And might look like this
If a sample of a pure element is
placed in the spectrometer, then all
the ions formed will have the same
charge. For example Cl → Cl-
Cl
Cl
Ionizer
The ions then pass through a
magnetic field that will change their
paths. Which will change direction
more, something heavy or something
light?
-
Cl
Cl
Cl
+
Magnetic Field
Detector
The computer attached to the
detector gives a readout like this:
The locations tell the masses – one
group of Cl had a mass of 35 amu’s
and the other had a mass of 37 amu’s.
The size of the bars indicate the
relative amounts of each isotope.
As the bar for 35Cl is 3 times bigger
than 37Cl (measuring with a ruler) thus
35Cl is about 75% abundant and 37Cl is
about 25% abundant.
Unstable Nuclei and
Radioactive Decay
39
1) When referring to nuclear
reactions people commonly
think of nuclear fission (the
splitting of large atoms into
smaller pieces)
40
1) and nuclear fusion (the
combining of small atoms into
one large one), but on earth
these reactions do not occur
naturally.
41
2) Naturally occurring nuclear
reactions result from the unusual
number of neutrons of an isotope
which makes it unstable
(unusually high in energy). This
often results in the isotope
changing from one element into
another element in an attempt to
become more stable (lower in
energy).
42
A)These reactions are called
nuclear reactions, as they
involve changes in the
nucleus.
B)During these nuclear
reactions, rays and
particles are given off,
which is called radiation.
43
C) Sometimes an unstable
nucleus will change into
many different elements as
it tries to become more
stable. This is called
radioactive decay.
44
3) When radioactive decay
occurs, there are three
different types of radiation
that can be given off. Each
type has a different mass,
and sometimes a charge.
45
A) The first type of radiation
to be discovered was called
alpha radiation and came
from alpha particles.
226 Ra → 222 Rn
88
86
+
?? ??
??
46
i) Because the mass
numbers must be equal,
226 = 222 + x. So the
mass of the alpha particle
must be 4.
47
ii) Because the atomic
numbers must be equal,
88 = 86 + x. So the atomic
number of the alpha
particle must be 2.
48
iii) The element Helium has a
mass of 4 and an atomic
number of 2, so the alpha
particle is just like a helium
atom without any electrons;
4
4
He or 
2
2
49
B) The second type of
radiation to be discovered
was called beta radiation
and came from beta
particles.
14 C → 14 N + ?? ??
6
7
??
50
i) Because the mass numbers
must be equal, 14 = 14 + x.
So the mass of the beta is
zero.
51
ii) Because the atomic
numbers must be equal,
6 = 7 + x. So the atomic
number must be -1.
52
iii) The electron has a mass
of zero and a charge of -1,
so the beta particle is just
like an electron;
0
β
-1
53
C) The last type of radiation
to be discovered was called
gamma radiation and came
from gamma particles.
238
92
U
→ 234
90
Th
+ 4
2
He
+ ??
??
??
54
i) Because the mass
numbers must be equal,
238 = 234 + 4 + x. So the
mass of the gamma particle
must be zero.
55
ii) Because the atomic
numbers must be equal,
92 = 90 + 2 + x. So the
atomic number must also
be zero.
56
iii) The gamma particle was
the last to be found
because it has no mass
and no charge;
0
0
γ
57
The Development of a
New Atomic Model
58
Previously, Rutherford reshaped our thoughts
of the atom by showing the protons were
located in the nucleus of the atom, but he could
not model for us where the electrons were,
other than outside the nucleus somewhere.
Fortunately, studies into the properties of light
and the effects of light on matter soon gave
clues to where electrons actually are.
59
Light is a small part of all the radiation
(something that spreads from a source) called
electromagnetic radiation. Electromagnetic
radiation is energy in the form of waves (of
electric and magnetic fields). Electromagnetic
radiation includes radio waves, microwaves,
infrared, visible light, X-rays, and Gamma
rays. All these together are considered the
Electromagnetic Spectrum.
60
61
As all the forms of electromagnetic radiation
are waves, they all have similar properties.
•All electromagnetic radiation travels at the
speed of light (c), 299,792,458 m/s (3 x 108) in a
vacuum
62
•The crest is the top of the waves, the trough is
the bottom of the waves, and the amplitude is a
measurement from the rest or zero line to a
crest or trough
63
64
•The wavelength (λ – lambda) is the distance
between successive crests/troughs and is
measured in meters (often nm = 1 x 10-9 m)
•The frequency (ν – nu) is the number of waves
that pass a point in one second and is measured
in
1
(per second – can be written as s-1) or
s
Hz (Hertz)
65
How many hertz is the first wave?
1 wave per second = 1 Hz
How many hertz is the second wave?
2 waves per second = 2 Hz
66
The speed of a wave is directly proportional to
the wavelength and the frequency; c = λν is the
formula
c
λ
ν
67
Example. A certain violet light has a
wavelength of 413 nm. What is the frequency
of the light?
68
Example. A certain violet light has a
wavelength of 413 nm. What is the frequency
of the light?
c
ν=
λ
ν=
3.00 × 108 m/s
413 nm
WAIT, This won’t work!
69
Example. A certain violet light has a
wavelength of 413 nm. What is the frequency
of the light?
c
ν=
λ
ν=
413 nm
1m
3.00 × 108 m/s
413 nm
=
70
Example. A certain violet light has a
wavelength of 413 nm. What is the frequency
of the light?
c
ν=
λ
ν=
413 nm
3.00 × 108 m/s
4.13 x 10-7 m
1m
-7 m
=
4.13
x
10
1 x 109 nm
71
Example. A certain violet light has a
wavelength of 413 nm. What is the frequency
of the light?
c
ν=
λ
ν=
3.00 × 108 m/s
4.13 × 10-7 m
ν = 7.26 × 1014 Hz
72
Unfortunately, thinking of light as waves lead
to a problem. It was noticed that if light strikes
a metal, then sometimes it could cause electrons
to be emitted (leave the atoms entirely – like in
a solar panel); called the photoelectric effect. If
light was a wave, then all amounts of light
energy should cause this to happen, but this
was not the case. It always took some
minimum amount of energy to get the electrons
to be emitted.
73
74
This lead Max Planck to theorize that light
must carry energy in basic minimum amounts
that he called quanta. Like a delivery person
cannot correctly deliver half a box, the
electrons in atoms cannot gain a fraction of a
quantum of energy (it has to be in whole
numbers).
75
He proposed that this energy was directly
proportional to the frequency of the
electromagnetic radiation and a constant, now
called Planck’s constant. E = h ν
E = energy in Joules (J)
h = Planck’s constant = 6.626 × 10-34 Js
ν = frequency in Hz or 1/s
E
h
ν
76
Example. What is the energy content of one
quantum of the light with a wavelength of 413
nm?
77
Example. What is the energy content of one
quantum of the light with a wavelength of 413
nm?
Note: wavelength is not in the energy equation,
but frequency is. So first, you must solve for
the frequency. As seen in the earlier example, a
wavelength of 413 nm gives a ν = 7.26 × 1014
Hz.
78
Example. What is the energy content of one
quantum of the light with a wavelength of 413
nm?
ν = 7.26 × 1014 Hz
E=h×ν
E = 6.626 × 10-34 Js × 7.26 × 1014 1/s
79
Example. What is the energy content of one
quantum of the light with a wavelength of 413
nm?
ν = 7.26 × 1014 Hz
E=h×ν
E = 6.626 × 10-34 Js × 7.26 × 1014 1/s
E = 4.81 × 10-19 J
80
In 1905 Einstein used Plancks work to propose
that electromagnetic radiation had a dual
wave-particle nature. As a particle,
electromagnetic radiation carries a quantum of
energy of energy, has no mass, and is called a
photon.
81
So to get an electron to emit from a metal, it
must be struck with a photon having quantum
energy big enough, or nothing will happen.
Each metal requires a different quantum
energy, thus each metal can be identified by the
frequency of light needed to emit electron.
82
This idea was expanded
upon to develop an idea
of where the electrons
were in an atom. It was
found that low pressure
gases could be trapped
in a tube and electrified,
and would then glow a
color particular to the
gas inside.
83
Furthermore this light could be passed into a
prism, and instead of getting the entire
spectrum (rainbow) of colors, only certain
wavelengths of light would be seen as small
bars of color, called a line-emission spectrum.
84
This would indicated that the electrons in an
atom were only absorbing specific amounts of
energy from the electricity, causing the
electrons to move from their ground state
(normal position close to the nucleus) to an
excited state (higher energy position further
away from the nucleus). The electrons do not
stay in the excited state for long and fall back
to their ground state, losing the energy equal to
what they gained.
85
Niels Bohr used this to develop a model of the
atom where the electrons could only be in
certain, specific energy level (n) orbits around
the nucleus. Just as you cannot go up half a
rung on a ladder, the electron could not go up a
partial energy level. The electrons gained or
lost enough energy to move a whole number
amount of energy levels (n) away from or closer
to the nucleus, or it did not move.
86
He calculated the amount of energy needed for
an electron of hydrogen to move between each
energy level (n) (which was not constant) and
his calculations agreed with experimental
results.
87
The Balmer series of hydrogen spectral lines
refer to the four lines seen in the visible light
region (the four colored bars). If the electron
was excited to energy level (n) 6, 5, 4, or 3 and
fell to energy level (n) 2, the resulting energy
given off would have a frequency in the visible
region of electromagnetic radiation. (One line
for dropping from 6 to 2, one for 5 to 2, one for
4 to 2, and one for 3 to 2).
88
However, there are other possibilities. If the
electrons drop from n=6, 5, or 4 to n=3, then
the energy given off is not big enough to be seen
as it is in the infrared region. These three lines
in the infrared region are referred to as the
Paschen series. If the electrons drop to n=1,
then the five lines given off are too high in
energy to be seen, as they are in the ultraviolet
region. These lines are referred to as the
Lyman series.
89
90
Model of Atom Review:
91
1. Thomson’s Plum Pudding Model – the atom
is a ball of evenly spread positive stuff with
random negative particles (electrons).
2. Rutherford’s Nuclear Model – the atom has
a central nucleus containing the positive
particles (protons) with the electrons outside.
3. Bohr’s Orbital Model – The electrons circle
the nucleus in specific energy orbits, like the
planets orbit the sun. Unfortunately this
only works for atoms with one electron…
92
4. Quantum Mechanical Model – electrons are
found in specific regions around the nucleus,
but the exact location of the electrons inside
the regions cannot be determined
93
The quantum mechanical model starts with a
Principal Quantum Number (n), which is the
basic energy level of an electron, and often
matches the period number. Possible values
(currently) are 1-7.
94
The quantum mechanical model starts with a
Principal Quantum Number (n), which is the
basic energy level of an electron, and often
matches the period number. Possible values
(currently) are 1-7.
Inside the principal quantum energy level are
sublevels that correspond to different cloud
shapes. The sublevels are designated as s
(sharp), p (principal), d (diffuse), and f
(fundamental).
95
Inside the sublevels are orbitals, specific
regions with a 90% probability of finding
electrons.
• s –orbitals are spherically shaped clouds
around the nucleus
• p -orbitals are bar-bell shaped clouds with
the nucleus between the lobes
• d and f are much more complex in shape
96
Each sublevel has room for a different amount
of electrons, because an orbital can hold two
electrons, then each sublevel has a different
amount of orbitals
97
• s –sublevel can hold 2 electrons, so it has 1
orbital (shape)
• p –sublevel can hold 6 electrons, so it has 3
orbitals (shapes)
• d –sublevel can hold 10 electrons, so it has 5
orbitals (shapes)
• f –sublevel can hold 14 electrons, so it has 7
orbitals (shapes)
98
The s sublevel is simply a sphere centered on the nucleus.
99
The p sublevel has three orbitals. These are often
referred to a dumbbell shape.
100
The d sublevel has five orbitals:
101
The f sublevel has seven orbitals
102
To know the maximum amount of electrons
that could be in any principal quantum level
(and the number of elements that could be
represented) use the formula 2n2
if n=1, then
103
To know the maximum amount of electrons
that could be in any principal quantum level
(and the number of elements that could be
represented) use the formula 2n2
if n=1, then 2 electrons will fit
if n=4,
104
To know the maximum amount of electrons
that could be in any principal quantum level
(and the number of elements that could be
represented) use the formula 2n2
if n=1, then 2 electrons will fit
if n=4, then 32 electrons will fit
105
In order to show on paper where electrons are
likely to be located in an atom, orbital filling
diagrams and electron configurations are
drawn or written. When this is done, three
rules must be followed:
106
1. Aufbau principle – electrons fill lower
energy levels first, thus 1 before 2 and s
before p, etc.
a. orbitals within a sublevel are equal in energy
(called degenerate)
b. the principal energy levels often overlap,
making them seem a little out of order
c. boxes are used to represent orbitals
107
Another way of writing the aufbau principle
diagram:
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d
7s 7p
108
2. Pauli Exclusion principle – an orbital (box)
can hold a maximum of two electrons
(arrows)
a. for two electrons to fit, they have to have
opposite spins
b.
c.
for one electron in the orbital
for two electrons in the orbital (opposite
spins)
109
3. Hund’s Rule – when electrons occupy
degenerate orbitals, one electron is placed
into each orbital with parallel spins before
doubling up
Ex. _____ _____ _____
3p
NOT _____ _____ _____
3p
110
Please navigate to
http://intro.chem.okstate.edu/WorkshopFold
er/Electronconfnew.html
And click through all the elements. Make note
of which ones are exceptions to the Aufbau
principle, and where they are located in the
periodic table.
When the d sublevel get close in energy to the s
sublevel the electrons from the s sublevel
repel each other and one ends up in the d,
even though it is slightly higher in energy.
111
We separate the
orbitals from
each other to be
able to talk
about the
electrons’
locations easier,
but in a real
atom all the
electrons and
orbitals exist at
once, which
might look like
this:
112
Orbital Notation shows the arrows in the boxes
to represent the electrons in an atom. To
shorten this process, an electron
configuration can be written. It leaves out
the information about the number of orbitals
in each sublevel, so it will be expect you
remember that information.
113
It has the general form nl°
n = principal quantum number (1-7…)
l = sublevel letter (s, p, d, or f)
° = number of e- in that orbital (1-14)
114
Ni = 28 e-
115
Ni = 28 e1s2 2s2 2p6 3s2 3p6 4s2 3d8
2 + 2 + 6 + 2 + 6 + 2 + 8 = 28
116
Sn = 50 e-
117
Sn = 50 e1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p2
118
If writing out the entire electron configuration is too
much, we can use the previous (in the periodic
table) noble gas to take the place of part of the
electron configuration:
Polonium:
1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p4
Xenon:
1s22s22p63s23p64s23d104p65s24d105p6
Polonium: [Xe] 6s24f145d106p4
119
When the electron configuration is written for an
element using the noble gas configuration the
electrons written after the noble gas are the ones that
appear on the outside of the atom, called valence
electrons..
120
When elements bond to form compounds, it is these
electrons that are involved. The amount of valence
electrons makes a big difference in how the element
will bond, so to make it easy to predict, we draw
electron dot diagrams.
A) In an electron dot diagram, we use the symbol of
the element and dots to represent the number of
valence electrons.
B) Only s and p electrons with the highest quantum
number count for dot diagrams, even if there are d
and f electrons after the noble gas.
121
Lithium =
So
2
1
1s 2s
Li
122
Beryllium =
So
2
2
1s 2s
Be
123
Boron =
So
2
2
1
1s 2s 2p
B
124
Carbon =
So
2
2
2
1s 2s 2p
C
125
Nitrogen =
So
2
2
3
1s 2s 2p
N
126
Oxygen =
So
2
2
4
1s 2s 2p
O
127
Fluorine =
So
2
2
5
1s 2s 2p
F
128
Neon =
So
2
2
6
1s 2s 2p
Ne
129
Periodic Law – properties of the
elements are a periodic function of
their atomic number
The moon’s phases
and magazine
subscriptions are also
periodic…
The earliest and most
successful arrangement of
the elements was by
Dmitri Mendeleev in the
1870s. He arranged the
elements known at that
time by their physical and
chemical properties into
groups.
His arrangements of the elements left
some gaps, which he claimed were
elements not yet discovered. With his
table he predicted the characteristics
of these missing elements, and was
correct.
During the 1910s, Henry
Moseley used x-ray
spectra to determine
the atomic number
(number of protons) for
elements and proved
that each element had a
different amount of
protons.
Moseley rearranged the elements based
on atomic number, as it is arranged
today, and found gaps which he also
claimed were undiscovered elements.
Some of these elements were found
quickly, while others were not found
until after his death due to their
highly radioactive (unstable) nature.
There are three key items in
understanding and explaining all the
trends (and exceptions) on the
periodic table.
There are three key items in
understanding and explaining all the
trends (and exceptions) on the
periodic table.
1. Effective Nuclear Charge
There are three key items in
understanding and explaining all the
trends (and exceptions) on the
periodic table.
1. Effective Nuclear Charge
2. Energy Levels
There are three key items in
understanding and explaining all the
trends (and exceptions) on the
periodic table.
1. Effective Nuclear Charge
2. Energy Levels
3. Coulombs Law
1. Effective Nuclear Charge – pull of
the protons in the nucleus on the
valence (outer) energy level electrons
1. Effective Nuclear Charge – pull of
the protons in the nucleus on the
valence (outer) energy level electrons
The greater the atomic number, the
greater the number of protons, and
the greater the effective nuclear
charge.
Effective Nuclear Charge has the
greatest effect moving from left to
right across a period (it increases).
Effective Nuclear Charge has the
greatest effect moving from left to
right across a period (it increases).
Which element has more effective
nuclear charge, P or S?
Effective Nuclear Charge has the
greatest effect moving from left to
right across a period (it increases).
Which element has more effective
nuclear charge, P or S?
Which element has more effective
nuclear charge, S or Se? (Why won’t
this work?)
2. Energy levels - the principal
quantum level of the electrons,
sometimes called shells
2. Energy levels - the principal
quantum level of the electrons,
sometimes called shells
As elements increase in atomic number
they also increase in the number of
electrons. These electrons occupy
higher and higher energy levels.
Higher numbered energy levels are
farther away from the nucleus.
1s
1s
2s
1s
2s
3s
Energy Levels have the greatest effect
moving from top to bottom within a
group (it increases).
Energy Levels have the greatest effect
moving from top to bottom within a
group (it increases).
Which element has more energy levels,
S or Se?
Energy Levels have the greatest effect
moving from top to bottom within a
group (it increases).
Which element has more energy levels,
S or Se?
Which element has more energy levels,
P or S? (Why won’t this work?)
3. Coulombs Law states that the force
of attraction between things is
directly proportional to the size of the
charge and inversely proportional to
the square of the distance between
them.
For the following trends, you must
know the general trends (memorize),
but using the three key items you
must be able to explain the trend as
well (understand).
1. Atomic Radius
size of the atom
a. increases from top to bottom within
a group
Why does it increase from top to
bottom?
1. Atomic Radius
size of the atom
b. decreases from left to right within a
period
Why does it decrease from left to right?
2. Ionization Energy
energy required to remove an
electron from a gaseous atom
a. increases bottom to top
Why is it easier to remove electrons
from atoms at the bottom of the
P.T.?
2. Ionization Energy
energy required to remove an
electron from a gaseous atom
b. increases left to right
Why is it easier to remove electrons
from atoms at the left of the P.T.?
2. Ionization Energy
energy required to remove an
electron from a gaseous atom
c. This is sometimes called metallic
character, as metals tend to lose
electrons easily.
Which element would exhibit the most
metallic character?
3. Electron Affinity
energy released when an electron is
added to a gaseous atom
a. increases bottom to top
Why is more energy released when
electrons are added to atoms at the
top of the P.T.?
3. Electron Affinity
energy released when an electron is
added to a gaseous atom
b. increases left to right
Why is more energy released when
electrons are added to atoms at the
right of the P.T.?
3. Electron Affinity
energy released when an electron is
added to a gaseous atom
c. This is sometimes called nonmetallic
character, as nonmetals tend to gain
electrons easily.
Which element would exhibit the most
nonmetallic character?
4. Electronegativity
measure of an atom’s pull on another
atom’s electrons
a. increases bottom to top
Why are atoms at the top of the P.T.
able to pull stronger on a different
atom’s electrons?
4. Electronegativity
measure of an atom’s pull on another
atom’s electrons
b. increases left to right
Why are atoms at the right of the P.T.
able to pull stronger on a different
atom’s electrons?
5. ions
charged atoms
a. cations – positive ion formed from
an atom losing an electron
The ionic radius is always smaller than
the original atom.
Why?
5. ions
charged atoms
b. anions – negative ion formed from an
atom gaining an electron
The ionic radius is always larger than
the original atom.
Why?
Let’s summarize the trends.
Remember you must know the
trend (memorize) but you also need
to be able to explain why the trend
exists.
Atomic Radius
Ionization Energy, Electron Affinity, Electronegativity
Energy Levels
Effective Nuclear Charge
Metallic Character
Nonmetallic Character