Roy Kennedy Massachusetts Bay Community College Wellesley Hills, MA Introductory Chemistry, 3 rd Edition Nivaldo Tro Chapter 9 Electrons in Atoms and the Periodic Table 2009, Prentice Hall Chapter Opening blimp
Roy Kennedy
Massachusetts Bay Community College
Wellesley Hills, MA
Introductory Chemistry, 3rd Edition
Nivaldo Tro
Chapter 9
Electrons in Atoms
and the
Periodic Table
2009, Prentice Hall
Chapter
Opening
blimp
9.1 Blimps, Balloons, and Models of the Atom
9.2 Light, Electromagnetic Radiation
9.3 The Electromagnetic Spectrum
9.4 The Bohr Model: Atoms with Orbits
9.5 The Quantum Mechanical Model: Atoms with Orbitals
9.6 Quantum Mechanical Orbitals
9.7 Electron Configurations and the Periodic Table
9.8 The Explanatory Power of the Quantum Mechanical
Model
9.9 Periodic Trends: Atomic Size, Ionization Energy, and
Metallic Character
2Tro's "Introductory Chemistry",
Chapter 9
OUTLINE
• Hydrogen blimps may burn up
• Helium blimps can not burn
• Why is H reactive, but He inert?
3Tro's "Introductory Chemistry",
Chapter 9
9.1 Blimps, Balloons, and
Models of the Atom
4Tro's "Introductory Chemistry",
Chapter 9
9.2 Light, Electromagnetic
Radiation
5Tro's "Introductory Chemistry",
Chapter 9
6Tro's "Introductory Chemistry",
Chapter 9
Classical View of the Universe• Uniuverse is made of matter or energy.
• Matter has mass and volume.
Energy doesn’t have mass and volume.
• Matter is composed of particles.
• Energy, should not be composed of particles.
• All energy has in common that it travels in waves.
7Tro's "Introductory Chemistry",
Chapter 9
The Nature of Light—Its Wave
Nature
• Light is a the forms of energy. Called
electromagnetic radiation made of waves.
• Electromagnetic radiation moves through
space like waves move across the surface of a
pond
8
Speed of Energy Transmission
Fireworks contrasting
the speed of sound = 340 m/s
With
The speed of light = 3.00 x 108 m/s
9Tro's "Introductory Chemistry",
Chapter 9
Electromagnetic Waves• Wave characteristics:
Wave speed,
Height (amplitude),
Wavelength,
Number of wave peaks that pass in a given time = frequency.
• All electromagnetic waves move through space at the same, constant speed.
3.00 x 108 meters per second in a vacuum = The speed of light, c.
10Tro's "Introductory Chemistry",
Chapter 9
Characterizing Waves• The amplitude is the height of the wave = intensity.
• The wavelength (l) is a measure of the distance of a repeat unit.
The distance from one crest to the next.
Usually measured in nanometers.
1 nm = 1 x 10-9 m
11Tro's "Introductory Chemistry",
Chapter 9
Electromagnetic Waves
Figure 9.1 showing a wave
12Tro's "Introductory Chemistry",
Chapter 9
Characterizing Waves
• The frequency (n) = number of waves that pass a point in a given period of time.
The number of waves = number of cycles.
Units are hertz (Hz), or cycles/s = s-1.
• The total energy is proportional to the amplitude and frequency of the waves.
The larger the wave amplitude, the more force it has.
The more frequently the waves strike, the more total force there is.
13
The Electromagnetic Spectrum• Light passed through a prism is separated into all its
colors called a continuous spectrum.
Figure 9.2 showing a continuous
spectrum
14Tro's "Introductory Chemistry",
Chapter 9
Color
• The color of light is determined by its wavelength.
Or frequency.
• White light is a mixture of all the colors of visible light.
A spectrum.
RedOrangeYellowGreenBlueViolet.
• The observed color is predominantly the colors reflected al other
are absorbed.
9.3 The Electromagnetic
Spectrum
15Tro's "Introductory Chemistry",
Chapter 9
16Tro's "Introductory Chemistry",
Chapter 9
Electromagnetic Spectrum
Figure 9.4 the electromagnetic
spectrum
17Tro's "Introductory Chemistry",
Chapter 9
Particles of Light
• Scientists in the early 20th century showed that
electromagnetic radiation was composed of
particles we call photons.
Max Planck and Albert Einstein..
• Each wavelength of light has photons that have
a different amount of energy.
18Tro's "Introductory Chemistry",
Chapter 9
The Electromagnetic Spectrum and
Photon Energy
• Short wavelength light has photons with high energy.
• High frequency light has photons with high energy.
• High-energy electromagnetic radiation can potentially damage biological molecules.
Ionizing radiation.
9.4 The Bohr Model: Atoms with
Orbits
19Tro's "Introductory Chemistry",
Chapter 9
20Tro's "Introductory Chemistry",
Chapter 9
Light’s Relationship to Matter
• Atoms can acquire extra energy, but
they must eventually release it,
usually in the form of light.
• However, atoms emit only very
specific wavelengths which can be
used to identify the element.
Figure 9.6 Neon light
Figure 9.7 emission
tubes
21Tro's "Introductory Chemistry",
Chapter 9
Emission Spectrum
Figure 9.8 (part 1) a line emission
spectrum.
22Tro's "Introductory Chemistry",
Chapter 9
Spectra
Figure 9.8 (part 2) comparison of a
continuous spectrum and the
emission spectrum of several
elements.
23Tro's "Introductory Chemistry",
Chapter 9
A figure showing the comlimentary
nature of emission (a few lines) an
an absorption spectum. Hole
appear in the absorption spectrum
At the same wavelengths lines
appear in the emission sprectum.
24Tro's "Introductory Chemistry",
Chapter 9
The Bohr Model of
the Atom• Neils Bohr developed a model of the atom to explain
how the atom changes when it undergoes energy transitions.
• Bohr’s major idea was that the energy of the atom was quantized, and energy in the atom depends of the electron’s position.
Quantized means only have very specific amounts of energy.
25
The Bohr Model of the Atom:
Electron Orbits• Electrons travel in orbits around the nucleus.
More like shells than planet orbits.
• The farther the electron is from the nucleus the more energy it has.
Figure 9.9 Bohr orbits
26Tro's "Introductory Chemistry",
Chapter 9
The Bohr Model of the Atom:
Orbits and Energy, Continued• Each orbit has a specific amount
of energy.
• The energy of each orbit is characterized by an integer, n, is called a quantum number. The larger the integer, the greater the distance from the nucleus, and the greater the energy.
27Tro's "Introductory Chemistry",
Chapter 9
The Bohr Model of the Atom:
Energy Transitions
• When the atom gains energy, the electron leaps to a
higher orbit.
• When the electron leaps from a higher energy orbit
to one that is closer to the nucleus, energy is emitted
as a photon of light—a quantum of energy.
28Tro's "Introductory Chemistry",
Chapter 9
The Bohr Model of the Atom
Figure 9.11 model of energy
transitions. Lowers to higher by
absorption of energy of a photon.
Higher to lower with emission of a
photon.
29Tro's "Introductory Chemistry",
Chapter 9
The Bohr Model of the Atom:
Ground and Excited States
• The lowest amount of energy corresponds to being in the n = 1 orbit called the ground state.
• Electrons in higher orbital are in an excited state.
• Excited state are unstable so they release energy to return to the ground state.
30Tro's "Introductory Chemistry",
Chapter 9
The Bohr Model of the Atom:
Hydrogen Spectrum
• Every hydrogen atom has identical orbits,
• Distances between the orbits differ so transitions do not have the same energy.
• The emission spectrum that has many lines that are unique to hydrogen.
• Visible lines are all transitions to level n=2
31Tro's "Introductory Chemistry",
Chapter 9
The Bohr Model of the Atom:
Hydrogen Spectrum, Continued
Figure 9.12 Correlation of orbital
transition and the wavelength of
light emitted
32Tro's "Introductory Chemistry",
Chapter 9
The Bohr Model of the Atom:
Success and Failure
• Bohr model accurately predicts the spectrum
of hydrogen.
• It fails when applied to multi-electron atoms.
• A better theory was needed.
9.5 The Quantum Mechanical
Model: Atoms with Orbitals
33Tro's "Introductory Chemistry",
Chapter 9
34Tro's "Introductory Chemistry",
Chapter 9
The Quantum-Mechanical Model
of the Atom
• Erwin Schrödinger applied the
mathematics of probability and the
ideas of quantizing energy to the
physics equations that describe
waves, resulting in an equation that
predicts the probability of finding
an electron with a particular
amount of energy at a particular
location in the atom.
Photo of
Schrödinger
35Tro's "Introductory Chemistry",
Chapter 9
The Quantum-Mechanical Model:
Orbitals
• The result is a map of regions in the atom
that have a particular probability for finding
the electron.
• An orbital is a region where we have a very
high probability of finding the electron
when it has a particular amount of energy.
9.6 Quantum Mechanical
Orbitals
36Tro's "Introductory Chemistry",
Chapter 9
37Tro's "Introductory Chemistry",
Chapter 9
Orbits vs. Orbitals
Exact Pathways vs. Probability Regions
Figure 9.13 a fixed path
Figure 9.14 A probability distribution
38Tro's "Introductory Chemistry",
Chapter 9
Wave–Particle Duality
• We’ve seen that light has the characteristics of waves and particles (photons).
• Electrons also have the characteristics of both particles and waves.
• It is impossible to predict the exact path of an electron.
39Tro's "Introductory Chemistry",
Chapter 9
The Quantum-Mechanical Model:
Quantum Numbers
• 3 integers, called quantum
numbers, that quantize the
energy.
• The principal quantum
number, n, specifies the main
energy level.
Figure 9.15
Atom Energy
Levels
40Tro's "Introductory Chemistry",
Chapter 9
The Quantum-Mechanical Model:
Quantum Numbers, Continued
• Each principal energy shell has one or more subshells.
The number of subshells = the principal quantum number.
• The subshell is often identified s,p,d,f, that
corresponds to a particular shape.
41Tro's "Introductory Chemistry",
Chapter 9
Shells and Subshells
Figure 9.19
Number of shells and subshells
In an atom
42Tro's "Introductory Chemistry",
Chapter 9
How Does the 1s Subshell Differ
from the 2s Subshell?
Figure 9.20 Sizes and shapes of the
1s and 2s orbitals
43Tro's "Introductory Chemistry",
Chapter 9
Probability Maps and Orbital Shape:
s Orbitals
44Tro's "Introductory Chemistry",
Chapter 9
Probability Maps and Orbital Shape:
p Orbitals
Figure 9.21 Sizes and shapes of the three
P orbitals
45Tro's "Introductory Chemistry",
Chapter 9
Probability Maps and Orbital Shape:
d Orbitals
Figure 9.22 d orbitals
46Tro's "Introductory Chemistry",
Chapter 9
Subshells and Orbitals
The subshells in a shell of H all have the same energy, but for multielectron atoms the subshells have different energies.
s < p < d < f.
• Each subshell contains one or more orbitals.
s =1 orbital.
p = 3 orbitals.
d = 5 orbitals.
f = 7 orbitals.
47Tro's "Introductory Chemistry",
Chapter 9
The Quantum-Mechanical Model:
Energy Transitions
• Atoms gain or lose energy as the electron
leaps between orbitals in different energy
shells and subshells.
• The ground state of the electron is the
lowest energy orbital.
• Higher energy orbitals are excited states.
48Tro's "Introductory Chemistry",
Chapter 9
The Bohr Model vs.
the Quantum-Mechanical Model
• Both predict the spectrum of hydrogen
accurately.
• Only the quantum-mechanical model
predicts the spectra of multi-electron atoms.
9.7 Electron Configurations and
the Periodic Table
49Tro's "Introductory Chemistry",
Chapter 9
50Tro's "Introductory Chemistry",
Chapter 9
Electron Configurations
• The distribution of electrons in the various energy
shells and subshells in its ground state is the
electron configuration.
• Maximum number of electrons each subshell can
hold.
s = 2, p = 6, d = 10, f = 14.
• We place electrons in the energy shells and
subshells in order of energy, from low energy up.
Aufbau principle.
• Figure 9.26 Energy levels of a multielectron
• atom
51Tro's "Introductory Chemistry",
Chapter 9
52Tro's "Introductory Chemistry",
Chapter 9
Filling an Orbital with Electrons
• Each orbital may have a maximum of 2 electrons.
Pauli Exclusion principle.
• Electrons spin on an axis and two electrons the same orbital must have opposite spins.
So their magnetic fields will cancel.
53Tro's "Introductory Chemistry",
Chapter 9
Orbital Diagrams
• orbital as a square and the electrons as arrows.
The direction of the arrow represents the spin of the
electron.
Figure 9.? Possible electron
occupancies of an orbital diagram
0 or 1 or 2(paired spins)
54Tro's "Introductory Chemistry",
Chapter 9
Order of Subshell Filling
in Ground State Electron Configurations
Draw the triangle of orbitals as
shown
Next, draw arrows through
the diagonals, looping back
to the next diagonal
each time.
Figure 9.27 orbital filling order
triangle
55Tro's "Introductory Chemistry",
Chapter 9
Filling the Orbitals in a Subshell
with Electrons
• Energy shells fill from lowest energy to highest.
• Subshells fill from lowest energy to highest.
s → p → d → f
• Orbitals in thesame subshell have the same
energy.
• When filling a subshell, place one electron in each
before completing pairs.
Hund’s rule.
56Tro's "Introductory Chemistry",
Chapter 9
Electron Configuration of Atoms
in their Ground State
• The electron configuration is a listing of the subshells in order with the number of electrons written as a superscript.
Kr = 36 electrons = 1s22s22p63s23p64s23d104p6
• A short-hand way of writing an electron configuration is to use the symbol of the previous noble gas in [] to represent all the inner electrons. Noble gases (except He) end with p subshells Ne (2p), Ar (3p), Kr (4p), Xe (5p)
Rb = 37 electrons = 1s22s22p63s23p64s23d104p65s1 = [Kr]5s1
57Tro's "Introductory Chemistry",
Chapter 9
Electron Configurations
Page 293 orbital diagrams for
Li, Be, B, and C
58Tro's "Introductory Chemistry",
Chapter 9
1. Atomic number from the periodic table = s
the number of protons and electrons
Mg Z = 12, so Mg has 12 protons and 12 electrons.
Example—Write the Ground State
Orbital Diagram and Electron
Configuration of Magnesium.
Electron configuration of any
atom or ion• Find the atomic number from periodic table = the number of electrons
for an atom.
• Fill the orbitals in the correct order, remember the maximum electron
per subshell. Use the triangle for the order.
• s = 2 e- p = 6 e- d = 10 e- f = 14 e-
• Place electron in orbital boxes, half fill boxes before pairing.
• When electrons are used up, write the electron configuration using a
subscript for the number of electron in a subshell
59Tro's "Introductory Chemistry",
Chapter 9
Example Mg
• Mg, Z = 12.
• We need 12 electrons
Need up to 3s
• Fill the correct number of electrons
Total = 2 4 10 12, stop
• Write configuration 1s22s22p63s2 = [Ne]3s2
• .
60
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
+2 = 4e−
+6 +2 = 12e−
2e−
1s 2s 2p 3s
1s 2s 2p 3s 3p
3p
Example Manganese
61Tro's "Introductory Chemistry",
Chapter 9
Mn, Z = 25, we need 25 electrons
If we fill to 4s = 20 and 3d = 30
So 3d is partially filled.1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
2 e−
+2 = 4e−
+6 +2 = 12e−
+10 = 30e−
+6 +2 = 20e−
Manganese continued
62Tro's "Introductory Chemistry",
Chapter 9
1s 2s 2p 3s 3p 4s
3 d
Lay out the required subshells in order
Manganese continued
63Tro's "Introductory Chemistry",
Chapter 9
1s 2s 2p 3s 3p 4s
3 d
Populate the orbitals with electrons. Notice reach 3d is half filled
2 4 10 12 18 20
(5 = 25, stop)
Write the electron configuration.
Mn = 1s22s22p63s23p64s23d5 or [Ar]4s23d5
.
64Tro's "Introductory Chemistry",
Chapter 9
Practice—Write the Full Ground State Orbital
Diagram and Electron Configuration of Potassium,
Continued.
K Z = 19, therefore 19 e−
1s 2s 2p 3s 3p 4s
Therefore the electron configuration is
K = 1s22s22p63s23p64s1 or K = [Ar]4s1
Ions - cations subtract and anions
add electrons
65Tro's "Introductory Chemistry",
Chapter 9
Example F- F = 9 e- and -1 means an extra e- 9 e- + 1 e- = 10 e-
Write configuration F- = 1s22s22p6 = [Ne]
1s 2s 2p
Ions Ca2+
66Tro's "Introductory Chemistry",
Chapter 9
Ca has 20 e- and (+2) means 2 e- lost 20 e- – 2 e- = 18 e-
Ca2+ = 1s22s22p63s23p6 or [Ar]
1s 2s 2p 3s 3p
67Tro's "Introductory Chemistry",
Chapter 9
Valence Electrons
• The electrons in all the subshells with the highest principal energy shells are called the valence electrons.
• Electrons in lower energy shells are called core electrons.
• Chemists have observed that one of the most important factors in the way an atom behaves, both chemically and physically, is the number of valence electrons.
68Tro's "Introductory Chemistry",
Chapter 9
Valence Electrons, Continued
Rb = 37 electrons = 1s22s22p63s23p64s23d104p65s1
• The highest principal energy shell the 5th, therefore,
Rb has 1 valence electron and 36 core electrons.
Kr = 36 electrons = 1s22s22p63s23p64s23d104p6
• The highest principal energy shell of Kr that contains
electrons is the 4th, therefore, Kr has 8 valence
electrons and 28 core electrons.
• Core electrons are always a noble gas number of
electrons
Practice—Determine the Number and Types of
Valence Electrons in an As Atom, Continued.
As Z = 33, therefore 33 e−.
Therefore, the electron configuration is 1s22s22p63s23p64s23d104p3.
The valence electrons are in level n = 4 4s24p3 2+3 = 5 valence
electrons
.
70Tro's "Introductory Chemistry",
Chapter 9
Electron Configurations and
the Periodic Table
Figure 9.28 Electron configuration
and the periodic table
71Tro's "Introductory Chemistry",
Chapter 9
Electron Configurations from
the Periodic Table (Main Group)
• Same period (row) have valence electrons in the same principal energy shell.
• Across a row valence electronsgoes from 1 to 8.
• Elements in the same group (column) have the same number of valence electrons.
72Tro's "Introductory Chemistry",
Chapter 9
Electron Configuration and the
Periodic Table
• Elements in the same column have similar chemical and physical properties because their valence shell electron configuration is the same.
• The number of valence electrons for the main group elements is the same as the group number.
73Tro's "Introductory Chemistry",
Chapter 9
Figure 9.29 illustration the s block,
P block, d block, and f blocks of the
Periodic table.
74Tro's "Introductory Chemistry",
Chapter 9
Electron Configuration from
the Periodic Table
• The inner electron configuration is the same as the
noble gas of the preceding period.
• To get the outer electron configuration from the
preceding noble gas, loop through the next period,
marking the subshells as you go, until you reach the
element.
The valence energy shell = the period number.
The d block is always one energy shell below the period
number and the f is two energy shells below.
• Figure 9.30, 9.31 – using the table to write
the electron configuration for P and As
75Tro's "Introductory Chemistry",
Chapter 9
76Tro's "Introductory Chemistry",
Chapter 9
Practice—Use the Periodic Table to Write the Short
Electron Configuration and Orbital Diagram for
Each of the Following and Determine the Number of
Valence Electrons.
• Na (at. no. 11).
• Te (at. no. 52).
77Tro's "Introductory Chemistry",
Chapter 9
Practice—Use the Periodic Table to Write the Short
Electron Configuration and Orbital Diagram for
Each of the Following and Determine the Number of
Valence Electrons, Continued.
• Na (at. no. 11). [Ne]3s1 1 valence electron
• Te (at. no. 52). [Kr]5s24d105p4 6 valence electrons
3s
5s 5p4d
78Tro's "Introductory Chemistry",
Chapter 9
The Explanatory Power of
the Quantum-Mechanical Model
• The properties of the elements are largely
determined by the number of valence electrons
they contain.
• Since elements in the same column have the same
number of valence electrons, they show similar
properties.
• Since the number of valence electrons increases
across the period, the properties vary in a regular
fashion.
Tro's "Introductory Chemistry",
Chapter 9
The Noble Gas
Electron Configuration• The noble gases have 8 valence electrons.
Except for He, which has only 2 electrons.
• non-reactive because the electron
configuration of the noble gases is especially
stable.
Figure 9.32
Similarity
In eletron
Configuration
Of the
Nobel gases
Tro's "Introductory Chemistry",
Chapter 9
Everyone Wants to Be Like a Noble Gas!
The Alkali Metals
• The alkali metals have one more
electron than the previous noble
gas.
• The alkali metals tend to lose their
extra electron giving electron
configuration of a noble gas.
Cation with a 1+ charge.
Figure 9.32
Similarity
In eletron
Configuration
Of the
Alkali metals
Group 1A
81Tro's "Introductory Chemistry",
Chapter 9
81
Everyone Wants to Be Like a Noble Gas!
The Halogens• Halogens all have one fewer electron than
the next noble gas.
• Reaction with Metals - Gain an electron and attain the electron configuration of the next noble gas forming an anion with charge 1−.
• Reactions with nonmetals - share electrons with the other nonmetal so that each attains the electron configuration of a noble gas.
Figure 9.32
Similarity
In electron
Configuration
Of the
Halogens
Group 7A
82Tro's "Introductory Chemistry",
Chapter 9
Everyone Wants to Be
Like a Noble Gas!
• Alkali metals are the most reactive metals.
• Halogens are the most reactive group of nonmetals.
• They are only one electron away from having a very
stable electron configuration.
The same as a noble gas.
83Tro's "Introductory Chemistry",
Chapter 9
Stable Electron Configuration
and Ion Charge• Metals losing valence
electrons = noble gas.
• Nonmetals gaining valence electrons = noble gas.
Atom Atom’s
electron
config
Ion Ion’s
electron
config
Na [Ne]3s1
Na+
[Ne]
Mg [Ne]3s2 Mg
2+ [Ne]
Al [Ne]3s23p
1 Al
3+ [Ne]
O [He]2s2p4 O
2- [Ne]
F [He]2s22p
5 F
- [Ne]
Periodic Trends in the
Properties of the Elements
85Tro's "Introductory Chemistry",
Chapter 9
Trends in Atomic Size
• Either volume or radius of sphere
• Down a column on the periodic table, the size of the atom increases.
Valence shell farther from nucleus.
Effective nuclear charge fairly close.
• Left to right across a period, the size of the atom decreases.
Adding electrons to same valence shell.
Effective nuclear charge increases.
Valence shell held closer.
86
Trends in Atomic Size, Continued
Figure 9.37 from book shows radii of main group atoms.
87Tro's "Introductory Chemistry",
Chapter 9
Graphical plot of atomic radius showing the trend
Similar to this
plot from webelements
88Tro's "Introductory Chemistry",
Chapter 9
Example 9.6 – Choose the
Larger Atom in Each Pair
• C or O
• Li or K
• C or Al
• Se or I?
Page 302 left margin
89Tro's "Introductory Chemistry",
Chapter 9
Practice—Choose the
Larger Atom in Each Pair.
1. N or F
2. C or Ge
3. N or Al
4. Al or Ge
Answers next slide
90Tro's "Introductory Chemistry",
Chapter 9
1. N or F
2. C or Ge
3. N or Al
4. Al or Ge? opposing trends
1. N or F, N is further left
2. Ge is lower
3. Al is down and to the left
Practice—Choose the
Larger Atom in Each Pair, Continued.
91Tro's "Introductory Chemistry",
Chapter 9
Ionization Energy
• Minimum energy needed to remove an electron
from an atom.
Gas state.
Endothermic process.
Valence electron.
M(g) + 1st IE M1+(g) + 1 e-
M+1(g) + 2nd IE M2+(g) + 1 e-
First ionization energy = energy to remove electron from
neutral atom; 2nd IE = energy to remove from +1 ion; etc.
92
Plot of 1st ionization energy trend from
schoolworkhelper.net
93Tro's "Introductory Chemistry",
Chapter 9
Trends in Ionization Energy
• As atomic radius increases, the ionization energy (IE) generally decreases.
Because the electron is closer to the nucleus.
• 1st IE < 2nd IE < 3rd IE …
• As you traverse down a column, the IE gets smaller.
Valence electron farther from nucleus.
• As you traverse left to right across a period, the IE gets larger.
Effective nuclear charge increases.
94Tro's "Introductory Chemistry",
Chapter 9
Trends in Ionization Energy, Continued
95
1. Al or S
2. As or Sb, Sb is further down
1. Al or S
2. As or Sb
3. N or Si, Si is further down and left
1. Al or S
2. As or Sb
3. N or Si
4. O or Cl, opposing trends
Example—Choose the Atom in Each Pair
with the Higher First Ionization Energy
1. Al or S, Al is further left
Page 305 left margin
Answers next page
96Tro's "Introductory Chemistry",
Chapter 9
Practice—Choose the Atom with the
Highest Ionization Energy in Each Pair
1. Mg or P
2. Cl or Br
3. Se or Sb
4. P or Se
97Tro's "Introductory Chemistry",
Chapter 9
Practice—Choose the Atom with the
Highest Ionization Energy in Each Pair,
Continued
1. Mg or P P is to the right
2. Cl or Br Cl is higher
3. Se or Sb Se is up and right
4. P or Se ? Opposing trends
98
Metallic Character• How well an element’s properties match the
general properties of a metal.• Metals:Malleable and ductile as solids.Solids are shiny, lustrous, and reflect light.Solids conduct heat and electricity.Most oxides basic and ionic.Form cations in solution.Lose electrons in reactions—oxidized.
• Nonmetals:Brittle in solid state.Solid surface is dull, nonreflective.Solids are electrical and thermal insulators.Most oxides are acidic and molecular.Form anions and polyatomic anions.Gain electrons in reactions—reduced.
99Tro's "Introductory Chemistry",
Chapter 9
Metallic Character, Continued
• In general, metals are found on the left of
the periodic table and nonmetals on the
right.
• As you traverse left to right across the
period, the elements become less metallic.
• As you traverse down a column, the
elements become more metallic.
100Tro's "Introductory Chemistry",
Chapter 9
Trends in Metallic Character
101Tro's "Introductory Chemistry",
Chapter 9
1. Sn or Te
2. P or Sb, Sb is further down
1. Sn or Te
2. P or Sb
3. Ge or In, In is further down & left
1. Sn or Te
2. P or Sb
3. Ge or In
4. S or Br? opposing trends
1. Sn or Te, Sn is further left
Example—Choose the
More Metallic Element in Each Pair
Page 306 left margin
102Tro's "Introductory Chemistry",
Chapter 9
Practice—Choose the
More Metallic Element in Each Pair
1. Sn or Te
2. Si or Sn
3. Br or Te
4. Se or I
103Tro's "Introductory Chemistry",
Chapter 9
Practice—Choose the
More Metallic Element in Each Pair,
Continued
1. Sn or Te Sn is farther left
2. Si or Sn Sn is lower
3. Br or Te Te is left and lower
4. Se or I ? Opposong trends