4B Introduction to Periodic Table and
Electrons
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A very Organized Man???
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Elements
• Arranged in a pattern called the Periodic
Table
• First arranged by Dmitri Mendeleev, a
Russian scientist in 1869
• Position on the table allows us to predict
properties of the element
• Called “Periodic” because as atomic number
increases, periodically, the properties of the
elements repeats
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Magnesium
burns in air
to give a
bright
white
flame.
Indonesian men carrying chunks of
elemental sulfur in baskets.
Source: API/Explorer/Photo Researchers, Inc.
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The Modern Periodic Table
• Elements with similar chemical and
physical properties are in the same
column
• Columns are called Groups or
Families
• Rows are called Periods
• Each period shows the pattern of
properties repeated in the next period
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Figure 4.11: The periodic table.
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1. Alkali metals 2. Alkaline Earth metals
3. Transition Metals 4. Metalloids
5. Non-Metals
6. Halogens
7. Noble Gases 8. Lanthanides
9. Actinides
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Alkali Metals Group 1
• Include elements Lithium,
Sodium, Potassium,
Rubidium, Cesium, and
Francium
– Very reactive, especially
with water
– Low density, soft metals
– Oxidize quickly, one
valence electron
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Alkaline Earth Metals Group 2
• Includes the elements of
Beryllium, Magnesium,
Calcium, Strontium,
Barium and Radium
– Also reactive, but not to
the degree of the Alkali
Metals
– 2 valence electrons
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• Transition Elements
– Hard metals with high melting point and
high boiling points.
– Relatively loose hold on valence shell
electrons allows for electrical conductivity.
– Compared to Groups 1A and 2A, transition
metals oxidize slower (if at all), e.g. Ag
tarnishes over months, iron also oxidizes
over time (rusts), but Au does not oxidize
easily.
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Metalloids
• share metallic and
non-metallic
properties.
•Si is a semiconductor
of electrical current (a
metallic property). It is
also shiny and brittle.
(a non-metal property
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Halogens
• Fluorine = F2
– pale yellow gas
• Group 7A =
Halogens
• Chlorine = Cl2
– pale green gas
• Bromine = Br2
– brown liquid that has lots of
brown vapor over it
– Only other liquid element at
room conditions is the
metal Hg
• Iodine = I2
Bromine liquid/gas – lustrous, purple solid
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Halogens
• Highly reactive
non-metals which
exist in different
forms at room
temperature
(gases, liquid,
solids)
• 7 valence
electrons
Chlorine gas
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Halogens
• Highly reactive
non-metals which
exist in different
forms at room
temperature
(gases, liquid,
solids)
• 7 valence
electrons
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Noble Gases
• Helium, Neon, Argon, Krypton, Xenon,
Radon
• all colorless gases at room temperature
• very non-reactive, practically inert
• 8 valence electrons
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Lanthanides/Actinides
Also known as rare earth elements
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Left to Right
• Metals
– about 75% of all the elements
– lustrous, malleable, ductile, conduct
heat and electricity
• Nonmetals
– dull, brittle, insulators
• Metalloids
– also known as semi-metals
– some properties of both metals &
nonmetals
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Figure 4.12: The elements classified
as metals and as nonmetals.
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4.9 Natural States of Elements
• Most “stuff” around us is a
mixture of pure substances, not
generally found in uncombined
form
• A few “native” elements can be
found, such as Au, Ag, Pt, Cu,
etc….
• Others include the elements of
Group 8, Noble Gases He, Ne,
Ar, Kr, Xe, Rn
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Air
• Air, as a mixture, is a mixture of
several pure substances, including……
– Argon—separate particles of Ar are
found
– Neon—separate particles of Ne are found
– Xenon—separate particles of Xe are
found
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Air
• ….but also to be found are O2 and N2 (and
others) which are found not as single
particles, but as
Diatomic molecules
A Diatomic molecule is made up of 2 atoms
H 2 O2
F2
Br2
N2
Cl2
I2
H2O2N2Cl2Br2I2F2
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So in a sample of air….you would
find some elements present as
atoms, some as molecules (of course, some
compounds will also be present…but we’re not talking about them )
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Other forms of elements in nature…..Allotropes
• Many solid nonmetallic elements can exist
in different forms with different physical
properties, these are called allotropes
• the different physical properties arise from
the different arrangements of the atoms in
the solid
• Allotropes of Carbon include
– diamond
– graphite
– buckminsterfullerene
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Graphite and diamond, two forms of carbon.
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Figure 4.18: The three solid elemental forms of
carbon (allotropes).
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Buckminsterfullerene
• These structures are
also known as
“buckyballs”
• Named after
Buckminster Fuller, a
great thinker and
industrial designer,
who designed the
“geodesic dome”
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Other allotropes
• Sulfur has three allotropes
– Orthorhombic form
(yellow)
– Monoclinic form (orange)
– Amorphous form (brown)
• Oxygen has two
allotropes
– O2 and O3
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States of Matter
• Note which elements are gases at room
temp.
• Note which elements are liquids at room
temp.
• All the rest are solids
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Modern Atomic Theory
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Rutherford’s Atom
The concept of a nuclear atom (charged electrons
moving around the nucleus) resulted from Ernest
Rutherford’s experiments.
• Rutherford showed:
– Nucleus is composed of protons (positive)
– The nucleus is very small compared to the size of the
entire atom.
– Nucleus surrounded by electrons
• Question left unanswered:
– What is the arrangement of the electrons? Some clues
come from “Light”
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?
Figure 10.1:
The
Rutherford
atom.
?
?
?
Electromagnetic Radiation
• Classical physics says matter made up of
particles, energy travels in waves
• Electromagnetic Radiation is radiant energy,
that travels in waves
• All waves have individual characteristics
of….
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Electromagnetic Waves
• velocity = c = speed of light
– 3.0x 108 m/s
– all types of light energy travel at the same speed
• amplitude = A = measure of the intensity of the
wave, “brightness”
• wavelength =  = distance between two
consecutive points in a wave
– EMR generally measured in nanometers (1 nm = 10-9 m)
• frequency =  = the number of waves that pass a
point in space in one second
– generally measured in Hertz (Hz),
– 1 Hz = 1 wave/sec
• c=x
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Types of Electromagnetic Radiation
•
•
•
•
Radiowaves = low frequency and energy
Microwaves
Infrared (IR)
Visible
– ROYGBIV
• Ultraviolet (UV)
• X-rays
• Gamma rays = high frequency and energy
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Figure 10.4: The different wavelengths of
electromagnetic radiation.
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Max Planck’s Revelation
• Stated that light came in particles called
quanta or photons, little packets of fixed
amounts of energy
– Basis of quantum theory
Nobel Prize 1918
• The energy of the photon is directly
proportional to the frequency of light
– Higher frequency = More energy in
photons
– Higher frequency light = bigger packets
(photons)
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Atomic Spectra
• It was observed that atoms which have gained extra
energy release that energy in the form of light
• The light atoms give off or gain is of very specific
wavelengths called a line spectrum
– light given off = emission spectrum
– light energy gained = absorption spectrum
Hydrogen
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Other Gases….
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Niels Hendrik David
Bohr (1885-1962)
Nobel Prize in 1922
for work on atomic
structure
Modified
Rutherford’s
model to
include specific
locations of
electrons,
based upon
spectral lines
Source:
Emilio Segre
Visual Archives
Bohr’s Model
• Explained spectral lines of hydrogen
• Energy of atom is related to the distance
electron is from the nucleus, or energy level.
• Energy of the atom is quantized
– atom can only have certain specific energy states
called quantum levels or energy levels (7 max)
– when atom gains energy, electron “moves” to a
higher quantum level
– when atom loses energy, electron “moves” to a
lower energy level
– Specific amount of energy is released….a photon
– lines in spectrum correspond to the difference in
energy between levels, or the different photons
released.
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• Animation
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Figure 10.17: The Bohr model of the hydrogen atom represented the
electron as restricted to certain circular orbits around the nucleus.
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When salts
containing
Li+, Cu2+,
and Na+
dissolved in
methyl
alcohol are
set on fire,
brilliant
colors result.
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A neon sign
celebrating
Route 66
Source:
Owaki-Kulla/Corbis
Figure 10.12: Hydrogen atoms have several
excited-state energy levels.
Or rest position
for electron
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Figure 10.13:
Each photon
emitted by an
excited
hydrogen
atom
corresponds
to a
particular
energy
change in the
hydrogen
atom.
Or rest position
for electron
Bohr’s Model
•
•
•
•
1st energy level can hold 2e-1
2nd energy level can hold 8e-1
3rd energy level can hold 18e-1
4th energy level can hold 32 e-1
– farther from nucleus = more space
– The highest ground state orbit (outermost
electrons) is called the valence shell
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Orbitals and Energy Levels
• Principal energy levels identify how much energy
the electrons in the orbital have
– N (n=1, n=2, n=3, etc….)
• Each principal energy level contains one or more
sublevels (called s, p, d, and f sublevels)
–
–
–
–
there are n sublevels in each principal energy level
1st energy level has one sublevel (s)
2nd energy level has two sublevels (s and p)
3rd energy level has three sublevels (s, p, and d)
–
–
–
–
s sublevel have 1 orbital,
p sublevels have 3 orbitals,
d sublevels have 5 orbitals
f sublevels have 7 orbitals
• Each sublevel contains one or more orbitals
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Pauli Exclusion Principle
• No orbital may have more than 2 electrons
• Electrons in the same orbital must have
opposite spins
• s sublevel holds 2 electrons (one orbital)
• p sublevel holds 6 electrons (three orbitals)
• d sublevel holds 10 electrons (five orbitals)
• f sublevel holds 14 electrons(seven orbitals)
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Brief Instructions
An electron configuration is a method of indicating the
arrangement of electrons about a nucleus. A typical electron
configuration consists of numbers, letters, and superscripts
with the following format:
•1. A number indicates the energy level (The number
is called the principal quantum number.)
•2. A letter indicates the type of orbital; s, p, d, f.
•3. A superscript indicates the number of electrons in
the orbital. Example: ls2 means that there are
two electrons in the ‘s’ orbital of the first energy
level. The element is helium.
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To write an electron configuration:
1. Determine the total number of electrons to
be represented.
2. Use the Aufbau process to fill the orbitals
with electrons, that is, fill the lowest energy
orbitals first. Build your electron
configuration “from the ground upwards”.
3. The sum of the superscripts should equal
the total number of electrons.
Example: 12Mg
ls2 2s2 2p6 3s2
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Electron Configuration
• Elements in the same family of the
Periodic Table have
– Similar chemical and physical properties
– Similar valence shell electron
configurations
• Same numbers of valence electrons
• Same orbital types
• Different energy levels
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s1
1
2
3
4
5
6
7
s2
p 1 p 2 p 3 p 4 p 5 s2
p6
d1 d2 d3 d4 d5 d6 d7 d8 d9 d10
f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14
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Hund’s Rule
• for a set of degenerate orbitals, half fill each orbital
first before pairing
• highest energy level called the valence shell
– electrons in the valence shell called valence electrons
– electrons not in the valence shell are called core electrons
– often use symbol of previous noble gas to represent core
electrons
1s22s22p6 = [Ne]
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4.10 Ions
• OBJECTIVES:
–Identify the charges on monatomic
ions by using the periodic table,
and name the ions.
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Ions
• All neutral atoms have same
number of protons as electrons,
and are neutral in charge
• Atoms sometimes can either
gain or lose electrons, and in
the process develop an overall
electrical charge. When they
do this, they become…Ions!
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Atoms and ions
• Atoms are electrically neutral.
– Because there is the same
number of protons and
electrons.
• Ions are atoms, or groups of
atoms, with a charge (positive
or negative)
– They have different numbers of
protons and electrons.
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I am Not
an
atom!!!
An Anion is…
• A negative ion.
• Has gained electrons.
• Nonmetals can gain electrons.
• Charge is written as a superscript on the right.
1F
Has gained one electron (-ide
is new ending = fluoride)
2O
Gained two electrons (oxide)
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A Cation is…
 A positive
ion.
 Formed by losing electrons.
 More protons than electrons.
 Metals can lose electrons
1+
K
Has lost one electron (no
name change for positive ions)
2+
Ca
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Has lost two electrons
Predicting Ionic Charges
Group 1A: Lose 1 electron to form 1+ ions
H1+ Li1+
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Na1+
K1+ Rb1+
Predicting Ionic Charges
Group 2A: Loses 2 electrons to form 2+ ions
Be2+ Mg2+ Ca2+ Sr2+ Ba2+
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Predicting Ionic Charges
Group 3A: Loses 3
electrons to form
3+ ions
B3+
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Al3+
Ga3+
Predicting Ionic Charges
Neither! Group 4A
elements rarely form
ions.
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Group 4A: Lose 4
electrons or gain
4 electrons?
Predicting Ionic Charges
Notice that the end of the element
name is dropped and –ide is added.
This is only done for ions formed when
an atom gains electron(s)
N3-
Nitride
P3-
Phosphide
As3- Arsenide
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Group 5A: Gains 3
electrons to form
3- ions
Predicting Ionic Charges
Group 6A: Gains 2
electrons to form
2- ions
O2-
Oxide
S2-
Sulfide
Se2- Selenide
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Predicting Ionic Charges
Group 7A: Gains
1 electron to form
1- ions
F1- Fluoride
Br1- Bromide
Cl1- Chloride
I1- Iodide
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Predicting Ionic Charges
Group 8A: Stable
noble gases do not
form ions!
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Predicting Ionic Charges
Group B elements: Many transition elements
have more than one possible oxidation state.
Note the use of Roman
Iron (II) = Fe2+
numerals to show charges
Iron (III) = Fe3+
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Naming cations
• Two methods can clarify when more
than one charge is possible:
1) Stock system – uses roman
numerals in parenthesis to indicate
the numerical value
2) Classical method – uses root word
with suffixes (-ous, -ic)
• Does not give true value
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Naming cations
• We will use the Stock system.
• Cation - if the charge is always the same
(like in the Group A metals) just write the
name of the metal.
• Transition metals can have more than
one type of charge.
• Indicate their charge with roman
numerals in parenthesis after the name of
the metal (Table 9.2, p.255)
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Predicting Ionic Charges
Some of the post-transition elements also
have more than one possible oxidation state.
Tin (II) = Sn2+
Lead (II) = Pb2+
Tin (IV) = Sn4+
Lead (IV) = Pb 4+
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Predicting Ionic Charges
Group B elements: Some transition elements
have only one possible oxidation state, such
as these three that are always:
Silver = Ag1+
Zinc = Zn2+
Cadmium = Cd2+
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Exceptions:
• Some of the transition metals have
only one ionic charge:
–Do not use roman numerals for
these:
–Silver is always 1+ (Ag1+)
–Cadmium and Zinc are always 2+
(Cd2+ and Zn2+)
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Practice by naming these:
•
•
•
•
•
•
•
Na1+
Ca2+
Al3+
Fe3+
Fe2+
Pb2+
Li1+
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Write Formulas for these:
• Potassium ion
• Magnesium ion
• Copper (II) ion
• Chromium (VI) ion
• Barium ion
• Mercury (II) ion
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Naming Anions
• Anions are always the same
charge
• Change the monatomic
element ending to – ide
1• F a Fluorine atom becomes
a Fluoride ion.
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Practice by naming these:
• Cl13•N
1• Br
• O23+
• Ga
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Write symbols for these:
• Sulfide ion
• Iodide ion
• Phosphide ion
• Strontium ion
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Polyatomic ions are…
• Groups of atoms that stay together and have
an overall charge, and one name.
• Usually end in –ate or -ite
• Acetate: C2H3O21• Nitrate: NO31• Nitrite: NO21• Permanganate: MnO41• Hydroxide: OH1- and Cyanide: CN1-?
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Note Table 9.3 on page 257
2-
• Sulfate: SO4
• Sulfite: SO32-
• Carbonate: CO32-
• Chromate: CrO42• Dichromate: Cr2O72-
• Phosphate: PO43• Phosphite: PO33• Ammonium: NH41+
(One of the few positive
polyatomic ions)
If the polyatomic ion begins with H, then combine the
word hydrogen with the other polyatomic ion present:
H1+ + CO32- →
HCO31hydrogen + carbonate → hydrogen carbonate ion
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Ions
• ions that have a positive charge are
called cations
–form when an atom loses electrons
• ions that have a negative charge are
called anions
–form when an atom gains electrons
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Ions
• ions with opposite charges attract
–therefore cations and anions
attract each other
• moving ions conduct electricity
• compound must have no total
charge, therefore we must balance
the numbers of cations and anions
in a compound to get 0 total charge
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Atomic Structures of Ions
• Metals form cations
• For each positive charge the ion has 1
less electron than the neutral atom
–Na = 11 e-, Na+ = 10 e–Ca = 20 e-, Ca+2 = 18 e-
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• Cations are named the same as the metal
sodium Na  Na+ + 1esodium ion
calcium Ca  Ca+2 + 2ecalcium ion
• The charge on a cation can be
determined from the Group number on
the Periodic Table for Groups IA, IIA,
IIIA
– Group 1A  +1, Group 2A  +2, (Al,
Ga, In)  +3
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Figure 4.16: (a) Sodium chloride (common table salt) can be decomposed
to the elements (b) sodium metal (on the left) and chlorine gas.
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Atomic Structures of Ions
• Nonmetals form anions
• For each negative charge the ion has 1 more electron
than the neutral atom
– F = 9 e-, F- = 10 e– P = 15 e-, P3- = 18 e-
• Anions are named by changing the ending of the name
to -ide
fluorine
F + 1e-  F- fluoride ion
oxygen
O + 2e-  O2oxide ion
• The charge on an anion can be determined from the
Group number on the Periodic Table
– Group 7A  -1, Group 6A  -2
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Figure 4.21:
(a) The
arrangement
of sodium
ions (Na+)
and chloride
ions (Cl-) in
the ionic
compound
sodium
chloride.
Figure 4.21: (b) Solid sodium
chloride highly magnified.
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Figure 4.19: The ions formed by selected
members of Groups 1, 2, 3, 6, and 7.
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Electrical Nature of Matter
• Most common pure substances are
very poor conductors of electricity
–with the exception of metals and
graphite
–Water is a very poor electrical
conductor
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Electrical Nature of Matter
• Some substances dissolve in water to
form a solution that conducts well these are called electrolytes
• When dissolved in water, electrolyte
compounds break up into component
ions
– ions are atoms or groups of atoms
that have an electrical charge
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Figure 4.20: (a) Pure water does not conduct a current. (b) Water
containing dissolved salt conducts electricity and the bulb lights.
Source: Tara Piasio/IFAS/University of Florida.
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Lenna Ma and Pteris
vittata–called the brake fern.
Source: Tara Piasio/IFAS/University of Florida
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Ancient Anasazi Indian cliff dwellings.
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