Transcript Chapter 5 Atomic Structure

Topic 3: Periodicity
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3.1: The periodic table
Understandings
• 3.1.1 The periodic table is arranged into four blocks
associated with the four sub-levels.
• 3.1.2 The periodic table consists of groups (vertical
columns) and period (horizontal rows).
• 3.1.3 The period number (n) is the outer energy level
that is occupied.
• 3.1.4 The number of the principal energy level and the
number of the valence electrons in an atom can be
deduced from its position on the periodic table.
• 3.1.4 The periodic table shows the positions of metals,
non-metals and metalloids.
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3.1.1 Describe the arrangement of elements in the
periodic table in order of increasing atomic number.
Development of the Periodic Table
• Johan Dobereiner
Grouped similar elements into groups of 3
(triads) such as chlorine, bromine, and
iodine. (1817-1829).
• John Newlands
Found every eighth element (arranged by
atomic weight) showed similar properties.
Law of Octaves (1863).
• Dmitri Mendeleev
Arranged elements by similar properties
but left blanks for undiscovered elements
(1869).
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Dmitri Mendeleev
1834 – 1907
• Russian chemist and
teacher
• Given the elements he
knew about, he organized
a “Periodic Table” based
on increasing atomic
mass (it’s now atomic #)
• He even left empty
spaces to be filled in later
At the time the elements gallium and germanium were not
known. These are the blank spaces in his periodic table. He predicted their
discovery and estimated their properties.
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Henry Moseley
1887 – 1915
• Arranged the elements
in increasing atomic
numbers (Z)
– Properties now
recurred periodically
IB prefers this one.
Classification of the Elements
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Design of the Table
• A Group (aka family) is a vertical column
– Elements have similar, but not identical,
properties
• Most important property is that
they have the same # of valence
electrons
Electron Arrangement
Core Electrons:
electrons that are in the
inner energy levels
Valence Electrons:
electrons that are in the
outermost (highest)
energy level
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3.1.4 Apply the relationship between the number of
electrons in the highest occupied energy level for an
element and its position in the periodic table.
Arrangement of the Periodic
Table
• Valence Electrons: electrons in
the outermost (highest)
energy level
–
–
–
–
–
Group 1 elements have 1 v.e.s
Group 2 elements have 2 v.e.s
Group 3 elements have 3 v.e.s
So on and so forth
Group 8 have 8 v.e. (except for
helium, which has 2)
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• Valence electrons: electrons in the
highest occupied energy level
• all elements have 1,2,3,4,5,6,7, or 8
valence electrons
Lewis Dot-Diagrams/Structures
• valence electrons are represented as dots
around the chemical symbol for the element
Na
Cl
2
1
3
2
5
8
3.1.4 Apply the relationship between the number of
electrons in the highest occupied energy level for an
element and its position in the periodic table.
Electron dot diagrams
Group 1A: 1 dot
X
Group 5A: 5 dots
X
Group 2A: 2 dots
X
Group 6A: 6 dots
X
Group 3A: 3 dots
X
Group 7A: 7 dots
X
Group 4A: 4 dots
X
Group 0: 8 dots (except He)
X
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Look, they are
following my
rule!
Electron Dot Diagram
Using the symbol for the element, place dots around the symbol corresponding
to the outer energy level s & p electrons (valence electrons). Will have from
one to eight dots in the dot diagram.
Draw electron dot diagrams for the following atoms
H
Be
H
Be
O
Al
Ca
Zr
O
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Electron Dot Diagram
Using the symbol for the element, place dots around the symbol corresponding
to the outer energy level s & p electrons. Will have from one to eight dots in
the dot diagram.
Draw electron dot diagrams for the following atoms
Al
Ca
Zr
Al
Ca
Zr
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4f
5f
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ns2np6
ns2np5
ns2np4
ns2np3
ns2np2
ns2np1
d10
d5
d1
ns2
ns1
Ground State Electron Configurations of the Elements
2.3.4 Deduce the electron arrangement for
atoms and ions.
Write electron configuration, orbital filling
diagrams, and electron dot diagrams.
Kr
Tb
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• B is 1s2 2s2 2p1;
– 2 is the outermost energy level
– it contains 3 valence electrons, 2
in the 2s and 1 in the 2p
• Br is [Ar] 4s2 3d10 4p5
How many valence electrons
are present?
3.1.3 Apply the relationship between the electron
arrangement of elements and their position in the periodic
table.
Arrangement of the Periodic
Table
• Group = Sum of electrons in
the highest occupied energy
level (s + p) = Number of
valence electrons
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2.3.4 Deduce the electron arrangement for atoms and
ions.
Valence electrons are electrons in the
outermost energy level of an atom
– The sum of electrons in the s & p orbitals in the
highest energy level
– Ex. Argon’s electron arrangement is
1s22s22p63s23p6. Since the highest energy level is 3,
we add the e-s in 3s2 + 3p6 = 8
– So, argon has 8 valence electrons
• The easy way is to look at its location on the periodic
table  (except for the transition metals)
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3.1.3 Apply the relationship between the electron
arrangement of elements and their position in the periodic
table.
Arrangement of the Periodic
Table
• Na = 1s22s22p63s1
• Since the sum of electrons in the
highest occupied energy level is 1, it
will be in the 1st group and have 1
valence electron
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3.1.3 Apply the relationship between the electron
arrangement of elements and their position in the periodic
table.
Arrangement of the Periodic
Table
• Na = 1s22s22p63s1
• It is in the 1st group because it
has 1 valence electron
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• B is 1s2 2s2 2p1;
– 2 is the outermost energy level
– it contains 3 valence electrons, 2
in the 2s and 1 in the 2p
• Br is [Ar] 4s2 3d10 4p5
How many valence electrons
are present?
• Periods are the horizontal rows
– do NOT have similar properties
– however, there is a pattern to their
properties as you move across the table
that is visible when they react with other
elements
3.1.3 Apply the relationship between the electron
arrangement of elements and their position in the periodic
table.
Arrangement of the Periodic
Table
• Period = The highest occupied
energy level (s and p) =
number of energy levels
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3.1.3 Apply the relationship between the electron
arrangement of elements and their position in the periodic
table.
Arrangement of the Periodic
Table
• Na = 1s22s22p63s1
• Sodium is in the 3rd period because
it has 3 energy levels  The
highest occupied energy level is 3
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3.1.3 Apply the relationship between the electron
arrangement of elements and their position in the periodic
table.
Arrangement of the Periodic
Table
• Write out the electron configuration
for selenium
• State the relationship between the
group that selenium is in and its
electron configuration
• State the relationship between the
period that selenium is in and its
electron configuration
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3.1.3 Apply the relationship between the electron
arrangement of elements and their position in the periodic
table.
Arrangement of the Periodic Table
Electron configuration for selenium:
Se = 1s22s22p63s23p64s23d104p4
• Group # is 16
• Since the highest energy level is 4, we add the e-s in
4s2 + 4p4 = 6
• Therefore, Se has 6 valence e-s
• Period # is 4
• The highest occupied s/p energy level is 4
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3.1.1 Describe the arrangement of elements in the
periodic table in order of increasing atomic number.
Metals
• Left side of the periodic table (except
hydrogen)
• Good conductors of heat and electricity
• Malleable: capable of being hammered
into thin sheets
• Ductile: capable of being drawn into
wires
• Have luster: are shiny
• Typically lose electrons in chemical
reactions
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3.1.1 Describe the arrangement of elements in the
periodic table in order of increasing atomic number.
Metals
• Alkali metals: Group 1 (1A)
• Alkaline earth metals: Group 2 (2A)
• Transition metals: Group B, lanthanide
& actinide series
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3.1.1 Describe the arrangement of elements in the
periodic table in order of increasing atomic number.
Nonmetals
•
•
•
•
Right side of the periodic table
Poor conductors of heat and electricity
Non-lusterous
Typically gain electrons in chemical
reactions
• Halogens: Group 17 (7A)
• Noble gases: Group 18 (0)
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3.1.1 Describe the arrangement of elements in the
periodic table in order of increasing atomic number.
Metalloids
• Between metals and non-metals, along
the stair step (except aluminum)
• Have properties of metals and nonmetals
• Some are semi-conductors
• Boron (B), Silicon (Si), Germanium
(Ge), Arsenic (As), Antimony (Sb),
Tellurium (Te), Astatine (At)
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Green = Metals
Blue = Metalloids
Yellow = Nonmetals
http://www.windows2universe.org/earth/geology/metals.html
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3.2: Periodic trends
Understandings
• 3.2.1 Vertical and horizontal trends in the
periodic table exist for atomic radius, ionic
radius, ionization energy, electron affinity and
electronegativity.
• 3.2.2 Trends in metallic and non-metallic
behavior are due to the trends above.
• 3.2.3 Oxides change from basic through
amphoteric to acidic across a period.
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Periodic Trend Definitions
• Atomic Radius: half the internuclear distance
between two atoms of the same element (pm)
• Ionic radius: the radius of an ion in the
crystalline form of a compound (pm)
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Periodic Trend Definitions
• First ionization energy: The energy required to
remove one electron from each atom in one
mole of gaseous atoms under standard
thermodynamic conditions (kJ mol-1)
• Electron Affinity: The energy released when
one electron is added to each atom in one mole
of gaseous atoms under standard
thermodynamic conditions (kJ mol-1)
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Periodic Trend Definitions
• Electronegativity: A measure of the tendency
of an atom in a molecule to attract a pair of
shared electrons towards itself
• Melting Point: The temperature at which a solid
becomes a liquid at a fixed pressure (degrees
Kelvin)
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Trends in the table
IB loves the alkali metals and
the halogens
Many trends are easier to understand if
you comprehend the following:
• The ability of an atom to “hang on to” or
attract its valence electrons is the result
of two opposing forces
– the attraction between the electron and the
nucleus
– the repulsions between the electron in
question and all the other electrons in the
atom (often referred to the shielding effect)
– the net resulting force of these two is
referred to effective nuclear charge
This is a simple, yet very good picture. Do you understand it?
Therefore…
• Periodic trends typically have to do with
an increase in nuclear charge
• Group trends typically have to do with
an increase in shielding effect (more
energy levels)
3.2.2 Describe and explain the trends in atomic
radii, ionic radii, first ionization energies,
electronegativities and melting points
Group 1A: Alkali Metals
•
•
•
•
•
Have 1 valence electron
Shiny, silvery, soft metals
React with water & halogens
Oxidize easily (lose electrons)
Reactivity increases down the group
Group 7A: Halogens
•
•
•
•
Have 7 valence electrons
Colored gas (F2, Cl2); liquid (Br2);
Solid (I2)
Oxidizer (gain electrons)
Reactivity decreases down the group
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3.2.2 Describe and explain the trends in atomic
radii, ionic radii, first ionization energies,
electronegativities and melting points
Atomic Radii
• Periodic trend (Period 3 Trend)
– Atomic radius decreases as you move across a
period.
– Number of protons in the nucleus increases
– Increase in nuclear charge increases the attraction to
the outer shell so the outer energy level progressively
becomes closer to the nucleus
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3.2.2 Describe and explain the trends in atomic
radii, ionic radii, first ionization energies,
electronegativities and melting points
Atomic Radii
• Group trend for Alkali metals &
Halogens
– Atomic radius increases as
you move down a group of the
periodic table.
– More energy levels are added
– More shielding
H
Li
Na
K
Rb
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Atomic Radii
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3.2.2 Describe and explain the trends in atomic
radii, ionic radii, first ionization energies,
electronegativities and melting points
Ionic Radii
The radius of an ion in the crystalline form of a
compound
• Atoms tend to gain or lose electrons in order to
have the electron configuration of a noble gas
• Most want 8 valence electrons and take the easiest
approach to obtaining that full “octet”
• Hydrogen and helium only want 2
• There are some other weird exceptions that we’re not
worried about yet
• Let’s assign charges!
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Common Ion Charges
Positive Ions (cations)
Group 1:
•
•
Lose 1 valence electron
Charge of +1: Li+, Na+, K+
Group 2
•
•
Lose 2 valence electrons
Charge of +2: Mg2+, Ca2+
Group 3
•
•
Lose 3 valence electrons
Charge of +3: Al3+
Negative Ions (anions)
Group 5:
•
•
Gain 3 electrons
Charge of -3: N3-, P3-
Group 6:
•
•
Gain 2 electrons
Charge of -2: O2-, S2-
• Group 7
•
•
Gain 1 electron
Charge of -1: F-, Cl-, Br-, I-
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Uncommon Ion Charges
In most transition elements, d electrons can
become involved in the reaction
Iron can lose 2 electrons (the 2 in the 4s) (Fe2+) or
3 electrons (the 2 in the 4s and 1 in the 3d) (Fe3+)
•
•
The name of the Fe2+ ion is iron(II) or ferrous
The name of the Fe3+ ion is iron(III) or ferric
Chromium can lose 2 electrons (the 2 in the 4s) (Cr2+) or
3 electrons (the 2 in the 4s and 1 in the 3d) (Cr3+)
•
•
The name of the Cr2+ ion is chromium(II) or chromous
The name of the Cr3+ ion is chromium(III) or chromic
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Uncommon Ion Charges
• These transition metals only form ONE ion:
Ag+1, Zn+2 and Cd+2
Label these on your periodic table!!
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3.2.2 Describe and explain the trends in atomic
radii, ionic radii, first ionization energies,
electronegativities and melting points
Ionic Radii
• Look at the ions compared to their parent atoms
• Do atoms become smaller or larger when they do
this?
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Cations (+ ions) are smaller than the parent
atom
• Have lost an electron and lost an entire
energy level!
• Therefore, have fewer electrons than protons
+
Li
0.152 nm
Li forming a
cation
Li+
.078nm
Anions (– ions) are larger than parent atom
• Have gained an electron to achieve noble gas
configuration
• Effective nuclear charge has decreased since
same nucleus now holding on to more
electrons
• Plus, the added electron repels the existing
electrons farther apart (kind of “puffs it out”)
F 0.064 nm
9e- and 9p+
F- 0.133 nm
10 e- and 9 p+
3.2.2 Describe and explain the trends in atomic
radii, ionic radii, first ionization energies,
electronegativities and melting points
Ionic Radii
• Periodic trend (Period 3 Trend)
– Decreases at first when losing electrons (+ ions)
– Suddenly increase when gaining electrons (– ions)
– Decreases again due to increased nuclear charge
• Group trend for Alkali metals & Halogens (same as
neutral atoms)
– Increase down a group
– More energy levels are added
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3.2.2 Describe and explain the trends in atomic
radii, ionic radii, first ionization energies,
electronegativities and melting points
First Ionization Energy
The energy required to remove one electron from each atom in one
mole of gaseous atoms under standard thermodynamic conditions
(kJ mol-1)
X(g)  X+(g) + eSecond ionization removes the second electron and so on.
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3.2.2 Describe and explain the trends in atomic
radii, ionic radii, first ionization energies,
electronegativities and melting points
First Ionization Energy
• Periodic Trend (Period 3 Trend)
– Increases as you move from left to right across a period.
– Number of protons in the nucleus increases.
– Effect of increasing nuclear charge makes it harder to remove an
electron.
• Group trend for Alkali metals & Halogens
– Decreases as you move down a group in the periodic table.
– Number of energy levels increases.
– Outer electrons are farther away from the nucleus and is easier
to remove.
– Inner core electrons “shield” the valence electrons from the pull
of the positive nucleus and therefore easier to remove.
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Filled n=1 shell
Filled n=2 shell
Filled n=3 shell
Filled n=4 shell
Filled n=5 shell
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3.2.2 Describe and explain the trends in atomic
radii, ionic radii, first ionization energies,
electronegativities and melting points
Electronegativity
A measure of the tendency of an atom in a molecule to
attract a pair of shared electrons towards itself.
Helps predict the type of bonding (ionic/covalent).
• Linus Pauling (1901 to 1994) came up with a
scale where a value of 4.0 is arbitrarily
given to the most electronegative element,
fluorine, and the other electronegativities
are scaled relative to this value.
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3.2.2 Describe and explain the trends in atomic
radii, ionic radii, first ionization energies,
electronegativities and melting points
Electronegativity
• Periodic Trend (Period 3 Trend)
– Increases as you move from left to right across a
period.
– Number of protons in the nucleus increases.
– Increasing nuclear charge makes it more likely to
want an electron.
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3.2.2 Describe and explain the trends in atomic
radii, ionic radii, first ionization energies,
electronegativities and melting points
Electronegativity
• Group trend for Alkali metals & Halogens
– Generally decreases as you move down a group in
the periodic table.
– Number of energy levels increases.
– Outer electrons are farther away from the nucleus
and aren’t as attracted to one another.
– Inner core electrons “shield” the valence electrons
from the pull of the positive nucleus and therefore
less attracted.
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Electronegativity
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3.2.2 Describe and explain the trends in atomic
radii, ionic radii, first ionization energies,
electronegativities and melting points
Electron Affinity
The energy released when one electron is added to each
atom in one mole of gaseous atoms under standard
thermodynamic conditions (kJ mol-1)
• In other words, the neutral atom’s likelihood of gaining
an electron.
• Example
• F (g) + e-  F-(g) will release 328 kJ mol-1 of energy
• The more negative the value, the greater the attraction
for the electron, the more affinity the atom has
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3.2.2 Describe and explain the trends in atomic
radii, ionic radii, first ionization energies,
electronegativities and melting points
Electron Affinity
• Periodic Trend (Period 3 Trend)
– Values decrease (become more negative) as you move from left
to right across a period….
– Energy released increases…
– Meaning affinity for electrons INCREASES
– Number of protons in the nucleus increases  increasing
nuclear charge makes it more likely to add an electron.
• Group trend for Alkali metals & Halogens
– Generally increase (become less negative) as you move down a
group in the periodic table…
– Meaning affinity for electrons DECREASES
– Number of energy levels increases  outer electrons are farther
away from the nucleus, adding to the shielding effect.
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3.2.2 Describe and explain the trends in atomic
radii, ionic radii, first ionization energies,
electronegativities and melting points
Electron Affinity
• Periodic Trend (Period 3 Trend)
– Values decrease (become more negative) as you move from left
to right across a period….
– Energy released increases…
– Meaning affinity for electrons INCREASES
– Number of protons in the nucleus increases  increasing
nuclear charge makes it more likely to add an electron.
• Group trend for Alkali metals & Halogens
– Generally increase (become less negative) as you move down a
group in the periodic table…
– Meaning affinity for electrons DECREASES
– Number of energy levels increases  outer electrons are farther
away from the nucleus, adding to the shielding effect.
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This one same as “IB textbook”
3.2.2 Describe and explain the trends in atomic
radii, ionic radii, first ionization energies,
electronegativities and melting points
Melting Points
The temperature at which a solid becomes a liquid at a
fixed pressure (degrees Kelvin)
The temperature at which a crystalline melts depends on
the strength of the attractive forces and on the way the
particles are packed in the solid state
• Requires understanding of concepts covered in later
topics (this year and next year)
• Know the type of bonding
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3.2.2 Describe and explain the trends in atomic
radii, ionic radii, first ionization energies,
electronegativities and melting points
Melting Points
• Don’t worry about the periodic trend!!!
• You will need to know the group trends for
Alkali metals and halogens… they’re
Element Melting
different!
Point (K)
• Alkali Metals: Melting point decreases
down the group
– Li (181 oC) to Cs (29 oC)
– As the atoms get larger the forces of attraction
between them decrease due to the type of
bonding (metallic)
– The “sea of electrons” is further away from the
metal ions
Li
453
Na
370
K
336
Rb
312
Cs
301
Fr
295
3.2.2 Describe and explain the trends in atomic
radii, ionic radii, first ionization energies,
electronegativities and melting points
Melting Points
• Halogens: Melting point
increases down the group
– F2 (-220 0C) to I2 (114 oC)
– Halogens molecules are held
together with weak van der
Waals’ attractive forces due to
their non-polar covalent nature
– Larger molecules have more
electrons, increasing the strength
of the IMF
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increases
increases
3.2.2 Describe and explain the trends in atomic
radii, ionic radii, first ionization energies,
electronegativities and melting points
Reactivity
The relative capacity of an atom to undergo a chemical
reaction with another atom, molecule or radical.
• Why do atoms react?
• Atoms that are good at becoming stable (getting a full
valence shell) are the most reactive
• Don’t worry about the periodic trend because elements on
opposite sides of the periodic table can be equally reactive…
but for different reasons!!!
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3.2.2 Describe and explain the trends in atomic
radii, ionic radii, first ionization energies,
electronegativities and melting points
Reactivity of Alkali metals
The relative capacity of an atom to undergo a chemical
reaction with another atom, molecule or radical.
• How many valence electrons do the alkali metals
have?
• Are alkali metals more likely to give up or get
electrons?
• What is the name of the property responsible for
doing that?
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3.2.2 Describe and explain the trends in atomic
radii, ionic radii, first ionization energies,
electronegativities and melting points
Reactivity of Alkali metals
• It all has to do with ionization energy since they
want to lose electrons!
• Lower ionization energy = more reactive
• Group trend for Alkali metals
– Increases as you move down group 1
– Since alkali metals are more likely to lose an electron, the
ones with the lowest ionization energy are the most
reactive since they require the least amount of energy to
lose a valence electron.
• Which alkali metal is the best at losing electrons?
• That’s the most reactive alkali metal!
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3.2.2 Describe and explain the trends in atomic
radii, ionic radii, first ionization energies,
electronegativities and melting points
Reactivity of halogens
The relative capacity of an atom to undergo a chemical
reaction with another atom, molecule or radical.
• How many valence electrons do the halogens have?
• Are halogens more likely to give up or get
electrons?
• What is the name of the property responsible for
doing that?
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3.2.2 Describe and explain the trends in atomic
radii, ionic radii, first ionization energies,
electronegativities and melting points
Reactivity of halogens
• It all has to do with electronegativity (or electron affinity)
since they want to gain electrons!
• Higher electronegativity = lower electron affinity = more
reactive
• Group trend for Alkali metals
– Decreases as you move down group 17 in the periodic table
– Since halogens are more likely to gain an electron, the ones
with the greatest electronegativity are the most reactive
since they are most effective at gaining a valence electron.
• Which alkali metal is best at getting electrons?
• That’s the most reactive halogen!
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least reactive
most reactive
IB Topic 3: Periodicity
3.3: Chemical properties
• Discuss the similarities and differences in the chemical
properties of elements in the same group.
• Discuss the changes in nature, from ionic to covalent
and from basic to acidic, of the oxides across period 3.
98
3.3.1 Discuss the similarities and
differences in the chemical properties of
elements in the same group.
Alkali Metals
React with water &
react with many
substances
because…
They have the same
number of valence
electrons
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3.3.1 Discuss the similarities and
differences in the chemical properties of
elements in the same group.
Alkali Metals
2Na(s) + 2H2O(l) 
2NaOH (aq) + H2(g)
In the reaction of alkali
metals and water, all
will:
• move around the
surface of the water,
• give off hydrogen gas,
• create a basic solution.
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3.3.1 Discuss the similarities and
differences in the chemical properties of
elements in the same group.
Alkali Metals
In the reaction of alkali metals and
water, the reactivity will increase
down the group because they
get better at getting rid of their
valence electron
(the 1st ionization energy
decreases)
So, alkali metals lower down will:
• React more vigorously
• React faster
• Give off a flame
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3.3.1 Discuss the similarities and
differences in the chemical properties of
elements in the same group.
Alkali Metals
Reaction with halogens
2M(s) + X2 (g)  2MX(s)
where M represents Li,Na,K,Rb, or Cs
Where X represents F,Cl,Br, or I
2Na(s) + Cl2(g)  2NaCl(s)
Reactivity decreases down the
group
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3.3.1 Discuss the similarities and
differences in the chemical properties of
elements in the same group.
Halogens
Halogens are diatomic as gases (two atoms
bond together) and called halides when
they form ions… These are BrINClHOF 
Halogens want to get one electron to fill its
outer shell.
Reactivity decreases down the group
because electronegativity decreases
Cl2 reacts with Br- and ICl2(aq) + 2Br-(aq)  2Cl-(aq) + Br2(l)
Cl2(aq) + 2I-(aq)  2Cl-(aq) + I2(s)
Br2 reacts with I-
Br2(aq) + 2I-(aq)  2Br-(aq) + I2(s)
I2 non-reactive with halide ions
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Reactivity of Elements… in action
Alkali Metals:
http://www.youtube.com/watch?v=m55kgy
ApYrY
Halogens:
http://www.youtube.com/watch?v=tk5xwS5b
ZMA&feature=related
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3.3.2 Discuss the changes in nature, from
ionic to covalent and from basic to acidic,
of the oxides across period 3.
Metallic Oxides in Period 3
Sodium oxide: Na2O
Magnesium oxide: MgO
Aluminum oxide: Al2O3
ionic
ionic
ionic
Metalloid oxide in Period 3
Silicon dioxide: SiO2
covalent
Nonmetallic oxides in Period 3
Tetraphosphorus decoxide: P4O10
Sulfur trioxide: SO3
Dichlorine heptoxide: Cl2O7
covalent
covalent
covalent
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3.3.2 Discuss the changes in nature, from
ionic to covalent and from basic to acidic,
of the oxides across period 3.
Acidic/Basic
Metallic oxides in Period 3 are basic
Sodium oxide:
Na2O + H2O  2 NaOH
Magnesium oxide:
MgO + H2O  Mg(OH)2
Aluminum oxide:
Al2O3 + H2O  2 Al(OH)3
basic
basic
amphoteric
Metalloid oxide in Period 3 is acidic
Silicon dioxide:
SiO2 + H2O  H2SiO3
acidic
Nonmetallic oxides in Period 3 are acidic
Tetraphosphorus decoxide: P4O10 + 6H2O  4H3PO4
Sulfur trioxide:
SO3 + H2O  H2SO4
Dichlorine heptoxide: Cl2O7 + H2O  2HClO4
Argon does not form an oxide
acidic
acidic
acidic
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Terms to Know
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Group
Period
Alkali metals
Halogens
Ionic radius
Electronegativity
First ionization energy
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Periodic Table of Video
• http://www.periodicvideos.com/
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