Transcript Periodic Table 2015

IB Topic 3: Periodicity
3.1: The periodic table
• 3.1.1 Describe the arrangement of elements in the
periodic table in order of increasing atomic number.
• 3.1.2 Distinguish between the terms group and period.
• 3.1.3 Apply the relationship between the electron
arrangement of elements and their position in the
periodic table up to Z = 20.
• 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.
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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
Design of the Table
• Groups are the vertical columns.
– elements have similar, but not identical,
properties
• most important property is that
they have the same # of valence
electrons
• valence electrons- electrons in the
highest occupied energy level
• all elements have 1,2,3,4,5,6,7, or 8
valence electrons
IB prefers this one.
3.1.1 Describe the arrangement of elements in the
periodic table in order of increasing atomic number.
3.1.2 Distinguish between the terms group and period.
Development of the Periodic Table
• Henry Mosley
Arranged the elements by increasing atomic
number instead of mass (1913)
• Glen Seaborg
Discovered the transuranium elements (93102) and added the actinide and lanthanide
series (1945)
Elements arranged by increasing
atomic number into
• periods (rows) and
• groups or families (columns),
which share similar characteristics
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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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• 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?
Classification of the Elements
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19
4f
5f
20
ns2np6
ns2np5
ns2np4
ns2np3
ns2np2
ns2np1
d10
d5
d1
ns2
ns1
Ground State Electron Configurations of the Elements
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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Lewis Dot-Diagrams/Structures
• valence electrons are represented as dots
around the chemical symbol for the element
Na
Cl
2
1
3
2
5
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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.
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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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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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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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 = 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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IB Topic 3: Periodicity
3.2: Physical properties
• 3.2.1 Define the terms first ionization energy and
electronegativity.
• 3.2.2 Describe and explain the trends in atomic radii,
ionic radii, first ionization energies, electronegativities
and melting points for the alkali metals (Li  Cs) and
the halogens (F  I).
• 3.2.3 Describe and explain the trends in atomic radii,
ionic radii, first ionization energies and
electronegativities for elements across period 3.
• 3.2.4 Compare the relative electronegativity values of
two or more elements based on their positions in the
periodic table.
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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?
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
• The radius of an atom, measured in pm (picometers)
• Periodic trend (Period 3 Trend)
– Atomic size decreases as you move across a period.
– The increase in nuclear charge increases the attraction to the
outer shell so the outer energy level progressively becomes
closer to the nucleus
• Group trend for Alkali metals & Halogens
– Atomic size increases as you move down a group of the
periodic table.
– Adding higher energy levels
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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 the ion form of atoms
(cations and anions)
•
Positive ions are smaller than their atoms.
•
Negative ions are larger than their atoms
•
Periodic trend (Period 3 Trend)
•
Group trend for Alkali metals & Halogens
– Ions get larger down a group
– More energy levels are added
–
–
–
–
–
–
Fewer electrons so nucleus attracts remaining electrons more strongly
One fewer energy level since valence electrons removed.
More electrons so nucleus has less attraction for them
Greater electron-electron repulsion
Decrease as you move across a period, then spike and decrease again
This 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
First Ionization Energies
The energy required to remove the first electron from a gaseous atom.
Second ionization removes the second electron and so on.
Can be used to predict ionic charges.
•
Periodic Trend (Period 3 Trend)
•
Group trend for Alkali metals & Halogens
– Generally decreases as you move down a group in the periodic table
– Since size increases down a group, the outermost electron is farther away from
the nucleus and is easier to remove.
– Increases as you move from left to right across a period.
– Effect of increasing nuclear charge makes it harder to remove an electron.
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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
Tendency for the atoms of the element to attract electrons
when they
are chemically combined with atoms of another element.
Helps predict the type of bonding (ionic/covalent).
• Periodic Trend (Period 3 Trend)
– Increases as you move from left to right across a period.
– Nonmetals have a greater attraction for electrons than metals & there is
a greater nuclear charge that can attract electrons
•
Group trend for Alkali metals & Halogens
– Generally decreases as you move down a group in the periodic table.
– For metals, the lower the number the more reactive.
– For nonmetals, the higher the number the more reactive.
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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
Reactivity
The relative capacity of an atom, molecule or radical to
undergo a chemical reaction with another atom, molecule or
radical.
• Don’t worry about the periodic trend!!!
•
Group trend for Alkali metals
– Increases as you move down group 1 in the periodic table
– Since alkali metals are more likely to lose an electron, the ones with the
lowest 1st ionization energy are the most reactive since they require the
least amount of energy to lose a valence electron.
•
Group trend for Halogens
– Decreases as you move down group 7 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.
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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
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
• Don’t worry about the periodic trend!!!
• 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)
• Halogens: Melting point increases down the group
– F2 (-220 0C) to I2 (114 oC)
– Weak attractive forces increase as the molecules get larger due to the
type of bonding (non-polar covalent)
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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.
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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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