Download Periodic Relationships Among the Elements

Periodic Relationships Among
the Elements
Chapter 8
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When the Elements Were Discovered
8.1
ns2np6
ns2np5
ns2np4
ns2np3
ns2np2
ns2np1
d10
d5
d1
ns2
ns1
Ground State Electron Configurations of the Elements
4f
5f
8.2
Classification of the Elements
8.2
Electron Configurations of Cations and Anions
Of Representative Elements
Na [Ne]3s1
Na+ [Ne]
Ca [Ar]4s2
Ca2+ [Ar]
Al [Ne]3s23p1
Al3+ [Ne]
Atoms gain electrons
so that anion has a
noble-gas outer
electron configuration.
Atoms lose electrons so that
cation has a noble-gas outer
electron configuration.
H 1s1
H- 1s2 or [He]
F 1s22s22p5
F- 1s22s22p6 or [Ne]
O 1s22s22p4
O2- 1s22s22p6 or [Ne]
N 1s22s22p3
N3- 1s22s22p6 or [Ne]
8.2
-1
-2
-3
+3
+2
+1
Cations and Anions Of Representative Elements
8.2
Effective nuclear charge (Zeff) is the “positive charge” felt
by an electron.
Zeff = Z - s
0 < s < Z (s = shielding constant)
Zeff  Z – number of inner or core electrons
Z
Core
Na
11
10
1
186
Mg
12
10
2
160
Al
13
10
3
143
Si
14
10
4
132
→Zeff
Radius (pm)
8.3
Periodic Properties
• Elements show gradual changes in certain
physical properties as one moves across a
period or down a group in the periodic table.
These properties repeat after certain
intervals. In other words they are PERIODIC
Periodic properties
include:
-- Ionization Energy
-- Electronegativity
-- Electron Affinity
-- Atomic Radius
-- Ionic Radius
.8
Atomic Size
• The electron cloud doesn’t have a definite
edge.
• They get around this by measuring more
than 1 atom at a time.
• Summary: it is the volume that an atom
takes up
• http://www.mhhe.com/physsci/chemistry/e
ssentialchemistry/flash/atomic4.swf
8.3
Atomic
Radius
• The radius increases on going down a group.
• Because electrons are added further from the
nucleus, there is less attraction. This is due to
additional energy levels and the shielding
effect. Each additional energy level “shields”
the electrons from being pulled in toward the
nucleus.
• The radius decreases on going across a
period.
.11
The Electron Shielding Effect
• Electrons
between the
nucleus and
the valence
electrons repel
each other
making the
atom larger.
.12
Group trends
• As we go down a
group (each atom
has another energy
level) the atoms get
bigger, because
more protons and
neutrons in the
nucleus
H
Li
Na
K
Rb
• The radius decreases across a period owing to
increase in the positive charge from the protons.
• Why? Stronger attractive forces in atoms (as you go
from left to right) between the opposite charges in the
nucleus and electron cloud cause the atom to be
'sucked' together a little tighter. Remember filling up
same energy level, little shielding occurring.
• Each added electron feels a greater and greater +
charge because the protons are pulling in the same
direction, whereas the electrons are scattered.
Large
All values are in
nanometers
Small
.14
Atomic Radius
.15
Atomic Radius
.16
8.3
Atomic Radii
8.3
Trends in Ion Sizes
Radius in pm
.19
Comparison of Atomic Radii with Ionic Radii
8.3
Ionic Size
• Cations form by losing electrons.
• Cations are smaller than the atom they
come from. The electron/proton attraction
has gone Up and so the radius Decreases.
• Cations like atoms increase as one
moves from top to bottom in a group.
• Metals form cations.
• Cations of representative elements have
noble gas configuration.
Ionic size
• Anions form by gaining electrons.
• Anions are bigger than the atom they
come from. The electron/proton attraction
has gone Down and so size Increases.
• Trends in ion sizes are the same as atom
sizes.
• Nonmetals form anions.
• Anions of representative elements have
noble gas configuration.
Periodic Trends
• Metals - losing from outer energy level, more
protons than electrons so more pull, causing it to
be a smaller species.
• Non metals gaining electrons in its outer energy
level, but there are less protons than electrons in
the nucleus, so there is less pull on the protons,
so found further out making it larger.
Li+1
B+3
Be+2
C+4
N-3
O-2
F-1
Size of Isoelectronic ions
• Positive ions have more protons so they
are smaller.
Al+3
Na+1
Mg+2
Ne
F-1
O-2
N-3
Cation is always smaller than atom from
which it is formed.
Anion is always larger than atom from
which it is formed.
8.3
Trends in Ionization Energy
Ionization energy is the energy required to
remove an electron from an atom
• Metals lose electrons
more easily than
nonmetals.
• Nonmetals lose electrons
with difficulty. (They like
to GAIN electrons).
• Ionization energy
increases across a period
because the positive
charge increases.
.26
Trends in Ionization Energy
• The ionization energy is
highest at the top of a
group. Ionization energy
decreases as the atom
size increases.
• This results from an
effect known as the
Shielding Effect
.27
Ionization Energies of the
Representative Groups
.28
Ionization energy is the minimum energy (kJ/mol) required
to remove an electron from a gaseous atom in its ground
state.
I1 + X (g)
X+(g) + e-
I1 first ionization energy
I2 + X+(g)
X2+(g) + e-
I2 second ionization energy
I3 + X2+(g)
X3+(g) + e-
I3 third ionization energy
I1 < I2 < I3
8.4
8.4
Variation of the First Ionization Energy with Atomic Number
Filled n=1 shell
Filled n=2 shell
Filled n=3 shell
Filled n=4 shell
Filled n=5 shell
8.4
Electronegativity
Electronegativity is a
measure of the ability
of an atom in a
molecule to attract
electrons to itself.
• How fair it shares.
• Big electronegativity
means it pulls the
electron toward it.
This concept was first proposed by Linus
Pauling (1901-1994). He later won the Nobel
Prize for his efforts.
.32
Periodic Trends:
Electronegativity
• In a group: Atoms with fewer
energy levels can attract
electrons better (less shielding).
So, electronegativity increases
UP a group of elements.
• In a period: More protons, while
the energy levels are the same,
means atoms can better attract
electrons. So, electronegativity
increases RIGHT in a period of
elements.
.33
Electronegativity is the ability of an atom to attract
toward itself the electrons in a chemical bond. (Desire
to gain electrons)
X (g) + e-
X-(g)
Electronegativity - relative, F is highest - or most
electronegative element
9.5
The Electronegativities of Common Elements
9.5
Variation of Electronegativity with Atomic Number
9.5
Group Trend
• The further down a group the farther the electron
is away and the more electrons an atom has.
• So as you go from fluorine to chlorine to bromine
and so on down the periodic table, the
electrons are further away from the nucleus
and better shielded from the nuclear charge and
thus not as attracted to the nucleus. For that
reason the electronegativity decreases as you
go down the periodic table.
Period Trend
• Electronegativity increases from left to
right across a period
• When the nuclear charge increases, so
will the attraction that the atom has for
electrons in its outermost energy level and
that means the electronegativity will
increase
Summary of Periodic Trends
Melting Point
-based upon types of intermolecular forces
-higher mp with metallic bonds (strong
intermolecular forces,
-network solid very high mp (very strong bonds
between atoms forming large molecules),
-covalent bonds lower mp (weak
intermolecular forces)
Melting Points of Group 1
Element
Melting Point (K)
Li
453
Na
370
K
336
Rb
312
Cs
301
Fr
295
Melting Points for halogens
Element
Melting Point (K)
Fluorine
85
Chlorine
238
Bromine
332
Iodine
457
Astatine
610
Melting Point
• the temperature at which a solid changes
to a liquid
• Trends within
• a. alkali metal: MP DECREASES down the
family/group
• b. halogens: MP INCREASES down the
family/group
Metallic bonding
• Collective bond, not a single bond
• Strong force of electromagnetic attraction
between delocalized electrons (move freely).
• This is sometimes described as "an array of
positive ions in a sea of electrons
Explanation
• Generally - MP depends upon the strength of forces
holding atoms or molecules together. The stronger the
IM (intermolecular) force the higher the MP (more energy
is needed to separate molecules from the solid to the
liquid phase)
• a. Alkali Metals have metallic bonds between atoms .
As the size of the atom increases (down the family) the
metallic bonding weakens and so the MP decreases.
• b. Halogens have Van der Waal’s forces between
diatomic molecules . The molecules are NP (nonpolar)
and have relatively weak IM forces. The strength of the
forces INC with INC molecular mass, so MM increases
down the family and therefore VdW forces increase, and
therefore MP increases.
Why does the melting point decrease
going down the alkali metals family?
• Atoms are larger and their outer electrons
are held farther away from the positive
nucleus.
• The force of attraction between the metal
ions and the sea of electrons thus gets
weaker down the group.
• Melting points decrease as less heat
energy is needed to overcome this
weakening force of attraction.
Why does melting point increase
going down the halogens?
• The halogens are diatomic molecules, so
F2, Cl2, Br2, I2
• As the molecules get bigger there are
more electrons that can cause more
influential intermolecular attractions
between molecules.
• The stronger the I.M. forces, the more
difficult it will be to melt. (more energy
needed to break the I.M. forces)
3.3 Chemical properties
• Reactions of elements within the same
family
• in general, if the electron arrangement
determines the chemical reactivity of an
element, then the members of the same
family/group can be expected to have
similar chemical reactivity.
Chemical properties
• Alkali Metals (Li and Na and K)
• most characteristic property is ability to lose an
electron - they have low ionization energies and
are very reactive and form ionic solids.
• Reactivity of alkali metals will increase down the
family as reactive electrons are farther from the
nucleus and easier to access and react with;
these elements tend to lose electrons and
become reducing agents (provide electrons for
oxidation rxns to occur)
Group 1 Elements (ns1, n  2)
M+1 + 1e-
2M(s) + 2H2O(l)
4M(s) + O2(g)
2MOH(aq) + H2(g)
2M2O(s)
Increasing reactivity
M
8.6
Group 1A Elements (ns1, n  2)
8.6
Alkali Metals
• Physical properties of alkali metals -soft
malleable metals with low mp and low
densities.
• Very reactive chemically -- including
exposure to air and water
Alkali Metals
• Reaction with water
Metal + water → H2(g) + metal hydroxide
• Due to the decrease in Ionization Energy
of metals moving down the Periodic Table
the reactivity of the metal INCREASES
down the table (Li reacts less violently
than does Na...)
Alkali Metals
• Reactions with halogens (ie. Cl2 and Br2)
• redox reactions to form ionic salts
2 Na + Cl2 → 2NaCl
• Note: oxidation/reduction reactions (redox)
• 1. oxidation is the increase in oxidation
number or the loss of electrons
• 2. reduction is the decrease in oxidation
number or the gain of e-.
Chemical properties
• Halogens (Cl2,Br2 and I2) – Group 17 (VII)
General Properties –
• diatomics, colored, phase changes as one goes
down the family. Cl2 is gas (green yellow), Br2 is
liquid (brown/red) and Iodine is a purple solid
• not soluble in water (non polar substance)
(hence use of oil in experiments-non polar to
dissolve halogens).
General Reactivity• highly reactive due to need for a single electron
to fill valence shell
Group 7A Elements (ns2np5, n  2)
8.6
Halogens
• Reactivity decreases as one goes down the
halogen family.
• Halogens will react by adding an electron to
themselves (they behave as oxidizing agents they are reduced - gain electrons). The smallest
and most electronegative element F is the most
reactive.
• Valence electrons that are farther from the
nucleus will have less attraction and are
therefore less reactive.
Group 7A Elements (ns2np5, n  2)
X2(g) + H2(g)
X-1
2HX(g)
Increasing reactivity
X + 1e-
8.6
Halide Ions (F-, Cl-, Br- and I-)
• Reactivity
oxidizing power of the ions decreases
going down the table (size of atom
increases and attraction for electrons
decreases) so Cl will oxidize I but I will not
oxidize Cl (higher halogen will displace a
lower halogen from its salts.)
Halide Ions
• Reactions : assume that the halogen is the one
reacting by removing electrons from the ion,
therefore if the halogen (diatomic) is higher on
the table than the ion , the reaction will take
place, but if the ION is higher on the table than
the HALOGEN the reaction will not take place.
• Cl2 + 2 I- → I2 + 2 Cl-
• Br2 + 2 I- → I2 + Br• I2 + 2 Br- → no rxn
Properties of Oxides Across a Period
basic
acidic
8.6
Metal oxide + water → metal hydroxide (base)
ie. Na2O(s) + H2O(l) → 2 NaOH(aq)
MgO(s) + H2O(l) → Mg(OH)2(aq)
Nonmetal oxide + water → acid
ie. SO3(g) + H2O(l) → H2SO4(aq)
P4O10(s) + 6 H2O(l) → 4 H3PO4(aq)
Properties of the Third Period
Oxides
Properties of the Third Period
Chlorides
The D Block Elements
• The d block elements
fall between the s
block and the p block.
• They share common
characteristics since
the orbitals of d
sublevel of the atom
are being filled.
The D Block Elements
•
•
The D block elements include the transition
metals. The transition metals are those d block
elements with a partially filled d sublevel in one
of its oxidation states.
Since the s and d sublevels are very close in
energy, the d block elements show certain
special characteristics including:
1. Multiple oxidation states
2. The ability to form complex ions
3. Colored compounds
4. Catalytic behavior
5. Magnetic properties
The D Block Elements
• The d electrons are close in energy to the s
electrons.
• D block elements may lose 1 or more d
electrons as well as s electrons. Hence they
often have multiple oxidation states
Some common D block oxidation states
Multiple Oxidation States
• There is no sudden sharp increase in ionization
energy as one proceed through the d electrons as
there would be with the s block.
• D block elements can lose or share d electrons
as well as s electrons, allowing for multiple
oxidation states.
• Most d Block elements have a +2 oxidation State
which corresponds to the loss of the two s
electrons.
• This is especially true on the right side of the d
block, but less true on the left.
---- For example Sc+2 does not exist, and
Ti+2 is unstable, oxidizing
in the presence of any
water to the +4 state.
Complex Ions
• The ions of the d block and the lower p block
have unfilled d or p orbitals.
• These orbitals can accept electrons either an
ion or polar molecule, to form a dative bond.
This attraction results in the formation of a
complex ion.
• A complex ion is made up of two or more ions
or polar molecules joined together.
• The molecules or ions that surround the metal
ion donating the electrons to form the complex
ion are called ligands.
Complex Ions
• Compounds that are formed with
complex ions are called coordination
compounds
• Common ligands
• Complex ions usually have either 4 or 6
ligands.
K3Fe(CN)6
Cu(NH3)42+
Complex Ions
• The formation of complex ions stabilizes
the oxidations states of the metal ion
and they also affect the solubility of the
complex ion.
»
»
»
»
»
The formation of a
complex ion often has
a major effect on the
color of the solution of
a metal ion.
The D Block Colored Compounds
• In an isolated atom all of the d sublevel electrons
have the same energy.
• When an atom is surrounded by charged ions or polar
molecules, the electric field from these ions or
molecules has a unequal effect on the energies of the
various d orbitals and d electrons.
• The colors of the ions and complex ions of d block
elements depends on a variety of factors including:
– The particular element
– The oxidation state
– The kind of ligands bound to the element
Various oxidation
states of Nickel (II)
Colors in the D Block
• The presence of a partially filled d sublevels in a
transition element results in colored compounds.
• Elements with completely full or completely empty
subshells are colorless,
– For example Zinc which has a full d subshell. Its
compounds are white
• A transition metal ion is colored, if it absorbs light in
the visible range (400-700
nanometers).
• If the compound absorbs a
particular wavelength of light its
color will be the composite of those
wavelengths that it does not absorb.
• In other words it shows its
complimentary color.
Colors and d Electron Transitions
• When ligands are attached to transition metal ions, the
d orbitals may split into two groups. Some of the
orbitals are at a lower energy than the others
• The difference in energy of these orbitals varies slightly
with the nature of the ligand or ion surrounding the
metal ion
• The energy of the transition: ∆E =hn may occur in the
visible region.
• When white light passes through a compound of a
transition metal, light of a particular frequency is
absorbed as an electron is promoted from a lower
energy d orbital to a higher one.
• The result is a colored compound
Magnetic Properties
• Paramagnetism --- Molecules with
one or more unpaired electrons are
attracted to a magnetic field. The
more unpaired electrons in the
molecule the stronger the attraction.
This type of behavior is called
• Diamagnetism --- Substances with
no unpaired electrons are weakly
repelled by a magnetic field.
• Transition metal complexes with
unpaired electrons exhibit simple
paramagnetism.
• The degree of paramagnetism
depends on the number of unpaired
electrons
Catalytic Behavior
• Many D block elements are
catalysts for various chemical
reactions
• Catalysts speed up the rate of a
reaction with out being consumed.
• The transition metals form complex
ions with ligands that can donate
lone pairs of electrons.
• This results in close contact
between the metal ion and the
ligand.
• Transition metals also have a wide
variety of oxidation states so they
gain and lose electrons in
oxidation- reduction reactions
Some Common D Block Catalysts
• Examples of D block elements that are
used as catalysts:
1. Platnium or
rhodium in a
catalytic converter
2. MnO2 decomposition
of hydrogen peroxide
3. V2O5 in the contact
process
4. Fe in Haber process
5. Ni in conversion of
alkenes to alkanes
The Periodic Table--Summary
The periodic table is a classification
system. Although we are most
familiar with the periodic table that
Seaborg proposed more than 60
years ago, several alternate designs
have been proposed.
Alternate Periodic Tables
Alternate Periodic Tables II