IB Topic 3 Periodicity
Download
Report
Transcript IB Topic 3 Periodicity
+
Topic 3
Periodicity
+
4.1 The Periodic Table: Learning
Goals
Understand
how the elements in the
periodic table are arranged
Understand
the terns ‘group’ and
‘period’
Understand
how the electronic
configuration of an element relates to
its position in the periodic table
+
The Periodic Table
Arranged
in order of increasing atomic
number
Vertical
Most
columns and horizontal periods
of the elements are metals
The
elements bordering between metals and
non-metals are sometimes referred to as semimetals or metalloids (Si, Ge, Sb)
Some
of the groups have specific names
+
+
+
History of Chemistry
In
1869 Russian scientist Dmitri
Mendeleev published the original
periodic table
He
arranged the elements based on
increasing atomic weight
He also left gaps for elements that had not
yet been discovered and predicted their
properties
The
modern periodic table is arranged
by increasing atomic number
+
The Periodic Table and Electron
Configurations
The
group number of an element indicates the
number of electrons in the element’s highest
main energy level (outer shell)
Group
1= 1 outer electron
Group 2= 2 outer electrons, etc.
The
period number indicates the number of
main energy levels (shells) in an atom
Period
1= 1 main energy level
Period 2= 2 main energy levels, etc.
+
4.2 Physical Properties: Learning
Goals
Define
first ionization energy and electronegativity of
an element
Understand
trends in atomic radius, first ionization
energy and electronegativity across a period
Understand
trends in atomic radius, first ionization
energy and electronegativity down group 1 and
group 7
Explain
the variation of melting points for elements
across period 3 and down group 1 and group 7
+
Electronegativity
In
covalent bonds between 2 different types of
atoms, the atoms do not attract the electron pair
equally
How strongly electrons are attracted depends on
the size of individual atoms and their nuclear
charge
Electronegativity
decreases down a group
Because, the size of atoms increase down a group
The larger an atom is, the less pull from the
nucleus thus, lower electronegativity
+
Example: HF vs. HCL
+
Electronegativity Across a Period
Increases
across a period
Because, increases
in nuclear charge with no
significant change in shielding
Shielding remains approximately constant
because atoms in the same period have the
same number of inner shells
Also, the number of protons attracting the
electrons increases
+
First Ionization Energy
The
energy required to remove the outermost
electron from a gaseous atom
Energy
Down
for: M(g) M+(g) + e-
a group:
I.E. decreases
Because,
size of the atom increases so the outer
electron is further from the attraction of the nucleus
Across
a period:
I.E. increases
Because,
the nuclear charge increases
+
Atomic Radius
Size
of an atom; mostly due to electron orbits
Trends:
Increases down a group
Because, there are more electrons when going down a group
Decreases across a period
Because increase in nuclear charge with no increase in shielding
(similar as electronegativity and ionization energy)
Across
period 3, Na is the larges and Cl is the smallest
Ar is not counted because it does not bond, thus the radius cannot be
measured
+
Ionic Radius
Measure
Positive
of the size of an ion
ions are smaller than their atomic
radii
Negative
radii
ions are greater than their atomic
+
Ionic Radii Trends
Not
as simple as neutral atomic radii
Due
to the type of ions changing from side to side
Ex. Na+ and Mg2+
They have the same number of electrons, but Mg2+ has
the highest nuclear charge (pull of nucleus on
electrons) and is smallest
Ex. P3- and S2 Equal electrons, but S2- has the stronger nuclear charge,
thus is smaller
Trend
is generally similar to atomic radii
+
Melting Point
Depends
Does
The
on the type of bonding in an element
not have a consistent trend
alkali metals are metallic m.p. decrease
Because, the
delocalized electron sea is further from the
pull of the nucleus
The
halogens increases down the family
Because
the relative atomic masses of X2 increases and
van der Waal’s forces increases
van der Waal’s Intermolecular forces resulting from
temporary dipole-dipole interactions
+
Variation in M.P. across period 3
Must look at bonding to understand changes
Increases from Na to Mg to Al all metallically bonded
Silicon increases sharply has a giant covalent structure which are
very difficult to break
There is a large decrease from Si to P (covalent) van der Waal’s
forces are weak, thus not a lot of energy is required to break them
Slight increase from P to S covalent compound, but has a higher
molecular mass and requires more E to break the bonds
Decrease from S to Cl to Ar also covalently bound, but much
smaller masses thus, easier to break
+
Graph the boiling points
Na 98
Mg 649
Al 660
Si 1410
P 44
S 119
Cl -101
Ar -189
+
4.3 Chemical properties of
elements in group 1 and group 7
Chemical
properties of elements in the
same group
Chemical
reactions depend on the amount
of outer electrons an atom contains
Thus, families have basically the same
properties
+
Reactions of elements in group 1
Very
reactive metals that react readily with oxygen,
water, halogens, and other things
Nearly
all of the reactions involve the single outer
electron being removed
Reactions
are more vigorous going down the
column due to lesser ionization energy
Reaction
React
with oxygen:
vigorously with oxygen
4M(s) + O2(g) 2M2O(s)
+
Con’t
Reaction
Rapid
with water
reactions with water
2M(s) + 2H2O(l) 2MOH(aq) + H2(g)
The alkali metal hydroxides are
strong bases
Reactions get more vigorous
descending down the column
+
Reactions of Group 7 Elements
Reactions
generally are due to the elements
gaining an electron to have an octet
Reactivity
Fluorine
decreases down the family
is the most reactive element known
Basically
reacts with every element on the periodic table
Variations
in reactivity are not as easily explained
as group 1
Several
factors must be considered when explaining
+
Halogen Reactions
React
with alkali metals to form salts
2M(s)
+ X2(g) 2MX(s)
Salts formed are white or colorless
Displacement
Reactions
Reactions
between a solution of a halogen
and a solution containing halide ions
Ex: Cl2(aq) + 2KBr(aq) 2KCl(aq) + Br2(aq)
+
4.4 Properties of the Oxides of
Period 3 Elements
Look
at table 4.3 p.156
Sodium
oxide, magnesium oxide,
aluminum oxide
Have
giant ionic structures
Ions held by strong electrostatic forces
causing high melting points and boiling
points
+
Covalent Oxides
Look
at 4.4 p. 158
+
Reactions of Period 3 Oxides with
Water
In
general, metallic oxides are
basic and non-metallic oxides are
acidic
Ex. Na2O(s)
2NaOH(aq)
P4O10(s)+
+ H2O(l)
6H2O(l) 4H3PO4(aq)
+
HL
first years
+
4.5 Properties of the Chlorides of
Period 3 Elements
Chlorides
of period 3
Table 4.6 p. 161
Bonding of aluminum chloride is complex and undergoes
changes of structure and bonding as it changes states
We assume bonding is covalent molecular in all states
In
the Lewis structure Al only has 6 electrons and will
bond with another AlCl3 to form a dimer
Overall
The
structure is non-planar
bonds are weak and easily broken, thus low
melting/ boiling points
+
Con’t
SiCl4
and PCl3 covalent molecular liquids
SiCl4
is tetrahedral and PCl3 is trigonal pyramidal
van der Waals’ forces are weak, so they have low melting/
boiling points
Neither conduct electricity
PCl5
Solid
complex bonding
at room temperature
Forms a covalent molecular liquid when melted
van der Waals’ are stronger than PCl3, has a slightly higher
melting/ boiling point
Does not conduct electricity
+
Con’t
Cl2
covalent molecular gas
Diatomic
molecule
Weak van der Waals’ forces in liquid
chlorine
Low molecular mass
Bonds can be broken at temperatures
below room temperature
Does not conduct electricity
+
Reactions of Period 3 Chlorides
with Water
Table 4.7 p. 162
NaCl(s) Na+(aq) + Cl-(aq)
[Mg(H2O)6]2+(aq) [Mg(H2O)5(OH)]+(aq) + H+(aq)
AlCl3(s) + 3H2O(l) Al(OH)3(s) +3HCl(g)
AlCl3(s) + 6H2O(l) [Al(H2O)6]3+(aq) + 3Cl-(aq)
SiCl4(l) + 4H2O(l) Si(OH)4(s) + 4HCl(aq)
PCl3(l) + 3H2O(l) H3PO3(aq) +3HCl(aq)
PCl5(l) +4H2O(l) H3PO4(aq) + 5HCl(aq)
Cl2(aq) + H2O(l) HCl(a) + HOCL(aq)
(with excess water)
+
4.6 The Transition Elements:
Learning Goals
Describe
the characteristic properties of transition metals
Explain
why transition metals have variable oxidation states
Explain
why formation and describe with shape of complex
ions
Explain
why transition metal complex ions are colored
Describe
some uses of transition metals and their
compounds and catalysts
+
The Transition Metals
The ‘d-block’ elements
3d
to 7d elements
Defined
as: an element that forms at least one
stable oxidation state (other than 0) with a
partially filled d subshell
Scandium
metals
and zinc excluded from transition
+
Properties of Transition Metals
All
typical metals; high m.p., b.p. and densities
Ionization
Radii
energies increase from Ti to Cu, but only slightly
decrease from Ti to Cu, but only slightly
Can
exhibit more than one oxidation number in compounds/
complexes
Form
complex ions
Usually
In
form colored complexes/ compounds
compound/ complex form can act as catalysts in many
reactions
+
Ionization of Transition Elements
The
transition elements form positive ions
Common transition metal ions
Element
Ion(s)
Cr
Cr2+ and Cr3+
Mn
Mn2+
Fe
Fe2+ and Fe3+
Co
Co2+
Cu
Cu+ and Cu2+
+
Variable Oxidation Numbers
All
transition metals show an oxidation
state of 2+
Because
they have a full 4s subshell and
removal of these electrons would result in
a 2+ charge
Table
4.25 Pg. 167
+
Cause of Multiple O.N.
4s
and 3d subshells are close in energy
Removing
electrons from either require
nearly the same amount of energy
Thus, lost electrons depend on several
factors
Lattice enthalpy
Ionization energy
Hydration of enthalpy
+
Complex Ions
Consist
of a central transition metal ion surrounded by
a ligand
Ligand
negative ions or neutral molecules that use
lone pairs of electrons to bond to a transition metal ion
to form a complex ion
Dative covalent bonds are formed between the ligand and the
transition metal ion
Look and Fe complex ion
Except
for Ti, al transition metals form an octahedral
complex ion to form [M(H2O)6]2+ in solution
+
Oxidation number of a Transition
Metal in a Complex Ion
May
be determined based on charge of ligand
Ligands
are either negative or neutral
Neutral Ligands
1- Ligands
H2O
Cl-
NH3
CN-
CO
Br-
Deduce
the charge of Ni in the complex ion
[Ni(CN)4]2-
+
Shapes of Complex Ions
Transition
metal complexes do not
obey the VSEPR theory rules
Six
coordinate (dative covalent bonds)
complexes are nearly always octahedral
Four coordinate complexes are either
tetrahedral or square planar
Table 4.9 P. 169
+
Formation of Complex Ions
May
undergo substitution reactions
Ex. Water
replaced with another ligand
[Cu(H2O)6]2+(aq) + 4Cl-(aq) [CuCl4]2-(aq) + 6H2O(l)
The
ligands are considered Lewis bases and
according to Le Chatelier’ principle, a strongly
acidic solution means that the concentration of H+
ions is high and the position of equilibrium is
shifted to the left-hand side
+
Formation of Colored Complexes
Complex
ions form colored solutions as a result of their 3d
orbitals being split into 2 groups by ligands
The
electrons are promoted from the group of lesser energy
to the group with higher energy, causing an emission of
colored light
Ions
with no electrons in the 3d subshell are colorless
Similarly, ions
with 10 electrons in the 3d subshell are
colorless because their electrons will not move between the
orbitals
+
Catalytic Ablility
The
elements and their compounds/ complexes are
able to act as catalysts
Ex.
Iron
is the catalyst in the Haber reaction to produce
ammonia
Vanadium (V) oxide can be used in the Contact process of
convert sulfur(IV) oxide to sulfuric(VI) acid
Manganese(IV) oxide is used in the decomposition of
hydrogen peroxide
Nickel is used in the hydrogenation of alkenes to form
alkanes
+
Catalysts of Haber and Contact
Processes
Catalysts
are used to speed up a reaction without
themselves being used up
With
a catalyst, less energy and time are required to
gain product
The
use of the catalysts in both the Haber and
Contact processes is to run the reaction at a lower
temperature and increase the yield of product
Improves
the overall cost effectiveness of these
processes in industry