IB Topic 3 Periodicity

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Transcript IB Topic 3 Periodicity

Topic 3
4.1 The Periodic Table: Learning
how the elements in the
periodic table are arranged
the terns ‘group’ and
how the electronic
configuration of an element relates to
its position in the periodic table
The Periodic Table
 Arranged
in order of increasing atomic
 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
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
modern periodic table is arranged
by increasing atomic number
The Periodic Table and Electron
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.
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
 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
 In
covalent bonds between 2 different types of
atoms, the atoms do not attract the electron pair
 How strongly electrons are attracted depends on
the size of individual atoms and their nuclear
 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
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
Ionic Radius
of the size of an ion
ions are smaller than their atomic
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
properties of elements in the
same group
reactions depend on the amount
of outer electrons an atom contains
Thus, families have basically the same
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)
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
with alkali metals to form salts
+ X2(g) 2MX(s)
Salts formed are white or colorless
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
at table 4.3 p.156
oxide, magnesium oxide,
aluminum oxide
giant ionic structures
Ions held by strong electrostatic forces
causing high melting points and boiling
Covalent Oxides
 Look
at 4.4 p. 158
Reactions of Period 3 Oxides with
general, metallic oxides are
basic and non-metallic oxides are
Ex. Na2O(s)
+ H2O(l) 
6H2O(l)  4H3PO4(aq)
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
 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
covalent molecular gas
 Diatomic
 Weak van der Waals’ forces in liquid
 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
 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
to 7d elements
as: an element that forms at least one
stable oxidation state (other than 0) with a
partially filled d subshell
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/
 Form
complex ions
 Usually
 In
form colored complexes/ compounds
compound/ complex form can act as catalysts in many
Ionization of Transition Elements
 The
transition elements form positive ions
Common transition metal ions
Cr2+ and Cr3+
Fe2+ and Fe3+
Cu+ and Cu2+
Variable Oxidation Numbers
transition metals show an oxidation
state of 2+
they have a full 4s subshell and
removal of these electrons would result in
a 2+ charge
4.25 Pg. 167
Cause of Multiple O.N.
and 3d subshells are close in energy
electrons from either require
nearly the same amount of energy
Thus, lost electrons depend on several
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
 Deduce
the charge of Ni in the complex ion
Shapes of Complex Ions
metal complexes do not
obey the VSEPR theory rules
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
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
 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
Catalysts of Haber and Contact
 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