Unit 3: Periodicity
IB Topics 3 & 13
AP Chapters (Zumdahl): 7.12-7.13; 19.2; 21.1-21.3; & pp. 906, 922
NOTES - Unit 3: Periodicity
PART 1: The Periodic Table and Physical Properties
_____________: elements with same number of valence electrons and therefore
similar chemical and physical properties; vertical column of elements (“family”)
_________________: elements with same outer shell; horizontal row of elements
____________________________________: regular variations (or patterns) of
properties with increasing atomic weight. Both chemical and physical properties
vary in a periodic (repeating pattern) ____________________________________. Another name for “metalloid” is “semi-metal”.
Periodic Groups
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alkali metals
alkaline earth metals
transition metals
halogens
noble gases
lanthanides
actinides
Alkali Metals
•
•
Group _____ on periodic table
Called alkali metals because they all react with water
to form an alkali solution of the metal hydroxide and
hydrogen gas.
– i.e. _________________________________
•
•
React by losing outer electron to form the metal ion
Good ________________________ agents (because
they can readily lose an electron)
Alkaline Earth Metals
• Group ______ on periodic table
• Abundant metals in the earth
•
•
Not as reactive as alkali metals
Higher density & melting pt. than alkali metals
Transition Metals
• Variable valency & ___________________ state
• Forms _______________________ compounds
•
•
Forms _________________________________
_______________________________ behavior
•
Good ____________________________ agents
(because they readily gain an electron)
Halogens
• Group _____ on periodic table
– group 17 or VIIA on other P.T.’s
• React by _______________________________
more electrons to from halide ions
Noble Gases
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Unit 3: Periodicity
•
•
IB Topics 3 & 13
Group ______ (or 8) on periodic table
– group 18 or VIIIA on other P.T.’s
Sometimes called rare gases or inert gases
Lanthanides
• Part of the “inner transition metals”
• Soft silvery metals
• Tarnish readily in air
Actinides
• _______________________________ elements
AP Chapters (Zumdahl): 7.12-7.13; 19.2; 21.1-21.3; & pp. 906, 922
•
•
Relatively _______________________________
_____________________ at room temperature
•
•
React slowly with water
Difficult to separate because all have two
effective valence electrons
•
Part of the “inner transition metals”
Periodic Trends: properties that have a definite trend as you move through the Periodic Table
• Valence Electrons
• Ionic radii
• Ionization energy
• Atomic radii
• Electronegativity
• Melting points
Valence Electrons:
The electrons in the outermost electron ____________________ (highest energy level) are called valence electrons.
“___________________”= Latin to be strong
Most chemical reactions occur as valence electrons seek a stable configuration.
(______________ energy level, and to a lesser extent _________________________)
Non Valence electrons are called “_________________ electrons”. Core electrons are relatively stable.
Example: Chlorine (Cl) – 17 total electrons
Example: Cadmium (Cd) – 48 total electrons
For the tall groups the
column number is the
number of valence electrons.
General Concepts: electron attraction to the nucleus depends on…
1. ___________________________________________________________________________________________
2. ___________________________________________________________________________________________
3. ___________________________________________________________________________________________
Atomic Radii
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Unit 3: Periodicity
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IB Topics 3 & 13
AP Chapters (Zumdahl): 7.12-7.13; 19.2; 21.1-21.3; & pp. 906, 922
The atomic radius is the distance from the nucleus to the outermost electron.
Since the position of the outermost electron can never be known precisely, the atomic radius is usually defined
as _____________________________________ between the nuclei of two bonded atoms of the same element.
Thus, values not listed in IB data booklet for ______________________________________________.
Trend: ________________________________________________________
o WHY???
 The atomic radius gets bigger because electrons are added to energy levels
___________________________ away from the nucleus.
 Plus, the inner electrons shield the outer electrons from the positive charge (“pull”) of the
nucleus; known as the ___________________________________________
Trend: _________________________________________________________
o WHY???
 As the # of protons in the nucleus increases, the positive charge increases and as a result, the
“pull” on the electrons increases.
Ionic Radii
Cations (+) are always ____________________________ than the metal atoms from which they are formed. (fewer
electrons than protons & one less shell of e’s)
Anions (-) are always ____________________________ than the nonmetal atoms from which they are formed. (more
electrons than protons)
•
•
Trend: For both cations and anions, _________________________________________________________
• WHY???
• Outer electrons are farther from the nucleus (more __________________________________)
Trend: For both cations and anions, _________________________________________________________
• WHY???
• The ions contain the same number of electrons (___________________________________),
but an increasing number of protons, so the ionic radius decreases.
First Ionization Energy: The energy to remove one outer electrons from a gaseous atom.
• Trend: _________________________________________________________
– WHY???
• Electrons are in _____________________________ energy levels as you move down a group;
they are further away from the positive “pull” of the nucleus and therefore easier to remove.
• As the distance to the nucleus increases, Coulomb force is reduced. (Remember from Physics
the ____________________________________________ nature of the force)
• Trend: _________________________________________________________
– WHY???
• The increasing charge in the nucleus as you move across a period exerts ___________________
“pull” on the electrons; it requires more energy to remove an electron.
3
Unit 3: Periodicity
IB Topics 3 & 13
AP Chapters (Zumdahl): 7.12-7.13; 19.2; 21.1-21.3; & pp. 906, 922
Electronegativity: a relative measure of the attraction that an atom has for a shared pair of electrons when it is covalently bonded to
another atom.
• Trend: _________________________________________________________
• WHY???
• Although the nuclear charge is increasing, the _________________________________________ produced by
the added energy levels means the electrons are farther away from the nucleus; decreased attraction, so
decreased electronegativity; plus, shielding effect
• Trend: _________________________________________________________
• WHY???
– ________________________________________________________, atomic radius is ___________________;
attractive force that the nucleus can exert on another electron increases.
Melting Point
Depends on both…
1. _______________________________________________________________________________________
2.
_______________________________________________________________________________________
Trend (using per 3 as an example):
• Elements on the ___________ exhibit metallic bonding (Na, Mg, Al),
which increases in strength as the # of valence electrons increases.
• Silicon in the _______________________ of the period has a
macromolecular covalent structure (network) with very strong
bonds resulting in a very high melting point.
• Elements in __________________________________ (P4, S8 and Cl2)
show simple molecular structures with weak van der Waals’ forces
(more on that next unit) of attraction between molecules (which
decrease with molecular size).
• The ____________________________ (Ar) exist as single individual
atoms with extremely weak forces of attraction between the atoms.
Within groups there are also clear trends:
 In __________________________ the m.p. ____________________________________________________ as the atoms
become larger and the strength of the metallic bond decreases.
 In __________________________ the van der Waals’ attractive forces between the diatomic molecules increase down the
group so the melting points ____________________________________.
4
Unit 3: Periodicity
IB Topics 3 & 13
AP Chapters (Zumdahl): 7.12-7.13; 19.2; 21.1-21.3; & pp. 906, 922
PART 1 Practice Problems: The Periodic Table and Physical Properties
1.
Fill in the table below with the period and group numbers as they are listed on the IB Periodic Table of Elements:
Element
Period (#)
Group (# and name if appropriate)
Helium
Chlorine
Barium
Francium
2.
How many valence electrons are present in the atoms of the element with atomic number 51?
3.
Describe and explain the trend in radii of the following ions: O2-, F-, Ne, Na+, Mg2+
4.
Explain why the magnesium ion is much smaller than the magnesium atom.
5.
Rank the following elements in order of increasing electronegativity: Sr, Cl, Al, P, Cs
6.
Provide an explanation for trend determined in question number 4.
7.
Explain why carbon has the highest melting point in its period
8.
Which of the following properties increase from sodium to argon?
a. Nuclear charge
b. Atomic radius
c. electronegativity
5
Unit 3: Periodicity
IB Topics 3 & 13
AP Chapters (Zumdahl): 7.12-7.13; 19.2; 21.1-21.3; & pp. 906, 922
PART 2: The Periodic Table and Chemical Properties of Groups 1 & 7
Group 1 – the alkali metals
 React by losing an electron (good reducing agents; tend to be oxidized)
 Reactivity increases down the group as the outer electron is in successively higher energy levels and less energy
is therefore required to remove it.
 Called “alkali metals” because all react with water to form an alkali solution of the metal hydroxide and
hydrogen gas.
o Alkali metal + water  metal hydroxide (base) + hydrogen gas
 Reactions you should know: Li, Na, K + H2O
1)
2)
3)

All react with chlorine, bromine and iodine to form ionic salts
o Alkali metal + halogen gas  metal halide (salt)
 Reactions you should know: Li, Na, K + I, Br, Cl
1)
2)
3)
Group 7 – the halogens
 React by gaining electrons (good oxidizing agents; tend to be reduced)
 Reactivity decreases down the group as the outer shell is increasingly at higher energy levels and further from
the nucleus. This, together with the fact that there are more electrons between the nucleus and the outer shell
(more shielding), decreases the attraction for an extra electron.
 Since chlorine is a stronger oxidizing agent than bromine, it can remove the electron from bromide ions in
solution to form chloride ions and bromine. Similarly, both chlorine and bromine can oxidize iodide ions to form
iodine.
o Reactions you should know: Cl, Br, I + Cl-, Br-, I- (none will react with Cl- since it is most reactive)
1)
2)
3)
6
Unit 3: Periodicity
IB Topics 3 & 13
AP Chapters (Zumdahl): 7.12-7.13; 19.2; 21.1-21.3; & pp. 906, 922
Test for halide ions
 The presence of halide ions in solution can be detected by adding silver nitrate solution.
 This works well because all silver halides are insoluble in water and have characteristic colors.
o Ag+ + halide ions (X-)  AgX(s), where X = Cl, Br or I
 AgCl = white
 AgBr = cream
 AgI = yellow
 Silver halides react with light to form silver metal. This is the basis of photography.
o AgX(s) + light Ag(s) + ½X2
PART 3: Oxides and chorides of the third period (sodium  argon)
Formula
Na2O
MgO
Oxides of period 3 elements
Al2O3
SiO2
P4O10
(P4O6)
SO3
(SO2)
Cl2O7
(Cl2O)
State at 25°C
Melting pt. (°C)
Boiling pt. (°C)
Electrical
conductivity in
molten state
Structure
Rxn w/ H2O
Nature of oxide
7
Unit 3: Periodicity
IB Topics 3 & 13
AP Chapters (Zumdahl): 7.12-7.13; 19.2; 21.1-21.3; & pp. 906, 922
Explanation of trends of per 3 oxides
Physical properties - melting points, boiling points & conductivity: (from left  right)
 Left side of PT: oxides of Na, Mg and Al = ionic (metal + nonmetal… large difference in electronegativity)
o Ionic solids have high melting points
o Ionic solids are capable of conducting electricity in molten state (moving charge = electricity)
 Middle of PT: oxide of Si (silicon dioxide) = macromolecular structure
o Strong diamond-like structure (covalently bonded network) accounts for high boiling point
 Right side of PT: oxides of P, S and Cl = simple covalent molecules
o difference in electronegativities between element and oxygen is small
o low melting and boiling points
Chemical properties - Acid-base nature of oxides: (from left  right)
 Oxides of electropositive (opposite of electronegative) elements are very basic and form alkaline solutions


o
Na2O(s) + H2O(l) 
o
MgO(s) + H2O(l) 
The amphoteric nature of aluminum oxide can be seen from its rxns w/ hydrochloric acid (a strong acid) and
sodium hydroxide (a strong base)
o
Acting as a base:
o
Acting as an acid:
Silicon dioxide behaves as a weak acid. It does not react with water (that would be weird… SiO2 is sand, so if
sand chemically reacted with water then our beaches would be very different places). However, silicon dioxide
will form sodium silicate with sodium hydroxide (meaning it reacts with a base).
o

SiO2(s) + 2NaOH(aq) 
The oxides of phosphorus, sulfur and chlorine are all strongly acidic (all form strong acids when added to water).
o
SO2(s) + H2O(l) 
o
P4O10(s) + H2O(l) 
o
Cl2O7(l) + H2O(l) 
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Unit 3: Periodicity
Formula
IB Topics 3 & 13
Chlorides of period 3 elements
NaCl
MgCl2
Al2Cl6
SiCl4
AP Chapters (Zumdahl): 7.12-7.13; 19.2; 21.1-21.3; & pp. 906, 922
PCl3
(PCl5)
S2Cl2
Cl2
State at 25°C
Melting pt. (°C)
Boiling pt. (°C)
Electrical
conductivity in
molten state
Structure
Rxn w/ H2O
Nature of
solution
Explanation of trends of per 3 chlorides
Physical properties - melting points, boiling points & conductivity: (from left  right)
 Structure affects physical properties of chlorides in the same way it did for oxides.
 NaCl and MgCl2 are ionic – conduct electricity in molten state – high melting points
 AlCl3 is covalent – poor conductor
 Unlike SiO2, SiCl4 has a simple molecular structure (not a macromolecular network)
 PCl3, PCl5, S2Cl2, Cl2 – all have simple molecular structures (covalently bonded molecules)
o held together by weak van der Waals’ forces, resulting in low melting and boiling points.
Chemical properties - Acid-base nature of chlorides: (from left  right)
 NaCl dissolves in H2O to give a neutral solution (more on why when we get to the acid/base unit).
 MgCl2 gives a slightly acidic solution with water (again, more on this later).
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Unit 3: Periodicity


IB Topics 3 & 13
AP Chapters (Zumdahl): 7.12-7.13; 19.2; 21.1-21.3; & pp. 906, 922
All other chlorides including AlCl3 react vigorously with water to produce acidic solutions of hydrochloric acid
together with fumes of hydrogen chloride.
o ___AlCl3(s) + ___H2O(l) 
o
SiCl4(l) + ___H2O(l) 
o
PCl3(l) + ___H2O(l) 
Chlorine itself reacts with water to some extent to form acidic solution.
o Cl2(g) + H2O(l) 
PART 4: d-block elements (first row)
The first row transition elements: electron configurations are [Ar]…
(Sc)
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
(Zn)
Note: for Cr and Cu it is more energetically favorable to half-fill or completely fill the d sub-level so they have only one 4s electron
Oxidation states of the first row transition series: (you need to be familiar with the ones in bold)
(Sc)
Ti
V
Cr
Mn
Fe
Co
Ni
+3
+2
+3
+4
+2
+3
+4
+5
+2
+3
+4
+5
+6
+2
+3
+4
+5
+6
+7
+2
+3
+4
+5
+6
+2
+3
+4
+5
+2
+3
+4
Cu
+1
+2
+3
(Zn)
+2
Transition element: an element that possesses an incomplete d sub-level in one or more of its oxidation states.
Based on the above definition, which of the elements above are not transition elements? Explain.
Characteristic properties of transition elements:
 _______________________________________
 _______________________________________
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_______________________________________
_______________________________________
10
Unit 3: Periodicity
IB Topics 3 & 13
AP Chapters (Zumdahl): 7.12-7.13; 19.2; 21.1-21.3; & pp. 906, 922
Variable oxidation states
 ______ and ______ sublevels are similar in energy
 Transition metals lose their ____-electrons first
 The _________________________ increase in ionization energy from losing the last 4s election to the first 3d
election explains the existence of ______________________________________________________________.
 All transition metals can show an oxidation number of ________.
 Some transition metals can form the _______ or _______ ion
o Fe3+
o Mn4+
 The M4+ ion is ______________ and in higher oxidation states the element is generally found not as the free
metal ion, but either ______________________________ bonded or as the oxyanion, such as MnO4-.
 Some common examples of variable oxidation states in addition to +2 follow:
Transition element Oxidation #
Formula
Name
Cr
CrCl3
Cr
Cr2O72-
Mn
MnO2
Mn
MnO4-
Fe
Fe2O3
Cu
Cu2O
Formation of complex ions
 Ions of the d-block elements attract species that are rich in electrons (ligands) because of their __________ size.
 The electron pair from a ____________________ can form ______________________________ covalent bonds
(a.k.a. “dative” - a bond in which both shared electrons are supplied by one species… remember Lewis Bases
from Honors Chem?) with the metal ion to form __________________________________________________.
 Ligand: _____________________________________________________________________________________
___________________________________________________________________________________________
o The word “ligand” is derived from _________________________, the Latin word for “bound”
o Most (but not all) transition metal ions exist as __________________ complex ions in aqueous solutions
 [Fe(H2O)6]3+
o Ligands can be replaced by other ligands (such as ___________ or __________).
 Coordination number: the number of ______________________________________ bonded to the metal ion.
o Examples: state the coordination numbers of the species below.
[Fe(CN)6]3[CuCl4]2[Ag(NH3)2]+
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Unit 3: Periodicity

IB Topics 3 & 13
AP Chapters (Zumdahl): 7.12-7.13; 19.2; 21.1-21.3; & pp. 906, 922
More examples: fill in the ligand, coordination number and oxidation number in the table below.
Complex
Ligand
Coordination
Oxidation number
Shape (more on
number
of central ion
this next unit)
3+
Fe(H2O)6]
octahedral
2[CuCl4]
tetrahedral
3+
Co(NH3)6]
octahedral
+
[Ag(NH3)2]
linear
MnO4
tetrahedral
Ni(CO)4
tetrahedral
PtCl2(NH3)2
square planar
Colored complexes
 In the free ion, the five ______________________________ are degenerate (of equal energy). However, in
complexes the d orbitals are split into _________________________________________________________.
 The energy difference between the levels corresponds to a specific ______________________________ and
_____________________________ in the ______________________ region of the electromagnetic spectrum.
 When the complex is exposed to light, energy of a specific wavelength is ______________________________
and electrons are excited from the _____________________ level to the _________________________ level.
 Cu2+(aq) appears blue because it is the ___________________________________________________________
to the wavelengths that have been absorbed.
 The energy separation between the orbitals and hence the color of the complex depends on the following
factors:
o _______________________________________________ (based on identity of the central metal ion)
o ___________________________________________________________________________________
 Ex: NH3 has a higher charge density than H2O and so produces a larger split in the d sublevel.
 [Cu(H2O)6]2+ absorbs red-orange light and appears __________ blue
 [Cu(NH3)4(H2O)2]2+ absorbs the higher energy yellow light and appears __________ blue
o Number of ____________________________ present (and hence the oxidation # of the central ion)
o ____________________________ of the complex ion
 Electric field created by the ligand’s lone pair of electrons depends on the geometry of the
complex ion
 If the d sublevel is completely empty, as in Sc3+, or completely full, as in Cu+ or Zn2+, no transitions within the d
sublevel can take place and the complexes are _________________________________.
o NOTE: it is important to distinguish between the words “clear” and “colorless.” Neither AP, nor IB, will
give credit for use of the word clear (which means translucent) when colorless should have been used.
Think about it, something can be pink and clear… colorless means something else.
Catalytic behavior
 Many transition elements and their compounds are very efficient catalysts.
 Catalysts __________________________________________________________________________________.
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Unit 3: Periodicity

IB Topics 3 & 13
AP Chapters (Zumdahl): 7.12-7.13; 19.2; 21.1-21.3; & pp. 906, 922
Examples: (need to be familiar with examples and economic significance of those in bold, but mechanisms will
not be assessed)
o Fe in the Haber Process (p. 906)
o
V2O5 in the Contact Process (p. 922)
o
Ni in the conversion of alkenes to alkanes
o
Pd and Pt in catalytic converters
o
MnO2 in the decomposition of hydrogen peroxide
o
Fe2+ in heme
o
Co3+ in vitamin B12
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