Periodicity Notes
3.1
Periodic table arranged in
a. periods -- horizontal - valence energy level
b. groups/ families - vertical - share valence electron configuration
alkali metals, alkaline earth metals, halogens, noble gases
c. metals/ non metals/ metalloids
d. transition metals
e. diatomic elements
3.2
Atomic radius - size of atom
- increases down group - more electrons in farther energy levels
- decreases across period - stronger attraction between e- and nucleus
Ionic radius - size of the ion of that element
- increases down group, same explanation as atom
--cations smaller than atom
--anions larger than atom
- decreases across period until group 4, peaks then decreases again (e- , nuclear
attractions)
Electronegativity -assigned attraction for electrons, max value 4.0 (F)
- decreases down group - bigger atom less nuclear attraction
- increases across period - nonmetals desire e- more than metals
Ionization energy - energy to remove the most loosely held electron
- decreases down group - atom size is bigger, attraction for valence e- is less,
therefore easier to remove electron
- increases across period (increased nuclear charge/ closer to stable octet)
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 Point
1. Definition: the temperature at which a solid changes to a liquid
2. Trends within
a. alkali metal: MP DECREASES down family
b. halogens: MP INCREASES down the family
c. period 3: MP shows increasing trend , then low mp (after Si)
3. Explanation
Generally - MP depends upon the strength of forces holding atoms or molecules
together. The stronger the 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 nonpolar and have relatively weak intermolecular forces. The
strength of the forces INCREASES with INCREASING molecular mass, so as
molecular mass increases down the family, vdW forces increase, and therefore
MP increases.
c. Group 3: MP depends on intermolecular forces
Across Period 3 metallic bonds change to macromolecules (network solids) to
simple molecules/ atoms with vdW forces between them. As metals decrease in
size MP increases. A large jump in MP occurs when network solid form, a large
decrease in MP when simple covalent molecules exist. (P4, S8, Cl2, Ar)
Valence electrons - electrons in the outermost energy level - those that can bond or be
removed to form ions
3.3
Chemical properties
A. Reactions of elements within the same family
1. In general, if we assume that the electron arrangement determines the chemical
reactivity of an element, then the members of the same family can be expected to have
similar chemical reactivities.
2. 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.
Metals behave as reducing agent (they are oxidized) and nonmetals behave as
oxidizing agents (they are reduced). Reactivity of alkali metals will increase down
the family as reactive electrons are farther from the nucleus and easier to access
and react and these elements tend to lose electrons and become reducing agents
(provide electrons for oxidation rxns to occur)
a. oxidation-reduction reactions
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-.
3. Assigning oxidation numbers: elements = 0...
4. When ox # increases, electrons are lost
5. Na + Cl2 → NaCl
Na0 + Cl20 → Na+1 + Cl-1 sodium is oxidized, chlorine reduced
6. Oxidizing agent - the part that allows oxidation to happen (itself is reduced)
7. Reducing agent = the part that allows reduction to happen (itself is oxidized)
b. physical properties of alkali metals
soft malleable metals with low mp and low densities. Very reactive chemically including exposure to air
c. rxn with water:
Metal + water → H2(g) + metal hydroxide
Note: sodium metal reaction. Due to the decrease in Ionization Energy of metals moving
down the Periodic Table, the reactivity of the metal INCREASES down the group. Li
reacts less violently than does Na.
d. with halogens (Cl2 and Br2)
Redox reactions to form ionic salts: 2Na + Cl2 → 2NaCl
3. Halogens (Cl2, Br2 and I2)
a. general properties –
diatomics, colored, phase changes as one goes down the family. Cl2 is gas (greenyellow), Br2 is liquid (brown-red) and Iodine purple solid
b. not soluble in water (non polar substance)
-hence use of oil in experiments-non polar to dissolve halogens.
c. general reactivity
-highly reactive due to need for a single electron to fill valence shell
d. 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 for electrons
and are therefore less reactive.
4. Halide ions (F-, Cl-, Br-, and I-)
a common test to identify halide ions is AgNO3 + HX → AgX + HNO3
-addition of HNO3 (dilute to prevent precipitation by other ions in solution) and
silver nitrate. Ppt indicates the presence of halide ions (Cl, Br or I - not F) in the
form of AgCl, AgBr or AgI (white, off white, pale yellow)
*** PbI2 is a bright yellow ppt and is easily used to identify the presence of iodide ions
b. 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.
c. 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 + 2I- → I2 + 2ClBr2 + 2I- → I2 + BrI2 + Br- → no rxn
B. Change in nature for metallic to non metallic of elements in period 3
- From lab recall change in physical appearance of elements metallic shiny to yellow to
green/yellow gas to Ar gas. Phosphorous is highly reactive in air (oxygen) and must be
stored under liquid (no access to this for observation purposes)
1. metallic behavior (Na, Mg and Al) (beginning of period)
a. to the left side of period (relatively low ionization energy and so these elements
may bond to other elements to form ionic compounds in which they have lost
their valence electrons and exist as cations) .
b. metallic properties : shiny, good conductors of heat and electricity
c. Chlorides of these ions dissolve in water to give neutral solutions.
NaCl + HOH → Na+ + Cl- + HOH
d. Oxides of these metals are ionic containing oxide ions that will neutralize acids
to produce a salt and water. (Metal oxides behave as bases and neutralize acid to
form salt and water)
Na2O + HCl → NaCl + HOH
e. Oxides will react with water to form hydroxides
Na2O + HOH → 2NaOH(aq)
f. reaction of metals with water
Na + HOH → NaOH + H2
2. metalloid behavior (in the middle of the period )
a. ionization energy becomes too great for cation formation → non metal behavior
b. Bonding is not ionic but covalent
c. semi-conductor
d. reaction of chloride of these elements with water produces acidic solutions with
the oxide or hydroxide of element
SiCl4 + 2 H2O → SiO2 + 4 H+ + 4 ClCO2 + H2O ⇌ H+ + HCO3e. reaction of oxides with water
SiO2 + HOH → NO REACTION
Al2O3 + HOH→ NO REACTION
these substances are insoluble in water
f. nature of oxide of metalloid - amphoteric (amphiprotic) may behave as either an
acid or a base (depending on what it is reacting)
metalloid oxide that react with acid → basic
metalloid oxide that reacts with base → acidic
3. non metal behavior (P, S, Cl and Ar) on the right side
a. the ionization increases further and gaining an electron becomes feasible. So
these ions may bond with metals
b. non metal - not conductors of electricity
c. bonding becomes covalent - molecular
d. reaction with chlorides
i. PCl3 + HOH → H3PO3 + 3 HCl
form acidic compounds and HCl (gaseous)
ii. Cl2 + HOH → HClO + HCl
e. reaction with oxides
i. P4O10 + HOH → H3PO4
ii. SO2 +HOH → H2SO3
iii. Cl2O7 + HOH → 2 HClO4 (perchloric acid)
f. nature of oxide - acidic ( forms acid in water solution)
13.1 Periodic trends Na → Ar : the third period
A. Physical properties
1. Physical states of compounds under standard conditions -in terms of bonding and
structure
a. oxides of third period
i. Ionic compounds are solids; covalent compounds vary from low mp solids to
liquids or gases
ii. Variety of oxides exist with different ox states for P and S compounds
iii. Na2O, MgO, Al2O3, SiO2, P4O6, P4O10, SO2, SO3, Cl20, and Cl2O7
b. chloride of the third period
i. First 2 are ionically bonded high mp crystalline solids; water soluble to form
neutral solutions
ii. Chlorides of non metals have covalent bonds. These molecules have weak
intermolecular forces and so have low mp and bp. When added to water a
hydrolysis rxn occurs and acidic soln’s are formed
iii. NaCl, MgCl2, Al2Cl6, SiCl4, PCl3, PCl5 and Cl2
**sulfur chloride will not be studied but the following compounds may
form S2Cl2 and SCl2 -- generally unstable and poorly understood
2. Electrical conductivity in the molten state - in terms of bonding and structure
electrical conductivity depends on free ions in the liquid phase, if they exist then a current
can flow through the substance as a liquid, if they don’t a current cannot flow. Ionically
bonded compounds will have ions in the liquid phase and will conduct a current.
Covalently bonded compounds will not have ions in the liquid phase and will not conduct
an electric current. Hence first two compounds are ionically bonded and will therefore
conduct a current when in the molten state. Others will not (since no ions will be
available to conduct current)
B. Chemical trends - include equations
1. acid base properties for oxides
a. formation by reaction with oxygen
b. basic →amphoteric → acidic behavior across the period
Na2O, MgO, Al2O3, SiO2, P4O10 (and P4O6), SO3 (SO2), Cl2O7 and Cl2O
c. Al and Si oxides are giant covalent molecules with very high melting and
boiling points. Giant molecules render them insoluble in water. Al2O3 is
amphoteric (acts as either an acid or base - depending on what it is with)
Al2O3 + HCl → AlCl3 + H2O
Al2O3 + 2 OH- + 3 H2O → 2 Al(OH)4
SiO2 is very weakly acidic (forms silicates in hot conc bases)
d. the non metal oxide react as acids ( form acids in aqueous solution)
2. rxn of chlorides with water
a. NaCl + water → Na+ + Cl- (dissolve to form neutral solution)
b. MgCl2 + water → Mg2+ + Cl- (dissolve) (slight acid solution)
c. AlCl3 - compounds show traits of nonmetal (due to small size and high charge
of Al+3 ion). Solid at room Temp but mp 178 (low) to give Al2Cl6 gas.
AlCl3 + H2O → Al(OH)3 + 3H+ + Cl- (acidic sol’n)
d. PCl3 + 3 H2O → H3PO3 (phosphorous acid) + 3 H+ +3 ClH3PO3 → H+ + H2PO3 SiCl4 + HOH → Si(OH)4 + HCl
e. Cl2 + H2O → HOCl + HCl (acidic and electrical conductor)
f. summary metal chlorides react with water to form neutral to slightly acid
solutions and non metals chlorides form acidic solutions
3. Rxn of oxides with water
a. Na2O + H2O → Na+ + OH- (basic)
b. MgO + H2O → Mg(OH)2 -- slightly soluble so slightly basic
(some OH- ions are in solution, but most are in the compound)
c. rxns of oxides and acid → salt + water
MgO + HCl → MgCl2 + H2O
d. in summary metal oxides behave as acids in water
e.Al2O3 + H2O → No reaction with metalloids as they are giant covalent
molecules
f. SiO2 + HOH → NR (water insoluble)
g. P and S oxides form acids which may dissociate
SO3 + H2O → H+ + HSO4h. in summary , metal oxides form bases in water , metalloid and Si oxides are
water insoluble and nonmetal oxides form acids in water
13.2 d-block elements - first row identify these elements
1. Characteristic properties of transition elements (due to the filling of the d orbital
shell)
a. variable oxidation states
b. complex ion formation
c. colored compounds
d. catalytic properties
2. Elements atypical of d-block
a. those that do not have characteristic properties include Sc and Zn
b. Sc does not exhibit those properties since as an ion (+3) it has no d electrons
c. Zn also does not exhibit properties as the (+2) ion has a full d sublevel
3. Variable oxidation states
a. all may show +2 oxidation states (for the s2 sublevel preceding them). The +2
cation is from the loss of the s2 electrons
b. Oxidation states of specific elements
i. Cr = +3 and +6
Cr has an electron configuration which differs from the rule [Ar] 4s13d 5 greater stability in half filled orbitals - max spin and the energy difference
between the 4s and 3d orbitals is small
+6 state is the loss of all d and s electrons in the sublevel
+3 state is another common one that does not follow the e rule
ii. Mn = + 4 and +7
iii. Fe = +3
iv. Cu = +1
v. explanation of those states; reason that oxidation states do not have big number
on the right side of the transition metals is due to the increasing nuclear charge
and attraction for the valence electrons.
vi. Note that electrons can be removed from the s sublevel to maintain the
integrity of the d sublevel. Electrons may be removed in such a way as to provide
MAXIMUM spin (all half filled orbitals)
4. Ligands
a. definition: species that can donate a pair of electrons to the central metal ion.
Most common examples are water, ammonia, and chloride ion (show that all these
have unbonded electron pairs that may be donated to a metal (electron poor
species)
b. d-block metals have low energy unfilled d-orbitals which can accept a pair of
electrons from a ligand. The ligand and the metal can form a coordinate covalent
bond (both electrons in the bond are donated by the same source)
c. compounds which have metals bonded to ligands are called complex ions
d. most complex ions have 6 ligands arranged octahedrally around the central
metal (usually water and ammonia) or 4 ligands arranged tetrahedrally around the
metal (usually Cl)
e. complexes often provide color or color changes in a solution
5. Formation of d-block element complexes
a. examples include: Co(H2O)62+ (cobalt hexahydrate), Co(NH3)62+ (cobalt
hexammonia), and CoCl4-2. Different complexes exhibit different colors
6. Reason for colored compounds
a. in a complex ion the energy in the d-orbitals is not equal
b. based on the orientation of the orbitals and the location of the ligand, some dorbitals will be energetically lower than others .
c. when white light passes through the complex, light of a specific energy
(wavelength) can be absorbed (moving d electrons from one orbital level to a
higher one).
d. the color that appears in the complex, then , is the remainder of the wavelengths
of light that have not been absorbed by the d electrons.(not the color absorbed). If
red and yellow are absorbed then the blues and greens remain visible in the
complex (recall that visible white light is the combination of all the colors of
light)
e. complexes that have no d electrons or all d-orbitals filled will not absorb the
light waves and will not appear colored.
7. catalytic behavior of d-block and compounds
a. definition - catalyst changes rate of rxn without itself becoming changed in the
rxn
b. reason that transition catalyst forms - complex ions can form (between
transition metal and ligands that donate a pair of electrons) The reactant can cling
temporarily (reversibly) to the complex ion and the complex ion can provide an
alternate pathway for the reaction (one that is faster). A variety of available
oxidation state allows different pathways to exist.
c. types of catalyst
i. homogeneous - of same phase as reactants
ii. heterogeneous - of different phase than reactants (site for rxn – often
solid catalyst in gas or liquid reactants
d. MnO2 in decomposition of H2O2 - solid MnO2
rxn is H2O2 ---MnO2--> O2 + H2O
e. V2O5 in Contact process: used to produce sulfuric acid industrially
S + O2 → SO2
SO2 + O2 → SO3
SO3 + H2O → H2SO4
This process uses V2O5 as a catalyst (contact since the sulfur dioxide and oxygen
molecules react in contact with the surface of the solid vanadium
f. Fe in the Haber process: used for the manufacture of ammonia
N2 + 3 H2 ---> 2NH3 (with Fe catalyst - still high T and P needed)
g. Ni in conversion of an alkene to an alkane (hydrogenation rxn)
C2H2 + H2 ---Ni---> C2H6