13.2.4 Define the term ligand.

Ligand- an ion or small polar molecule that is attracted to transition metal ions because it has an electron pair that it can donate to the central metal ion.

Examples)

NH3
    _
Cl
     _
CN
     _
OH

H2O

13.2.3 Explain the existence of variable oxidation number in ions of transition elements.

• all transition elements can show an oxidation state of +2

• But you also need to know these: 

Cr (+3, +6)
Mn (+4, +7)
Fe (+3)
Cu (+1)

13.2.2 Explain why Sc and Zn are not considered to be transition elements.

• they both do not have partially filled d-subshells
• only have one oxidation state; where as transition metals have variable oxidation states

13.2.1 List the characteristic properties of transition elements.

• higher melting points 
• harder and denser than group 1 and 2 metals 
• variable oxidation number
• form complex ions
• most compounds are colored 
• are able to act as catalysts
• no significant change in atomic radii due to repulsion between 4s and 3d orbitals 

13.1.2 Describe the reactions of chlorine and the chlorides referred to in 13.1.1 with water.

Simply:

• NaCl/MgCl2 dissolves into its ions when water is added.
• AlCl3undergoes hydrolysis.
• SiCl4 and PCl3/PCl5 form acidic solutions:

In more detail and also extra notes:

CHLORIDES 

Reactions with water:

·     Sodium chloride dissolves readily due to polar water molecules being able to pull oppositely charged ions out of the lattice; NaCl forms a neutral solution;

·     magnesium chloride and aluminium trichloride form acidic solutions as their small highly charged cations (high charge density) attract water molecules (to form complexes)  some of which give up their hydrogen ions to other water molecules surrounding the complex.

               [Mg (H₂O)₆]²⁺+  H₂O    ¬®       [Mg(H₂O)₅ OH] ⁺ + H₃O⁺

               [Al (H₂O)₆]³⁺  + H₂O    ¬®         [Al (H₂O)₅ OH] ²⁺  + H₃O⁺

    

      in addition, AlCl₃ also forms HCl when it reacts with water:

                AlCl₃ (s)   + 3H₂O    ¬®       Al (OH)₃     +  3HCl (=fumes)  (exothermic  reaction)

·     simple molecular chlorides all react rapidly with water to form hydrochloric acid.

                           

                SiCl₄ (l)  +4H₂O (l)   ®    Si(OH)₄ (s)  +   4HCl  (g)                 

                PCl₅  (l)  +  4H₂O    ®    H₃ PO₄ (aq)   + 5HCl  (g)      (= HCl fumes)(exo)

                Cl₂ (g)  +   H₂O    ®    HClO   +   HCl  

OXIDES

Reactions with water

• ionic oxides are soluble and react with water to form alkaline solutions (hydroxides) with a decreasing pH when going to the right of the period (basic oxides);

         Na₂O (s)  +   H₂O (l)     ¾®  2Na⁺   (aq) +2OH^-  (aq)     /2NaOH (aq)

         MgO (s)  +   H₂O (l)     ¾®    Mg²⁺   (aq)  + 2OH^-  (aq)     /Mg(OH)₂ (aq)

•  aluminium oxide is amphoteric:

         reaction with acid:      Al₂ O₃ (s)  + 6HCl (aq)   ¾®  2AlCl₃(aq) + 3H₂O (l)

         reaction with alkali:     Al₂ O₃ (s)  + 2NaOH (aq)  ¾®  2NaAlO₂ (aq) +H₂O (l)

                                                                                                                   (sodium aluminate)

• silicon dioxide does not dissolve in or react with water ( water remains neutral) but it can react with sodium hydroxide which is why it is considered an acidic oxide;

           SiO₂ (s)  +   2NaOH (aq)¾®  Na₂SiO₃(s) +H₂O (l)

                             

• simple molecular oxides are soluble in water and react to form strong acidic solutions like phosphoric and sulphuric acid. HClO, hypochlorous acid, is a weak acid.

            P₄ O₁₀ (s)  + 6H₂O (l)     ¾®  4H₃PO₄   (aq) 

            SO₂ (g)  +H₂O (l)     ¾®  H₂SO₃   (aq) 

            Cl₂O (l)  +    H₂O     ¾®  2HClO  (aq)  

13.1.1 Explain the physical states (under standard conditions) and electrical conductivity (in the molten state) of chlorides and oxides of the elements in Period 3 in terms of their bonding and structure.

Physical Trends:

• Overall decrease in melting and boiling points as the structures change from ionic to simple molecular.  This happens between Mg and Al; metallic chlorides (NaCl, MgCl₂, AlCl₃ ) are solids, non-metal/molecular 
• chlorides are gases (covalent bonding between atoms within molecules but Van Der Waal’s between molecules).  No giant covalent chloride as silicon tetrachloride is a gas;

Conductance:

• good in molten ionic chlorides while poor in the simple molecular chlorides.

Chemical trends

·     the type of bonding formed changes from ionic to covalent;  aluminium trichloride however is now covalently bonded instead of ionic;

·     no chloride for argon as it does not form any compounds as it is a noble gas with full outer shell.

·     there is a regular trend in the empirical formula of the highest chlorides because of the increase in oxidation state from +1 to +5;

     

            NaCl      MgCl₂     Al₂ Cl₆ (=dimer)       SiCl₄         PCl₅        

·     some elements show variable oxidation states; e.g. PCl₅  and PCl₃.

Physical state:

• melting and boiling points: rise to a maximum at SiO₂ and then decline
• ionic compounds (Na₂O, MgO and Al₂O₃) on the left side have relatively high melting and boiling points (high lattice energy); solids at room temperature;
• giant covalent molecular structure like silicon dioxide has the highest melting and boiling points (lattice held together by strongest bond i.e. covalent); solids at room temperature;
• lowest melting and boiling points are for the simple molecular structures (starting from P₄O₆); they are gases at room temperature (weaker dipole-dipole attractions).

Electrical conductivity in molten state:

• changes from good to semi-conductor to poor; this is the case because their structure changes from ionic to simple molecular. 
• None of the conductors will conduct as solids (because no free moving electrons).

Chemical trends:

• the type of bonding between element and oxygen changes from ionic to covalent;  aluminium oxide is still ionic but has quite a strong covalent character whilst silicon dioxide is mostly covalent;
• type of oxide:

_o    basic:   Na₂O      MgO    
_o    amphoteric:  Al₂ O_3 
o    acidic:  SiO₂         P₄ O₁₀        SO₃     Cl₂ O₇ 

• some elements show variable oxidation states; e.g.  P₄ O₁₀ and P₄ O₆ and  SO₃ and  SO₂;
• trend in empirical formula of highest oxide: ratio of  element : oxygen  (X : O) increases from 0.5 to 3.5.

ALL NOTES FOR THIS POST HAVE BEEN TAKEN FROM https://www.google.com.sa/url?sa=t&rct=j&q=&esrc=s&source=web&cd=2&cad=rja&uact=8&ved=0CCMQFjAB&url=https%3A%2F%2Fchemistryatdulwich.wikispaces.com%2Ffile%2Fview%2Ftopic13periodicityHLnotes.doc&ei=mqpUVMeqBdCP7AaSj4DwAg&usg=AFQjCNGPC4QKZuvvru8tbYbUYfF3jOX_Qw&sig2=Gprca7QzPG8q9VDN70vqXQ

3.3.2 Discuss the changes in nature, from ionic to covalent and from basic to acidic, of the oxides across period 3.

As you go across period 3, nature of oxides changes from basic to acidic.

Oxide	Acid or base
Na2O	Basic Giant Ionic
MgO
Al2O3 - giant ionic	Amphoteric
SiO2 - giant covalent
P4O10	Acidic Molecular covalent
SO3/SO2
Cl2O7

Amphoteric: 

• posses properties of both acids and bases.
• when it reacts with an acid, it behaves as a base, and vice-versa. 
• tested with aluminum oxide
• When aluminum oxide reacts with an acid such as hydrochloric, aluminum oxide acts as a base.
• Al₂O₃ (s) + 6HCl (aq) → 2AlCl₃ (aq) + 3H₂O (l)
• Now, when aluminum oxide reacts with a base such as Sodium hydroxide, it acts as an acid.
• Al₂O₃ (s) + 2NaOH (aq) + 3H₂O (l) → 2NaAl(OH)₄

Reactions between Na₂O, MgO, P₄O₁₀  and SO₃  and water.

• Na₂O (s) + H₂O (l) →2NaOH (aq)
• Sodium oxide + water → Sodium hydroxide
• MgO (s) + H₂O (l)→ Mg(OH)₂ (aq)
• Magnesium oxide + water → Magnesium oxide
• P₄O₁₀ (s) + 6H₂O (l) → 4H₃PO₄
• Phosphorous (v) oxide + water → Phosphoric (V) acid
• SO₃(l) + H₂O (l) → H₂SO₄(aq)
• Sulfur trioxide + water → Sulfuric (VI) acid.

Across a period 3, there is a change from metallic to non-metallic. 

• Left. ex) Sodium, Magnesium and Aluminium all have the typical properties of metallic substances (E.g. good conductors of electricity), 
• Righ. ex) such as Chlorine and Sulfur are non-metals.

Metals will react to form ionic compounds. 

Non-metals will react to form covalent compounds. 

3.3.1 Discuss the similarities and differences in the chemical properties of elements in the same group.

Group 1

• Group 1 metals = Alkali Metals
• share similarities and differences in their chemical properties.
• reactive, especially with air and water 
• often stored in paraffin to avoid contact with air
• react quite vigorously with water
• ex) lithium with water will most likely lead to a slow reaction
• ex) caesium with water is almost an explosion
• reactivity increases down the group because of decreasing electrostatic attraction to outer shells, therefore electrons can be lost easily.
• good reducing agents because they donate electrons
• the reaction is exothermic (releases energy) 
• metal hydroxide and hydrogen
• first two float
• might ignite and produce a flame (ex) potassium = violet flame 

Alkali Metals reactions with water

• A = generic alkali
• 2A (s) + 2H2O (l) → 2AOH (aq) + H2(g)
• Alkali metal + Water →Metal hydroxide + hydrogen
• We can also write this reaction in its ionic form, as alkali metals form ionic compounds:
• 2A (s) + 2H2O (l) →2A+ (aq) + 2OH- (aq) + H2 (g)

Group 7: Halogens

• 7 outer electrons
• reactivity decreases down the group because the distance between the nucleus and outer-electron increases as you go down the group
• are good oxidizing agents because they accept electrons easily

Halogens reaction with alkalis

• A = generic alkali
• Ha = generic halogen
• Halogen reaction with alkali's
• 2A (s) + HA2→2AH2 (s) = 2KBr +Cl2 → 2KCl +Br2 
• chlorine = gaseous state
• bromine = liquid state
• iodine = solid state 
• after reaction known as a halide
• produce a salt
• we often call this “Displacement reactions”

Group 0/8 : Nobel Gases

• colorless gases
• monoatomic 
• very unreactive because they do not gain or lose electrons
• stable octect (full outer shell)

3.2.4 Compare the relative electronegativity values of two or more elements based on their positions in the Periodic Table.

The most electronegativie element is on the top right of the Periodic Table and the least electronegative element is on the bottom left.

• Increases from left to right across a period because of increase in nuclear charge, resulting in an increased attraction between the nucleus and the bond electrons. 
• Electronegativity decreases down a group. The bond electrons are furthest from the nucleas and so there is reduced attraction. 

• Fluorine has the highest value of 4.0.
• Ceasium has the lowest value of 0.7.

3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points for the alkali metals (Li to Cs) and the halogens (F to I). 3.3.3 Describe and explain the trends in atomic radii, ionic radii, first ionization energies and electronegativities for elements across Period 3.

Effective nuclear charge

• Nuclear charge - Number of protons an atom contains. 
• 
  
• The inner electrons reduce the attraction of the nucleus for the outer electrons, therefore the outer electrons are shielded from the nucleus and repelled by the inner electrons. 
• 
  
• The effective charge experienced by the outer electrons is less than the full nuclear charge. 

Element	Na	Mg	Al	Si
Nuclear Charge	11	12	13	14
Electron Arrangement	2, 8, 1	2, 8, 2	2, 8, 3	2, 8, 4

From left to right, across a periods, one proton is added to the nucleus, there fore its added to the outer most shell, the effective charge increases with nuclear charge because there is not change in the number of inner electrons.

Element	Nuclear charge	Electron arrangement
Li	3	2, 1
Na	11	2, 8, 1
K	19	2, 8, 8, 1

As we descend the group, the increase in the nuclear charge is largely balanced by the increase in the number of inner electrons, both increase by eight between successive elements. The effective nuclear charge stays the same down the group. 

3.2.1 Define the terms first ionization energy and electronegativity.

First ionization energy - The first ionization energy is the energy required to remove the most loosely held electron from one mole of gaseous atoms to produce 1 mole of gaseous ions each with a charge of 1+. This is more easily seen in symbol terms. 

In simpler terms the first ionization energy is the energy required to remove the first electron from an atom, thus producing an ion with a 1+ charge. 

Electronegativity - Electronegativity is the ability of an atom to attract electrons to itself in a covalent bond. The Pauling scale is the most commonly used. Fluorine (the most electronegative element) is assigned a value of 4.0, and values range down to caesium and francium which are the least electronegative at 0.7.

definitions from www.chemguide.co.uk/

3.1.4 Apply the relationship between the number of electrons in the highest occupied energy level for an element and its position in the Periodic Table.

The block structure of the Periodic Table is based on the electron sub-levels of the atom.

The position of an element in the Periodic Table is based on the sub levels of the highest-energy electron in the ground-state atom.

Element	Period	Group	Electron Arrangement	Electron configuration
Helium	1	0	2	1s2
Lithium	2	1	2,1	1s2
Carbon	2	4	2,4	1s22s22p2
Aluminum	3	3	2,8,3	1s22s22p63s23p1
Chlorine	3	7	2,8,7	1s22s22p63s23p5
Potassium	4	1	2,8,8,1	1s22s22p63s23p64s1
Calcium	4	2	2,8,8,2	1s22s22p63s23p64s2

The number of electrons in the outer shell of elements with higher atomic numbers can be deduced from the group number of the element.