+

Unit 2 – 2014-2015

Periodicity

2

+Lesson 1: Review of Periodic Table StructureThursday, October 2nd

3+Periodicity

IB Understandings

The periodic table is arranged into four blocks associated with the four sublevels – s, p, d and f

The periodic table consists of groups (vertical columns) and periods (horizontal columns)

The period number (n) is the outer energy level that is occupied by electrons

The number of the principal energy level and the number of the valence electrons in an atom can be deduced from its position on the periodic table

The periodic table shows the positions of metals, non-metals and metalloids

4+3.1 Applications and Skills

Applications and skills:

Deduction of the electron configuration of an atom from the element’s position on the periodic table, and vice versa.

Guidance:

The terms alkali metals, halogens, noble gases, transition metals, lanthanoids and actinoids should be known.

The group numbering scheme from group 1 to group 18, as recommended by IUPAC, should be used.

+ 5

Periodic Table• The development of the periodic table brought a

system of order to what was otherwise an collection of thousands of pieces of information.

• The periodic table is a milestone in the development of modern chemistry. It not only brought order to the elements but it also enabled scientists to predict the existence of elements that had not yet been discovered .

6+Early Attempts to Classify Elements Dobreiner’s Triads (1827)

• Classified elements in sets of three having similar properties.

• Found that the properties of the middle element were approximately an average of the other two elements in the triad.

6

7

Dobreiner’s Triads

Element Atomic Mass(amu)

Average Density(g cm-3)

Average

ClBrI

35.579.9126.9

81.21.563.124.95

3.25

CaSrBa

40.187.6137.3

88.71.552.63.5

2.53

Note: In each case, the numerical values for the atomic mass and density of the middle element are close to the averages of the other two elements

8+Newland’s Octaves -1863John Newland attempted to

classify the then 62 known elements of his day.

He observed that when classified according to atomic mass, similar properties appeared to repeat for about every eighth element

His attempt to correlate the properties of elements with musical scales subjected him to ridicule.

In the end his work was acknowledged and he was vindicated with the award of the Davy Medal in 1887 for his work.

.8

9+Dmitri Mendeleev

Dmitri Mendeleev is credited with creating the modern periodic table of the elements.

He gets the credit because he not only arranged the atoms, but he also made predictions based on his arrangements His predictions were later shown to be quite accurate.

.9

10+Mendeleev’s Periodic Table

• Mendeleev organized all of the elements into one comprehensive table.

• Elements were arranged in order of increasing mass.

• Elements with similar properties were placed in the same row.

.10

11+ Mendeleev’s Periodic Table

12+Mendeleev’s Periodic Table

Mendeleev left some blank spaces in his periodic table. At the time the elements gallium and germanium were not known. He predicted their discovery and estimated their properties.

13+Periodic PropertiesElements show gradual changes in certain

physical properties as one moves across a period or down a group in the periodic table. These properties repeat after certain intervals. In other words they are PERIODIC

Periodic properties include:

-- Ionization Energy-- Electronegativity-- Electron Affinity-- Atomic Radius-- Ionic Radius .

13

14+REVIEW: Periodic Table

Elements are arranged by increasing atomic number, Z

Groups, or vertical columns, group elements with the same number of valence electrons and therefore similar chemical properties

Periods, or rows, are horizontal groups; the period number is equal to the principal quantum number, n, of the highest occupied energy level of the elements in that period

15+Important Group To Remember

Group Number

Recommended name

Characteristics

1 Alkali metals Very reactive metals

2 Alkaline metals

Less reactive metals

15 Pnictogens

16 Chalcogens

17 Halogens Very reactive non-metals

18 Noble Gases Non-reactive non-metals; stable octet

16+Arrangement

Metals make up most of the periodic table and are found on the left side

Metalloids include B, Si, Ge, As, Sb and Te and separate metals and non-metals (Po and At are sometimes considered metalloids); have characteristics of both metals and nonmetals

Non-Metals are on the right side of the periodic table after the metalloids

17+Metals and Nonmetals

NONMETAL

S METALS

METALS

Transition metals

Metalloids

18+Additional Groupings in the Periodic Table

Nonmetals, Metals, Metalloids, Noble gases

19+IB Goodness!

You have a large number of periodic tables in your data booklet

All groups are numbered 1-18

The position of an element is related to its electron configuration

Know the s-block, p-block, d-block and f-block; those blocks represent which valence electrons are getting filled

20+Blocks

S

dp

f

+ 21

Let’s Practice!

+ 22

Let’s Practice

23

+Lesson #2 – Review of Periodic Table FamiliesFriday, October 3rd

24+3.2 Periodic Trends

IB Understandings

Vertical and horizontal trends in the Periodic Table exist for atomic radius, ionic radius, ionization energy, electron affinity, and electronegativity.

Guidance

Only examples of general trends across periods and down groups are required. For ionization energy the discontinuities in the increase across a period should be covered. Trends in metallic and non-metallic behaviour are due to the trends

above. Oxides change from basic through amphoteric to acidic across a

period.

25+Topic 3.2 Applications and Skills

Applications and skills:

Prediction and explanation of the metallic and non-metallic behaviour of an element based on its position in the Periodic Table.

Discussion of the similarities and differences in the properties of elements in the same group, with reference to alkali metals (Group 1) and halogens (Group 17).

Guidance

Group trends should include the treatment of the reactions of alkali metals with water, alkali metals with halogens and halogens with halide ions.

Construction of equations to explain the pH changes for reactions of Na2O, MgO, P4O10, and the oxides of nitrogen and sulfur with water.

26+Periodicity

Properties of elements repeat themselves periodically

The Periodic Table is arranged to show these trends

27+Effective Nuclear Charge

Nuclear charge is the number of protons in the nucleus of an atom and increases steadily from left to right on the Periodic Table

HOWEVER, the outer electrons are shielded from the nuclear charge by the inner electrons and feel less of this positive pull

The effective nuclear charge experience by the outer electrons is less than the full nuclear charge

28+Effective Nuclear Charge - Trends

Effective nuclear charge INCREASES from left to right across the Periodic Table (no change in number of inner electrons)

Effective nuclear charge REMAINS THE SAME as we go down a group

29+Atomic Radius

Atomic radius is measured as ½ the distance between neighboring nuclei

Atomic radii DECREASE from left to right across the period and INCREASE down a group

30+Atomic Radius (cont.)

Atomic radii increase down a group because we are adding Principle Energy Levels

Atomic radii decrease from left to right as we increase effective nuclear charge

31+Ionic Radius

Positive ions (cations) have a smaller radius than the parent atom; lose outer valence shell

Negative ions (anions) have a larger radius than the parent atom; adding electrons to outer shell which increases electron repulsion

Atomic radii decrease from Groups 1 to Group 14 for the positive ions as we increase effective nuclear charge

Atomic radii decrease from Groups 14 to Group 17 for the negative ions as we increase effective nuclear charge

Ionic radii increase down the group as number of energy shells increases

32+Let’s Practice

Describe and explain the trend in radii of the following atoms and ions: O2–, F–, Ne, Na+, and Mg2+.

The ions and the Ne atom have 10 electrons and the electron configuration 1s22s22p6. The nuclear charges increase with atomic number:

O: Z = +8, F: Z = +9, Ne: Z = +10, Na: Z = +11, Mg: Z = +12

The increase in nuclear charge results in increased attraction between the nucleus and the outer electrons. The ionic radii decrease as the atomic number increases.

33+Ionization Energies

34+Ionization Energy Trends

Ionization energy increases across a period as we increase effective nuclear charge

Ionization energy decreases down a group as we increase the shielding effect

Exceptions to these trends can be explained in terms of electron configuration stability (i.e. it takes more energy to remove an electron from a full sub-level or shell)

35+Electron Affinity

The first electron affinity of an element is the energy change when one mole of electrons is added to one mole of gaseous atoms to form one mole of gaseous ions:

X(g) + e– → X–(g)

Values are tabulated in Table 8 of the IB data booklet.

1st electron affinity is general exothermic as the electron is attracted to the positive nucleus; 2nd and 3rd electron affinities become endothermic as the atom already has a negative charge!

36+Electron Affinity Trends Group 17 have the highest electron affinity

Group 1 has the lowest electron affinity

Group 2 and Group 15 have the highest electron affinities; here you are adding the first electron to a half-filled sub-level; electrostatic repulsion

37+Electronegativity

The electronegativity of an element is a measure of the ability of its atoms to attract electrons in a covalent bond

Electronegativity increases from left to right across a period owing to the increase in nuclear charge, resulting in an increased attraction between the nucleus and the bond electrons.

Electronegativity decreases down a group. The bonding electrons are furthest from the nucleus and so there is reduced attraction.

38+Metals vs. Non-Metals

The ability of metals to easily conduct electricity is due to their low ionization energy and ability to move their electrons away from the nucleus

As you move from left to right on the Periodic Table, there is a slow transition from metal to semi-metal to non-metal

39

+Lesson #3 – Review of Periodic Table FamiliesTuesday, October 7, 2014

40+Nota Bene

Remember when studying that the group of an element on the Periodic Table also tells you the number of valence electrons for that element!

+The Electron Shielding Effect

Electrons between the nucleus and the valence electrons repel each other making the atom larger.

.41

+How is atomic radius measured?

Half the distance between neighboring nuclei. OR

Distance from the nucleus to the outermost electron.

+Atomic Radius Across a Period

Why does atomic radius decrease across a period?

Think about the effective nuclear charge!

How many occupied energy levels does each atom have?

Fluorine’s radius is almost half of Lithium’s.

+Atomic Radius

45+Atomic Radius

Extra Info:

We typically measure atomic radius at bonding atomic radius where we look at atoms in chemical bonds with other atoms of the same element and take ½ the diameter between the two nuclei

There is also nonbonding atomic radius or van der Waals’ radius for things like Noble Gases that do not bond; take a look at these atoms in the solid phase and measure the distance between two nuclei

46+Ionic Radius – Some More Info

The factors that effect ionic radius include nuclear charge, number of filled energy shells, electrostatic repulsion and overall charge

At the HL level, the IB likes to ask questions about ionic radius that are not clear-cut and require you to think

Which ion would have a larger ionic radius, Na+ or Mg2+?

Let’s think about this…we cannot directly compare these thinking about trends from left to right because they have different charges! But, we can recognize that both now have the electron configuration of Neon. However, Mg has one more proton in the nucleus meaning electrons are pulled in tighter!

+Atomic Radius

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+ Trends in Ion SizesRadius in pm

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49+Effective Nuclear Charge

Extra Info (way beyond IB!)

The formula for effective nuclear charge is:

Zef = Z – S

Where s= screening or shielding constant

50+Melting Points

Comparing melting points is complex because it depends on bonding as well as nuclear charge

Trends down Groups 1 through 17 can be explained as elements in each group bond in the same way!

Melting points DECREASE down Group 1 as these metallic bonds have delocalized electrons and as you move down the group the attraction decreases

Melting points INCREASE down Group 17 because these diatomic elements are held together by London Dispersion forces which get stronger as the electron cloud gets bigger

Melting points generally rise from left to right until Group 14 and then fall from Group 14 to Group 18

51+Melting Points

52+Let’s Practice

1 (a) Explain what is meant by the atomic radius of an element.(b) The atomic radii of the elements are found in Table 9 of the IB data booklet. (i) Explain why no values for ionic radii are given for the noble gases. (ii) Describe and explain the trend in atomic radii across the Period 3

elements.

2 Si4+ has an ionic radius of 4.2 × 10–11 m and Si4– has an ionic radius of 2.71 × 10–10 m. Explain the large difference in size between the Si4+ and Si4– ions.

53+Answers

1 (a)Half the distance between the nuclei of neighbouring atoms of the same element.

(b) (i) The noble gases do not form stable ions and engage in ionic bonding so the distance between neighbouring ions cannot be defined.

(ii) The atomic radii decrease from Na to Cl. This is because the number of inner, shielding, electron is is constant (10) but the nuclear charge increases from +11 to +17. As we go from Na to Cl, the increasing effective nuclear charge pulls the outer electrons closer.

2. Si4+ has an electronic configuration of 1s22s22p6 where Si4– has an electronic configuration of 1s22s22p63s23p6. Si4+ has two occupied energy levels and Si4− has three and so Si4− is larger.

54+Let’s Practice

More practice!

55+Even more practice

Yay!

56+Answers

(a) The electron in the outer electron energy level (level 4) is removed to form K+. The net attractive force increases as the electrons in the third energy level experience a greater effective nuclear charge.

(b) P3− has electronic configuration of 1s22s22p63s23p6 whereas Si4+ has an electronic configuration of 1s22s22p2. P3– has one more principal energy level than Si4+ so its valence electrons will be further from the nucleus and it will have a larger ionic radius.

(c) The ions have the same electron configuration, 1s22s22p63s23p6: both have two complete shells; the extra proton in Na+ attracts the electrons more strongly

57+More Practice

Cl- < Cl < Cl+

Phosphorus exists as molecules with four atoms: P4. Sulfur exists as molecules with eight atoms: S8. There are stronger London dispersion forces between the larger S8 molecules as there are more electrons.

58+More

59+More Practice

60

+Lesson #4 – Vertical and Horizontal Periodic TrendsWednesday, October 8, 2014

61+Warm-Up

Why does I2 have a higher melting point than F2?

Why does aluminum have a higher melting point than sodium?

62+Metals

Metallic properties are related to ionization energy; the lower the ionization energy, the more metallic an element is

A metallic structure consists of a regular lattice of positive ions in a sea of delocalized electrons

Metallic character decreases from left to right and increases from top to bottom

63+Metals

1. Are good conductors of heat and electricity

2. Are malleable (capable of being hammered into thin sheets)

3. Are ductile (capable of being draw into wires)

4. Have lustre (they are shiny)

5. Typically lose electrons; i.e. they like to be oxidized

** Note: Mercury is the only metal that is liquid at room temperature! It can actually dissolve other metals and solutions formed this way are called amalgams

64+Non-Metals + Metalloids

Non-Metals

Are poor conductors of heat and electricity

Typically gain electrons; i.e. they like to be reduced

Metalloids

Are semiconductors meaning they are able to conduct electricity only at high temperature which has widespread applications in electronics (e.g. Silicon Valley)

65+Group 1 Alkali Metals

These metals are highly reactive, soft and have low melting points

Tend to form +1 ions

Held together by metallic bonds; as you go down Group 1 melting point decreases because electrons are held less tightly meaning you can disrupt the lattice more easily

66+Group 1 Reactions

Reaction With Oxygen

All alkali metals react vigorously with oxygen to form an oxide

4Na(s) + O2(g) 2Na2) (s)

As you go down the group, the reactions become more vigorous because of their lower ionization energy.

Reaction With Water

All alkali metals react rapidly with water; more reactive down the group

2Na(s) + 2H2O(l) 2NaOH(aq) + H2(g)

The products are highly basic!

67+Group 1 Reactions With Water

Lithium floats and reacts slowly. It releases hydrogen but keeps its shape.

Sodium reacts with a vigorous release of hydrogen. The heat produced is sufficient to melt the unreacted metal, which forms a small ball that moves around on the water surface.

Potassium reacts even more vigorously to produce sufficient heat to ignite the hydrogen produced. It produces a lilac coloured flame and moves excitedly on the water surface.

Cesium just EXPLODES!

68+Group 17 Elements Group 17 are known as the halogens

They all exist as diatomic elements; Cl2, I2, etc.

Reactivity decreases down the group (opposite of Group 1)

Reactions with Group 1:

All halogens react with Group I to form a salt (halides)

2Na(s) + Cl2(g) 2NaCl(s)

69+Salts

Just as a reminder, salts are ionic compounds, positive and negative ions held together by the electrostatic attraction

Salts have high melting points and only conduct electricity in the liquid phase, gas phase or dissolved in water

More on this in the future!

70+Displacement Reactions With Halogens More reactive halogen with displace a less reactive halogen

(REMEMBER from last year – single displacement reactions!)

Reactivity: F2 > Cl2 > Br2 > I2

Example:

Cl2(g) +2KBr(aq) 2KCl(aq) + Br2(g)

All of these reactions are redox reactions, meaning one atom gains elements and another atom loses electrons

You need to know that solutions of chlorine are green, bromine are orange and iodine are violet (red-brown); solutions of the halides are colorless

71+Silver + Halides

Halogens form an insoluble (i.e. not able to be dissolved in water) solution when combined with silver

Adding these together produces a precipitate

72+Group 18 Noble Gases

Group 18 consists of the Noble Gases which are all extremely stable because they have a stable octet They are colorless They are monatomic; they exist as single atoms They are very unreactive

Other elements all want to be the Noble Gases! Elements in Groups 1, 2, and 13 lose electrons to adopt the

arrangement of the nearest noble gas with a lower atomic number. Elements in Groups 15 to 17 gain electrons to adopt the electron

configuration of the nearest noble gas on their right in the Periodic Table.

The metalloids in the middle of the table show intermediate properties.

+Chemical Properties

1. Why are alkali metals called alkali?

2. How are halides formed?

3. Why does Cl- displace Br- in a displacement reaction?

+Chemical Properties in Groups 1,2,and 7

4. How do the reactivities of the alkali metals and the halogens vary down a group?

5. Which property of the halogens increases from fluorine to iodine?

A. ionic charge

B. electronegativity

C. melting point of the element

D. chemical reactivity with metals

+ 75

Let’s Practice

+ 76

Let’s Practice

77

+Lesson 5 – Construction of Oxide Equations and Explaining pH ChangesThursday, October 9th

+Warm up: Explain why each trend occurs below in terms of effective nuclear charge:

+Oxides

What is an oxide?

Oxides are formed from the combination of an element with oxygen

80+pH of Oxides

Metal oxides are basic; they react with water to form metal hydroxides

CaO(s) + H2O(l) Ca(OH)2(aq)

Non-metallic oxides are acidic: they react with water to form acidic solutions

CO2(g) + H2O(l) H2CO3(aq) Carbonic Acid

If an oxide can act both as an acid and a base, it is classified as amphoteric!

Al2O3(s) + 2NaOH(aq) + 3H2O(l) 2NaAl(OH)4 (acts as an acid)

Al2O3(s) + 6HCl(aq) 2AlCl3(aq) + 3H2O(l) (acts as a base)

81+Group 3 Oxides – Know the Trends!

Formula of Oxide

Na2O MgO Al2O3 SiO2 P4O10 SO3 and SO2

Metal/Non-Metal

Metal Metal Metal Nonmetal

Nonmetal

Nonmetal

Nature of Oxide

Basic Basic Amphoteric

Acidic Acidic Acidic

+Metal Oxides

The transition from metallic to non-metallic character is illustrated by the bonding of the Period 3 oxides. Ionic compounds are usually formed between metals and non-metal elements, so the oxides of elements Na to Al have what we call super ionic structures.

+Metalloid Oxides

Covalent Bonds

Giant Molecular Structure

Ex: SiO2

Silicon dioxide does not dissolve in water but it is still classified as acidic because it can react with NaOH to form Na2SiO3(aq)

+Non-metal Oxides

Molecular covalent bonds

Ex: P4O10, SO3, Cl2O7

+Oxides of Period 3

Oxide Formula

Na2O (s)

MgO (s) Al2O3 (s)

SiO2 (s) P4O10(s)/P4O6 (s)

SO3(l)/SO2(g)

Cl2O7(l)/Cl2O (g)

Oxidation number

+1 +2 +3 +4 +5/+3 +6/+4 +7/+1

Electrical conductivity in molten state

high high high very low

none none none

Structure giant ionic giant covalent

molecular covalent

+Oxide Trends

As you go across period 3, the ionic character decreases.

As you go down a group of oxides, the ionic character increases.

How can you explain these trends in terms of electronegativity differences within the oxide?

Li2O

Na2

O

K2O

Rb2

O

Cs2O

MgO Al2O3 SiO2 P4O10 SO3 Cl2O7

87

+Acid-Base Character

Formula of oxide

Na2(s) MgO(s) Al2O3(s) SiO2(s) P4O10 (s) /P4O6 (s)

SO3(l)/SO2(g)

Cl2O7(l)/Cl2O (g)

Acid-Base character

basic amphoteric acidic

+Basic Character

Sodium and magnesium oxides will dissolve in water to form alkaline solutions:

Na2O (s) + H2O (l) 2NaOH (aq)

MgO (s) + H2O (l) Mg(OH)2 (aq)

+Basic Oxides

Basic oxides react with an acid to form a salt and water:

MgO (s) + HCl (aq) MgCl2 (aq) + H2O (l)

Li2O (s) + HCl (aq) LiCl (aq) + H2O(l)

+Acidic Oxides

Non-metallic oxides react with water to produce acidic solutions:

P4O10(s) + 6H2O 4H3PO4(aq)

P4O6(s) + 6H2O 4H3PO3 (aq)

+Acidic OxidesSulfur trioxide reacts with water to produce

sulfuric acid:

SO3 (l) + H2O (l) H2SO4 (aq)

Sulfur dioxide reacts with water to produce sulfic acid

Dichlorine heptoxide reacts with water to produce chloric acid

Dichlorine monoxide reacts with water to produce chlorous acid

Silicon dioxide does not react with water, but will react with

alkalis to form silicates.

+

+

95

96

97

+Lesson 7 – Intro to Transition MetalsTuesday, October 14

98+Mole Day Is Coming!Let’s plan a party!

99+IB Understandings

Transition elements have variable oxidation numbers, form complex ions with ligands, have coloured compounds, and display catalytic and magnetic properties.

Guidance

Common oxidation numbers of the transition metal ions are listed in the IB Data booklet in sections 9 and 14.

Zn is not considered to be a transition element as it does not form ions with incomplete d orbitals.

Transition elements show an oxidation number of +2 when the s electrons are removed.

100+Applications and Skills

Explanation of the ability of transition metals to form variable oxidation states from successive ionization energies.

Explanation of the nature of the coordinate bond within a complex ion.

Deduction of the total charge given the formula of the ion and ligands present.

Explanation of the magnetic properties in transition metals in terms of unpaired electrons.

101+Periodic Trends The d-block elements show a lull in the periodic patterns we

have seen in the s and p block

The 10 d-block elements all show similar characteristics to each other (as an example look at the small range of atomic radii)

102+Electron Configuration Any of the d-block elements in the same group as chromium (Cr)

and Copper (Cu) have electron configurations that are exceptions to Aufbau’s principal because s electrons get promoted to either half-fill or fully fill the d-sublevel

103+Effective Nuclear Charge

The small decrease in atomic radii across the d-block is due to the fact that there is only a small increase in the effective nuclear charge as you move left to right

Why is this?

Remember that the d-sublevel is not the valence shell! It is actually an inner shell so each time you add a proton, you are also adding an electron to an inner shell which provides some shielding effect!

Why is this amazing?

Since most of the d-block metals have similar radii, they can easily form alloys or metal mixtures because mixing two different elements does not disrupt the metallic bond structure!

+ 104

Let’s Practice

+ 105

Answer!

106+Physical Properties

High electrical and thermal conductivity

High melting point

Malleable – they are easily beaten into shape

High tensile strength – they can hold large loads without breaking • ductile – they can be easily drawn into wires

Iron, cobalt, and nickel are ferromagnetic.

Properties can be explained by the fact that the s and d sublevel electrons all delocalize to form a sea of electrons! Sea of electrons hold metal lattice together.

107+Chemical Properties Form compounds with more than one oxidation number • form a

variety of complex ions

Form coloured compounds

Act as catalysts when either elements or compounds.

+Zinc

Why is zinc not considered to be a transition element? Discuss.

109+Zinc

Zinc only forms a +2 ion AND in both its natural state and as an ion it has a completely filled d-sublevel

Scandium3+ (Sc3+) also forms a colorless solution because it has NO d sublevel electrons but it is still a transition metal because it has an incomplete d sublevel. It can also sometimes for the +2 ion!

110+Oxidation States One of the key features of transition metals is their ability to

form multiple ions with various oxidation states!

This is different than alkali metals which only form +1 ions and alkaline metals which only form +2 ions

Look at the different in ionization energies between Calcium, which only forms a +2 ion, and Titanium, a transition metal. Ti is lacking the big jump Ca has from the 2nd to 3rd Ionization Energy

111+Multiple Oxidation States

This is because in Calcium, once you remove 2 electrons, you get to a noble gas configuration; removing one more electron makes that ion very unstable!

In Titanium, the 3d and 4s electrons are very close in energy so it can lose from both sublevels until they are all gone; the jump doesn’t happen until the next electron lost would have to be from the 3p sublevel!

+Oxidation States

What is the most common oxidation state of the transition elements? Look at the red!

Sc Ti V Cr Mn Fe Co Ni Cu Zn

+1

+2 +2 +2 +2 +2 +2 +2 +2 +2 +2

+3 +3 +3 +3 +3 +3 +3 +3 +3

+4 +4 +4 +4 +4 +4 +4

+5 +5 +5 +5 +5

+6 +6 +6

+7

113+Important Points For IB!!

All the transition metals show both the +2 and +3 oxidation states. The M3+ ion is the stable state for the elements from scandium to chromium, but the M2+ state is more common for the later elements. The increased nuclear charge of the later elements makes it more difficult to remove a third electron.

The maximum oxidation state of the elements increases in steps of +1 and reaches a maximum at manganese. These states correspond to the use of both the 4s and 3d electrons in bonding. Thereafter, the maximum oxidation state decreases in steps of –1.

114+Important Points For IB!! Oxidation states above +3

generally show covalent character. Ions of higher charge have such a large charge density that they polarize negative ions and increase the covalent character of the compound (see Figure 3.13).

Compounds with higher oxidation states tend to be oxidizing agents. The use of potassium dichromate(VI) (K2Cr2O7), for example, in the oxidation of alcohols.

+ 115

Let’s Practice!

+ 116

Answers!

117+Next Week

Quiz on Tuesday!!! Short + review

Complex Ions – Wednesday

D-Sublevel Split – Thursday

Colored Ions – Friday

Exam the following week!!! Thursday, October 23rd tentatively unless we need one more day of review!

118

+Lesson 8 – Complex Ions and Coordinate ComplexesTuesday, October 14th, 2004

119+Complex Ions

Transition metal ions in solution have a high charge density

These ions attract polar water molecules which form coordinate bonds

REMBEMBER: A coordinate bond is a bond where one atom provides both of the electrons to the covalent bond

120+Ligands

When a complex is formed where a central ion is surrounded by molecules or ions which possess a lone pair of electrons that can enter into a coordinate bond, the surrounding species are called ligands

Ligands have to have a lone pair of electrons to donate!

The number of coordinate bonds from the ligands to the central atom is called the coordination number

6 ligands attached to the central ion! Coordination number = 6

121+Shapes of Complex Ions These complex ions form a variety of 3-dimensional shapes

Think back to VSEPR (electron cloud repulsion) and the shapes of covalent molecules; electron clouds want to be as far apart as possible

122+Shapes of Complex IonsCoordination Number

Oxidation Number of Central Ion

Shape

6 +3 Octahedral4 +7 Tetrahedral4 +2 Square Planar4 +2 Tetrahedral2 +1 Linear

Yikes! Why?! There is no easy way to predict whether it will be square planar or tetrahedral. Sad!Most are tetrahedral…that’s all I can tell ya!

123+Oxidation Number of Metal and Charge of Ion Don’t cheat! Let’s try to figure out the oxidation state of the

metal in each of these ions…Let’s look at Fe(H2O)6

3+ first… The overall charge is +3. H2O does not have a charge itself so the charge of the iron must be +3.

+3 + 0 = +3

Now, let’s look at [CuCl4]2-. The chlorine here is chlorine ion which always has a charge of -1. We have four of them so they contribute an overall charge of -4. The total charge on the ion is -2 so…

X + (-4) = -2, X= +2, the oxidation state of the copper

Let’s try the rest!

HINT: CN-1 is a polyatomic ion with a charge of -1.

124+Oxidation Number of Central Ion

1. Ligands can either be neutral (H2O, NH3, etc.) or negatively charged (Cl-, CN-, etc.). Decide the ligands’ charges

2. Add up all the charges on the ligands

3. Set up the following equation:

X (oxidation # of trans. Metal) + Y (sum of charges of ligands) = Z (overall charge)

HINT: The complex ion will always be kept in brackets [ ]! [Fe(H2O)6]2+ Fe has a +2 oxidation #

[Ni(CN)4]2- Ni has a +2 oxidation #

Hint 2: These can enter into ionic compounds! Reverse criss-cross Li2[Ni(CN)4] I know Lithium forms a +1 ion so the complex ion has to

have a -2 charge (+2 + -2 = 0)

125+Examples

126+Determining Charge of Complex Ion When a complex ion bonds with another ion, we can use the

reverse criss-cross method to determine the charge on that ion!

EXAMPLE: What is the charge for the complex ion in compound? [Cr(H2O)6]Cl3?

So, we know that chlorine always has a charge of -1. Since the overall charge for this compound is zero (I don’t see a charge on the outside), I know that the charge of the ion has to be +3 to balance out the 3 -1 charges from the chlorine. Let’s try two more!

What is the charge for the complex ion in the following compounds? [CrCl(H2O)5]Cl2.H2O and [CrCl2(H2O)4]Cl.2H2O

+2 and +1

127+Let’s Practice

Platinum (II) can form a complex ion with 1 ammonia and 3 chloride ligands. What is the overall charge and formula for this complex ion?

Platinum forms a +2 ion. Ammonia is neutral and chloride ions each have a -1 charge. So, +2 + (3*-1) = -1 overall charge

[Pt(NH3)Cl3]-1

128+Monodentate vs. Polydentate

Monodentate ligands contain a single donor atom and have one lone pair contributing to the coordinate bond in a complex E.g. Cl- and H2O

Polydentate (chelate) ligands contain two or more donor atoms that form coordinate bonds with a transition metal center. E.g. 1,2-ethanediamine (en) Ethanedioate (ox)

129+Chelating Agents If something has more than one lone pair of electrons and can

coordinate bond with a central metal ion in more than one place, it is called a polydentate ligand

EDTA4– (old name ethylenediaminetetraacetic acid) is an example of a polydentate ligand as it has six atoms (two nitrogen atoms and four oxygen atoms) with lone pairs available to form coordinate bonds.

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130+Chelating Agents

Since compounds like EDTA4+ can grip the central atom in more than one place, it is called a chelate

The resulting complex ions formed with chelates are very energetically stable, more stable than if individual ligands are bound to the central ion

131+EDTA Usage

1. Removal of heavy metals Example: lead poisoning

2. Chelating Therapy Reduces calcium ions in the blood and can remove gunk from

hardened arteries

3. Water Softening Getting rid of metal ions in water

4. Food preservation Gets rid of metal ions that would otherwise spoil food

5. Restoring metal scultures

6. Cosmetics - presevative

132+Use In Medicine and Food

Chelating agents are use widely in food because they remove transition metals by forming very energetically stable complex ions; these transition metals would otherwise catalyze oxidation reactions and spoil the food

Chelating agents are also used in medicine to remove toxic metals in the body

+ 133

Let’s Practice

+ 134

Answers

135

+Lesson 9 – d-Block Elements as Catalysts and MagnetsWednesday, October 15, 2014

136+Catalysts A catalyst alters the rate of reaction by providing an alternate

pathway with a lower activation energy

Catalysts are extremely important as they allow reactions to proceed faster

In a homogeneous catalyst, the catalyst is in the same state as the reactants

In a heterogeneous catalyst, the catalyst is in a different state as the reactants

Transition metals are super effective heterogeneous catalysts since they can use their s and d electrons to weakly interact with reactant molecules and arrange them in the correct orientation (remember that in order for a reaction to occur, the reactants must combine with enough energy in the correct orientation!)

137+Examples of Transition Metals as Heterogeneous Catalysts Iron – Used in the Haber process to make ammonia

N2(g) + 3H2(g) 2NH3(g)

Nickel – Converts alkenes to alkanes

Palladium and Platinum (and Nickel) – Used in cars’ catalytic converters

2CO(g) + 2NO(g) 2CO2(g) + N2(g)

MnO2 – Decomposition of H2O2 (we did this last year in lab!)

2H2O2(l) 2H2O(l) + O2(g)

V2O5 – Used in Contact process - 2SO2(g) + O2(g) 2SO3(g)

Fe

Pd or Pt or Ni

MnO2

V2O5

138+Hydrogenation of Oil

Nickel can be used to add hydrogen to unsaturated oils

RCH=CHR + H2 RCH2CH2R

This allows these oils to be solid at room temperature and is useful for cooking

However, this process forms unhealthy trans fats which are not metabolized correctly in the body and leads to cardiovascular problems!

139+Examples of Transition Metals as Homogeneous Catalysts The ability of transition metals to show various oxidation states

makes them effective homogeneous catalysts in redox reactions

Many of these reactions are important in Biochemistry!

Fe2+ in heme Oxygen gets transported throughout the body by forming a weak bond the the iron ion held in the center of the hemoglobin protein! Why would you want a weak bond here?

Co3+ in Vitamin B12 Part of the vitamin B12 molecule consists of an octahedral Co3+ complex. Vitamin B12 is needed for the production of healthy red blood cells and for a healthy nervous system

In a biological setting, catalysts are called enzymes!

140+Catalytic Converters in Cars

In a running car engine, nitrogen gas in the air and oxygen can react in the high heat to form nitrogen monoxide

When NO is released into the air, it combines with more oxygen to form NO2

This NO2 pollutes the air forming smog and can also react with water to form acid rain! (Think about our oxide trends!!!)

Also, carbon monoxide can be formed – CO

Therefore these gases pass through a catalytic converter – often Pd, Pt or Rd, which helps convert these dangerous high energy molecules into more stable ones!

141+Magnetism

Every spinning electron can behave like a tiny magnet

However, when electrons are paired and have opposite spins, their spins cancel out this magnetic effect

Some transition metals are unusual because of the number of unpaired electrons they have and when aligned create highly magnetic substances!

142+Types of Magnetism

Diamagnetism is a property of all materials and produces a very weak opposition to an applied magnetic field.

Paramagnetism, which only occurs with substances which have unpaired electrons, is stronger than diamagnetism. It produces magnetization proportional to the applied field and in the same direction.

Ferromagnetism is the largest effect, producing magnetizations sometimes orders of magnitude greater than the applied field.

* Note: all substances with paired electrons exhibit diamagnetism but the effect is much smaller than paramagnetism and certainly much smaller than ferromagnetism!**

143+Transition Metals

Iron, nickel, and cobalt are ferromagnetic; the unpaired d electrons in large numbers of atoms line up with parallel spins in regions called domains.

Paramagnetism increases with the number of unpaired electrons so generally increases from left to right across the Periodic Table, reaches a maximum at chromium, and decreases. Zinc has no unpaired electrons and so is diamagnetic.

144+Let’s Practice Fe2+

145+Let’s Practice

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+

QuizThursday, October 16, 2014

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+Lesson 10 – d-Sublevel SplitsFriday, October 17, 2014

148+IB Understandings

Understandings:

The d sub-level splits into two sets of orbitals of different energy in a complex ion.

Complexes of d-block elements are coloured, as light is absorbed when an electron is excited between the d orbitals.

The colour absorbed is complementary to the colour observed.

Guidance

The relation between the colour observed and absorbed is illustrated by the colour wheel in the IB Data booklet in section 17.

149+Applications and Skills

Explanation of the effect of the identity of the metal ion, the oxidation number of the metal, and the identity of the ligand on the colour of transition metal ion complexes.

Guidance

Students are not expected to recall the colour of specific complex ions.

Explanation of the effect of different ligands on the splitting of the d orbitals in transition metal complexes and colour observed using the spectrochemical series.

The spectrochemical series is given in the IB data booklet in section 15. A list of polydentate ligands is given in the data booklet in section 16. Students are not expected to know the different splitting patterns and their relation to the coordination number. Only the splitting of the 3-d orbitals in an octahedral crystal field is required.

150+Crystal Field Theory (CFT)

The d-sublevel consists five orbitals – dxy, dyz, dxz, dx2-y2, dz2

151+Colored Ions The color of the ions is related to the presence of partially filled

d-sublevel orbitals

Any ions that do not have any electrons in the d sublevel will not be colored!

152+Visible Spectrum

The visible spectrum ranges from about 400nm to 700nm. The color we see depends on the wavelength

The color depends on which colors are absorbed and which colors are transmitted or reflected!!

153+LightWhite light contains all

wavelengths in the visible spectrum

Transition metals absorb some light and transmits others

The light emitted is the composite of the wavelengths

154+D-Level Splits

But, why do these transition metals absorb light?!

In an isolated transition metal, the d-orbitals are said to be degenerate or have the same energy

However, when ligands bond with the central atom, an electric field is produced which causes the d sublevel to split into two – a higher and a lower energy level

When light passes through these d sublevels, the d electrons jump from the lower energy d orbitals to the higher energy d orbitals

155

156+Energy Level SplitsThe energy separation between the orbitals is ∆E and hence the color of the complex depends on the following factors:

1. the nuclear charge and the identity of the central metal ion

2. the charge density of the ligand

3. the geometry of the complex ion (the electric field created by the ligand’s lone pair of electrons depends on the geometry of the complex ion)

4. the number of d electrons present and hence the oxidation number of the central ion.

157+Nuclear Charge

The strength of the bond between the ligand and the ion depends on the electrostatic attraction

Ligands interact more efficiently the higher the nuclear charge

Example: [Mn(H2O)6]2+and [Fe(H2O)6]3+ both have the same electron configuration but the iron nucleus has a higher nuclear charge and so has a stronger interaction with the water ligands

Mn absorbs in the green region and therefore look pink in solution

Fe absorbs in the higher energy blue region and therefore look yellow/brown

158+Charge Density of Ligand

The greater the charge density of the ligand, the greater the energy level split between the d orbitals and the higher energy that is absorbed

The spectrochemical series arranges the ligands according to the energy separation, ∆E, between the two sets of d orbitals

The wavelength at which maximum absorbance occurs, λmax, decreases with the charge density of the ligand, as shown in the table below and in Section 15 in your IB Booklet!!

159+Geometry of the Complex

The coordination number and geometry of the complex ion also affects the color

160+Number of d electrons and oxidation state of central ion The strength of the interaction between the ligand and the

central metal ion and the amount of electron repulsion between the ligand and the d electrons depends on the number of d electrons and hence the oxidation state of the metal.

EXAMPLE: [Fe(H2O)6]2+ absorbs violet light and so appears green/yellow, whereas [Fe(H2O)6]3+ absorbs blue light and appears orange/brown.

161+Let’s Practice!

State the formula and the shape of the complex ion formed in the following reactions.(a) Some iron metal is dissolved in sulfuric acid and then left exposed to air until a yellow solution is formed.

(b) Concentrated hydrochloric acid is added to aqueous copper sulfate solution to form a yellow solution.

(c) A small volume of sodium chloride is added to aqueous silver nitrate solution. The white precipitate dissolves to form a colourless solution when ammonia solution is added.

162+Answer

(a)  [Fe(H2O)6]3+

The oxidation number is +3 as the complex is left exposed to air. The shape is octahedral as the coordination number = 6 (see left).

(b)  The complex [CuCl4]2– is yellow.The shape is tetrahedral as the coordination number = 4 (see left).

(c) NaCl(aq) + AgNO3(aq) → AgCl(s) + NaNO3(aq)The complex [Ag(NH3)2]+ is linear as the coordination number is 2. [H3N—Ag—NH3]+

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Lesson 11 – Colored IonsWednesday, October 22nd

+The D Block Colored Compounds In an isolated atom all of the d sublevel electrons

have the same energy.

When an atom is surrounded by charged ions or polar molecules, the electric field from these ions or molecules has a unequal effect on the energies of the various d orbitals and d electrons.

The colors of the ions and complex ions of d block elements depends on a variety of factors including:

The particular element The oxidation state The kind of ligands bound to the element

Various oxidation states of Nickel (II)

+Colors in the D Block

The presence of a partially filled d sublevels in a transition element results in colored compounds.

Elements with completely full or completely empty subshells are colorless, For example Zinc which has a full d subshell. Its compounds

are white

A transition metal ion exhibits color, if it absorbs light in the visible range (400-700 nanometers)

If the compound absorbs a

particular wavelengths of light its

color is the composite of those

wavelengths that it does not absorb.

It shows the complimentary color. .168

+Colors and d Electron Transitions The d orbitals may split into two groups so that two orbitals

are at a lower energy than the other three

The difference in energy of these orbitals varies slightly with the nature of the ligand or ion surrounding the metal ion

When white light passes through a compound of a transition metal, light of a particular frequency is absorbed as an electron is promoted from a lower energy d orbital to a higher one.

When the energy of the transition: ∆E =hn may occur in the visible region, the compound is colored

+Magnetic Properties Paramagnetism --- Molecules with one

or more unpaired electrons are attracted to a magnetic field. The more unpaired electrons in the molecule the stronger the attraction. This type of behavior is called

Diamagnetism --- Substances with no unpaired electrons are weakly repelled by a magnetic field.

Transition metal complexes with unpaired electrons exhibit simple paramagnetism.

The degree of paramagnetism depends on the number of unpaired electrons

.170

+Catalytic Behavior Many D block elements are catalysts

for various reactions

Catalysts speed up the rate of a chemical reaction with out being consumed.

The transition metals form complex ions with species that can donate lone pairs of electrons.

This results in close contact between the metal ion and the ligand.

Transition metals also have a wide variety of oxidation states so they gain and lose electrons in redox reactions

.

+Some Common D Block CatalystsExamples of D block elements that

are used as catalysts:

1. Platinum or rhodium is used in a catalytic converter2. MnO2 catalyzes the

decomposition of hydrogen peroxide 3. V2O5 is a catalyst for

the contact process 4. Fe in Haber process 5. Ni in conversion of alkenes to alkanes

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Lesson 12 – ReviewTuesday, October

174+Hints

Know the EXACT definition of electronegativity

Remember that silicon forms network solids; these have VERY high melting points!

Any elements in Group 12 (where Zinc lives) are not considered transition metals