Topic 3: Periodicityhttp://www.youtube.com/watch?v=nsbXp64YPRQ

http://www.youtube.com/watch?v=nsbXp64YPRQ

● Chemistry is not study of a random collection of elements.

● Periodicity = repeating patterns of chemical and physical properties.

● The development of the periodic table took many years and involved scientists from several countries.

● Mendeleev grouped the known elements into families on the basis of their relative atomic masses and their chemical properties.

● He left gaps where no known elements fitted in and predicted the physical and chemical properties of these missing elements.

3.1 The periodic table

Dimitri Mendeleev (1834-1907)

http://www.youtube.com/watch?v=157KVUQXRJA

http://www.youtube.com/watch?v=157KVUQXRJA

● In the modern periodic table the elements are arranged in order of increasing atomic number , Z.

http://education.jlab.org/itselemental/

http://www.rsc.org/periodic-table/

http://education.jlab.org/itselemental/http://www.rsc.org/periodic-table/

Groups

● Elements with similar chemical and physical properties are placed underneath each other in vertical columns called groups.

● The groups are numbered 1-18.

● The number of valence electrons of the elements in group 1-2 and 13-18 can be found from the group number .

● Certain groups have their own names:- alkali metals- alkaline earth metals- pnictogens- chalcogens- halogens- noble gases

Periods

● The period number is equal to the principal quantum number, n, of the highest occupied energy level in the elements of the period.

Metals, non-metals and metalloids

http://www.ptable.com/

http://www.ptable.com/

● Shiny● Malleable and ductile● Good conductors of heat and

electricity

MetalsMetals

Non-metals

Selenium

Phosphorus

Sulphur

Carbon

Metalloids

Si GeB

● Both metallic and non-metallic properties.● They resemble metals physically and non-metals

chemically.● Silicon and germanium are semiconductors.

Transition metals

● They have very similar chemical and physical properties:– Relatively high mp, high bp and high densities– They form more than one stable cation– They often have coloured compounds and coloured

solutions.

Lanthanoides● Elements from Z=57 to Z=71● Ex. Europium (Eu) is a hard silvery-white metallic

element that is used in the security marking of euronotes.

Actinoids

● Elements from Z=89 to Z=103

● Ex. Naturally occuring uranium is a silvery-white metal. It is a mixture of U-238 (99.27%) U-235 (0.72%) and U-234 (0.006%).

● Uranium decays slowly by emitting an alpha particle. The half-life of U-238 is about 4.47 billion years and that of U-235 is 704 million years.

A collection of uranium glassware

3.2 Periodic trends

Effective nuclear charge

● In every atom there is a balance between the attraction of the positively charged nucleus for the negatively charged electrons and repulsion between the electrons.

● The outer electrons (valence electrons) are shielded from the nucleus and repelled by the inner electrons.

● The outer electrons do not experience the full attraction of the positive nucleus because of the presence of inner electrons.

● The effective charge experienced by the outer electrons is less than the full nuclear charge.

Atomic radius

● Atomic radii increase down a group and decrease across a period.

● The effective charge increases as a period is crossed from the left to the right: – One proton is added to the nucleus and one

electron to the outermost electron shell.– There is no change in the number of inner

electrons.

● The effective charge remains almost the same down a group, because:– The increase in the nuclear charge is offset by the

increase in the number of inner electrons.

Ionic radius

Ionic radii● Positive ions are smaller than their parent atoms

(because of loss of the outer shell).

● Negative ions are larger than their parent atoms (because of increased electron repulsion by addition of electrons).

● The ionic radii decrease as a period is crossed from the left to the right (because of increased attraction between the nucleus and the electrons).

● The ionic radii increase down a group as the number of electron shells increase.

Ionization energies

● The energy required to remove one mole of electrons from one mole of gaseous atoms.

Electronegativity● The ability of an atom to attract the shared pairs of

electrons in a covalent bond.

Electron affinity

● The energy change when an electron is added to an isolated atom in the gaseous state.

X(g) + e- → X- (g)

Ex. F (g) + e- → F- (g) Ea = -328 kJ mol-1

Melting points

● The nature of the bonding between the particles of an element determines its melting point.

● Strong bonds require higher energy to break:

- The stronger the bonds, the higher the melting point.

Chemical properties

http://www.webelements.com/

http://www.webelements.com/

Metallic and non-metallic character

● Metals have a tendency to lose electrons during a chemical reaction = they tend to be oxidized.

● Non-metals have a tendency to gain electrons during a chemical reaction = they tend to be reduced.

Group 18: the noble gases

● Colourless gases● They exist as single atoms ● They are very unreactive

http://www.periodicvideos.com/

http://www.periodicvideos.com/

Group 1: the alkali metals● Shiny, silvery, soft● Good conductors of electricity and heat● Low densities● Too reactive to be found in nature

Reaction with water

● The alkali metals react with water to produce hydrogen and a metal hydroxide = the resulting solution is alkaline.

● Reactivity increases down the group.Li:Na:K:

https://www.youtube.com/watch?v=uixxJtJPVXk

https://www.youtube.com/watch?v=uixxJtJPVXk

Reaction with halogens2 Na(s) + Cl

2 (g) → 2 NaCl (s)

2 K(s) + Br2 (l) →2 KBr(s)

● The most vigorous reaction occurs between the most reactive alkali metal and the most reactive halogen: francium with fluorine.

Group 17: the halogens

● Diatomic molecules

Reaction between halogens and alkali metals

● The outer electron moves from the alkali metal to the halogen atom and an ionic halide is formed.

2 Na (s) + Cl2 (g) → 2NaCl (s)

● The most vigorous reaction occurs between elements which are furthest apart in the Periodic Table (Fr and F).

https://www.youtube.com/watch?v=Mx5JJWI2aaw

https://www.youtube.com/watch?v=Mx5JJWI2aaw

Reactions between halogens and halides

● Reactivity decreases down a group as the atomic radius increases and the attraction for outer electrons decreases.

● A solution of a more reactive halogen, X2(g), will react with

a solution of halide ions of a less reactive halogen, X-(g).

● Ex. Which of the following chemical reactions are possible?

Silver halide precipitate

● The halogens form insoluble salts (= precipitate) with silver:

AgNO3 (aq) + NaI (aq) → AgI (s)

Metal oxides

Metal oxides (Na2O, MgO) are ionic compounds:

- solids in room temperature - high mp & bp - conduct electricity when molten (or in aqueous solutions) - basic in aqueous solutions: Na

2O (s) + H

2O (l) → 2 NaOH (aq)

MgO (s) + H2O (l) → Mg(OH)

2 (s)

Aluminium oxide

● Al2O

3 is an ionic oxide with some covalent character.

● It is amphoteric as it acts as a base when it reacts with acids and acts as an acid when it reacts with bases:

Silicon oxide

● Silicon oxide, SiO2, has a giant covalent structure with

very high melting and boiling points.

● It idoes not dissolve in water.

● It is classified as an an acidic oxide, because it reacts with NaOH at temperatures above 350° C.

Non-metallic oxides

● Nonmetal oxides are covalently bonded because of the small difference in the elements' electronegativity values.

● sulfur: SO2, SO

3

● chlorine: Cl2O, Cl

2O

7

● phosphorus: P4O

6, P

4O

10

● They have low mp and bp and do not conduct electricity.

● Non-metallic oxides are acidic in aqueous solutions: P

4O

10 + 6 H

2O (l) → 4 H

3PO

4 (aq)

phosphoric(V)acid

SO3 (g) + H

2O (l) → H

2SO

4 (aq)

sulfuric(VI)acid

Chemical properties of elements in the Period 3

● Na, Mg and Al are metals. They are shiny and good conductors of heat and electricity.

● Si is a semi-conductor and is called a metalloid since it has some of the properties of a metal and some of a non-metal.

● P, S, Cl and Ar are all non-metals and do not conduct electricity.

Formula of oxide

Na2O(s) MgO(s) Al

2O

3(s) SiO

2(s) P

4O

10(s) SO

3(l) and

SO2(g)

Nature of oxide

basic basic amphoteric acidic acidic acidic

Acidic rain

• Rain is naturally acidic, because the water molecules react with the CO

2 in the air and form the weak acid

H2CO

3.

• Acid rain: precipitation (rain, snow) with pH lower than 5.6.

● The main acids present in

acid rain are sulfuric acid

nitric acid.

● The sulfuric acid in the rain reacts with calcium carbonate (in limestone or marble) to create calcium sulfate, which then flakes off.

CaCO3(s) + H

2SO

4(aq) → CaSO

4(aq) + CO

2(g) +

H2O(l)

13.1 First row d-block elements

d-block elements:

Transition elements

● An element whose atoms have an incomplete d sub-shell, or which can give rise to cations with an incomplete d sub-shell.

● Zinc is a d-block element but not a transition metal.

Physical properties

● Small atomic radii compared to the neighbouring s-block elements

● Only a small increase in atomic radii across the period● Metallic bonding:high electrical and thermal

conductivity, high mp, malleable, ductile, high density● Magnetic properties

Trends in the first IE

● The rate of increase in the first IE across the period is much lower for transition elements compared to that for the main-group elements.

Chemical properties

Oxidation number (oxidation state)

● The oxidation number of an element keeps track of the number of electrons it has lost or gained.

● Some elements always have the same oxidation number:

Variable oxidation states

● When transition metals lose electrons they lose the 4s electrons first.

● All transition elements (except for Cr and Cu) contain two 4s electrons, which means that they all have an oxidation state of +2

● In an positive ion all energy levels are closer to the nucleus (more p than e) and as the energy levels move down in energy the order of the 4s and 3d sub-levels are reversed.

● The M3+ ion is the most stable for scandium to chromium.

● The M2+ ion is the most stable for Mn to Zn (the increased nuclear charge makes it more difficult to remove a third electron).

● In the higher oxidation states the elements usually not exist as a free metal ions, but covalently bonded or as a oxyanions (MnO

4-).

Complexes

● Because of their small size and relatively high charge, the transition metal ions have a high charge density.

● They attract species that are rich in electrons: ligands.

● A ligand has at least one atom with a lone pair of electrons, e.g. H

2O, NH

3, Cl-, CN-

● These electron pairs form co-ordinate covalent bonds with the metal ion and a complex ion is formed.

● The number of co-ordinate covalent bonds from ligands to the central metal ion is called the coordination number.

● In a co-ordinate bond the shared pair of electrons orginates from the same atom.

● The charge on a complex ion is the sum of the charge of the d-block metal ion and the charge of the ligand (if they are ions).

Ligand exchange (replacement)

● In aqueous solutions the water molecule usually act as a ligand, but it can be replaced with another ligand.

Catalytic behaviour

● A catalyst increases the rate of a reaction without participating in the reaction itself.

● A catalyst does not become chemically changed at the end of the reaction and can be reused.

● Many transition elements and their compounds are very efficient catalysts.

Heterogeneous catalysts● The catalyst is in a different state from the reactants.

● Ex. Iron (Fe) in the Haber-Bosch process:N

2(g) + 3H

2 (g) 2NH

3 (g)

● Nickel (Ni) in the conversion of alkenes to alkanes

● Vanadium (V) oxide in the contact process

● Manganese (IV) oxide with hydrogen peroxide

Homogenous catalysts

● Homogenous catalysts are in the same state of matter as the reactants and products.

● Transition metals are effective homogenous catalysts in redox reactions since they can relatively easily be oxidized and reduced due to their variable oxidation states.

Magnetic properties

● Every spinning electron can behave as a tiny magnet (= a magnetic dipole).

● Electrons with opposite spins cancel each other out. Most substances have all electrons paired up and so are non-magnetic.

● Materials ca be classified as diamagnetic, paramagnetic or ferromagnetic based on their behaviour in an external magnetic field.

a) Diamagnetism

● A property of all materials that do not contain unpaired electrons.

● In an external magnetic field the paired electrons orientate themselves such that the field created by their spin opposes the applied field → They are weakly repelled by an external magnetic field.

Ex. Ne

b) Paramagnetism

● Stronger than diamagnetism.● Increases with the number of unpaired electrons (from

left to right in the periodic table).● The spins of unpaired electrons in an atom or ion can

temporarily be aligned in an external magnetic field.● These unpaired electrons behave as tiny magnets and

are attracted by an external magnetic field.

c) Ferromagnetism

● The largest effect. ● Only occurs in materials where unpaired d electrons in

a large number of atoms can line up with parallell spins to form regions or domains of parallell spins.

● In an external magnetic field the unpaired electrons in the domains align themselves such that the magnetic field created by their spin is aligned with the applied field.

● After the external magnetic field is removed, the domains remain aligned making a permanent magnet.

https://www.youtube.com/watch?v=hK_Yi5nxuKQ

https://www.youtube.com/watch?v=hK_Yi5nxuKQ

Colour of transition metal complexes

● White light is the combination of equal brightness of red, green and blue light (they are primary colours since all other colours can be formed from them).

● When white light is shone on a substance, some light is absorbed and some is reflected.

www.youtube.com/watch?v=EHMH0uQDEOU

● If all light is absorbed, the substance appears black.● If only certain wavelengths are absorbed, the

compound appears coloured.● If all light is reflected, the compound appears white.

● When white light falls on an aqueous solution of a transition metal complex, the ions absorb some colours.

● The Fe3+ ion appears yellow because it absorbs light in the blue region of the spectrum. Yellow is the complementary colour to blue (= they are opposite each other in the colour wheel.)

Transition metals absorb visible light

The colour of transition metal ions

Splitting of the d-orbitals

● In an isolated gaseous transition metal atom, the five 3 d orbitals are degenerate since they all have the same energy.

● However, since the 3 d sub-shells all have different orientations in space they will be orientated differently relative to the ligands in a complex ion.

● The 3d electrons close to a ligand will experience repulsion and be raised in energy.

● The 3d electrons further from the ligand will be reduced in energy.→ The 3d sub/shell splits into two energy levels.

http://www.chemguide.co.uk/inorganic/complexions/whatis.html

● The amount the orbitals are split (∆E) and hence the colour of the complex depends on:

– The nature of the transition metal– The oxidation state– The shape of the complex– The nature of the ligand

● If the orbital is completely empty (Sc3+) or completely full (Cu+, Zn2+), the complexes are colourless.

Nuclear charge

● Ions with a higher nuclear charge form stronger coordinate bonds with the ligands lone pair of electrons (= because of stronger electrostatic attraction).

● This stronger attraction causes a larger split in the 3d orbitals (larger ΔE = higher energy light is absorbed).

● Mn2+ is pale pink (green light absorbed)● Fe3+ is yellow (blue light is absorbed)

Charge density of the ligand

● Ligands can be arranged in order of their ability to split the d orbitals in octahedral complexes.

● Spectrochemical series:

● Ammonia has greater charge density than water and produces a larger split in the d-orbitals.

UV-vis absorbtion spectrum of some copper complexes

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