Section 2 Periodicity Bonding in the Elements 1-20 (a)

L.I. To learn about Bonding in the Elements 1-20
S.C. By the end of this lesson you should be able to describe the metallic bond explain what is meant by the term monatomic explain what London dispersion forces are and how they arise explain what happens to the strength of LDF as the atom size increases explain the difference between covalent network and covalent molecular in terms of bpt and mpt give examples of metallic, covalent molecular, covalent network and monatomic elements

Periodic Pattern Johan Wolfgang Dobereiner – triads
the atomic mass of the central element was approximately the mean of the other two.

What does the periodic table and the sound of music have in common?
John Newlands – octaves based on atomic mass (musical notes). o Every eighth element showed similarities

fun Period I II III IV V VI VII 1 H 2 Li Be B C N O F 3 Na K Mg Ca Al
The modern Periodic Table is based on the work of Dimtri Mendeleev in 1869 He arranged the elements based on: atomic mass, similar properties He left gaps and made predictions for missing elements Period I II III IV V VI VII 1 H 2 Li Be B C N O F 3 Na K Mg Ca Al * Si Ti P S Cr Cl Mn Fe Co Ni 4 Cu Rb Zn Sr Y Zr As Nb Se Mo Br Ru Rh Pd 5 Ag Cd In Sn Sb Te fun activity in workbook

Bonding in metals activity in workbook Metallic bonding is the electrostatic attraction between the positively charged ions and the delocalised electrons.

Bonding in metals The outer electrons are delocalised and free to move throughout the lattice, making metals good conductors of electricity. The greater the number of electrons in the outer shell the stronger the metallic bond. So the melting point of Al>Mg>Na

Bonding in Monatomic elements
He ++ Noble gases have full outer electron shells They do not need to combine with other atoms. The noble gases occur as single atoms, they are said to be monatomic. Since they can be liquefied and solidified there must be some weak attraction between the atoms.

Bonding in monatomic elements
++ The electrons in an atom “wobble” and become unevenly distributed causing one side of the atom becomes slightly negative while the other side becomes slightly positive. These slight charges are given the symbol δ ‘delta’. δ- δ+ A temporary dipole is therefore formed. A dipole can induce other atoms to form dipoles, resulting in weak attractions between particles. δ- δ+ London dispersion forces activity in workbook

Bonding in monatomic elements
London dispersion forces are a type of van der Waal force. They are very weak attractive forces.

Group 8 element Electron arrangement Boiling Point oC Helium Argon
Comparing Boiling points The melting and boiling point of a substance gives an indication of the strength of the forces of attraction holding atoms or molecules together. activity in workbook Group 8 element Electron arrangement Boiling Point oC Helium Argon Krypton Xenon

Noble gases b.p.’s 20 40 60 80 100 120 140 160 180 166 121 Helium Neon b.p / K 87 Argon Krypton Xeon 27 4 activity in workbook B.p.’s increase as the size of the atom increases This happens because the London dispersion forces increases with increasing size of atoms.

Covalent Molecular Elements
Many non-metals exist as discrete covalent molecules held together by covalent bonds. Discrete molecules have a definite number of atoms bonded together. 9+ Fluorine atom Fluorine molecule F2 diatomic activity in workbook

Examples of discrete molecules:
Cl activity in workbook - table

weak London dispersion forces
strong covalent bonds Melting point low – why? activity in workbook - paragraph

Halogen Boiling point (oC)
Comparing Boiling points activity in workbook Halogen Boiling point (oC)

Halogens b.p.’s 50 100 150 200 250 300 350 400 450 500 457 332 Fluorine b.p./ K 238 Chlorine Bromine Iodine 85 activity in workbook As the size of the halogen molecule increases the boiling point increases. The bigger the molecule the stronger the London dispersion forces between the halogen molecules.

Fullerenes, molecules of carbon
Fullerenes exists as large covalent molecules with a definite formula. Fullerenes were discovered in 1985 by Buckminster Fuller. Fullerenes are spherical in shape and usually contain sixty or seventy carbons. C60 is known as Buckminster fullerene activity in workbook

Covalent Network Elements
Carbon - diamond Diamond has a covalent network structure Each of the outer electrons in a carbon atom can form a covalent bond with another carbon atom. So every C bonds to 4 others. m.p.’s C > 3642oC It is high because many covalent bonds have to be broken.

Carbon - Graphite Carbon bonded to only 3 other Carbons
The spare (4th) electron is delocalised and so free to move. Graphite is a conductor of electricity. Van der Waals forces between the layers allows layers to slide over each other. Graphite can be used as a lubricant

Properties of graphite and diamond
Property Diamond Graphite Appearance Colourless transparent solid Conduction No Feel Smooth, not slippery Hardness Very hard

Other Network Structures
In the first 20 elements, only Boron, Carbon and Silicon have covalent network structures. m.p.’s B 2300oC, C > 3642oC and Si 1410oC activity in workbook

BONDING IN ELEMENTS - A SUMMARY
activity in workbook

Bonding patterns of the 1st 20 elements
Covalent Molecular Si C B Metallic lattice Li Li Be Be B N N C O O F F Ne Monatomic Na Na Mg Mg Al P P Si S S Cl Cl Ar Covalent Network K K Ca Ca C , in the form of fullerenes, is covalent molecular

animations/bonding_structure.asp This interactive animation provides a visual representation of the bonding and structure of the first twenty elements in the periodic table, taking into account both the intra- and inter-molecular forces involved.

Questions on elements – bonding and structure
Explain why the covalent network elements have high melting and boiling points. Explain why the discrete molecular and monatomic elements have low melting and boiling points. Does diamond conduct electricity? Explain. Does graphite conduct electricity? Explain. How does the hardness of diamond compare with graphite? Explain. Give a use for both diamond and graphite. Complete the following table:

Questions on elements – bonding and structure
7. Complete the following table: Type of bonding and structure Properties Metallic solids ……………. of electricity Covalent network solids ……….. …. melting points exception ………………. Covalent molecular solids ………….. melting points …………… of electricity Covalent molecular (diatomic) gases and monatomic gases …………… boiling points

Patterns in the Periodic Table (b)
Section 2 Periodicity Patterns in the Periodic Table (b)

L.I. To learn about covalent radius
S.C. By the end of this lesson you should be able to describe the term covalent radius explain the changes in covalent radius down a group explain the changes in covalent radius across a period explain why there is no stated covalent radius for the noble gases

Covalent Radius The size of an atom is indicated by its covalent radius. (Page 7 of data booklet). There is no definite edge to an atom. 266pm However, bond lengths can be worked out. The covalent radius of an element is half the distance between the nuclei of 2 of its bonded atoms From above the covalent radius would be 133 pm. Covalent radius – picometres (pm) 1pm =1 X 10 – 12 m

Trends in covalent radius - Across a period
activity in workbook Na 154 pm, Mg 145 pm, Al 130 pm, Si 117 pm, P 110 pm, S 102 pm Going across a period the covalent radius (atomic size) decreases. Why? Going across a period the nuclear charge increases. The attraction between the outer electrons and the positive nucleus increases. Thus the outer electrons are more strongly attracted and so the atom size is smaller.

Trends in covalent radius – Down a group
activity in workbook Li 134 pm Na 154 pm K pm Rb 216 pm Going down a group the covalent radius (atomic size) increases. On moving down a group from one element to the next the number of electron shells increases. So the outer electrons are further from the nucleus and the atom size increases.

Why is there no covalent radius value for the noble gases?

L.I. To learn about ionisation energies
S.C. By the end of this lesson you should be able to describe the term 1st ionisation energy write equations for the 1st ionisation energy explain the trend in 1st ionisation energy down a group explain the trend in 1st ionisation energy across a period describe the term 2nd ionisation energy carry out calculations involving ionisation energy

Ionisation energies Na(g) Na+ (g) + e
The first ionisation energy of an element is the energy required to remove one mole of electrons from one mole of atoms in the gaseous state. Units are kJmol-1. This is an endothermic process. (Page 11 of the data booklet.) Na(g) Na+ (g) e

Cl(g) Cl+ (g) e

Trends in 1st ionisation energy – Across a period
activity in workbook Going across a period the ionisation energy increases. Going across a period the nuclear charge increases. The attraction between the negative electrons and the positive nucleus increases. Thus the electrons are more tightly held and so more energy is needed to remove the outer electrons.

Trends in 1st ionisation energy – Down a group
activity in workbook Going down a group the ionisation energy decreases. The explanation for this is (i) on moving down a group from one element to the next the number of electron shells increases and so the outer electron is further from the nucleus and less tightly held. (ii) the inner shells provide a screening effect which also decreases the attractive forces between the outer electrons and nucleus.

The 2nd Ionisation Energy
The second ionisation energy of an element is the energy required to remove the second mole of electrons. First Ionisation Second Ionisation Mg(g)  Mg+(g) + e- Mg+(g)  Mg2+ + e- ΔH = +738 kJ mol-1 ΔH = kJ mol-1 Third Ionisation Mg2+(g)  Mg3+ + e- ΔH = kJ mol-1 activity in workbook

L.I. To learn about electronegativity
S.C. By the end of this lesson you should be able to describe the term electronegativity explain the trend in electronegativity down a group explain the trend in electronegativity across a period explain why there are no quote values of electronegativity for the noble gases

Electronegativity The electronegativity is a measure of the attraction an atom involved in a bond has for the shared pair of electrons. Electronegativity values are based on the Pauling Scale, devised by Linus Pauling an American Chemist. Values on the Pauling Scale range from 0 to 4. A list of these values can be found in the data booklet on page 11. The higher the number on the Pauling scale is, the greater the attraction an atom has for the bonding electrons.

Electronegativity – Across a Period
Electronegativity values can be useful in predicting which type of bonding is most likely between two elements. (More about this later) Electronegativity – Across a Period activity in workbook On crossing a period, electronegativity values increase. This is caused by an increase in nuclear charge as you move across a period from left to right. Electronegativity – Down a Group activity in workbook As you go down a group, electronegativity values decrease. This is caused by the addition of another energy level of electrons as you go down a group which shields the bonded electrons from the nucleus; therefore they are not attracted as strongly.

Electronegativity - The Monatomic Gases
Why no values for group 8 elements?

Section 3 Structure and Bonding Bonding in Compounds

L.I. To learn about bonding in compounds (a)
S.C. By the end of this lesson you should be able to describe the bonding and structure in ionic compounds explain the melting point of ionic compounds describe the bonding and structure in covalent network compounds explain the melting point of covalent network compounds describe the bonding and structure in covalent molecular explain the melting point of covalent molecular compounds

Three different types of compound - ionic, covalent molecular or covalent network.
Ionic Bonding Na Cl Na Cl- 2)8)1 2)8)7 2)8 2)8)8 In ionic compounds atoms achieve a full outer shell by either losing or gaining electrons and so form charged particles called ions.

Complete for sodium, chlorine, bromine, oxygen, aluminium and
Element Atom electron arrangement Ion electron arrangement Ion symbol Mg 2)8)2 2)8 Mg2+ Complete for sodium, chlorine, bromine, oxygen, aluminium and nitrogen. Metal atoms always lose electrons to form positive ions e.g Na+ Non-metal atoms always gain electrons to form negative ions e.g F- Glow: ionic bonding ionic compounds

Sodium chloride Lithium fluoride Magnesium oxide Aluminium nitride Calcium chloride Now write ionic formula for the above. On show me boards – work out how these elements form an ionic compound

The electrostatic attraction between positive and
Ionic Bonding NaCl 3D lattice – regular repeating pattern of ions The attraction between positive and negative ions holds the compound together. The electrostatic attraction between positive and negative ions is an ionic bond. + - ionic bond

Ionic Compounds

Ionic Compounds NaCl as with all ionic compounds have many strong ionic bonds which are broken on melting thus the melting points are high (801 0C) Complete workbook

shared pair of electrons
COVALENT BONDING In covalent bonding the atoms share electrons. Covalent bonding is the electrostatic attraction between the shared electrons and the positive nuclei. shared pair of electrons positive nuclei

COVALENT NETWORK Silcion dioxide - SiO2 Silcion carbide - SiC Mpt = 1610oC Mpt = 2700oC

Silicon dioxide and silicon carbide exist as a covalent network.
All network structures have very high melting and boiling points. It is the strong covalent bonds that are broken on melting.

Molecular Compounds Write formula for the following compounds:
carbon monoxide, sulphur trioxide, carbon tetrafluoride, dinitrogen tetraoxide, phosphorus trifluoride

Draw electron dot cross diagrams for the following molecules and structural formula
CH4 SCl2 CO2

Weak intermolecular forces
COVALENT MOLECULAR Weak intermolecular forces Strong covalent bonds Covalent molecules tend to have low melting and boiling points as it is the weak intermolecular forces that are broken on melting. Complete workbook

Plot a graph of melting points of the carbon tetrahalides against the covalent radius of the halogen in each molecules (see data book) CF4 = -184oC CCl4 = -23oC CBr4 = 90oC CI4 = 171oC

m.p.’s of the carbon halides
COVALENT MOLECULAR m.p.’s of the carbon halides -183 90 -23 171 Temp/ oC increasing size of molecules CF4 CCl4 CI4 CBr4 What happens to the melting point as the size of the molecule increase? Why? As the molecule size increases the m.pt.s increase. This is because the strength of the London dispersion forces increase, so more energy is needed to separate molecules. Complete workbook

L.I. To learn about polar covalent bonds (b)
S.C. By the end of this lesson you should be able to use electronegativities to explain the difference between pure covalent and polar covalent bonds explain the term permanent dipole use the data book to assign δ+ and δ+ partial charges on atoms

POLAR COVALENT BONDS Two types of covalent bond can be formed: Pure covalent (or non-polar covalent) Polar covalent.

Covalent Bonding Picture a tug-of-war:
If both teams pull with the same force the mid-point of the rope will not move.

Pure Covalent Bond H e H This even sharing of the rope can be compared to a pure covalent bond, where the bonding pair of electrons are held at the mid-point between the nuclei of the bonding atoms.

Covalent Bonding What if it was an uneven tug-of-war?
The team on the right are far stronger, so will pull the rope harder and the mid-point of the rope will move to the right.

Polar Covalent Bond A polar covalent bond is a bond formed when the shared pair of electrons in a covalent bond are not shared equally. This is due to different elements having different electronegativities.

I Polar Covalent Bond δ- δ+ H e.g. Hydrogen Iodide e
If hydrogen iodide contained a pure covalent bond, the electrons would be shared equally as shown above. This makes iodine slightly negative and hydrogen slightly positive. This is known as a dipole. However, iodine has a higher electronegativity and pulls the bonding electrons towards itself (winning the tug-of-war)

H - H PURE COVALENT (OR NON-POLAR COVALENT)
A pure covalent bond is formed when the atoms involved in the bond have an equal share of the bonding electrons. They have the same electronegativity. 2.2 2.2 H - H Complete workbook

POLAR COVALENT When atoms with different electronegativity values join together, a polar covalent bond is formed. The dipole produced is permanent.

A polar covalent bond is a bond where the electrons are not shared equally, one atom in the bond has a greater attraction than the other for the bonded electrons. Complete workbook

L.I. To learn about the bonding continuum (c)
S.C. By the end of this lesson you should be able to explain the relationship between differences in electronegativities and type of bonding use data from the properties of compounds to deduce the type of bonding and structure

Electronegativity difference Actual difference in electronegativity
BONDING CONTINUUM Electronegativity Difference and Bond Type: Electronegativity difference Bond type Example Actual difference in electronegativity covalent (non polar) H-H 0.0 covalent (polar) H-Cl 0.9 H2O 0.7 covalent (very polar) H-F 1.9 .2.0 ionic NaCl 2.1

The greater the difference in electronegativity the greater the polarity between two bonding atoms and the more ionic in character. A bonding continuum can be used to help us understand the differences in bonding. Complete workbook

Bonding Continuum “Covalent compounds are formed by non-metals only”
IS NOT AN ABSOLUTE LAW! Some compounds break this rule….

Tin(IV)iodide – covalent or ionic?
Predict its melting point. Complete workbook Melting point of tin(IV)iodide is 143oC. Tin electronegativity of 1.8 Iodine has electronegativity of 2.6 Molecule contains polar covalent bonds, but the symmetry cancels out the dipoles, therfore only weak London’s forces so low melting an boiling point.

L.I. To learn about intermolecular forces (d)
S.C. By the end of this lesson you should be able to explain the difference between intramolecular and intermolecular forces name the three types of van der Waals forces explain how London dispersion forces arise

INTERMOLECULAR FORCES
Intramolecular bonds are bond between atoms within a molecular – covalent bond. Intermolecular bonds are bonds which occur between molecules. Intermolecular bonds are called van der Waals’ forces. They are named after the Dutch Chemist Johannes Diderik van der Waals.

There are three types of van der Waals’ forces:
London dispersion forces Dipole-dipole interactions (permanent dipoles) Hydrogen bonding

1. London Dispersion Forces
Electrons ‘wobble’ and temporary dipole occur. These cause induced dipoles on other atoms. The attraction between atom resulting from the temporary dipoles are known as London dispersion forces.

London Dispersion Forces
London dispersion forces are very weak attractive forces.

L.I. To learn about intermolecular forces
S.C. By the end of this lesson you should be able to explain how dipole-dipole interactions arise describe a test that can be used to determine if a molecule is polar explain the connection between symmetry and polarity describe how most hydrocarbons are classified in terms of polarity

2. Dipole-Dipole Interactions
Polar Molecules Are all molecules with polar bonds polar? H O -- + Water has a polar covalent bonding between O and H. Is water polar?

see scholar animation on polarity test
Complete activity – testing polarity, and complete the table see scholar animation on polarity test

Asymmetrical molecules e.g. H2O are POLAR
Symmetry and polarity Asymmetrical molecules e.g. H2O are POLAR In an asymmetrical molecule there is a permanent dipole workbook activity

Symmetrical molecules e.g. CCl4 are NON-POLAR
In a completely symmetrical molecule the polarities cancel each other out so there is no permanent dipole. workbook activity

Most hydrocarbons are non-polar

H – H H - H Dipole-Dipole Interactions
Dipole-dipole interactions are intermolecular forces which occur between polar molecules. polar molecule Dipole – dipole interactions H – H H - H non - polar molecule London dispersion forces

Dipole-dipole interactions are stronger than London dispersion forces.
Polar molecules have higher melting and boiling points than non-polar molecules. b.p. 56 o C b.p. -1 o C polar molecule non - polar molecule

L.I. To learn about intermolecular forces
S.C. By the end of this lesson you should be able to explain how H-bonds arise what is necessary in a molecule to allow H-bonds to arise

Hydrogen Bonding Hydrogen bonding is a special type of dipole-dipole interaction involving; H-N, H-O or H-F bonds workbook activity For H-bonds to exist between molecules the molecules must have a strong polar covalent bond 2. the polar covalent bond must be between a hydrogen atom and either nitrogen, fluorine and oxygen (NOF)

Hydrogen bonds are stronger than normal dipole-dipole interactions and London dispersion forces.
Molecules which contain hydrogen bonding have much higher melting and boiling points than those with dipole- dipole interactions or London dispersion forces.

L.I. To learn about relating properties of compounds to
intermolecular forces (e) S.C. By the end of this lesson you should be able to explain the connection between size of molecule and strength of London dispersion forces describe the evidence that proves the existence of permanent dipole-dipole interactions. H-bonds explain why ice is less dense that water explain how intermolecular forces affect bpts, mpts, viscosity and solubility explain the term “like dissolves like”

RELATING PROPERTIES TO INTERMOLECULAR BONDING
1. Melting and boiling points give an indication of the amount of energy needed to overcome the van der Waal’s forces between molecules. London dispersion forces – between non polar molecules workbook activity Alkane name Alkane formula boiling point (oC) pentane hexane heptane octane

From the table we can see that as the molecular size increases (number of electron shells) the boiling point increase and so the strength of the London dispersion forces increase.

Dipole-dipole interactions – between polar molecules
b.p. 56 o C b.p. -1 o C polar molecule non - polar molecule Formula mass 58 Formula mass 58 workbook activity For polar molecules, the melting and boiling points are higher than those of non-polar molecules. More energy is needed to overcome the dipole-dipole interactions between polar molecules. Molecules with a similar mass are used to allow us to ignore the London Dispersion Forces.

Hydrogen bonding For polar molecules that contain H-N, H-O or H-F the melting points are related to the strong hydrogen bonds between molecules. workbook activity For polar molecules which contain H-N, H-O or H-F bonds the melting point and boiling points are higher than those of other polar and non-polar molecules. More energy is required to overcome the strong hydrogen bond.

NH3, has a higher boiling point than expected.

H2O has a higher boiling point than expected.

HF has a higher boiling point than expected.

Evidence of Hydrogen Bonding
H2O HF NH3

Strength of van der Waals’ forces:
Hydrogen > Dipole-dipole > London dispersion forces

“like dissolves like” 2. SOLUBILITY
polar solvents, such as water, will dissolve polar and ionic solutes Na+Cl- (s) Na+(aq) + Cl-(aq)

Dissolving in Water

3. VISCOSITY scholar animation

The stronger the van der Waals’ forces between a liquid are, the more viscous it will be.
Liquids containing hydrogen bonding will be more viscous than molecules containing dipole-dipole interactions or London dispersion forces The most viscous liquid is propane-1,2,3-triol.

Why does the ice float in Mrs Brown’s gin and tonic?
Why do icebergs float? Why do fish in ponds not die in winter when the water freezes?

DENSITY OF WATER/ICE look at candle wax and ice The density of ice is unusual. Normally solids sink in their liquids. On cooling, water contracts but at 4oC it expands. At freezing point an open structure exists as a result of H-bonds. So ice floats!!

SUMMARY

Section 2 Periodicity Bonding in the Elements 1-20 (a)

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Section 2 Periodicity Bonding in the Elements 1-20 (a)

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