The origin of chemistry can be traced to alchemy, or the art of converting metals like copper to

gold.

Chemistry

Ionic Bonding

Bonds

What is a chemical bond? Electrostatic forces of attractions (between 2 atoms)

between the nuclei of one atom and the electrons of the other atom.

Why do atoms bond?Three kinds of bonding:

Ionic Covalent Metallic

Noble Gas

Also called the inert gases or rare gases: He, Ne, Ar, Kr, Xe and Rn.

Noble gases are unreactive.Exist as individual atoms, monatomic.Is there a need to bond?Are there noble gas compounds?

Noble gas compounds

Examples: Xenon tetrafluoride (XeF4) Xenon tetroxide (XeO4) Krypton difluoride (KrF2) Radon difluoride (RnF2) Xenon trioxide (XeO2)

Notice what these compounds contain?

Noble Gas Structure

Duplet or octet configurations are most stable.

Also known as a noble gas configuration.Common feature – fully filled valence

electron shell

He Ne

Why atoms combine

Atoms WANT to achieve the noble gas configuration.

How do atoms achieve the noble gas structure? Transferring or sharing electrons.

Recap – ions

Normally, an atom is electrically neutral number of protons = number of electrons

An ion is formed when an atom loses or gains electrons.

An ion is a charged particle formed from an atom or a group of atoms by the loss or gain of electrons.

Cations

Positive ions (cations) are formed by removing/ losing electrons from atoms.

Loss of electrons tends to occur in atoms with few valence electrons (e.g. 1, 2 & 3)

Notice that these are METALS.

Cations

Electronic configuration: 1s22s22p63s1 1s22s22p6

Na Na++ e-

Na atom Na+ ion

Number of protons

11 11

Number of electrons

11 10

Na Na + + e-

Anions

Negative ions (cations) are formed by gaining electrons from atoms.

Gain of electrons tends to occur in atoms with larger number of valence electrons (e.g. 5, 6 & 7)

Notice that these are NON-METALS.

Anions

Electronic configuration: 1s22s22p63s23p5 1s22s22p63s23p6

Cl + e- Cl-

Cl + e-

Cl atom Cl- ion

Number of protons

17 17

Number of electrons

17 18

Cl-

Discuss

Consider sodium atom and sodium ion. Which do you think is bigger and why?

Na Na+

106 pm186 pm

Consider chlorine atom and chloride ion. Which do you think is bigger and why?

Cl-Cl

100 pm 181 pm

Trends in radii

Atoms are always larger than any of their cations.

Atoms are always smaller than any of their anions.

http://chewtychem.wiki.hci.edu.sg/Ionic+Bonding

Quickcheck

Is Mg+ or Mg2+ bigger?Is O- or O2- bigger?

Ionic Bond

An ionic bond is a chemical bond formed by the electrostatic attraction between the positive and negative ions.

Formation of an ionic bond can be viewed as a transfer of electrons from a metallic atom to a non-metallic atom.

Both will gain a duplet or octet configuration.

Video

Link to video

Ionic bond

Na Na + + e-

Cl + e- Cl-

Two processes occuring:

Na

Na1s22s22p63s1

Cl1s22s22p63s23p5

Cl

+

Na+

1s22s22p6

-

Cl-

1s22s22p63s23p6

Dot and Cross Diagram

Neon:1s22s22p6

Argon:1s22s22p63s23p6

Isoelectronic with

Isoelectronic with

Dot and Cross Diagram

Cl-

Cl-

Mg2+

Al3+

O2-

2 3

Magnesium Chloride

Aluminium Oxide

Is this possible?

Na + 7e- Na7-

Cl + 7e-Cl7+

1s22s22p63s1 1s22s22p63s23p6 (Ar)

1s22s22p63s23p5 1s22s22p6 (Ne)

Chemical formulae

Take for example the following: An ionic compound made of Magnesium and

Fluorine.

Mg2+

F-

Mg2+

F-F-

Chemical formulae

For Magnesium fluoride, the ions present are Mg2+ and F-.

Mg2+ has 2 positive chargesF- has 1 negative charge. To make the overall compound electrically

neutral there must be two F- to balance one Mg2+.

The formula is MgF2.

Chemical formulae

Mg2+ F-

Mg1 F2 MgF2

Na+ OH-

Na1 OH1 NaOH

Charge on ion Name of ion Formula+1 Sodium Na+

Potassium K+

Silver Ag+

Ammonium NH4+

Hydrogen H+

+2 Magnesium Mg2+

Calcium Ca2+

Iron(II) Fe2+

Zinc Zn2+

+3 Aluminium Al3+

Iron (III) Fe3+

-1 Chloride Cl-

Fluoride F-

Hydroxide OH-

Nitrate NO3-

-2 Carbonate CO32-

Oxide O2-

Sulfate SO42-

Find the chemical formula

1. Magnesium oxide2. Zinc chloride3. Calcium hydroxide4. Iron(II) fluoride5. Iron(III) sulfate6. Ammonium nitrate7. Silver chloride8. Potassium iodide9. Manganese(IV) oxide

General knowledge

General knowledge正离子 = cation负离子 = anion离子键 = ionic bonding共价键 = covalent bonding金属键 = metallic bondingNow you can hao4 lian4 to your friends!

Ionic lattices

In ionic compounds, ions are held in fixed positions in an orderly arrangement by strong electrostatic forces (or ionic bonds) between the cations and anions.

Attractions are maximised in this structure while repulsions are minimised.

How many Cl- surround one Na+ and vice versa?The coordination number is the number of nearest

neighbours (atoms, ions or molecules). What are the coordination numbers of Na+ and Cl-?

Ionic lattices

Ionic Lattices

Different ionic lattice structures exist.http://www.avogadro.co.uk/structure/

chemstruc/ionic/g-ionic.htm

Properties

High Melting and Boiling PointsHard and brittleConducts electricity when dissolved in water

or when moltenMany ionic compounds are soluble in water

or polar solvents like alcohol, but insoluble in most non-polar solvents like hexane

WHY?

High Boiling and Melting Point

due to breaking of strong ionic

bonds

Conductor of electricity in liquid state and when in

solution due to presence of free moving charges.

Is brittle due to repulsion between similarly charged ions

Properties - NaCl

Linus Pauling

Only person to have won two unshared Nobel Prizes (Chemistry and Peace)

The Nature of the Chemical Bond and the Structure of Molecules and Crystals

Pauling Electronegativity Scale

Electronegativity

Every atom has an attraction for the electrons shared in a bond. Why?

Click this link for the electronegativity table.http://chemwiki.ucdavis.edu/@api/deki/files/4756/=electronegativity_chart.png

Degree of attraction can be related to the electronegativity of the atom.

The higher the electronegativity, the more the electrons in a chemical bond are attracted to the atom.

Electronegativity

What patterns/ trends do you see?Metals generally have low electronegativity values, while non-metals have higher electronegativity values.Electronegativity values generally decrease down the group.Electronegativity generally increases across a period.

The high affinity for electrons of fluorine leads it to direct reactions with all other elements in which the reaction has been attempted, except for helium and neon.

Predicting Ionic compounds

What do you notice about ionic compounds?What are they made up of?

METALLIC BONDS

Metallic Bonding

Metallic names

Ever wondered why some metals have weird symbols in the Periodic table?

E.g. Au (aurum) and Hg (hydrargyrum)Aurum actually means ‘shining dawn’.Hydragyrum means watery silver.

Comparing Ionic and Metallic Bonding

Li Cl Li Li?+ -

Metallic bonding (I)

Metals exist as giant structures too.In metals, atoms are packed closely together

in regular three-dimensional patterns to form a giant lattice.

Metallic bonding (II)

Image originally created by IBM Corporation.

Platinum atoms

Metallic bonding (III)

Li+

e

Li+ Li+

Li+ Li+ Li+

Li+ Li+ Li+

Li+

Li+

Li+

e e e

e e e e

e e e e

Mobile/ delocalised electrons.

‘Sea of electrons’ surrounding the positively charged metal cations.

Opposites attract i.e. every positive ion is attracted to the ‘sea of electrons’.

Metallic bonding (III)

Forces of attraction between positively charged ions and negatively charged electrons – metallic bonding.

Only found in metals! Not ionic or covalent bonding!

PROPERTIES OF METALS

Metallic Bonding

Malleability and Ductility (I)

A malleable substance is one which can be bent or hammered out of shape without breaking.

A ductile substance is one which can be stretched or drawn into thin wires without breaking.

Malleability and Ductility (II)

The following video shows the malleability of metals.

Metals can be forged into different shapes and sizes by beating and hammering

Especially useful in making tools and machines.

Malleability and Ductility (III)

The following video shows the ductility of metals.

Notice that an iron strip can be twisted many times before it finally breaks.

Useful in making wires and cables.Why are metals malleable and ductile?

Malleability and Ductility (IV)

Metal atoms are arranged in orderly layers.Application of a force causes metal atoms to

slide over each other easily. Why does the metal not break easily?Look at the following animation to

understand better.

Malleability and Ductility (V)

What is this property useful for?

Malleability Ductility

Sculpting

Architecture Machinery

Tubing

Aircraft

Wire

Melting and Boiling Points (I)

Metals generally have high melting and boiling points.

Most metals are packed closely together and the strong forces of attraction between the positively charged metal ions and the ‘sea of electrons’ result in strong metallic bonding.

A lot of energy is required to separate the metal atoms.

Melting and Boiling Points (II)

Can you name some exceptions?Just for fun: The following metals might melt

at Singapore’s room temperature! Francium (27oC), Caesium (28oC) and Gallium (30oC).

Mercury

Melting and Boiling Points (III)

What is this property useful for?

High melting point

Lights

Electrical appliances

Aircraft

Electrical Conductivity (I)

‘Sea of electrons’ surround metal cations.Mobile or delocalised electrons in the metal

structure allows conduction of electricity.

Electrical Conductivity (II)

Metals conduct electricity due to mobile delocalised electrons.

Can they conduct electricity in the molten state?

Can ionic compounds conduct electricity in the solid state?

Using electrical conductivity, how can we determine whether a substance is an ionic compound or metal?

Electrical Conductivity (IV)

Sodium chloride

Iron

Solid Does not conduct electricity

Conducts electricity

Molten Conducts electricity

Conducts electricity

Aqueous Conducts electricity

-(Insoluble)

Electrical Conductivity (V)

What is this property useful for?

Good conductor

of electricity

WiresElectrical appliances

Lightningrod

Heat Conductivity (I)

Metals conduct heat well due to mobile delocalised electrons as well.

When you heat one end of a metal, delocalised electrons gain energy, move faster, and collide with neighbouring electrons, thus transferring heat from one end to the other.

Look at the following animation to observe what happens when metal is heated.

Heat Conductivity (II)

Compare the movement of free electrons when a metal conducts heat and electricity. What are the differences?

Heat Conductivity (III)

What is this property useful for?

Good conductor

of heatAircon

Heating coil

Refrigerator

Metals as a whole

In general, metals… Reason(s)

Have high melting and boiling points.

Strong metallic bonding.

Are good conductors of electricity. Mobile/ delocalised electron that can carry electric charges.

Are good conductors of heat. Mobile/ delocalised electrons that collide and transfer heat.

Are malleable and ductile. Atoms arranged in orderly layers that can slide past one another.

Have high densities. For you to find out.

Are shiny in appearance. For those who are interested.

Covalent Bonding

Covalent Bonding

Consider the following (chlorine):We know that chlorine is often written as Cl2.Is this possible?

Cl-

Cl-

Covalent Bonding

Covalent Bonding is the mutual electrostatic attraction between the nuclei of atoms and their shared electrons.

Normally occurs between 2 non-metals only.For example, when two Hydrogen atoms meet, they

will each share one electron to get a duplet configuration.

Covalent Bonding

Let’s take a look at Hydrogen.As two H atoms approach, the electron on

each atom is attracted to the nucleus of the other, i.e. there are forces of attraction (between what?).

What happens if they are too close?Because of repulsion between nuclei and

attractions between electrons and nuclei, there is a distance between the two atoms where the molecule is most stable. This distance is called the bond length.

Lennard-Jones Potential

H H

Distance between two atoms

H H

Bond lengt

h

http://upload.wikimedia.org/wikipedia/commons/5/5a/12-6-Lennard-Jones-Potential.png

Covalent Bonding

Bond length or bond distance is the average distance between nuclei of two bonded atoms in a molecule.

Covalent Bonding

Now take a look at two chlorine atoms. How do they share electrons?

Cl Cl

How many electrons are there around each chlorine atom now?

Covalent bonding

Covalent bonds exist for compounds too.What about the water molecule?

O H

H

Covalent bonding

How many electrons does an oxygen atom require?

O O

In this case, oxygen requires 2 electrons each. They can achieve the noble gas configuration by sharing 2 electrons each. The resulting bond is a double bond.

Nitrogen

Try drawing the covalent molecule, nitrogen.

N N

In this case, nitrogen requires 3 electrons each. They can achieve the noble gas configuration by sharing 3 electrons each. The resulting bond is a triple bond.

Try these!

NH3

CO2

SiCl4

CH3I

Structural Formula

Molecule Chemical Formula Structural Formula

Hydrogen H2 H – H

Oxygen O2 O = O

Nitrogen N2 N N

Water H2O H – O – H

Methane CH4

H – C – H

H

H

Polyatomic ions

O- N+

O

O-

Nitrate (NO3-)

O- CO

O-

Carbonate (CO32-)

SO

O-

Sulfate (SO42-)

O

O-

O- H

Hydroxide (OH-)

Sigma Bonds

Head-on overlap of orbitals (s, p or d-orbitals)The resultant electron cloud is called a sigma

bondStrongest kind of covalent bond

1s 1s

H atom H atomH2

molecule

pp

Pi bond

Side-on overlap of orbitals (p or d orbitals)Weaker than sigma bond because of less

overlap.

Can a molecule with only single bonds have pi bonds? E.g. Cl2?

Sigma bond vs Pi bond

Pi bonds can only form after a sigma bond is formed.

Sigma bonds can be found in all covalent compounds.

Pi bonds are only found in double bonds or triple bonds.

Identify the bonds in the given molecules

C C

H

HH

H

F F

O O

N N

C CH H

Hydrogen

Let’s take a look at the hydrogen molecule.

H H = 2.1 =

2.1

Each H atom has the same attraction for the shared electrons .

Thus the electron density is evenly distributed over the whole molecule.

The molecule is said to be non-polar.

Hydrogen chloride

Now consider the hydrogen chloride molecule.

H Cl = 2.1 =

3.0

Chlorine is much more electronegative than hydrogen, hence, it attracts the bonding pair of electrons more strongly.

Thus the electron density of the bond in HCl is pulled towards the Cl end of the molecule.

Hydrogen chloride

Now consider the hydrogen chloride molecule.

H Cl = 2.1 =

3.0

This results in a separation of charge.The molecule is said to be polar and the

bond is a polar covalent bond.The molecule is said to have a permanent

dipole moment.

+ -

Polar bonds

What happens if one atom is very much more electronegative than the other?

+ -

+ -

If one atom is very much more electronegative than the other, it pulls away the bonding electrons such that it becomes an ionic compound.

Fluorine

Fluorine is so electronegative that it can bond to almost all atoms, even some noble gas atoms.

It does so because it has very strong attractions for electrons, such that even the electrons of noble gas compounds can be attracted by fluorine, thus forming a bond!

Electronegativity and bond polarity

As a rule of thumb: If difference in electronegativity ranges from 0 to 0.5,

it is considered a non-polar covalent bond. If difference in electronegativity ranges from 0.6 to

1.6, it is considered a polar covalent bond. If difference in electronegativity is above 2.0, it is

considered an ionic bond. For 1.7 to 1.9, if a metal and non-metal is present, it is

considered ionic; if two non-metals are involved, it is considered polar covalent.

Try these!

Predict whether the following are covalent or ionic. sodium bromide hydrogen fluoride aluminium oxide aluminium chloride beryllium chloride caesium fluoride

Overall dipole

Vector sum of all dipole moments on the molecule

O

HH-

+ +

Individual dipoleOverall dipole

Polar or non-polar?

Not all molecules with a polar bond has a dipole!

For example, carbon dioxide.

O=C=O+ -

Cancel each other

-

Try these!

Predict whether the following are polar molecules. NH3

H2S CO CH4

CH3Cl

Bond strength

The strength of the covalent bond between atoms is known as the bond strength.

The higher the bond strength, the more energy is required to break the bond.

Which molecule has the highest bond strength, nitrogen, oxygen or hydrogen?

Bond strength

一根竹竿容易弯，一把筷子难折断Strengthtriple bond>Strengthdouble bond>Strength

single bond

What is the relationship between bond length and bond strength?

Bond Bond Length Bond Energy

C – C 154 pm 348 kJ/mol

C = C 134 pm 614 kJ/mol

C C 120 pm 839 kJ/mol

266 kJ/mol

225 kJ/mol

Bond strength

In general, the shorter the bond length, the greater the bond strength.

A number of factors affect bond strength. These include, number of bonds, size of atoms, electronegativity of atoms etc.

Bond Bond Length

Bond Energy

C – F 135 pm

488 kJ/mol

C – O 143 pm

360 kJ/mol

C – N 147 pm

308 kJ/mol

Across a period: Down a group:Bond Bond

Length

Bond Energy

C – F 135 pm

488 kJ/mol

C – Cl 177 pm

330 kJ/mol

C – Br 194 pm

288 kJ/mol

C – I 214 pm

216 kJ/mol

Intermolecular Forces

So far the attractive forces we have covered hold atoms or ions together (intramolecular). What holds MOLECULES together (intermolecular)? van der Waals’ Forces hydrogen-bonding

van der Waals’ forces

3 kinds Permanent dipole-permanent dipole Permanent dipole-induced dipole Instantaneous dipole-induced dipole

Permanent dipole-permanent dipole

Consider the hydrochloride molecule

H Cl H Cl

+ - + -

If two HCl molecules approach each other, they will tend to arrange themselves such that the positive end of one molecule is close to the negative end of the other.

dipole-dipole attractions

Permanent dipole-permanent dipole

Found in molecules with a permanent dipole moment i.e. a polar molecule.

Stronger the dipole moment, stronger the intermolecular force. Why?

Much weaker than a covalent bond.

H Cl H Cl

+ - + -

Weak dipole-dipole attractions(intermolecular)

Strong covalent bond

(intramolecular)

Polar molecules

When exposed to an electric field, polar molecules align themselves to the positive and negative terminals

http://witcombe.sbc.edu/water/chemistrystructure.html

Instantaneous dipole-induced dipole

What then holds non-polar molecules together?

Instantaneous dipole-induced dipole, also known as London dispersion forces.

Exists between ALL molecules and atoms.Only kind of intermolecular attraction

possible between non-polar molecules.

Instantaneous dipole-induced dipole

Electron cloud distribution is symmetrical

Electron cloud distribution

becomes unsymmetrical for

an instant

Neighbouring electron cloud experiences an induced dipole

+- +- +-

NOTE: This kind of attraction is very short-lived because electrons are always moving and the dipoles will disappear very quickly!

Instantaneous dipole-induced dipole

What does this London dispersion forces depend on? Number of electrons/ Electron cloud size Shape/ Surface area

B.P. vs Number of electrons

0 10 20 30 40 50 60 70 80 90 100

-300

-250

-200

-150

-100

-50

0

Boiling point/ oC vs Number of electrons

Number of electrons

Boil

ing p

oin

t/ o

C

He

Ne

Ar

Kr

Xe

Rn

Helium

Has a boiling point of -268.93oC. Lowest among all elements!

Remains at a liquid even at -273.15oC (absolute zero)

Can only become a solid at 25 bar pressure at -272.2oC

Because of the weak van der Waals’ forces!

Recap

3 kinds Permanent dipole-permanent dipole Permanent dipole-induced dipole Instantaneous dipole-induced dipole

Recap

How does pd-pd and id-id work?Under what conditions does pd-pd happen?Does id-id work on only non-polar molecules?

Instantaneous dipole-induced dipole

Instantaneous dipole-induced dipole becomes more important than permanent dipole-permanent dipole in determining the boiling point as the molecule or atom’s electron cloud increases.

Topics

van der Waals’ in actionHydrogen-bondingProperties of Hydrogen-bonded moleculesHydrogen-bonding in real lifeProperties of simple covalent molecules(Properties of giant covalent structures)

Geckos and Spiders

How do they stay on walls and even ceilings?Let’s take a look at this video.Then look at this website.

Geckos

Geckos - Geckos have millions of setae--microscopic hairs on the bottom of their feet. These tiny setae are only as long as two diameters of a human hair. That's 100 millionth of a meter long. Each seta ends with 1,000 even tinier pads at the tip. These tips, called spatulae, are only 200 billionths of a meter wide-below the wavelength of visible light.

A single seta can lift the weight of an ant. A million setae, which could easily fit onto the area of a dime, could lift a 45-pound child. If a gecko used all of its setae at the same time, it could support 280 pounds.

Geckos cannot stick to teflon (non-stick coating on cooking pans). Go and find out why!

http://www.sciencedaily.com/releases/2002/08/020828063412.htm

Spiders

On each of the spider's feet there are hair-like tufts, called scopulae, … it was discovered that a single scopula is itself composed of many, many, much smaller, single hairs… The number of setules per foot is estimated to be 78,000 each, and since spiders have eight feet, they have upwards of 600,000 individual points of contact with any given surface.

The total adhesive force is extremely powerful, up to 170 times the weight of the spider, if all eight legs are in contact.

http://www.istl.org/05-summer/article3.html

Hmmm…

If scientists manage to come up with materials that are like a gecko’s legs…,

What uses can they be used for?What issues will crop up?

Intermolecular

Are the only intermolecular forces van der Waals’?

Group IV hydrogen compounds

CH4

SiH4

GeH4

SnH4

0 20 40 60 80 100 120 140

-190

-140

-90

-40

10

60

110

Group IV

Group IV

Molecular mass

Boil

ing

Poin

t/ o

C

CH4 has the lowest boiling point as it is the smallest molecule.

Larger molecules have greater number of electrons and thus, greater intermolecular forces. As a result, they have higher boiling points.

Group V hydrogen compounds

NH3

PH3

AsH3

SbH3

0 20 40 60 80 100 120 140

-140

-90

-40

10

60

110

Group V

Group V

Molecular Mass

Boil

ing p

oin

t/ o

C

NH3 is the smallest molecule and is expected to have the lowest boiling point. Some additional force must be present!

Group VI hydrogen compounds

0 20 40 60 80 100 120 140

-70

-50

-30

-10

10

30

50

70

90

110

Group VI

Group VI

Molecular Mass

Boil

ing p

oin

t/ o

C

H2O

H2S

H2Se

H2Te

H2O is the smallest molecule and is expected to have the lowest boiling point. Some additional force must be present!

Group VII hydrogen compounds

HF

HCl

HBr

HI

0 20 40 60 80 100 120 140

-140

-90

-40

10

60

110

Group VII

Group VII

Molecular Mass

Boil

ing

poin

t/ o

C

HF is the smallest molecule and is expected to have the lowest boiling point. Some additional force must be present!

Hydrogen compounds

0 20 40 60 80 100 120 140

-190

-140

-90

-40

10

60

110

Hydrogen compounds

Group IVGroup VGroup VIGroup VII

Molecular mass

Boil

ing P

oin

t/ o

C

CH4

SiH4

GeH4

SnH4

NH3

PH3

AsH3

SbH3

H2O

H2S

H2Se

H2Te

HF

HCl HBr

HI

The additional forces existing between NH3, H2O and HF molecules are called hydrogen bonds.

Hydrogen bonding

Hydrogen bonding occurs only when molecules contain an H atom covalently bonded to a very small, highly electronegative atom with lone pairs of electrons, i.e. F, O and H.

How H-bonding works

When H is covalently bonded to an extremely electronegative atom, F, O or N, the electronegative atom will attract the electron cloud strongly, leaving the H nucleus almost bare.

Thus when another molecule containing an F, O or N with lone pair of electrons approaches, it can get very close to the H atom, thus the intermolecular force is much stronger.

This accounts for the high boiling points of water, ammonia and hydrogen fluoride.

Conditions for H-bonding

The H-atom must be covalently bonded to either N, O or F, the 3 most electronegative elements.

There must be a lone pair on N, O or F of the neighbouring molecule which can attract the partial positive charge on the H-atom.

Hydrogen bonding

HO

H

H

OH

H

OH

H

OH

H

OH

H

OH

Strong (intramolecular

) covalent bonds

Weaker (intermolecular

) Hydrogen bonds

Why must it be hydrogen???

Strength of H-bonds

Strength of H bonds: H – F > H – O > H – N Why does water have a much higher boiling

point than HF when HF forms stronger H-bonds?

Hydrogen Bonding in water

Why does ice float on water?When most substances freeze, the particles

are closer to one another as they are in the solid state as compared to the liquid state.

Water, however, has maximum density at 4oC. Recall that density = mass/volume. When water cools down from 4oC to 0oC, the formation of solid ice actually forces the water molecules to be fixed in a tetrahedral shape.

Click here for animation.

Hydrogen Bonding in water

Similarly when ice melts, the solid structure actually collapses, so that the water molecules are closer together. They are closest at the temperature of 4oC. Why is this important for freshwater fish in winter?

Ice (00C)

Water (10C)

Water (20C)

Water (30C)

Water (40C)

http://v.ku6.com/show/SXoq6mSJK2Ysgaga.html

Solubility in water

Polar molecules are able to dissolve in water, which is a highly polar molecule, due to H-bonding.

C C

H

H

H

H

H

O H

Ethanol

Solubility in water

Sugar (sucrose) is highly polar and can dissolve very well in water.

Carboxylic acids like ethanoic acid (vinegar) are able to form H-bonds as well in water.

C C

H

H

HO

OH

Sucrose Ethanoic acid

Solubility in water

Some ionic compounds dissolve in water as well.

See animation.

H-bonds in life

DNA (deoxyribonucleic acid) consists of two strands of polymers (very long molecules).

The 2 strands are held together by H-bonds.H-bonds can be broken by heating to high

temperatures.

http://en.wikipedia.org/wiki/File:DNA_chemical_structure.svghttp://en.wikipedia.org/wiki/File:DNA_orbit_animated_static_thumb.png

H-bonds in Life

Proteins are made up of amino acids. Primary, secondary, tertiary and quaternary

structure. Secondary – alpha helix and beta-pleated

sheet.

Alpha Helix Beta-pleated Sheet

Properties of covalent molecules

We have learnt that van der waals’ forces and hydrogen bonds hold molecules together. These forces are relatively weak as compared to covalent bonds.

What are the properties of such compounds?

Melting point/boiling point

Simple covalent molecules have high volatility, i.e. they have low boiling point.

Note: No breaking of covalent bonds required!

Reasons for low melting/boiling point: Strong covalent bonds within the molecules

but weak van der Waals’ forces between the molecules

Little energy is required to overcome the weak intermolecular forces.

Melting point/boiling point

In an iodine molecule, the two atoms are held by a strong covalent bond.

Weak van der Waals’ forces hold the iodine molecules together.

When heat is supplied, the weak intermolecular forces break and iodine sublimes.

Iodine molecule,

I2

Strong covalent bond

Weak van der Waals’ forces between molecules of iodine

Solubility

Simple covalent molecules are generally insoluble in water or polar solvents UNLESS they are able to form hydrogen bonds or can dissociate (later).

Rule: “Like dissolves like.”For instance if you try to dissolve a non-polar

molecule like oil in water, water molecules will prefer to bond to water molecules (they have H-bonds) whereas the oil molecules will prefer clump together because energy is needed to break the H-bonds. Hence oil does not dissolve in water.

Electrical conductivity

Simple covalent molecules do not conduct electricity.

Absence of free moving electrons/ ions.

N N

What about water?

Electrical conductivity

Some simple covalent molecules (often acids or bases) when dissolved in water produce free moving ions in the solution which can conduct electricity.

H Cl

H ClH+ Cl-

Checklist – Covalent or not?

Iodine? Sand?Are they covalent?

Do they have similar physical properties?

Commonly formed between non-metals Form bonds by sharing electrons After bonding, each atom achieves noble

gas configuration

Note: Sand is made up mostly of silicon dioxide, SiO2 and is a main component of glass.

Iodine vs Glass

What do you think happens to glass (75% SiO2)?What about iodine?

Watch the Youtube Video on heating iodine in glass: http://www.youtube.com/watch?v=E-fs9OwE9Y0&NR=1

After gentle heating, Physical form of glass remains intact. Iodine sublimed (Solid → Gas)