Properties of MatterPart 4

Bonding

● investigate the role of electronegativity in determining the ionic or covalent nature of bonds between atoms

● investigate the differences between ionic and covalent compounds through:

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– using nomenclature, valency and chemical formulae (including Lewis dot diagrams) (ACSCH029)

– examining the spectrum of bonds between atoms with varying degrees of polarity with respect to their constituent elements’ positions on the periodic table

– modelling the shapes of molecular substances (ACSCH056, ACSCH057)● investigate elements that possess the physical property of allotropy ● investigate the different chemical structures of atoms and elements, including but not

limited to:– ionic networks– covalent lattices (including diamond and silicon dioxide)– covalent networks– metallic structure

● explore the similarities and differences between the nature of intermolecular and intramolecular bonds and the strength of the forces associated with each, in order to explain the:– physical properties of elements– physical properties of compounds (ACSCH020, ACSCH055, ACSCH058)

● investigate the role of electronegativity in determining the ionic or covalent nature of bonds between atoms

● examining the spectrum of bonds between atoms with varying degrees of polarity with respect to their constituent elements’ positions on the periodic table

investigate the role of electronegativity in determining the ionic or covalent nature of bonds between atoms

Chemical Bonds

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Noble gas configuration is the electron configuration of noble gases. The basis of

all chemical reactions is the tendency of chemical elements to acquire stability.

Atoms form bonds to achieve “the Noble Gas Configuration”.

Electronegativity

The electronegativity of an element is a measure of the ability of the atom of that

element to attract bonding electrons towards itself when it forms compounds.

Electronegativity increases across a period and decreases down a group.

If the difference in electronegativity between two elements is greater than 1.5 the

compound will be ionic. If it is less than 1.5 it will be covalent.

Electronegativity Difference Type of Bond Formed

0.0 to 0.2 non-polar covalent (pure covalent)

0.3 to 1.4 polar covalent

> 1.5 ionic

Pauling’s Electronegativities for some common elements.

Element Electronegativity

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(Pauling Scale)

H 2.2

He 0

Li 1.0

Be 1.6

B 2.0

C 2.6

N 3.0

O 3.4

F 4.0

Ne 0

Na 0.9

Mg 1.3

Al 1.6

Si 1.9

P 2.2

S 2.6

Cl 3.2

Ar 0

K 0.8

Ca 1.0

Fr 0.7

● investigate the differences between ionic and covalent compounds through:– using nomenclature, valency and chemical formulae (including Lewis dot

diagrams) (ACSCH029)

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Nomenclature:Ionic Compounds: As previously demonstrated.

Covalent binary compounds:

Prefixes

mono = 1; di = 2; tri = 3; tetra = 4; penta = 5; hexa = 6

For example:

Compound Chemical Formula

Carbon Monoxide CO

Carbon Dioxide CO2

Sulfur Trioxide SO3

Carbon Tetrachloride CCl4

Sulfur Hexafluoride SH6

Check your understanding 5.7, 5.8, 5.9 p103 - 104

Lewis Electron DOT Structures

Count Electrons

Lewis electron dot structures show the valence electrons for each atom.

(You don't need to worry about the total number of electrons, only those

in the outer shells).

Lewis Electron Dot Structure for water

The following diagram shows the formation of the Lewis electron dot

diagram for water.

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The diagram on the right hand side shows the Lewis electron dot structure

for water. The ‘x’ shows the electron from hydrogen that is being shared

with the oxygen electron.

Lewis Dot Structure for Ammonia

Lewis Dot Structure for Hydrogen Sulfide

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– modelling the shapes of molecular substances

Valence Shell Electron-Pair Repulsion Theory (VSEPR)

It is a method for predicting the shape of a molecule from the knowledge of the

groups of electrons around a central atom.

Electron pairs (bonding and nonbonding electrons) repel one another, as a result,

the electron pairs remain as far apart as possible from another as possible to

minimize the repulsion.

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● investigate the different chemical structures of atoms and elements, including but not limited to:– ionic networks– covalent lattices (including diamond and silicon dioxide)– covalent networks– metallic structure

On the basis of their melting points and electrical conductivity, substances can be

classified into four classes with the following properties:

1. Covalent molecular substances – low melting point, non-conducting in the solid

and liquid states.

2. Covalent network substances – high melting points, non-conducting in the solid

and liquid states.

3. Ionic substances – high melting point, non-conducting in the solid state and

conducting in the molten state.

4. Metallic substances – high melting point, high conductivity in the solid and

molten states.

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Classification of Substances

Solid

Melting point

High Low

Conductivity ofSolid and molten state

Non-conductingConducting Non-conducting

Conductivity of Molten state

Conducting Non-conducting

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Covalent molecular substance

Not covalent molecular substance

Metallic substance

Covalent network solid

IonicSubstance

Covalent Molecular Substances

Covalent bonds form when atoms share electrons.

In most cases the shared pair(s) of electrons are provided by both of the atoms.

However, in some cases the shared pair of electrons are provided by one atom.

These covalent compound are called coordinate covalent bonds (or dative

covalent bond).

An example of this is the ammonium ion, NH4+:

(Draw the structure of the ammonium ion)

Examples of covalent molecular structures are oxygen, carbon dioxide and water.

Covalent molecular substances have the following properties:

They have low melting and boiling points.

Many are liquids or gases at room temperature.

They are non-conductors of electricity both in the liquid and solid states.

They form solids that are generally quite soft.

Covalent network Substances (Covalent Lattices)

These compounds are based on the elements carbon and silicon and include

substances such as diamond and silicon dioxide (quartz/sand).

Their properties come from the three dimensional network of strong covalent

bonds.

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Covalent network substances have the following properties:

Very high melting and boiling points

Non-conductors of electricity in the solid and liquid states

Extremely hard and brittle

Chemically inert

Insoluble in water and most other solvents

Structure of Diamond

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Structure of silicon

Crystalline silicon has the same structure as diamond.

As covalent compounds do not have ‘free’ ions or ‘electrons’ they do not conduct

electricity.

Ionic Compounds

Ionic substances consist of oppositely charged ions arranged in a three-dimensional

array.

The physical properties of ionic compounds are:

Hard and brittle

Non-conductors in the solid state

Good conductors when molten or in aqueous solutions

High melting and boiling points

The strong electrostatic attraction between oppositely charged ions results in ionic solids

being hard and brittle.

This comes from the fact that if a layer of ions is forced to slide past one another, then

electrostatic repulsion comes into effect and the crystal fractures.

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For electrical conductivity to occur, ions must be mobile, in a solid ionic compound this

is not possible, but in an aqueous solution the ions can freely migrate through the

solution and current can flow.

Structure of sodium chloride

Metallic Bonding

Metals consist of positive ions surrounded by a ‘sea’ of mobile electrons. The outer

electrons are said to be delocalised as they are not associated with a particular metal

ion and can move throughout the lattice of metal ions.

The properties of metal ions are:

Relatively high densities

Good conductors of heat and electricity

Malleable

Ductile

Lustrous

Relatively high melting and boiling points.

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Structure of Metals

Chapter Review Questions p114

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● explore the similarities and differences between the nature of intermolecular and intramolecular bonds and the strength of the forces associated with each, in order to explain the:– physical properties of elements– physical properties of compounds

Intermolecular ForcesForces exist between molecules that are not chemical bonds.

1. Dipole-Dipole interactions

Molecules in which the bonding electrons are unevenly shared between the bonded

atoms are called polar molecules and the bonds are called polar (covalent) bonds.

A pair of equal and opposite charges separated in space is called a dipole.

The positive and negative charges interact between molecules to form an inter-

molecular attraction.

Molecules which have atoms of significantly different electronegativity’s form

dipoles. Theses dipoles are permanent.

For example Hydrogen Chloride, HCl is a polar molecule.

Because Chlorine has a higher electronegativity than Hydrogen, most

electrons will be around the Chlorine end of the molecule.

As electrons are negatively charged, the chlorine end of the molecules has a

small net negative charge.

Therefore, the hydrogen end of the molecule has a small net positive

charge.

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This is shown by the following diagrams:

1. Electron distribution cloud diagram – shows most electrons are around the Cl atom.

2. The following diagram shows the net charge on the molecule.

The δ (small Greek letter delta) means “small”. The δ+ means a small positive charge.

The ‘dashed’ line between the chlorine and the hydrogen shows the dipole – dipole

interaction caused by the electrostatic attraction between positive and negative

charges. This is the intermolecular force.

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Hydrogen Bonding

Hydrogen bonds are specific dipole-dipole intermolecular attractions that form

between H – O, H – N and H – F.

These intermolecular forces are stronger than other dipole - dipole interactions.

NOTE: They are called Hydrogen bonds but they are not chemical bonds. They

are intermolecular forces.

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Hydrogen bonding in water:

Dispersion Forces

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The dispersion force is the weakest intermolecular force.

The dispersion force is a temporary attractive force that results when the electrons

in two adjacent atoms occupy positions that make the atoms form temporary

dipoles.

This force is sometimes called an induced dipole-induced dipole attraction.

Dispersion forces are the attractive forces that cause nonpolar substances to

condense to liquids and to freeze into solids when the temperature is lowered

sufficiently.

Because of the constant motion of the electrons, an atom or molecule can develop a

temporary (instantaneous) dipole when its electrons are distributed unsymmetrically

about the nucleus.

A second atom or molecule, in turn, can be distorted by the appearance of the dipole in the

first atom or molecule (because electrons repel one another) which leads to an electrostatic

attraction between the two atoms or molecules.

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Dispersion forces are present between any two molecules (even polar molecules) when

they are almost touching.

Molecular SizeDispersion forces are present between all molecules, whether they are polar or

nonpolar.

Larger and heavier atoms and molecules exhibit stronger dispersion forces than

smaller and lighter ones.

In a larger atom or molecule, the valence electrons are, on average, farther from

the nuclei than in a smaller atom or molecule. They are less tightly held and can

more easily form temporary dipoles.

The ease with which the electron distribution around an atom or molecule can be

distorted is called the polarizability.

Dispersion forces tend to be:

stronger between molecules that are easily polarized.

weaker between molecules that are not easily polarized.

● investigate elements that possess the physical property of allotropy

Allotropy is the property of some chemical elements to exist in two or more different

forms, in the same physical state, known as allotropes of these elements.

Allotropes are different structural modifications of an element; the atoms of the element

are bonded together in a different manner.

Allotropes have the same chemical properties but different physical properties.

For example, the allotropes of carbon include diamond (the carbon atoms are bonded

together in a tetrahedral lattice arrangement), graphite (the carbon atoms are bonded

together in sheets of a hexagonal lattice), graphene (single sheets of graphite), and

fullerenes.

Check your understanding: 6.1, 6.2, 6.3 p122

Chapter Review Questions p 135,136,

Module 1 Review Questions p137 & 138

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