Atomic Structure - Science Websitersgscience.weebly.com/uploads/8/0/6/0/8060490/bonding.doc · Web viewCompounds form when atoms are joined together by forces called chemical bonds

Bonding and Structure


document.doc Page 1 of 40 © Rob Ritchie, 09/05/23


CHEMICAL BONDINGStable atomsThe noble gases exist as single atoms that are very stable. All electrons are paired and the bonding shells are full.

helium neon argon

BondingCompounds form when atoms are joined together by forces called chemical bonds.

When bonds form, unpaired electrons often pair up to form a noble gas electron structure. This is often referred to as the ‘octet rule’.

the bonding shells are full

the electron structure is very stable.

Types of bondingBonds are classified into two main types: ionic and covalent.

As a general rule,

Ionic bonding occurs between a metal and a non-metal.

Covalent bonding occurs between a non-metal and a non-metal.

QuestionsPredict the type of bonding in the following compounds.

(a) sodium chloride ………………….

(b) zinc oxide ………………….

(c) hydrogen chloride ………………….

(d) silver bromide ………………….

(e) nitrogen bromide ………………….

(f) sulphur dioxide ………………….



IONIC BONDSIonic bonds are present in a compound of a metal and a non-metal.

Electrons are transferred from a metal atom to a non-metal atom forming ions.

The metal ion is positive

The non-metal ion is negative

An ionic bond is the electrical attraction between oppositely charged ions.

Sodium chloride, NaCl

Na atom Cl atom Na+ ion Cl ion

noble gas neon argon

The ions that are formed often have stable noble gas electron structures with full outer electron shells.

Magnesium chloride, MgCl2

Mg atom Cl atoms Mg2+ ion Cl ions

noble gas neon argon



Problems(a) Draw ‘dot-and-cross’ diagrams to show how atoms can form the four ionic

compounds below.

(b) Write down the noble gas that would have the same electron structure as each ion that you draw.

(i) magnesium oxide,

….. atom ….. atom ….. ion ….. ion

noble gas ……… ………..

(ii) sodium oxide,



(iii) aluminium oxide.





IONS AND THE PERIODIC TABLEThe charge on an ion can easily be predicted from its position in the Periodic Table.

Complete the table below.

Group 1 2 3 4 5 6 7 8number of outer shell electrons

1 2 3 4 5 6 7 8

element Li Be B C N O F Ne

2,1 2,2 2,3 2,4 2,5 2,6 2,7 2,8

ion Li+ F

2 2,8

element Na Mg Al Si P S Cl Ar

2,8,1 2,8,2 2,8,3 2,8,4 2,8,5 2,8,6 2,8,7 2,8,8

ion

Multiple chargesSometimes elements can form ions with different charges. The charge is shown by a Roman numeral.

iron(II) for Fe2+ and iron(III) for Fe3+;

copper(I) for Cu+ and copper(II) for Cu2+.

Predicting ionic formulae

calcium chloride: ion charge ion charge

Ca2+ 2+ Cl 1–

equalise charge Ca2+ 2+ 2 Cl 2–

formula CaCl2



Predicting formulae1. Predict the formula of the ionic compounds below.

(a) lithium chloride ………………….

(b) sodium oxide ………………….

(c) beryllium oxide ………………….

(d) lithium nitride ………………….

(e) aluminium fluoride ………………….

(f) aluminium sulphide………………….

(g) aluminium phosphide……………….

(h) magnesium nitride………………….

2. Predict the formula of the ionic compounds below.

(a) copper(II) oxide ………………….

(b) copper(I) oxide ………………….

(c) iron(II) fluoride ………………….

(d) iron(III) chloride ………………….

(e) cobalt(II) oxide ……..………….

(f) chromium(III) oxide………………..

(g) manganese(VI) oxide……………….

(h) vanadium(V) nitride………………….



Molecular ions

common molecular ions

ion 1+ 1– 2– 3–

ammonium NH4+ hydroxide OH– carbonate CO3

2– phosphate PO43–

nitrate NO3– sulphate SO4

2–

nitrite NO2– sulphite SO3

2–

Ionic compounds from molecular ionssodium carbonate ion charge ion charge

Na+ 1+ CO32– 2–

equalise charge 2 Na+ 2+ 2 CO32– 2–

formula Na2CO3

calcium nitrate ion charge ion chargeCa2+ 2+ NO3

– 1–

equalise charge Ca2+ 2+ 2 NO3– 2–

formula Ca(NO3)2

Predicting more formulaePredict the formula of the ionic compounds below.

(a) lithium nitrate ………………….

(b) sodium carbonate ………………….

(c) aluminium sulphate………………….

(d) calcium hydroxide ………………….

(e) iron(III) sulphite ………………….

(f) chromium(III) nitrite ……………….

(g) ammonium phosphate …………….

(h) vanadium(V) sulphate……………….



COVALENT BONDSA covalent bond is formed between atoms of non-metals with a similar attraction for electrons.

A covalent bond is a shared pair of electrons.

A covalent bond forms when two atoms attract the same pair of electrons.

H2 molecule Cl2 molecule

Molecules with different atoms (a) Draw ‘dot-and-cross’ diagrams to show a molecule of the following compounds

of hydrogen.

(b) Write down the formula of each compound.

(c) What is the common name given to each compound?

hydrogen chloride hydrogen oxideformula: …………. formula: ………….common name: ………………….. common name: …………………..

hydrogen nitride hydrogen carbide

formula: …………. formula: ………….

common name: ………………….. common name: …………………..



Multiple bondsDouble and triple bonds are possible.

O2 N2

Drawing ‘dot-and-cross’ diagrams 1 Draw dot-and-cross diagrams of the following molecules of (i); (ii); (iii); (iv),

F2 HF

SiF4 SCl2

2 Draw dot-and-cross diagrams of the following molecules, each containing at least one multiple covalent bond.

CO2 C2H4

HCN H2CO



DATIVE COVALENT BONDSA dative covalent bond forms when the shared pair of electrons comes from just one of the atoms.

The formation of the ammonium ion, NH4+.

NH3 H+ NH4+

More covalent questions1. Draw dot-and-cross diagrams to show the formation of a dative covalent bond in

the oxonium ion, H3O+.

2. (a) Draw a dot-and-cross diagram for a molecule of boron trifluoride.

(b) How many electrons surround the boron atom?

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(c) In what way, is the electron structure different to those met previously?

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(d) Why do you think this structure is still stable?

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Bonding and Structure3. When covalent bonds form, unpaired electrons often pair up to form a noble gas

electron structure. This is often referred to as the ‘octet rule’. Many molecules, however, form electron structures that do not form an octet.

Draw ‘dot-and-cross’ diagrams for the following molecules which do not obey the octet rule.

AlCl3 PCl5

SF6 SO3

SO2 SO3



SHAPES OF MOLECULESthe shape of a molecule depends upon the number of electron pairs surrounding the central atom.

The shapes of simple molecules can be explained using electron-pair repulsion.

The pairs of electrons that surround an atom repel one another.

When this happens the electron pairs become as far apart as possible.

The electron-pairs naturally repel into the following shapes around a central atom.

molecule

number of electron

pairs around central atom

dot-and-cross diagram shape bond

angle

BCl2

BF3

CH4

SF6



Molecules with lone pairsA lone pair of electrons repels more than a bonded pair. This is because a lone pair is closer to the atom:

In general:

lone-pair/lone-pair repulsion > bonded-pair/lone-pair repulsion

bonded-pair/lone-pair repulsion > bonded-pair/bonded-pair repulsion

Lone pairs will distort the shape of a molecule and reduce the bond angle:

molecule number of lone pairs

dot and cross diagram shape bond angle

CH4

NH3

H2O



Some shape questionsFor each of the following molecules, draw a dot-and-cross diagram, and predict the shape and bond-angles.

molecule number of lone pairs

dot and cross diagram shape bond angle

BeF2

AlCl3

SiH4

H2S

PH3



Double bondsA double bond is treated in the same way as a bonded pair. Each double bond and pair of electrons should be treated as a bonding ‘area’.

molecule number of bonding areas ‘dot and cross’ diagram shape bond

angle

CO2

Now try these…

moleculenumber of bonding areas

‘dot and cross’ diagram shape bond

angle

CS2

C2H4

SO3

SO2



IONIC OR COVALENT?An ionic bond with 100% ionic character would require the complete transfer of an electron from a metal atom to a non-metal atom.

In practice, this never completely happens!

Between the extremes of ionic and covalent bonding, there is a whole range of intermediate bonds, which have both ionic and covalent contributions.

Ionic bonds with covalent characterMany ionic compounds have a degree of covalency resulting from incomplete transfer of electrons.

This occurs in ionic compounds with

a small + ion (cation) with a high charge density and

a large – ion (anion) with a low charge density.

charge density =

The electric field around a small cation distorts the electron shells around a large anion.

This effect is called polarisation.

As a result, electrons are attracted towards the cation, giving a degree of covalency (electron sharing).

In Al2O3, polarisation takes place between Al3+ and O2– ions.



ELECTRONEGATIVITYThe nuclei of the atoms in a molecule attract the electrons pair in a covalent bond.It is this attraction which is responsible for the covalent bond.

Electronegativity is a measure of the attraction of an atom in a molecule for the pair of electrons in a covalent bond.

The most electronegative atoms attract bonding electrons most strongly.

Elements with small atoms have the most electronegative atoms.

Highly reactive non-metallic elements (such as O, F and Cl) have the most electronegative atoms.

Reactive metals (such as Na and K) have the least electronegative atoms.

How is electronegativity measured?

Fluorine is the most electronegative

element with an electronegativity of

4.0;

Fluorine has small atoms which attract the pair of electrons in a covalent bond more strongly than larger atoms.

The greater the difference between electronegativities, the greater the ionic character of the bond.


Li1.0

Be1.5

B2.0

C2.5

N3.0

O3.5

F4.0

Na0.9

Cl3.0

K0.8

Br2.8

electronegativity increases


The greater the similarity in electronegativities, the greater the covalent character of the bond

POLAR AND NON-POLAR MOLECULESNon-polar bondsA covalent bond is non-polar when

the bonded electrons are shared equally between both atoms

the bonded atoms are the same

the bonded atoms have similar electronegativities.

A covalent bond must be non-polar if the bonded atoms are the same.

H H Cl Cl

H2 molecule Cl2 molecule



Polar bondsA covalent bond is polar when

the electrons in the bond are shared unequally making a polar bond;

the bonded atoms are different, each has a different electronegativity.

In hydrogen chloride the H–Cl bond is mainly covalent with one pair of electrons shared between the two atoms.

However, the chlorine atom is more electronegative than the hydrogen atom.

The chlorine atom attracts the bonded pair of electrons more than the hydrogen atom making the covalent bond polar.

The electrons in the bond are shared unequally making a polar bond.

+ H Cl

HCl molecule

The HCl molecule is polarised with a small positive charge + on the H atom and a small negative charge – on the Cl atom.

The hydrogen chloride molecule is polar with a permanent dipole.

Symmetrical and unsymmetrical moleculesIn symmetrical molecules, dipoles cancel and there are no permanent dipoles.

dipoles cancel

CHCl3 - polar CCl4 - non-polar

In CCl4

each C–Cl bond is polar but the dipoles act in different directions

the overall effect is for the dipoles to cancel each other

CCl4 is a non-polar molecule.



Some polar questions1. Draw simple diagrams for each of the following molecules. Decide whether each

molecule is polar. If so, use the periodic table to predict the dipole present and shows these clearly as + and on your diagrams.

Br2 H2O O2

HBr NH3 CF4

2. (i) Predict the shape of a molecule of BF3 and of PF3.

BF3 PF3

(ii) Explain why BF3 is non-polar whereas PF3 is polar.

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3. (i) Predict the shape of a molecule of H2O and of CO2.

H2O CO2

(ii) Explain why H2O is polar whereas CO2 is non-polar.

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INTERMOLECULAR FORCESStrength of bonds and forces

Ionic, covalent and metallic bonds are of comparable strength.

Intermolecular forces are much weaker:

type of bond bond enthalpy/ kJ mol1

covalent bond 200-500hydrogen bond 5-40van der Waals’

forces2

VAN DER WAALS' FORCES Van der Waals’ forces (induced dipole-dipole interactions) exist between all molecules whether polar or non-polar. Without these forces, non-polar molecules could never form a liquid or a solid.

Van der Waals’ forces are

weak intermolecular interactions

caused by attractions between very small dipoles in molecules.

movement of electrons produces an oscillating dipole

oscillating dipole induces a dipole in a neighbouring molecule which is induced onto further molecules

What affects the strength of van der Waals’ forces? Van der Waals' forces result from interactions of electrons between molecules.

The greater the number of electrons in each molecule

the larger the oscillating and induced dipoles

the greater the attractive forces between moleculesdocument.doc Page 21 of 40 © Rob Ritchie, 09/05/23

dipole oscillates and continually changes with time

- + -+

induced dipoles attract one another

- +- +

- +- + - +

- +

Bonding and Structure the greater the van der Waals' forces.



Van der Waals’ forces and electronsMolecules with more electrons will generate larger oscillating and induced dipoles. These produce larger attractive forces between molecules.

The table below shows the boiling points of the hydrogen compounds of group 4.

compound boiling point /K

number of electrons

CH4 112SiH4 161GeH4 178SnH4 221

1. Complete the table above with the number of electrons in each compound.

2. Explain why the boiling point increases down this group.

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3. Compare these boiling points with those for the hydrogen compounds of Group 6 (page 4). Explain the similarities and differences between the two sets of data.

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PERMANENT DIPOLE-DIPOLE INTERACTIONSPermanent dipole-dipole interaction are weak intermolecular forces.

They are simply attractions between weak dipole charges on different molecules.

Permanent dipole-dipole interactions are ‘non-directional’.

Intermolecular forces between HCl molecules

Between HCl molecules, there will be both

weak intermolecular forces:

van der Waals’ forces and

permanent dipole-dipole interactions.

permanent dipole-dipole interactions > van der Waals’ forces.

Now do this…..Draw a diagram to show the dipole-dipole interactions that exist between molecules of nitrogen monoxide, NO.



HYDROGEN BONDSA hydrogen bond is a special type of permanent dipole-dipole interaction found between molecules containing the following groups:

Hydrogen bonds act between:

a lone pair HO:, HN: or HF: and

a hydrogen atom HO, HN or HF on a different molecule.

Hydrogen bonding occurs between molecules such as H2O.

Hydrogen bonding is important in organic compounds containing –OH or –NH bonds: e.g. alcohols, carboxylic acids, amines, amino acids.

When water change state, the covalent bonds between the H and O atoms in an H2O molecule are strong and do not break - the much weaker intermolecular forces break



Some hydrogen bonding questions1. Put a tick below the following molecules that have hydrogen bonding:

molecule H2O H2S CH4 HOCH3 NH3 NO2 H2NCH3

hydrogen bonding?

2. Each molecule that you draw should show relevant lone pairs and dipoles.

Draw diagrams showing hydrogen bonding between

(a) 2 molecules of ammonia;

(b) 2 molecules of hydrogen fluoride;

(c) 2 molecules of ethanol, C2H5OH;

(d) 1 molecule of water and 1 molecule of ethanol.



SPECIAL PROPERTIES OF WATERA hydrogen bond has only about one-tenth the strength of a covalent bond.

However, hydrogen bonding is strong enough to have significant effects on physical properties, resulting in some unexpected properties for water.

The solid (ice) is less dense than the liquid (water) Particles in solids are usually packed closer together than in liquids.

Hydrogen bonds hold water molecules apart in an open lattice structure.

ice is less dense than water.

The diagram below shows how the open lattice of ice collapses on melting.

H2O has a relatively high melting point and boiling point There are relatively strong hydrogen bonds between H2O molecules.

The hydrogen bonds are extra forces, over and above van der Waals’ forces.

These extra forces result in higher melting and boiling points than would be expected from just van der Waals’ forces.

When the ice lattice breaks, hydrogen bonds are broken.

Other propertiesThe extra intermolecular bonding from hydrogen bonds also explains the relatively high surface tension and viscosity in water.



How do hydrogen bonds affect the properties of water?Hydrogen bonds increase the attraction between molecules. This will result in an increase in the boiling point of a compound. The table below shows the boiling points of the hydrogen compounds of group 6.

compound boiling point /K

relative molecular mass

H2O 373H2S 213H2Se 231H2Te 270

1. Complete the table above with the relative molecular mass of each compound.

2. Plot a graph of boiling point against relative molecular mass. Join each point with a line.

3. The boiling point of water is higher than expected owing to hydrogen bonding. Use your graph to estimate what the boiling point of water would be if there were no hydrogen bonding. Show this clearly on your graph.

4. Refer to Chapter 3 of the Foundation Chemistry textbook. Water has several peculiar properties that can be explained by hydrogen bonds. The unexpected high boiling and melting point is just one of these. Explain how hydrogen bonding is responsible for three more unusual properties of water.

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BONDING, STRUCTURE AND PROPERTIESThe properties of a substance depend upon its bonding and structure.

Ionic and covalent bonds are of comparable strength - Intermolecular forces are far weaker.

Ionic bonds break when a giant ionic lattice melts or boils

Simple molecular structure – weak forces break - low melting point

Covalent bonds break when a giant molecular lattice melts or boils

Giant structure – strong forces break - high melting point

GIANT IONIC LATTICESIonic compounds form giant ionic lattices with each ion surrounded by ions of the opposite charge.

A giant ionic lattice is held together by strong electrostatic attraction between positive and negative ions.

Each ion is surrounded by oppositely-charged ions, forming a giant ionic lattice.

Part of the sodium chloride lattice,

Each Na+ ion surrounds 6 Cl ions

Each Cl ion surrounds 6 Na+ ions

Although it is convenient to look at ionic bonding between two ions only, each ion is able to attract oppositely charged ions in all directions.

This results in a giant ionic lattice structure with hundreds of thousands of ions (depending upon the size of the crystal).

This arrangement is characteristic of all ionic compounds.



Properties of Giant ionic latticesHigh melting point and boiling pointIonic compounds are solids at room temperature.

The attraction is strong between + and ions.

High temperatures are needed to break the strong electrostatic forces holding the ions rigidly in the solid lattice.

ionic compounds have high melting and boiling points.

Electrical conductivityIn the solid lattice,

the ions are in a fixed position and there are no mobile charge carriers.

an ionic compound is a non-conductor of electricity in the solid state.

When melted or dissolved in water,

the solid lattice breaks down,

the ions are now free to move as mobile charge carriers.

an ionic compound is a conductor of electricity in liquid and aqueous states.

Solubility The ionic lattice dissolves in polar solvents (e.g. water).

The polar water molecules break down the lattice and surround each ion in solution as shown below for sodium chloride.



PROPERTIES OF COVALENT COMPOUNDSElements and compounds with covalent bonds have either of two structures:

a simple molecular structure,

a giant molecular structure.

SIMPLE MOLECULAR STRUCTURESSimple molecular structures have small molecules, such as Ne, H2, O2, N2.

A simple molecular structure is held together by weak forces between molecules

The atoms within each molecule are bonded strongly together by covalent bonds

Properties of simple molecular structuresLow melting point and boiling point Low temperatures provide sufficient energy to break the intermolecular forces.

simple molecular structures have low melting and boiling points.

When the simple molecular structure of I2 is broken,

only the weak van der Waals’ forces between the I2 molecules break;

the covalent bonds, I–I, are strong and do not break.

Electrical conductivity There are no free charged particles.

simple molecular structures are non-conductors of electricity.

Solubility Van der Waals' forces form between a simple molecular structure and a non-polar

solvent, such as hexane.

These interactions weaken the structure.

simple molecular structures are soluble in non-polar solvents (e.g. hexane).



GIANT MOLECULAR STRUCTURESDiamond, graphite and SiO2 are common examples of giant molecular lattices.

This type of structure is known by a variety of names: a giant molecular lattice, a giant covalent lattice, a giant atomic lattice and a macromolecular lattice.

A giant molecular structure is held together by strong covalent bonds between atoms.

Properties of giant molecular structuresHigh melting point and boiling point High temperatures are needed to break the strong covalent bonds in the lattice.

giant molecular structures have high melting and boiling points.

Electrical conductivity Except for graphite (see below), there are no free charged particles.

giant molecular structures are non-conductors of electricity.

Solubility The strong covalent bonds in the lattice are too strong to be broken by either polar

or non-polar solvents.

giant molecular structures are insoluble in polar and non-polar solvents.



PROPERTIES OF DIAMOND AND GRAPHITEproperty diamond graphite

structure

tetrahedral hexagonal layers

symmetrical structure held together by strong covalent bonds throughout lattice

strong layer structure but with weak bonds between the layers

electrical conductivity

poor conductivity

There are no delocalised electrons

All outer shell electrons are used for covalent bonds.

good conductivity

Delocalised electrons between layers.

Electrons are free to move parallel to the layers when a voltage is applied.

hardness

hard

Tetrahedral shape enables external forces to be spread throughout the lattice.

soft

Strong bonding within each layer weak forces between layers

easily allow layers to slide.



BONDING AND STRUCTURE IN METALSMetallic bondingA metallic bond holds atoms together in a solid metal or alloy.

In solid metals, the atoms are ionised.

The positive ions occupy fixed positions in a lattice;

The outer shell electrons are delocalised – they are spread throughout the metallic structure and are able to move freely throughout the lattice.

A metallic bond is the electrostatic attraction between the positive metal ions and delocalised electrons.

In the metallic lattice, each metal atom exists as + ions by releasing its outer shell electrons to the ‘sea of electrons’.

The ‘sea of electrons’

Delocalised and localised electronsIn a metallic bond,

the delocalised electrons in metals are spread throughout the metal structure;

the delocalised electrons are able to move throughout the structure;

it is impossible to assign any electron to a particular positive ion.

In a covalent bond,

the localised pair of electrons is always positioned between the two atoms involved in the bond;

the electron charge is concentrated between the bonded atoms.



PROPERTIES OF GIANT METALLIC LATTICESAll metals form giant metallic lattices in the solid state

A giant metallic lattice is held together by strong electrostatic attractions between positive ions and negative electrons.

A metallic bond acts between the positive metal ions and delocalised electrons.

High melting point and boiling point Generally high temperatures are needed to separate the ions from their rigid

positions within the lattice.

most metal have high melting and boiling points.

Good thermal and electrical conductivity The existence of mobile, delocalised electrons allows metals to conduct heat and

electricity well, even is the solid state.

The electrons are free to flow between positive ions.

The positive ions do not move.

When a metal conducts electricity, only the electrons move.

Solubility Metals are insoluble.

Solvents such as water are unable to form strong enough forces with the ions and electrons to pull the lattice apart.



COMPARISON OF STRUCTURE, BONDING AND PROPERTIESThe different types of bonds and forces Covalent bonds act between atoms.

Ionic bonds act between ions.

Metallic bonds act between positive ions and electrons.

Hydrogen bonds act between polar molecules.

Dipole-dipole interactions act between polar molecules.

Van der Waals' forces act between induced dipoles of molecules.

structure m pt/ b pt

reason electrical conductivity

reason solubility reason

giant ionic

simple molecular

giant molecular

hydrogen bonded

giant metallic



Properties questions1. Magnesium oxide forms a similar ionic lattice to that of sodium chloride.

(a) Draw a clear diagram of the magnesium oxide lattice.

(b) Predict, with reasons, the following properties of magnesium oxide:

(i) melting and boiling points

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(ii) electrical conductivity

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(iii) solubility.

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(c) Suggest, with reasons, why magnesium oxide has a higher melting point than sodium chloride,

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Bonding and Structure2. (a) Describe what is meant by hydrogen bonding,

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(b) Describe how hydrogen bonding influences the properties of water.

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Bonding and Structure...............................................................................................................................


Bonding and Structure3. The data below gives some properties of four substances, A, B, C and D.

compound solubility in water

solubility in hexane

electrical conductivity boiling

point /Ksolid aqueous solution

A good insoluble poor good 1738B poor good poor poor 456C poor poor poor - 2503D poor poor good - 3160

(a) Explain how these data suggest different structures for the four substances shown.

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(b) Suggest an identity for each substance.

A ..................................................................................................................

B ..................................................................................................................

C ..................................................................................................................

D ..................................................................................................................document.doc Page 40 of 40 © Rob Ritchie, 09/05/23

Documents

Atomic Structure - Science Websitersgscience.weebly.com/uploads/8/0/6/0/8060490/bonding.doc · Web viewCompounds form when atoms are joined together by forces called chemical bonds