Ionic bondingIONIC BONDWhen atoms combine they do so by trying to achieve an inert gas configuration. Ionic compounds are formed when electronsare transferred from one atom to another to form ions with complete outer shells of electrons. In an ionic compound the positiveand negative ions are attracted to each other by strong electrostatic forces, and build up into a strong lattice. Ionic compoundshave high melting points as considerable energy is required to overcome these forces of attraction.

The classic example of an ionic compound is sodium chloride Na+CI-, formed when sodium metal burns in chlorine. Chlorine isa covalent molecule, so each atom already has an inert gas configuration. However, the energy given out when the ionic latticeis formed is sufficient to break the bond in the chlorine molecule to give atoms of chlorine. Each sodium atom then transfers oneelectron to a chlorine atom to form the ions.

2.B.1 I[NeI3s') 2.B.71[NeI3s'3p')11 protons 1 7 protons11 electrons 17 electrons

2.81[Nell 2.8.81[ArD11 protons 1 7 protons10 electrons 18 electrons

The charge carried by an ion depends on the number of electrons the atom needed to lose or gain to achieve a full outer shell.

Cations Anions

Group 1 Group 2 Group 3 Group 5 Group 6 Group 7+1 +2 +3 -3 -2 -1

l.i" Na" K+ Mg2+Ca1+ A13+ N3- p3- 02-S1- F- CI- Br

Thus in magnesium chloride two chlorine atoms each gain one electron from a magnesium atom to form Mg1+CI-1. Inmagnesium oxide two electrons are transferred from magnesium to oxygen to give Mg1+02-. Transition metals can form morethan one ion. For example, iron can form Fe1+ and Fe3+ and copper can form Cu" and Cu2+.

FORMULAS OF IONIC COMPOUNDSIt is easy to obtain the correct formula as the overall charge of the compound must be zero.

lithium fluoride Li+F- magnesium chloride Mg1+CI-2 aluminium bromide AI3+Br-3sodium oxide Na+20

2- calcium sulfide Ca1+S2- iron(1I1) oxide Fe3+202-3potassium nitride K+3N3- calcium phosphide Ca2+3p3-1 iron(ll) oxide Fe2+02-

Note: the formulas above have been written to show the charges carried by the ions. Unless asked specifically to do this it iscommon practice to omit the charges and simply write LiF, MgCI2, etc.

IONS CONTAINING MORE THAN ONE ELEMENTIn ions formed from more than one element the charge is often spread (delocalized) over the whole ion. An example of a positiveion is the ammonium ion NH/, in which all four N-H bonds are identical. Negative ions are sometimes known as acid radicalsas they are formed when an acid loses one or more H+ ions.

hydroxide OH-nitrate NO)sulfate SO/-hydrogensulfate HS04-

(from nitric acid, HN03)

{ from sulfuric acid, H1S04 }

carbonate CO/-hydrogencarbonate HC03-ethanoate CH3COO-

{ from carbonic acid, H1C03 }

(from ethanoic acid, CH3COOH)

The formulas of the ionic compounds are obtained in exactly the same way. Note: brackets are used to show that the subscriptcovers all the elements in the ion.

sodium nitrate Na+N03-ammonium sulphate (NH/)2S0/-

calcium carbonate Ca2+CO/-magnesium ethanoate Mg2+(CH3COO-)2

aluminium hydroxide AP+(OH-)3

IONIC OR COVALENT?Ionic compounds are formed between metals on the left of the Periodic Table and non-metals on the right of the Periodic Table;that is, between elements in groups 1, 2, and 3 with a low electronegativity (electropositive elements) and elements with a highelectronegativity in groups 5,6, and 7. Generally the difference between the electronegativity values needs to be greater thanabout 1.8 for ionic bonding to occur.

AI F AI ° AI CI AI BrElectronegativity 1.5 4.0 1.5 3.5 1.5 3.0 1.5 2.8~ '---v----' '-------v--' ~Difference in electronegativity 2.5 2.0 1.5 l.3Formula AIF3 AI103 A12C16 AI2Br6Type of bonding ionic ionic intermediate between covalent

ionic and covalent

M. pt / °C 1265 2050 Sublimes at 180 97

Bonding 19

Covalent bondingSINGLE COVALENT BONDSCovalent bonding involves the sharing of one or more pairs of electrons so that eachatom in the molecule achieves an inert gas configuration. The simplest covalentmolecule is hydrogen. Each hydrogen atom has one electron in its outer shell. Thetwo electrons are shared and attracted electrostatically by both positive nucleiresulting in a directional bond between the two atoms to form a molecule. When onepair of electrons is shared the resulting bond is known as a single covalent bond.Another example of a diatomic molecule with a single covalent bond is chlorine, C12.

LEWIS STRUCTURESIn the Lewis structure (also known as electron dot structure) all the valence electrons are shown. There are various differentmethods of depicting the electrons. The simplest method involves using a line to represent one pair of electrons. It is alsoacceptable to represent single electrons by dots, crosses or a combination of the two. The four methods below are all correctways of showing the Lewis structure of fluorine.

IF-FIxx xx~nnxx xx

: F: F:.. ..xx ••

~HF:xx ••

Sometimes just the shared pairs of electrons are shown, e.g. F-F. This gives information about the bonding in the molecule, but itis not the Lewis structure as it does not show all the valence electrons.

SINGLE COVALENT BONDSIFI

- I _IF-C-FI- I -

1£1tetrafluoro-

methane

HI

H-C-HI

Hmethane

H-FIH-N-HI

Hammonia

H-OII

Hwater hydrogen

fluoride

The carbon atom (electronic configuration 2.4) has fourelectrons in its outer shell and requires a share in four moreelectrons. It forms four single bonds with elements that onlyrequire a share in one more electron, such as hydrogen orchlorine. Nitrogen (2.5) forms three single bonds withhydrogen in ammonia leaving one non-bonded pair ofelectrons (also known as a lone pair). In water there aretwo non-bonded pairs and in hydrogen fluoride threenon-bonded pairs.

BOND LENGTH AND BOND STRENGTHThe strength of attraction that the two nuclei have for theshared electrons affects both the length and strength of thebond. Although there is considerable variation in the bondlengths and strengths of single bonds in differentcompounds, double bonds are generally much stronger andshorter than single bonds. The strongest covalent bonds areshown by triple bonds.

Length StrengthInm IkJ mol-1

Single bonds CI-CI 0.199 242C-C 0.154 348

Double bonds C=C 0.134 6120=0 0.121 496

Triple bonds C=C 0.120 837N=N 0.110 944

e.g. ethanoic acid: 0.1240nm "II-0 C 0.143nm,/~

,

Shapes of simple molecules and ionsVSEPR THEORY

No. of charge Shape Name of BondThe shapes of simple molecules and ions can be centres shape angle(s)determined by using the valence shell electron pair 2 0-0-0 linear 180°repulsion (VSEPR) theory. This states that pairs ofelectrons arrange themselves around the central 0atom so that they are as far apart from each other as 1

trigonalpossible. There will be greater repulsion between 3 0 120°non-bonded pairs of electrons than between

0/ ""0 planar

bonded pairs. Since all the electrons in a multiple 0bond must lie in the same direction, double and 4 1 tetrahedral 109.5°0triple bonds count as one pair of electrons. Strictly Q-j'/ ""0speaking the theory refers to negative charge 0

centres, but for most molecules this equates to pairs 00'-,1 trigonalof electrons. 5 90°, 120°, 180°'0-0 bipyramidalThis results in five basic shapes depending on the 0/1

number of pairs. 0

00'-.1 octahedral 90°, 180°6 0-'0-0

I~o0

WORKING OUT THE ACTUAL SHAPETo work out the actual shape of a molecule calculate the number of pairs of electronsaround the central atom, then work out how many are bonding pairs and how manyare non-bonding pairs. (For ions the number of electrons which equate to the chargeon the ion must also be included when calculating the total number of electrons.)

2 NEGATIVE CHARGE CENTRES

CI-Be-CIO=C=O

'Uble bondcounts as one pair

H-C==C-H H-C==N

~ triple bond ~counts as one pair

3 NEGATIVE CHARGE CENTRES3 bonding pairs - trigonal planar

F1

[0 ]2-

O/~""Ocarbonate ion

2 bonding pairs, 1 non-bonded pair - bent or V-shaped

nitrite ion

5 AND 6 NEGATIVE CHARGE CENTRES5 and 6 negative charge centres

F FF", 1 F'"I /F;P-F ;s~F 1 F 1 F

F F

CICI", \

'S:)C( I

CIdistorted

tetrahedraltrigonal

bipyramidaloctahedral square planar

non-bonding pairs asfar apart as possible

above and below plane

even greaterrepulsion by twonon-bondingpairs so bond angleeven smaller

4 NEGATIVE CHARGE CENTRES

4 bonding pairs - tetrahedral

CI1

,C"/,

Intermolecular forces and allotropes of carbonMOLECULAR POLARITYWhether a molecule is polar, or not, depends both on the relative electronegativities of the atoms in the molecule and on itsshape. If the individual bonds are polar then it does not necessarily follow that the molecule will be polar as the resultant dipolemay cancel out all the individual dipoles.

0- 20+ 0-o=c=onon-polar

(resultant dipole zero)

resultant

70-dipole0+ __W, 0+H

polar

0- 1resultantCI dipole

0+1H'--'C""

, HHpolar

0-CI140+

°CI/C"", 0-Clo-CI

non-polar (resultant dipole zero)

Van der Waals' forcesEven in non-polar molecules the electrons can at anyone moment beunevenly spread. This produces temporary instantaneous dipoles. Aninstantaneous dipole can induce another dipole in a neighbouring particleresulting in a weak attraction between the two particles. Van der Waals' forcesincrease with increasing mass.

increasing van der Waals' forces

CI270.9-34.0

Br216058.0

F2M, 38.0b. pt I °C -188

increasing van der Waals' forces

CH4M, 16.0b. pt/ °C -162

C3HS44.0-42.2

INTERMOLECULAR FORCESThe covalent bonds between the atoms within a moleculeare very strong. The forces of attraction between themolecules are much weaker. These intermolecular forcesdepend on the polarity of the molecules.

Hydrogen bondingHydrogen bonding occurs when hydrogen is bonded directly to asmall highly electronegative element, such as fluorine, oxygen, ornitrogen. As the electron pair is drawn away from the hydrogen atomby the electronegative element, all that remains is the proton in thenucleus as there are no inner electrons. The proton attracts a non-bonding pair of electrons from the F, N, or 0 resulting in a muchstronger dipole:dipole attraction. Water has a much higher boilingpoint than the other group 6 hydrides as the hydrogen bondingbetween water molecules is much stronger than the dipole:dipolebonding in the remaining hydrides. A similar trend is seen in thehydrides of group 5 and group 7. Hydrogen bonds between themolecules in ice result in a very open structure. When ice melts themolecules can move closer to each other so that water has itsmaximum density at 4 "C.

boilingtemperature / K

400

300

200

100

the ice lattice

~= hydrogen~ bond

3period

22 Bonding

254183

Dipole:dipole forcesPolar molecules are attracted to each otherby electrostatic forces. Although stillrelatively weak the attraction is strongerthan van der Waals' forces.

non-polar

H H H HI I I I

H-C-C-C-C-HI I I IH H H H

polar0-

H ~ HI C IH-9/0+"9-HH H

butane Mr = 58b. pt -0.5°C

propanone Mr = 58b. pt 56.2°C

identical masses(different intermolecular forces)

ALLOTROPES OF CARBONAllotropes occur when an element can exists indifferent crystalline forms. In diamond each carbonatom is covalently bonded to four other carbonatoms to form a giant covalent structure. All thebonds are equally strong andthere is no plane of weakness in diamondthe molecule so diamond isexceptionally hard and becauseall the electrons are localized itdoes not conduct electricity.Both silicon and silicondioxide, Si02, form similargiant tetrahedral structures.

In graphite each carbon atom has very strong bonds tothree other carbon atoms to give layers of hexagonalrings. There are only very weak bonds between thelayers. The layers can slide over each other so graphiteis an excellentlubricant andbecause theelectrons aredelocalizedbetween the layersit is a goodconductor ofelectricity.

A third allotrope of carbon is buckminsterfullerene.Sixty carbon atoms are arranged in hexagons andpentagons to give a geodesic spherical structuresimilar to a football. Following the initial discoveryof buckminsterfullerene many other similar carbonmolecules have been isolated. This has led to anew branch of science called nanotechnology.

Metallic bonding and physical properties related tobonding type

METALLIC BONDINGThe valence electrons in metals become detachedfrom the individual atoms so that metals consist ofa close packed lattice of positive ions in a sea ofdelocalized electrons. A metallic bond is theattraction that two neighbouring positive ions havefor the delocalized electrons between them.Metals are malleable, that is, they can be bent andreshaped under pressure. They are also ductile,which means they can be drawn out into a wire.

delocal izedelectrons

Metals are malleable and ductile because the close-packed layers of positive ions can slide over eachother without breaking more bonds than are made.

••88888888888888888+

Impurities added to the metal disturb the lattice andso make the metal less malleable and ductile. This iswhy alloys are harder than the pure metals they aremade from.

nucleus andinner shells

TYPE OF BONDING AND PHYSICAL PROPERTIESMelting and boilinq pointsWhen a liquid turns into a gas the attractive forces between the particles arecompletely broken so boiling point is a good indication of the strength ofintermolecular forces. When solids melt the crystal structure is broken down, butthere are still some attractive forces between the particles. Melting points areaffected by impurities. These weaken the structure and result in lower meltingpoints.

SOlubility'Like tends to dissolve like'. Polarsubstances tend to dissolve in polarsolvents, such as water, whereas non-polar substances tend to dissolve innon-polar solvents, such as heptaneor tetrachloromethane. Organicmolecules often contain a polar headand a non-polar carbon chain tail. Asthe non-polar carbon chain lengthincreases in an homologous series themolecules become less soluble inwater. Ethanol itself is a good solventfor other substances as it containsboth polar and non-polar ends.

Covalent macromolecular structures have extremely high melting and boilingpoints. Metals and ionic compounds also tend to have relatively high boilingpoints due to ionic attractions. Hydrogen bonds are in the order of fa th thestrength of a covalent bond whereas van der Waals' forces are in the order of lessthan ,bo of a covalent bond. The weaker the attractive forces the more volatile thesubstance.

CH30H

C2HsOH

C3H70H

C4HgOH

NaCIdecreasingsolubility in water

tDiamond (melting point over 4000 °C)All bonds in the macromolecularstructure covalent

Sodium chloride (melting point 801°C)Ions held strongly in ionic lattice

H toI II

H-C-C-HI b+

H

H H HI I I

H-C-C-C-HI I IH H H

H HI I b-Oo+OH-C-C-O-HI IH H

CompoundMrM. pt I °CPolarityBonding type

ethanal4420.8polardipole:dipole

ethanol4678.5polarhydrogen bonding

propane44

-42.2non-polarvan der Waals' Ethanol is completely miscible with water as it

can hydrogen-bond to water molecules.

ConductivityFor conductivity to occur the substance must possess electrons or ions that are free to move. Metals (and graphite) containdelocalized electrons and are excellent conductors. Molten ionic salts also conduct electricity, but are chemicallydecomposed in the process. Where all the electrons are held in fixed positions, such as diamond or in simple molecules,no electrical conductivity occurs.

When an ionic compound melts, the ions are free to move to oppositely chargedelectrodes. Note: in molten ionic compounds it is the ions that carry the charge,not free electrons.When a potential gradient is applied to the metal, the delocalized

electrons can move towards the positive end of the gradientcarrying charge. 8 88 + 8

heat 8 88 8NaCI 8 88 8

+

Bonding 23

Molecular orbitals and hybridization (1)

COMBINATION OF ATOMIC ORBITALS TO FORM MOLECULAR ORBITALSAlthough the Lewis representation is a useful model to represent covalent bonds it does make the false assumption that all thevalence electrons are the same. A more advanced model of bonding considers the combination of atomic orbitals to formmolecular orbitals.

a bondsA a (sigma) bond is formed when two atomic orbitals on differentatoms overlap along a line drawn through the two nuclei. Thisoccurs when two s orbitals overlap, an s orbital overlaps with a porbital, or when two p orbitals overlap 'head on'.

Jt bondsA Jt (pi) bond is formed when two p orbitals overlap'sideways on'. The overlap now occurs above andbelow the line drawn through the two nuclei. A Jt bondis made up of two regions of electron density.

rt bond a bond

~ ~o,'=""\§below line of centress p

([X)t

CDt

HYBRIDIZA TlON (1)

Sp3 hybridizationMethane provides a good example of Sp3 hybridization. Methane contains four equal C-H bonds pointing towards thecorners of a tetrahedron with bond angles of 109.5°. A free carbon atom has the configuration 1s22s22p2. It cannot retain thisconfiguration in methane. Not only are there only two unpaired electrons, but the p orbitals are at 90° to each other and willnot give bond angles of 109.5° when they overlap with the s orbitals on the hydrogen atoms.

When the carbon bonds in methane one of its 2s electrons is promoted to a 2p orbital and then the 2s and three 2p orbitalshybridize to form four new hybrid orbitals. These four new orbitals arrange themselves to be as mutually repulsive aspossible, i.e. tetrahedrally. Four equal a bonds can then be formed with the hydrogen atoms.

lower in energy than2p orbitals, so moreenergetically favourable

\11 1 11 11 11 1

~ 2p CD 2p25

electron25 25 and three 2p.

promoted orbitals hybridiz~

0 free carbon ~15 atom 15

11 11 11 1

sp3

o15

24 Bonding

Molecular orbitals and hybridization (2)

HYBRIDIZA TlON (2)

Sp2 hybridizationSp2 hybridization occurs in ethene. After a 2s electron on the carbon atom is promoted the 2s orbital hybridizes with two ofthe 2p orbitals to form three new planar hybrid orbitals with a bond angle of 1200 between them. These can form 0 bondswith the hydrogen atoms and also a 0 bond between the two carbon atoms. Each carbon atom now has one electronremaining in a 2p orbital. These can overlap to form a it bond. Ethene is thus a planar molecule with a region of electrondensity above and below the plane.

11 11 12p: ,

one 2p orbitalremains

it bond (above and below plane)

1L~oo0~CCJDc~ond

&O~O~

CD2s

2s and two2p orbitalshybridize

ethene

sp hybridizationsp hybridization occurs when the 2s orbital hybridizes with just one of the 2p orbitals to form two new linear sp hybridorbitals with an angle of 1800 between them. The remaining two p orbitals on each carbon atom then overlap to form two itbonds. An example is ethyne.

Two it bonds at90° to each other

1l

CD2s

11 11 1:2p: two 2p orbitals

remain

2s and one2p orbitalshybridize

RELA TlONSHIP BETWEEN TYPEOF HYBRIDIZA TION, LEWIS Hybridization Regular bond angle Examples

HSTRUCTURE, AND MOLECULAR Sp3 109S I 0. 0. 0.SHAPES ,C N ,OJ

,N____N-'H/J "H H/~ "H H/~ H/ ~ / ~Molecular shapes can be arrived at either

H H H H H H •by using the VSEPR theory or by knowing hydrazinethe type of hybridization. Hybridization

H" /H H" l\i0 °can take place between any sand p orbital Sp2 IIin the same energy level and is not just

120°/C=C" N= /C"(::. "

restricted to carbon compounds. If the H H H H H

shape and bond angles are known fromusing Lewis structures then the type ofhybridization can be deduced. Similarly if sp 180° H-C=C-H (N=N)the type of hybridization is known theshape and bond angles can be deduced.

Bonding 25

Delocalization of electrons

RESONANCE STRUCTURESWhen writing the Lewis structures for some moleculesit is possible to write more than one correct structure.For example, ozone can be written:

These two structures are known as resonance hybrids.They are extreme forms of the true structure, which liessomewhere between the two. Evidence that this is truecomes from bond lengths, as the bond lengths betweenthe oxygen atoms in ozone are both the same and areintermediate between an 0=0 double bond and an0-0 single bond. Resonance structures are usuallyshown with a double headed arrow between them.Other common compounds which can be written usingresonance structures are shown here.

DELOCALlZA TION OFELECTRONSResonance structures can also beexplained by the delocalization ofelectrons. For example, in theethanoate ion the carbon atom and thetwo oxygen atoms each have a porbital containing one electron afterthe 0 bonds have been formed. Insteadof forming just one double bondbetween the carbon atom and one ofthe oxygen atoms the electrons candelocalize over all three atoms. This isenergetically more favourable thanforming just one double bond.

Delocalization can occur wheneveralternate double and single bondsoccur between carbon atoms. Thedelocalization energy in benzene isabout 150 k] mol:", which explainswhy the benzene ring is so resistant toaddition reactions.

26 Bonding

o -0II I

CH3 - C - o-+----; CH3 - C = 0

H HI I

H.....C

CC/H H

IB QUESTIONS - BONDING

1. Which compound contains both covalent and ionicbonds?

A. sodium carbonate, Na2C03B. magnesium bromide, MgBr2C. dichloromethane, CH2CI2D. ethanoic acid, CH3COOH

2. Which pair of elements is most likely to form a covalentlybonded compound?

A. Li and CI

B. P and °C. Ca and S

D. Zn and Br

3. Given the following electronegativities,H: 2.2 N: 3.0 0: 3.5 F: 4.0

which bond would be the most polar?

A. O-H in H20

B. N-F in NF2

C. N-O in N02D. N-H in NH3

4. What is the correct Lewis structure for methanal?

A. H:C:::O:H C. HC::O:H

D. :C:O:HH

B. HH:C::O:

5. When CH4, NHy H20, are arranged in order of increasingbond angle, what is the correct order?

A. CH4, NHy H20 C. NHy CH4, H20

B. NHy H20, CH4 D. H20, NHy CH4

6. When the H-N-H bond angles in the species NH2, NH3,NHr are arranged in order of increasing bond angle(smallest bond angle first), which order is correct?

A. NH2 < NH3 < NH~ C. NH3 < NH2 < NH~

B. NH~ < NH3 < NH:; D. NH:; < NHr < NH3

7. In which of the following pairs does the second substancehave the lower boiling point?

A. F2, CI2B. H20, H2S

C. C2H6, C3HaD. CH30CHy CH3CH20H

8. In which of the following substances would hydrogenbonding be expected to occur?

I. CH4

II. CH3COOH

III. CH30CH3

A. II only

B. I and III only

C. I and III only

D.I, II and III

9. Which one of the following statements is correct?

A. The energy absorbed when liquid ammonia boils isused to overcome the covalent bonds within theammonia molecule.

B. The energy absorbed when sol id phosphorus (P4) meltsis used to overcome the ionic bonds between thephosphorus molecules.

C. The energy absorbed when sodium chloride dissolvesin water is used to form ions.

D. The energy absorbed when copper metal melts is usedto overcome the non-directional metallic bondsbetween the copper atoms.

10. A solid has a melting point of 1440 "C. It conducts heatand electricity. It does not dissolve in water or in organicsolvents. The bond between the particles is most likely tobe

A. covalent.

B. dipole:dipole.

C. ionic.

D. metallic.

11.What are the types of hybridization of the carbon atoms inthe compound

H2CIC-CHrCOOH ?123

3.Sp2

sp

12. Which molecule or ion does not have a tetrahedral shape?

A. XeF4

B. SiCI4

c. BF:!D. NH4

13.When the substances below are arranged in order ofincreasing carbon-carbon bond length (shortest bond first),what is the correct order?

I. H2CCH2

A. I < II < IIIB. I < III < II

III.©II. H3CCH3

C. II < I < III

D.III < II < I

14.Which of the following species is considered to involve sp-'hybridization?

I. BCI3

A. lonly

B. II only

II. CH4 III. NH3

C. I and III only

D. II and III only

15. How many J1 bonds are present in CO2?

A. OneB. Two

C. Three

D. Four

16. When the following substances are arranged in order ofincreasing melting point (lowest melting point first) thecorrect order is

A. CH3CH2CH3, CH3COCH3, CH3CH2CH20H

B. CH3CH2CHy CH3CH2CH20H, CH3COCH3C. CH3COCHy CH3CH2CH20H, CH3CH2CH3D. CH3CH2CH20H, CH3CH2CHy CH3COCH3

IB questions - Bonding 27

Wt &

Enthalpy changes

enthalpy, H

EXOTHERMIC AND ENDOTHERMIC REACTIONSEnergy is defined as the ability to do work, that is, move a force through a distance. It is measured in joules.

Energy = force x distance(J) (N x m)

In a chemical reaction energy is required to break the bonds in the reactants, and energy is given out when new bonds areformed in the products. The most important type of energy in chemistry is heat. If the bonds in the products are stronger thanthe bonds in the reactants then the reaction is said to be exothermic, as heat is given out to the surroundings. Examples ofexothermic processes include combustion and neutralization. In endothermic reactions heat is absorbed from the surroundingsbecause the bonds in the reactants are stronger than the bonds in the products.

The internal energy stored in the reactants is known as its enthalpy,H. The absolute value of the enthalpy of the reactants cannot beknown, nor can the enthalpy of the products, but what can bemeasured is the difference between them, !1H. By convention !1Hhas a negative value for exothermic reactions and a positive valuefor endothermic reactions. It is normally measured under standardconditions of 1 atm pressure at a temperature of 298 K. Thestandard enthalpy change of a reaction is denoted by !1H".

TEMPERATURE AND HEATIt is important to be able to distinguish between heat andtemperature as the terms are often used loosely.

• Heat is a measure of the total energy in a given amount ofsubstance and therefore depends on the amount ofsubstance present.

• Temperature is a measure of the 'hotness' of a substance.It represents the average kinetic energy of the substance,but is independent of the amount of substance present.

\\ (r-- 50°C (' Two beakers of water.

r- 50°C ~Both have the sametemperature, but the100 cm-' of water contains

~50 cm-'twice as much heat as the

~ 100 ern-50 crn '.

simple calorimeter

polystyrene

reactionmixture

reactants

!1H = Hproducts - Hreactants(value negative)

products (more stable than reactants)

Representation of exothermic reaction usingan enthalpy diagram.

enthalpy, H

products (less stable than reactants)

!1H = Hproducts - Hreactants(value positive)

reactantsRepresentation of endothermic reaction usingan enthalpy diagram.

CALORIMETRYThe enthalpy change for a reaction can be measuredexperimentally by using a calorimeter. In a simplecalorimeter all the heat evolved in an exothermic reaction isused to raise the temperature of a known mass of water. Forendothermic reactions the heat transferred from the water tothe reaction can be calculated by measuring the lowering oftemperature of a known mass of water.

To compensate for heat lost by the water in exothermicreactions to the surroundings as the reaction proceeds a plotof temperature against time can be drawn. By extrapolatingthe graph, the temperature rise that would have taken placehad the reaction been instantaneous can be calculated.

Compensatingfor heat lostTo = initial temperature of

reactantsTl = highest temperature

actually reachedT2 = temperature that would have

been reached if no heat:- - - - ___ lost to surroundings, --

extrapolation atsame rate ofcooling

!1Tfor reaction = T2 - To

reactants mixed

time

28 Energetics