class04 chemistryg12 notes and homework

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GRADE 12 CHEMISTRY

OLYMPIAD SCHOOL

‐. .

FRIDAY 4.45‐6.45



Structures and

Properties of Substances

em ca on ng• Of more than one hundred elements that occur in nature or that have

een pro uce syn e ca y, on y e no e gases ex s na ura y assingle, uncombined atoms.

• The atoms of all other elements occur in some combined form,

on e oge er.• Chemical bonds are electrostatic forces that hold atoms together incompounds.

• Why, however, do atoms form bonds at all? The answer involvesenergy. In nature, systems of lower energy tend to be favoured over

systems of higher energy.• Said in a different way, lower-energy systems tend to have greater

stability than higher-energy systems .

, , ,

uncombined atoms.



Using Lewis Structures to Represent

Atoms in Chemical Bonding• Chemical bonding involves the interaction of valence electrons

• Valence Electrons: the electrons that occupy the outermost principal

energy level of an atom.•electrons of atoms.

• To draw the Lewis structure of an atom, replace its nucleus and inner .

symbol to represent the valence electrons (Many chemists place thedots starting at the top and continue adding dots clockwise).

•molecules and the simplest formula unit of an ionic solid. Drawing aLewis structure for a molecule lets you see exactly how manyelectrons are involved in each bond which hel s to kee track of the

number of valence electrons. Some chemists use dots only. Other chemists show the bonding pairs as lines between atoms. Dots arereserved for re resentin a lone air a non-bondin air of electrons.



Ionic Bonding• The force of attraction between oppositely

charged ions, cations {+} and anions {-},.

occurs between atoms of elements that

have large differences in electronegativity—electronegativity and a non-metal withhigh electronegativity.

• e un s o on c compoun s suc assodium chloride and magnesium fluoridecannot be separated easily by directea ng o e crys a sa s. e ons a

make up the ionic solid are arranged in a

specific array of repeating units. In solidso um c or e, or examp e, e ons arearranged in a rigid lattice structure. Insuch systems, the cations and anions arearranged so that the system has theminimum possible energy.



Because of the large differences in electronegativity, the atoms in

p block non-metals.For example, magnesium in Group 2 and fluorine in Group 7com ne o orm e on c compoun magnes um uor e, g .Through bonding, the atoms of each element obtain a valenceelectron confi uration like that of the nearest noble as. In thiscase, the nearest noble gas for both ions is neon.



• Magnesium fluoride has many properties that are characteristic. ,

properties:

• Cr stalline with smooth shin surfaces

• Hard but brittle

• Non-conductors of electricity and heat

• High melting points• Most ionic solids are also soluble in water.



Covalent Bonding• ova en on s are orces o a rac on orme w en a oms o a

molecule share electrons. Covalent bonding involves a balancebetween the forces of attraction and repulsion that act between thenuclei and electrons of two or more atoms.

• This idea is represented with a molecule of hydrogen, H2. There is anelectron attractions, nucleus-nucleus repulsions, and electron-electron repulsions achieve this balance. This optimum separation

,covalent bond between the two hydrogen atoms. Unlike ionicbonding, in which electrons behave as if they are transferred fromone a om o ano er, cova en on ng nvo ves e s ar ng o pa rs oelectrons.



• The quantum mechanical model, which we applied to theatom, can be extended to explain bonding.

• A covalent bond ma form when two half-filled atomic orbitals

from two atoms overlap to share the same region of space. Acovalent bond involves the formation of a new orbital, causedby the overlapping of atomic orbitals.

• The new orbital has energy levels that are lower than those

of the original atomic orbitals. Since electrons tend to occupy,more energetically favourable configuration than the twoatoms had before they interacted.



Characteristics of Covalent Bondin• Electron-sharing enables each atom in a

covalent bond to acquire a noble gas. ,

filled valence level like He by treating theshared electrons like part of its owncomposition.

• s ng e s are pa r o e ec rons - a on ngpair - fills the valence level of both hydrogenatoms at the same time. The period 2 non-

, ,

three 2p orbitals to acquire a noble gasconfiguration like Ne. Covalent bonding thatinvolves these elements obeys the octet rule

.• In the formation of the diatomic fluorine

molecule, F2, for example, the bondingatom a complete valence level. Each fluorineatom also has three unshared pairs of electrons. These pairs of electrons, called lone

, .



The covalent bond that holds molecules of hydrogen, fluorine, and

.

bonding pair of electrons. Some molecules are bonded together

with two shared airs of electrons. These are called double bonds. Carbon dioxide is an example of a covalent molecule that consists

of double bonds. Molecules that are bonded with three shared

pairs of electrons have triple bonds. Nitrogen, N2, another diatomic molecule, is a triple-bonded molecule.



Bond Ener• In chemistry, bond energy (E ) is a measure of bond strength in a chemical

bond. It is the heat required to break Avogadro's number of molecules intoe r n v ua a oms. on energy s e energy requ re o rea e orce

of attraction between two atoms in a bond and to separate them. Thus, bond

energy is a measure of the strength of a bond.• ou m g t expect t at t e on energy wou ncrease more e ectrons

were shared between two atoms because there would be an increase incharge density between the nuclei of the bonded atoms. In other words, youm g pre c a ou e on s are s ronger an s ng e on s, an atriple bonds are stronger than double bonds.



Pr i in I ni n v l n B n • What if you know the formula of a compound but don’t know

how atoms with hi her electrone ativit values attract electronsmore strongly than atoms with the lower electronegativity

values. [ELECTRONEGATIVITY: the ability of an atom to attract.

• In sodium chloride, for instance, chlorine (EN = 3.16) attracts anelectron much more stron l than sodium EN = 0.93 .

Therefore, sodium’s valence electron has a very highprobability of being found near chlorine, and a very low

.

• A high electronegativity difference is characteristic of ionic

com ounds.



- It is common for chemists todescribe the formation of ioniccompounds in terms of a transferring

of electrons from one atom to another.For example, in forming sodiumchloride, you often read that sodiumloses its valence electron and thatchlorine gains an electron.

- Terms such as these - electronrans er, os ng e ec rons, ga n ng

electrons - make it easier to discussand model the behaviour of atoms

formation of ionic compounds.- The figure provides a guide to

c ass y ng c em ca on s ase on

their electronegativity difference.

Notice that this “bonding continuum,”as it is sometimes called, avoids

definitive classifications. Instead, the

character of chemical bonds chan es

gradually with a change in

electronegativity.



• Chemists consider bonded atoms with a difference in EN between 0 and 0.4 as being mostly covalent; electrons areshared equally or nearly equally.

• For bonded atoms that have a difference in EN between 0.4. , .bond is a covalent bond with an unequally shared pair of electrons between two atoms.

• This unequal sharing results in a bond that has partially positiveand partially negative poles. Bonded atoms with a difference in

. .so unequally that one bonded atom has a strong negative

charge and the other has a strong positive charge.



How To Draw Lewis Structures for

Sim le Molecules and Ions with a Central Atom

• Step 1: Position the least electronegative atom in the centre of

the molecule or polyatomic ion. Write the other atoms around,a single bond.

• Step 2: (a) Determine the total number of valence electrons ine mo ecu e or on. or po ya om c ons, pay c ose a en on o

the charge. For example, if you are drawing a polyatomic anionsuch as CO3 2-, add two electrons to the total number of va ence e ec rons ca cu a e or e s ruc ure . or apolyatomic cation such as NH4+, subtract one electron from the

total number of valence electrons calculated for the structure.

• (b) Once you have the total number of valence electrons,

to achieve a noble gas electron configuration.



• the second total (number of electrons needed to satisfy theoctet rule) to get the number of shared electrons.

• Then divide this number by 2 to give the number of bonds.

Double or triple bonds may be needed to account for this. .bonds count as three bonds.

•

number of valence electrons to get the number of non-bondingelectrons.

• Add these electrons as lone pairs to the atoms surrounding thecentral atom so that you achieve a noble gas electron

.



Example: Lewis structure of the nitrite ion

The formula of the nitrite ion is NO2−.

• , .

•Step two: Count valence electrons. Nitrogen has 5 valence electrons; each oxygenhas 6, for a total of (6 × 2) + 5 = 17. The ion has a charge of −1, which indicates an

extra electron so the total number of electrons is 18.

•Step three: Place ion pairs. Each oxygen must be bonded to the nitrogen, which uses

four electrons — two in each bond. The 14 remaining electrons should initially be placed

as 7 lone pairs. Each oxygen may take a maximum of 3 lone pairs, giving each oxygen

8 electrons including the bonding pair. The seventh lone pair must be placed on the

nitrogen atom.

•Step four: Satisfy the octet rule. Both oxygen atoms currently have 8 electrons

assigned to them. The nitrogen atom has only 6 electrons assigned to it. One of the

lone pairs on an oxygen atom must form a double bond, but either atom will work

equally well. We therefore must have a resonance structure.

• tep ive: Tie up loose ends. Two Lewis structures must be drawn: one with each

oxygen atom double-bonded to the nitrogen atom. The second oxygen atom in each

structure will be single-bonded to the nitrogen atom. Place brackets around eachs ruc ure, an a e c arge − o e upper r g ou s e e rac e s. raw a

double-headed arrow between the two resonance forms.



Co-ordinate Covalent Bonds• A covalent bond involves the sharing of a pair of electrons between two

atoms; each atom contributes one electron to the shared pair.

• , , ,of the electrons to the shared pair. The bond in these cases is called a co-ordinate covalent bond. In terms of the quantum mechanical model, a co-

ordinate covalent bond forms when a filled atomic orbital overlaps with anemp y a om c or a . nce a co-or na e on s orme , e aves n esame way as any other single covalent bond.

• Note that the Lewis structure for NH4+ does not indicate which atomprov es eac s are pa r o e ec rons aroun e cen ra n rogen a om.

• However, the quantum mechanical model of the atom can explain thebonding around this nitrogen atom. The condensed electron configuration for nitrogen is [Ne] 2s 2 .

• Each nitrogen atom has only three unpaired 2p electrons in three half-filled

orbitals available for bonding.• Since there are four covalent bonds shown around nitrogen in the Lewis

structure, electrons in one of the bonds must have come from the filledorbitals of nitro en.

• Therefore, one of the bonds around the central nitrogen atom must be a co-ordinate covalent bond.



Resonance Structures: More Than

One Possible Lewis Structure•

sulfur dioxide, SO2. A typical answer would look like this

• This Lewis structure suggests that SO2 contains a singleon an a ou e on .

• However, experimental measurements of bond lengthsindicate that the bonds between the S and each O are

.• The two bonds have properties that are somewhere between

a single bond and a double bond. In effect, the SO2 molecule

contains two “one-and-a-half” bonds.• To communicate the bonding in SO2 more accurately,

chemists draw two Lewis structures and insert a double-headed arrow between them.

E h f th L i t t i ll d t t



• Each of these Lewis structures is called a resonance structure.

• Resonance structures are models that give the same relative positionof atoms as in Lewis structures but show different laces for their

bonding and lone pairs.

• Many molecules and ions—especially organic ones—requireresonance structures to represent their bonding. It is essential to bear in mind that resonance structures do not exist in reality.

• For example, SO2 does not shift back and forth from one structure tothe other. An actual SO2 molecule is a combination—a hybrid—of itstwo resonance structures. It is more properly referred to as a

.• You could think of resonance hybrids as being like “mutts.” For

example, a dog that is a cross between a spaniel and a beagle (a

“s ea le” is art s aniel and art bea le.• It does not exist as a spaniel one moment and a beagle the next.

Instead, a speagle is an average, of sorts, between the two breeds.

• ,

more resonance structures.



Central Atoms with an Expanded

a ence eve• The octet rule allows a maximum of four bonds (a total of eight

e ec rons o orm aroun an a om.

• Based upon bond energies, however, chemists have suggested

best explained by a model that shows more than eight electronsin the valence energy level of the central atom.

• This central atom is said to have an expanded valence energylevel. One example of a molecule with an expanded valence, , ,

pale-yellow powder that is used in various industries.

U til tl h i t l i d b di i l l h



• Until recently, chemists explained bonding in molecules such asPCl5 by assuming that empty d orbitals of the central atom

.

• However, experimental evidence does not support this idea.Current thinkin su ests that lar er atoms can accommodateadditional valence electrons because of their size.

• Assume that the octet rule usually applies when drawing Lewiss ruc ures or s mp e mo ecu es.

• Sometimes, however, you must violate the rule to allow for .



Shapes and Polarity of Molecules

n ro uc ng a ence- e-

Theory• The fundamental principle of the Valence-Shell Electron-

Pair Re ulsion theor is that the bondin airs and lone

non-bonding pairs of electrons in the valence level of anatom repel one another.

• As you know electron pairs of atoms are localized in orbitals



As you know, electron pairs of atoms are localized in orbitals,which are shapes that describe the space in which electronsare most likel to be found around a nucleus. The orbital for

each electron pair is positioned as far apart from the other orbitals as possible.

between electron pairs. A lone pair (LP) will spread out more

than a bond pair. Therefore, the repulsion is greatest between.between the atomic nuclei, so they spread out less than lonepairs. Therefore, the BP-BP repulsions are smaller than the

.

lone-pair (BPLP) has a magnitude intermediate between theother two.

• In other words, in terms of decreasing repulsion:LP - LP > LP - BP > BP - BP.

• ,the trigonal planar shape around a carbon atom with onedouble bond, and the linear shape around a carbon atom with a

bonding pairs of electrons.

• The figure shows the five basic geometrical arrangements that



• The figure shows the five basic geometrical arrangements thatresult from the interactions of lone pairs and bonding pairs

.electron groups. An electron group is usually one of thefollowing:

• a single bond

• a double bond• a triple bond

• a lone pair

• ac o t e e ectron-group arrangements resu ts n a m n mumtotal energy for the molecule. Variations are possible,de endin on whether a bondin air or a lone air occu ies a specific position.

• It is important to distinguish between electron-group arrangement an mo ecu ar s ape .

• Electron-group arrangement refers to the ways in which groups.

Molecular shape refers to the relative positions of the atomicnuclei in a molecule.

When all the electron groups are bonding pairs a molecule will have one of the five



When all the electron groups are bonding pairs, a molecule will have one of the five

geometrical arrangements shown . If one or more of the electron groups include a lone

air of electrons , variations in one of the five eometric arran ements result .



• Each of the molecules has four pairs of electrons around the central



patom.

• Observe the differences in the number of bondin and lone airs in

these molecules. Methane, CH4, has four BPs. Ammonia, NH3, hasthree BPs and one LP. Water, H2O, has two BPs and two LPs. Noticethe effect of these differences on the shapes and bond angles of themolecules.

• Methane, with four BPs, has a molecular shape that is the same as

the electron- rou arran ement in which four electron rou ssurround a central atom: tetrahedral. The angle between any twobonding pairs in the tetrahedral electron-group arrangement is109.5°. This angle corresponds to the most favourable arrangemento e ec ron groups o m n m ze e orces o repu s on among em.

• When there are 1 LP and 3 BPs around a central atom, there are twotypes of repulsions: LP - BP and BP - BP. Since LP - BP repulsionsare grea er an - repu s ons, e on ang e e ween ebond pairs in NH3 is reduced from 109.5° to 107.3°. When you drawthe shape of a trigonal pyramidal molecule, without the lone pair, you

triangular base. In a molecule of H2O, there are two BPs and twoLPs. The strong LPLP repulsions, in addition to the LPBP repulsions,

104.5.

The table summarizes the molecular shapes that commonly occur. The VSEPR



p y

notation used for these shapes adopts the letter “A” to represent the central atom,

the letter “X” to represent a bonding pair, and the letter “E’ to represent a lone pair

of electrons. For example, the VSEPR notation for NH3 is AX3E. This indicatesthat ammonia has three BPs around its central atom, and one LP.



• 1. Draw a preliminary Lewis structure of the molecule based on

.

• 2. Determine the total number of electron groups around the

central atom (bonding pairs, lone pairs and, where applicable,account for the charge on the ion). Remember that a doublebond or a triple bond is counted as one electron group.

.

accommodate this total number of electron groups.• .

by the bonding pairs and lone pairs.



The Relationship Between



p

Polarit• A molecule’s shape and polarity are directly related. As you

know, the electronegativity difference between two atoms is thepr nc pa ac or a e erm nes e ype o c em ca on ngbetween them. When atoms of two different elements have adifference in EN between 0.5 and 1.7 the bond between them

is polar covalent. The electrons are shared unequally betweenthe two types of atoms.

• or examp e, n e mo ecu e, e a om s moreelectronegative than the H atom. Thus, the Cl atom exerts a

stron er force of attraction on the shared air of electrons. As aresult, the Cl end of the bond develops a partial negativecharge (delta neg) and the H end develops a partial positive

.

dipole .

Compare two molecules with the same molecular shape: CCl4



and CCl3H. Since the polarity of the CCl bond is different from

that of the CH bond, the polarities of these two molecules aredifferent. The CCl4 molecule is symmetrical about any axis

.

bonds counteract one another. Thus, the carbon tetrachloride

molecule is non- olar. In the case of CCl3H the olarit of theHCl bond is different from that of the three CCl bonds. Thus, the

molecule is polar.



Intermolecular Forces



Intermolecular Forces

in Liquids and Solids• a er an y rogen su e ave e same mo ecu ar s ape: en .

Both are polar molecules. However, H2O, with a molar mass of 18 g,

is a li uid at room tem erature while H2S with a molar mass of 34is a gas.

• Water’s boiling point is 100°C, while hydrogen sulfide’s boiling point

s - . ese proper y erences are cu o exp a n us ng w a

we know thus far.

•solids, between individual atoms in covalent and polar covalentcompounds, and between positive metal ions and free electrons in

. —molecule or polyatomic ion. During a chemical change, intramolecular forces are overcome and chemical bonds are broken. Therefore,

.

that influence the physical properties of substances are calledintermolecular forces.

• These are forces of attraction and repulsion that actbetween molecules or ions



between molecules or ions.

• n ermo ecu ar orces were ex ens ve y s u e y eDutch physicist Johannes van der Waals (1837–1923). Tomark his contributions to the understandin of intermolecular attractions, these forces are often calledvan der Waals forces.

• ey are ca egor ze n o e o ow ng groups: po e-dipole forces, ion-dipole forces, induced dipole forces,dis ersion London forces and h dro en bondin .

Di ole Di ole Forces



Di ole-Di ole Forces• In the liquid state, polar molecules (dipoles) orient themselves so that

oppositely charged ends of the molecules are near to one another.

• -forces.

• As a result of these dipole-dipole forces of attraction, polar molecules will

non-polar molecules would.

• The energy required to separate polar molecules from one another istherefore reater than that needed to se arate non- olar molecules of similar molar mass.

• This is indicated by the extreme difference in melting and boiling points of these two types of molecular substances.

Ion Di ole Forces



Ion-Di ole Forces• Sodium chloride and other soluble ionic solids dissolve in polar

solvents such as water because of ion-dipole forces.

• n on- po e orce s e orce o a rac on e ween an on an apolar molecule (a dipole).

• For example, NaCl dissolves in water because the attractionsbetween the Na+ and Cl- ions and the water molecules provideenough energy to overcome the forces that bind the ions together.

• The fi ure shows how ion-di ole forces dissolve an t e of solubleionic compound.

Dis ersion London Forces



Dis ersion London Forces• The shared pairs of electrons in covalent bonds are constantly

vibratin . While this a lies to all molecules it is of articular interest for molecules that are non-polar. The bond vibrations,which are part of the normal condition of a non-polar molecule,

cause momentar uneven distributions of char e.• In other words, a non-polar molecule becomes slightly polar for

an instant, and continues to do this on a random but on-going. -

polar condition, it is capable of inducing a dipole in a nearbymolecule. An intermolecular force of attraction results. This.

commonly refer to it as a London force, in honour of theGerman physicist, Fritz London, who studied this force.

• n e o er n ermo ecu ar orces, spers on on on orcesact between any particles, polar or otherwise. They are themain intermolecular forces that act between non-polar molecules.

Two factors affect the magnitude of dispersion forces.One is the number of electrons in the molecule and the other is the shape of the



molecule. Vibrations within larger molecules that have more electrons thansmaller molecules can easily cause an uneven distribution of charge.The dispersion forces between these larger molecules are thus stronger, whichhas the effect of raising the boiling point for larger molecules.

A molecule with a spherical shape has a smaller surface area than a straight

chain molecule that has the same number of electrons. The smaller surface

area allows less o ortunit for the molecule to induce a char e on a nearbmolecule. Therefore, for two substances with molecules that have a similar

number of electrons, the substance with molecules that have a more spherical

.

H dro en Bondin



H dro en Bondin• A dipole-dipole interaction that is very significant in many

,H2O(l), is hydrogen bonding.

• Hydrogen bonding is a particularly strong form of dipole-dipole-bonded molecule that contains bonds such as HO, HN, or HF,and an unshared pair of electrons on another small,e ec ronega ve a om suc as , , or .

• The small, electronegative atom can be on its own, but isusuall bonded in a molecule.

• When hydrogen is bonded to oxygen, fluorine, or nitrogen, thestrong force of attraction exerted by any of these three very

hydrogen atom in the polar-bonded molecule, leaving thehydrogen atom with a partial positive charge.

•

on a nearby electronegative atom: N, O, or F.

• You can see the effect of hydrogen bonding clearly in the boiling pointdata of the binary hydrides of Groups 14 to 17 (IVA to VIIA). Theincrease in boiling point from methane CH4 to tin(IV) hydride SnH4



increase in boiling point from methane, CH4, to tin(IV) hydride, SnH4,.

• Therefore, there are larger dispersion forces, since more electronsare able to temporarily shift from one part of the molecule to another.In SnH4 this results in reater intermolecular forces of attraction.There is no hydrogen bonding in these substances, because thereare no lone pairs on the molecules.

• However, the hydrides of Groups 15, 16, and 17 (VA to VIIA) only mass hydrides.

• The smallest mass hydrides, NH3, H2O, and HF, have relatively high,

hydrogen bonding.• Why is there no hydrogen bonding in H2S and H2Se, which have

unshared electron airs on their central atoms? Neither S nor Se issma enoug or e ec ronega ve enoug o suppor y rogenbonding.

• Thus, H2O (with its small and very electronegative O atom) has a.

• A hydrogen bond is only about 5% as strong as a single covalentbond. However, in substances that contain numerous hydrogen

.example, the double-helix structure of DNA occurs because of hydrogen bonds.

class04 chemistryg12 notes and homework

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