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Unit 02

Structure &BondingAL Chemistry

NIVANTHA SATHARASINGHE



CHEMISTRY Advanced Level Unit 2

NIVANTHA SATHARASINGHE Page 2

What is a chemical bond?

An electrostatic attraction between atoms

How chemical bonds are formed?

Atoms try to achieve more stable electronic configurations by exchanging or sharing electrons

Primary Interactions

1. COVALENT BOND-

• consists of a shared pair of electrons, each atom supplies one electron

• atoms are held because their nuclei are attracted to the shared electrons

Formation

• between atoms of the same element; (N2, O2, diamond and graphite)

• between atoms of different elements on RHS of the periodic table; (CO2, SO2).

• when one of the elements is in the middle of the table; (e.g. C, Si)

• head-of-the-group elements with high ionisation energies, (e.g. Be in BeCl2)

ChemicalBonds

PrimaryInteractions

Covalent Bond

Non-polar

Polar

Co-ordinate

Ionic Bond

Metallic Bond

SecondaryInteractions

Hydrogen Interactions

Dipole-dpole Interactions

Ion-dipole Interactions

Ion-induced dipole Interactions

Dipole-induced dipoleInteractions

Dispersion Interactions





• atoms share electrons in order to complete their „octet‟ of electrons

• some don‟t achieve an „octet‟ as they haven‟t enough electrons - Al in AlCl3

• others share only some - if they share all their „octet‟ is exceed - NH3 and H2O

• atoms of elements in the 3rd period onwards can exceed their „octet‟ because they are not restricted to eight

electrons in their „outer shell‟ - S in SF6

Covalent bonds can be divided into 3 groups;

A.

Non-polar-B.

C.

B. Polar- Polar covalent bonding is the process of unequal sharing of electrons. It is considered the middle

ground between ionic bonding and covalent bonding. It happens due to the differing electronegativity values

of the two atoms. Because of this, the more electronegative atom will attract and have a stronger pulling force

on the electrons. Thus, the electrons will spend more time around this atom.

The symbols above indicate

that on the flourine side it is

slightly negitive and the

hydrogen side is slightly

positive.





C. Co-ordinate-

• differs from a covalent bond only in its formation

• both electrons of the shared pair are provided by one species (donor) and it shares the electrons with the

acceptor

• donor species will have lone pairs in their outer shells

• acceptor species will be short of their “octet” or maximum.

Donor group- (Lewis Base) a lone pair donor

Acceptor group- (Lewis Acid) a lone pair acceptor

e. g. 1

ammonium ion, NH4+

The lone pair on N is used to share with the hydrogen ion which needs two electrons to fill its outer shell. The N

now has a +ive charge as it is now sharing rather than owning two electrons.

e. g. 2

Boron trifluoride-ammonia NH3BF3Boron has an incomplete shell in BF3 and can accept a share of a pair of electrons donated by ammonia. The B

becomes -ive as it is now shares a pair of electrons (i.e. it is up one electron) it didn‟t have before.

Fajan’s Rules• not all ionic compounds have high melting points

• some covalently bonded compounds have higher than expected boiling points this is due to dipoles in their

structure

• reason :- in many substances the bonding is not 100% ionic or covalent

The ideal ionic compound has completely separate, spherical ions and the electron densities are apart

from each other.

If the positive ion has a high charge density it can distort the negative ion by attracting the outer shell electrons

to give an area of electron density between the two species ... a bit like a covalent bond

The feasibility of formation of covalent bonds is predicted using Fajan’s Rules.





The rules A compound is more likely to be covalent if the ...

CATION small size

high charge

“highly polarising” attracts electrons in the

anion

ANION large size

high charge

“highly polarisable” will be easily distorted

N.B. Just because a substance is less likely to be covalent according to Fajan’s Rules doesn’t mean it will be

ionic; it will remain covalent but have some ionic character.

Examples Changes in bond type of chlorides as the positive charge density increases due to higher charge (across

Period 3) or larger size (down Group 1)

What is Electronegativity?

„The ability of an atom to attract the pair of electrons in a covalent bond to itself.‟

Pauling Scale

• a scale for measuring electronegativity

• values increase across periods

• values decrease down groups

• fluorine has the highest value





How Electronegativity differences affect to determine the bond type?

When Electronegativity differences;

1.7 <

1.7 - 0.4 -

0.4 >

Polar molecules

• some molecules are polar if they contain polar bonds

• the molecules will be polar if they have a NET DIPOLE MOMENT

• it is a bit like balanced forces

• non-polar molecule dipoles in bonds within the molecule „cancel each other‟

• polar molecule dipoles do not „cancel each other out‟

I denti fying polar molecules

• place a liquid in a burette

• allow a narrow stream to run out

• place a charged rod next to the flow

• polar molecules will be attracted





Structures contain ing covalent bonds

1. SIMPLE MOLECULES

Bonding - Atoms are joined together within the molecule by covalent bonds.

Electri cal conductivity - Don‟t conduct electricity as they have no mobile ions or electrons.

solubili ty - Tend to be more soluble in organic solvents than in water; some are hydrolysed

boil ing pt - Low - the forces between molecules (intermolecular forces) are weak known as van der Waals

forces

Attractions between molecules increases as the molecules get more electrons.

e.g. CH4 -161°C C2H6 - 88°C C3H8 -42°C

as forces are weak, little energy is required to to separate molecules from each other

so boiling points are low

2. GIANT COVALENT LATTI CES (covalent networks) - DI AMOND, GRAPHI TE and SI LI CA

Bonding - Many atoms joined together in a regular array by large numbers of covalent bonds

Diamond - each carbon atom is joined to four others - Co-ordination No. = 4

Graphite - each carbon atom is joined to three others - Co-ordination No. = 3

melting point - Very hi gh - structures are made up of a large number of covalent bonds, all of which need to

be broken if the atoms are to be separated.

strength –

Diamond and sil ica (SiO2)

hard - exists in a rigid tetrahedral structure





Graphite

soft - consists of layers which are attracted by weak van der Waals’ forces

layers can slide over each other

it used as a lubricant and in pencils

electrical conductivity - Do not conduct electricity as they have no mobile ions or electrons.

But Graphite conducts electr icity

• each atom only uses three of its outer shell electrons for bonding to other atoms

• remaining electron can move through layers allowing the conduction of electricity

• carbon atoms in diamond use all four electrons for bonding so have no free ones

3. MOLECULAR SOLI DS

Iodine

At room temperature, iodine is a grey solid. However, on gentle warming it produces a purple vapour. This is

because iodine is composed of diatomic molecules (I 2) existing in an ordered molecular crystal. Each

molecule is independent and attracted by weak van der Waals’ forces. Therefore, little energy is required to

separate the iodine molecules.





• IONIC BOND [ELECTROVALENT ]-

Formation of ions from atoms

Positive ions [cations]

• formed when electrons are removed from atoms

• are smaller than the original atom

• the energy associated with the process is known as the ionisation energy (IE).

1st I.E. The energy required to remove one mole of electrons (to infinity) from the one mole of gaseous atoms

to form one mole of gaseous positive ions.

e.g. Na(g) —— > Na+(g) + e¯ or Mg(g) —— > Mg+(g) + e¯

There are as many ionisation energy steps as there are electrons in the atom.

2nd I.E. Mg+(g) —— > Mg2+(g) + e¯ and so on

Notes • successive ionisation energies get larger as the proton : electron ratio increases.

3rd ionisation energy > 2nd ionisation energy > 1st ionisation energy

• big jumps in value occur when electrons are removed from shells nearer the

nucleus - less shielding so more energy is needed to overcome the attraction.

1st I.E 500 kJmol-1 2nd I.E 900 kJmol-1 3rd I.E 6000 kJmol-1

The 3rd electron must have been in a shell nearer the nucleus - In Group 2

• if the IE values are very high, covalent bonding is favoured (e.g. beryllium).

Negative ions [anions]

• larger than the original atom due to electron repulsion in outer shell

• formed when electrons are added to atoms

• energy is released as the nucleus pulls in an electron

• this energy is the electron affinity.

Electron Aff inity - The energy change when one mole of gaseous atoms acquires one mole of electrons (from

infinity) to form one mole of gaseous negative ions.

e.g. Cl (g) + e¯ —— > Cl ¯ (g) and O(g) + e¯ —— > O¯ (g)

The greater the effective nuclear charge (ENC) the easier an electron is pulled in.







• 1 electron is transferred from the 3s orbital of sodium to the 3p orbital of chlorine

• both species end up with an „octet‟ of electrons in their outer shell

• the resulting ions are held together in a crystal lattice by electrostatic attraction

GIANT IONIC LATTICES

bonding • oppositely charged ions held in a regular 3-d lattice by electrostatic attraction

• ions pack together in the most efficient way so there is little wasted space

• the arrangement of ions in a lattice depends on the relative sizes of the ions

The Na+ ion is small enough relative to the Cl ¯ ion to fit in the spaces so that both ions occur in every plane.

Each Na+ is surrounded by 6 Cl ¯ (co-ordination number = 6) and each Cl ¯ is surrounded by 6 Na+ (co-ordination

number = 6).

Physical properties of ionic compounds

1. melting pt

Very high A large amount of energy must be put in to overcome the strong electrostatic attractions and separate the

ions.

2. strength

Very brittle Any dislocation leads to layers moving and similarly charged ions being next to each other. The

repulsion splits the crystal.





3. Electrical conductivity

• do not conduct electricity when solid - ions are held strongly in the lattice

• conduct electricity when molten or in aqueous solution - the ions become mobile and conduction takes place.

4. solubility

• insoluble in non-polar solvents

• soluble in water as it is a polar solvent and stabilises the separated ions

• energy is needed to overcome the electrostatic attraction and separate the ions

• stability is achieved by polar water molecules surrounding the ions

3. METALLIC BOND-

Metal atoms achieve stability by “off -loading” electrons to attain the electronic structure of the nearest noble

gas. These electrons join up to form a mobile cloud which prevents the newly formed positive ions from

flying apart due to repulsion between similar charges.

Atoms arrange themselves in regular close packed 3-dimensional crystal lattices.

The outer shell electrons of each atom leave to join a mobile “cloud” or “sea” of electrons which can roam

throughout the metal. The electron cloud binds the newly-formed positive ions together.

Metallic bond strength depends on

• number of outer electrons donated

• the size of the metal atom/ion.

The melting point is a measure of the attractive forces within the metal.

electrical conductivity - Conduct electricity as there are mobile electrons.

strength - The delocalised electron cloud binds the “ions” together making metals ...

Solid State Liquid/ molten or aqueous State

Ions are attracted in electrostatically & packed

in definite pattern;

Ions don’t move vibrate only

Doesn’t conduct electricity

Ions exist separately;

Ions can move

conduct electricity





• malleable can be hammered into sheets

• ductile can be drawn into rods

melting pt.- High. Ease of separation depends on the - density of the electron cloud and ionic size/charge.

Metallic bonds get stronger when;

Size of metal ion increased

charge of metal ion increased

the no of electrons contributed to the bond increased

Determining the shape of molecules and ions

1. Lewis structure

Determining the "best" Lewis structure

Example; NO3-

1. Determine the total number of valence electrons in a molecule

2. Draw a skeleton for the molecule which connects all atoms using only single bonds. In simple molecules, the atomwith the most available sites for bondng is usually placed central. The number of bonding sites is detemined byconsidering the number of valence electrons and the ability of an atom to expand it's octet. As you become better, youwill be able to recognise that certain groups of atoms prefer to bond together in a certain way.

3. Of the 24 valence electrons in NO3-

, 6 were required to make the skeleton. Consider the remaining 18 electrons and place them so as to fill the ocets of as many atoms as possible (start with the most electronegative atoms first then proceed to the more electropositive atoms).







Octet Rule

Most elements follow the octet rule in chemical bonding, which means that an element should have contactto eight valence electrons in a bond or exactly fill up its valence shell. Having eight electrons total ensuresthat the atom is stable. This is the reason why noble gases, a valence electron shell of 8 electrons, arechemically inert; they are already stable and tend to not need the transfer of electrons when bonding with

another atom in order to be stable. On the other hand, alkali metals have a valance electron shell of oneelectron. Since they want to complete the octet rule they often simply lose one electron. This makes themquite reactive because they can easily donate this electron to other elements. This explains the highlyreactive properties of the Group IA elements.

Some elements that are exceptions to the octet rule include Aluminum(Al), Phosphorus(P), Sulfur(S), andXenon(Xe).

Hydrogen(H) and Helium(He) follow the duet rule since their valence shell only allows two electrons. Thereare no exceptions to the duet rule; hydrogen and helium will always hold a maximum of two electrons.

2.

Valence Shell Electron Pair Repulsion (VSEPR) Theory

The VSEPR Model

The electrons around the atoms in a molecule repel each other. They move to be as far apart as possible while stillmaintaining the bonding within the molecule. The procedure for using the model is as follows:

1) Determine the correct Lewis structure for the molecule. If it is a diatomic (has only two atoms) it is linear andfalls into the AX category. If it has 3 or more atoms continue with step 2.

2) Count the number of electron groups around the central atom. A group of electrons is a bond, a nonbonding

electron pair, or occasionally an unpaired nonbonding electron. Each triple or double bond counts as only onegroup for the purposes of this model.

3) Based on this number of groups around the central atom the molecule falls into one of six basic categories.Within each category there are a number of different names for the shapes depending upon the number of atomsand nonbonding groups around the central atom.





Electron Pairs How thepairs

arranged

Shape of the molecule

Total Bond unpaired Examples

Linear Linear

CO2

Trigonal planner

Trigonal planner

BF3

Angular

SO2

Tetrahedral

Tetrahedral

CH4

Trigonal pyramid

NH3

Angular

H2O





Trigonal bipyramidal

Trigonal bipyramidal

PCl5

See-saw shape

SF4

T-shape

ClF3

Linear

XeF2

Octahedral

Octahedral

SF6

Square pyramid

BrF5

Square planner

XeF4





Hybridization

Hybrid orbitals are the result of a model which combines atomic orbitals on a single atom in

ways that lead to a new set of orbitals that have geometries appropriate to form bonds in the

directions predicted by the VSEPR model. The VSEPR model predicts geometries that are very

close to those seen in real molecules

How to find the type of hybridisation

If you can assign the total electron geometry (geometry of all electron domains, not just bonding domains) on the

central atom using VSEPR, then you can always automatically assign hybridization. Hybridization was invented to

make quantum mechanical bonding theories work better with known empirical geometries.

Linear - sp - the hybridization of one s and one p orbital produce two hybrid orbitals oriented 180∘apart.

Trigonal planar - sp2 - the hybridization of one s and two p orbitals produce three hybrid orbitals

oriented 120∘ from each other all in the same plane.

Tetrahedral - sp3 - the hybridization of one s and three p orbitals produce four hybrid orbitals oriented toward the points of a regular tetrahedron, 109.5∘ apart.

Trigonal bipyramidal - dsp3 or sp3d - the hybridization of one s, three p, and one d orbitals produce five hybrid

orbitals oriented in this weird shape: three equatorial hybrid orbitals oriented 120∘ from each other all in the

same plane and two axial orbitals oriented 180∘ apart, orthogonal to the equatorial orbitals.

Octahedral - d2sp3 or sp3d2 - the hybridization of one s, three p, and two d orbitals produce six hybrid orbitals

oriented toward the points of a regular octahedron 90∘ apart.

E.g. 1 SP Hybridization





E.g. 1. SP2 Hybridization

E.g. 3. SP3 Hybridization





Table 1: Relationship of the VSEPR geometries to atomic hybridization.

Number of Groups

around the atom

Sub-Shapes

A = central atom

X = atom attached to central atom

E = nonbonding electron group oncentral atom

Hybridization

2 AX2, Linear sp

3 AX3, Trigonal Planar

AX2E, Bentsp2

4 AX4, Tetrahedral

AX3E, Trigonal Pyramidal

AX2E2, Bent

sp3

5 AX5, Trigonal Bipyramidal

AX4E, See-Saw

AX3E2, T-Shaped

AX2E3Linear

sp3d

6 AX6, Octahedral

AX5E, Square Pyramidal

AX4E2, Square Planar

sp3d2





Resonance Structures

Resonance theory is one of the most important theories that helps explain many interesting aspects of chemistry

ranging from differences in reactivity of related compounds to physical properties such a the absorption of light by

molecules. The electron-dot formula for many of the compounds and ions presented us a choice when we placed 4

electrons between 2 of the atoms in the formation of double bonds.

E.g. 1; O3

These structures are called resonance forms. A resonance form does not represent what the truestructure actually looks like. It is merely one of the electron-dot formulas that is possible for thecompound. In ozone, both of the oxygen-oxygen bonds have double bond and single bondcharacter. The negative charge is spread equally between both of the outer oxygen atoms. Wecan write a mechanism for this phenomenon showing how 2 of the electrons jump back and

forth.

This movement of electrons is constantly happening and we think of the real structure as anaverage of the 2 equal resonance forms.

E.g. 2; NO3-

The charge is evenly distributed among the 3 oxygen atoms in the nitrate anion. This makes the negative charge less

available for the reverse reaction and helps explain why nitric acid is a fairly strong acid. The bond order for thenitrogen-oxygen bonds in the nitrate anion is 1.33.

Bond Order

Bond order i s a measure of the number of bonding electron pair s between atoms. Single bonds have a bond order

of 1, double bonds have a bond order of 2 and tr ipl e bonds (the maximum number) have a bond order of 3. A

fractional bond order i s possibl e in molecules and ions that have resonance structures. I n the example of ozone, the

bond order would be the average of a double bond and a single bond or 1.5 (3 divided by 2). As the bond order

becomes larger, the bond length becomes smal ler.





Nature of bonds in molecules/ions

Bonds- sigma bonds are covalent bonds formed by direct overlapping between two

adjacent atom's outermost orbitals. The single electrons from each atom's orbital combine to form an electron

pair creating the sigma bond.

Bonds- A pi bond is a covalent bond formed between two neighboring atom's unbonded p-orbitals.

An unbound p-orbital electron in one atom forms an electron pair with a neighboring atom's unbound, parallel

p-orbital electron. This electron pair forms the pi bond. Double and triple bonds between atoms are usually

made up of a single sigma bond and one or two pi bonds.

Polarity & dipole moment

The polarity of a bond is determined by determining the difference in the electronegativities. If a molecule is diatomic

(2 atoms) there is often only one bond and that will determine whether the molecule is polar.

• E.g. the H−F bond is polar with fluorine being the more electronegative. A partial negative charge resides on the

fluorine atom and partial positive charge.

HOW TO DETERMINE MOLECULAR POLARITY (EXCLUDING DIRECTION OF DIPOLE MOMENT

• Molecules that are not totally symmetrical are polar molecules. In a polar molecule, electron density accumulates

toward one side of the molecule giving that side a slight negative charge δ-, and the other side a slight positive charge

of equal value δ+. Polar molecules are said to possess a dipole moment which means that it has 2 poles (+ and -). A

polar molecule is a dipole.

• Polarity is due to the polarity of the bonds and the lone pairs on the central atom. One lone pair on the central atom

makes the molecule polar. This only works for one lone pair, not 2, 3, 4, etc. If more than one lone pair, determine the

polarity from the bonded atoms.

To determine if a molecule is polar (has a dipole moment):





1. Draw an acceptable Lewis dot structure,

2. Predict the electron-pair geometry

3. Determine whether the molecule is totally symmetrical.

4. An analogy for polarity is to imagine that an object is being pulled in directions determined by the

electronegativities of the atoms. If the forces are equal, the object will not move (nonpolar).

Criteria for polarity:

• If a molecule is diatomic (2 atoms) and the atoms are different, it is polar.

• A molecule having just one lone pair of electrons is polar.

• If all of the terminal atoms are the same and there are no lone pairs of electrons around the central atom, the

molecule is totally symmetrical and nonpolar.

• If the molecule is not symmetrical, it is polar. The terminal atoms are different and the dipole moments do not cancel

each other out. (Pulling moves the object).

HOW TO DETERMINE THE DIRECTION OF THE DIPOLE MOMENT

Polarizability

Polarizability allows us to better understand the interactions between nonpolar atoms and molecules and other

electrically charged species, such as ions or polar molecules with dipole moments.

Neutral nonpolar species have spherically symmetric arrangements of electrons in their electron clouds. When in the presence of an electric field, their electron clouds can be distorted (Figure 1). The ease of this distortion is defined as

the polarizability of the atom or molecule. The created distortion of the electron cloud causes the originally nonpolar





molecule or atom to acquire a dipole moment. This induced dipole moment is related to the polarizability of themolecule or atom and the strength of the electric field by the following equation:

μind=α′E(1)

Where

E denotes the strength of the electric field and α′ is the polarizability constant with units of C m2V-1.

Figure 1: A neutral nonpolar species's electron cloud is distorted by A.) an Ion and B.) a polar molecule to induce adipole moment.

In general, polarizability correlates with the interaction between electrons and the nucleus. The amount of electrons ina molecule affects how tight the nuclear charge can control the overall charge distribution. Atoms with less electrons

will have smaller, denser electron clouds, as there is a strong interaction between the few electrons in the atoms‟

orbitals and the positively charged nucleus. There is also less shielding in atoms with less electrons contributing to the

stronger interaction of the outer electrons and the nucleus. With the electrons held tightly in place in these smalleratoms, these atoms are typically not easily polarized by external electric fields. In contrast, large atoms with manyelectrons, such as negative ions with excess electrons, are easily polarized. These atoms typically have very diffuseelectron clouds and large atomic radii that limit the interaction of their external electrons and the nucleus.

Factors that Influence Polarizability

The relationship between polarizability and the factors of electron density, atomic radii, and molecular orientation isas follows:

1. The greater the number of electrons, the less control the nuclear charge has on charge distribution, and thus the

increased polarizability of the atom.2. The greater the distance of electrons from nuclear charge, the less control the nuclear charge has on the charge

distribution, and thus the increased polarizability of the atom.3. Molecular orientation with respect to an electric field can affect polarizibility (labeled Orientation-dependent),

except for molecules that are: tetrahedral, octahedral or icosahedral (labeled Orientation-independent). This factoris more important for unsaturated molecules that contain areas of electron dense regions, such as 2,4-hexadiene.





Greatest polarizability in these molecules is achieved when the electric field is applied parallel to the moleculerather than perpendicular to the molecule.

Polarizability Influences Dispersion Forces Dispersion Forces

The dispersion force is the weakest intermolecular force. It is an attractive force that arises from surroundingtemporary dipole moments in nonpolar molecules or species. These temporary dipole moments arise when there areinstantaneous deviations in the electron clouds of the nonpolar species. Surrounding molecules are influenced by these

temporary dipole moments and a sort of chain reaction results in which subsequent weak, dipole-induced dipoleinteractions are created. These cumulative dipole- induced dipole interactions create the attractive dispersion forces.Dispersion forces are the forces that make nonpolar substances condense to liquids and freeze into solids when thetemperature is low enough.

Polarizability affects dispersion forces in the following ways:

As polarizability increases dispersion forces also become stronger . Thus, molecules attract one anothermore vigorously and melting and boiling points of covalent substances increase with larger molecular mass.

Polarazibility also affects dispersion forces through the molecular shape of the affected molecules. Elongated

molecules have electrons that are easily moved increasing their polarizability and thus strengthening the dispersion

forces, see the example in Figure 2. In contrast, small, compact, symmetrical molecules are less polarizableresulting in weaker dispersion forces, see the example in Figure 3.

Figure 3: An example of a compact less polarizable molecule.





2. Secondary Interactions

•

Hydrogen Interactions

• Dipole-dipole Interactions

• Ion-dipole Interactions

•

Ion-induced dipole Interactions

• Dipole-induced dipole Interactions

• Dispersion Interactions





Physical properties of substances

Physical propertiesType of compoundI onic Co valent Metallic

Melting point

Electrical conductivity

Thermal conductivity

Hardness

Q: Both H2S & H2O are covalent compounds. H2S is a gas but H2O is a liquid. Explain this statement using the

knowledge of intermolecular forces.

Q: Explain endo-thermic and exo-thermic chemical reactions in terms of energy transfers associated with breakin &

making chemical bonds.





Different types of Lattice

The substances in the solid state, have particles are arranged in a lattice.

A lattice is a regular arrangement of particles, whether these are atoms, ions or molecules

Type Particle Example Type of bond Solubility inWater

Melting Point

Homoatomic Atoms DiamondGraphite

Heteroatomic Molecules SiO2 Poor

Non-polarmolecularlattices

Molecules I2 CovalentBonding betweenatomsforming

molecules; Non-polarintermolecularBonding betweenseparatemolecules





polarmolecularlattices

Molecules Ice CovalentBonding betweenatomsformingmolecules; polar

intermolecularBonding betweenseparatemolecules

Ionic lattices Ions NaCl Ionic Bonding(throughout)

Generallygood, but may be low owingto high latticeenthalpy

Metalliclattices

Metal ionswithdelocalisedelectrons

Cu, Zn MetallicBonding(throughout)

Poor, but somemetals reactwith water

Exercise;

1. Draw Lewis structures for;

BeCl2, BF3, CH4, NH3, H2O, PCl3, NH4+, SF6, SF4, ClF3, XeF2, LF5, XeF4, NO3

-, CO32-, SO4

2-

2. Determine the hybridization of the central atom of,

BeCl2, BF3, CH4, NH3, H2O, PCl3, NH4+

3.

Draw resonance structures for,O3, NO3

-, CO32-, SO4

2-, N2O



Chem Bond Local Al

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