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Bonding. Chemical bonding. Chemical Bonding. Graphite. Diamond. Silicon dioxide. PreDicting the type of bonding from electronegativity values. Ionic bonding typically occurs between metal and non-metal. E.g. Barium fluoride, BaF 2

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Bonding

Bonding


Chemical bonding

Chemical bonding


Chemical bonding1

Chemical Bonding

Graphite

Diamond

Silicon dioxide


Bonding

PreDicting the type of bonding from electronegativity values

  • Ionic bonding typically occurs between metal and non-metal. E.g. Barium fluoride, BaF2

  • The reactivity of metals and non-metals can be assessed using electro-negativity

    Electro-negativity

  • ability of an atom in a covalent bond to

    attract ___________________________ to itself.

shared paired of electrons


Bonding

  • Metals generally have low electronegativity values, while non-metals have relatively high electronegativity values.

  • Fluorine, which has the greatest attraction for electrons in bond-forming situations (highest E value).


Bonding

Electrons are not shared.

E.g. Na+ Cl- , electron is transferred.

Polar covalent bonds are covalent bonds with ionic character.

Ionic bond

Electrons are equally shared.

E.g.Cl-Cl

Non polar

(pure covalent)

bond

Electrons are not equally shared.

E.g.

Atoms have different

electronegativity values

Polar

covalent bond


Bonding

Example

What type of bond is the following?

  • N (3.0) and H (2.1)

  • H (2.1) and H(2.1)

  • Ca(1.0) and Cl(3.0)

  • Al (1.5) and Cl(3.0)

  • H (2.1) and F(4.0)


Dipole

Dipole

  • No bond is purely ionic or covalent..they have a little bit of both characters.

  • When there is unequal sharing of electrons a dipole exists.

    Dipole

    - is a molecule that has 2 poles or regions with opposite charges.

    - is represented by a dipole arrow towards the more negative end.


Attractions between molecules

Attractions Between Molecules

  • Besides ionic, metallic, and covalent bonds, there are also attractions between molecules

  • Intermolecular attractions are weaker than ionic, covalent, and metallic bonds

  • There are 2 main types of attractions between molecules: Van der Waals and Hydrogen


Van der waals forces

Van der Waals Forces

  • Van der Waals forces consists of the two weak attractions between molecules

2. dispersion forces – caused by the motion of electrons (weakest of all forces)

1. dipole interactions – polar molecules attracted to one another


Hydrogen bond

Hydrogen Bond

  • Hydrogen Bonds are forces where a hydrogen atom is weakly attracted to an unshared electron pair of another atom


Hydrogen bond1

Hydrogen Bond

  • This other atom may be in the same molecule or in a nearby molecule, but always has to include hydrogen.

INTERMOLECULAR HYDROGEN BONDING


Bonding

INTRAMOLECULAR HYDROGEN BONDING

  • Hydrogen Bonds have about 5% of the strength of an average covalent bond

  • Hydrogen Bond is the strongest of all intermolecular forces


Intermolecular attractions

Intermolecular Attractions

  • A few solids that consist of molecules do not melt until the temperature reaches 1000ºC or higher called network solids (Example: diamond, silicon carbide)

  • A Network Solid contains atoms that are all covalently bonded to each other

  • Melting a network solid would require breaking bonds throughout the solid (which is difficult to do)

http://library.thinkquest.org/C006669/data/Chem/bonding/inter.html


Bonding

  • The bonding pair of electrons spends most of its time between the two atomic nuclei.

  • screening the positive charges from one another and enabling the nuclei to come closer together.

  • Negative charge on the electron pair attracts both nuclei and holds them together in a covalent bond.

  • When two atoms are chemically bonded, the two atoms close together have less energy and therefore are more stable than when separated.

  • Energy is given off by the atoms to form a bond, and energy must be supplied (absorbed) to break the bond.

  • A covalent bond is the result of electrostatic attraction between the nuclei of the 2 atoms and the pair of shared electrons.


Electron dot lewis structure

Electron Dot (LEWIS) Structure


Valence electrons for elements

Valence electrons for Elements


Valence electrons and number of bonds

H

a

l

o

g

e

n

s

F

,

B

r

,

C

l

,

I

C

a

l

c

o

g

e

n

s

O

,

S

N

i

t

r

o

g

e

n

N

,

P

C

a

r

b

o

n

C

,

S

i

Valence electrons and number of bonds

#

C

o

v

a

l

e

n

t

B

o

n

d

s*

F

a

m

i

l

y

Number of bonds elements prefers depending on the number of valence electrons. In general -

X

1 bond often

2 bond often

O

3 bond often

N

4 bond always

C


Lewis structure octet rule guidelines

.

.

Lewis Structure, Octet Rule Guidelines

When compounds are formed they tend to follow the Octet Rule.

Octet Rule: Atoms will share electrons (e-) until it is surrounded by eight valence electrons.

Rules of the (VSEPR) game-

i) O.R. works mostly for second period elements.

Many exceptions especially with 3rd period elements (d-orbitals)

ii) H prefers 2 e- (electron deficient)

iii) :C:N::O::F:

4 unpaired3unpaired2unpaired1unpaired up = unpaired e-

4 bonds3 bonds2 bonds1 bond

O=C=ONNO = OF - F

iv) H & F are terminal in the structural formula (Never central)

.

.

.

.

.

.


Atomic connectivity

N

O

O

Cl

H

O

O

H

Se

O

O

F

H

C

H

H

H

O

O

H C

C

O H

H O

S

O H

H

O

Atomic Connectivity

The atomic arrangement for a molecule is usually given.

CH2ClFHNO3CH3COOHH2SeH2SO4O3

In general when there is a single central atom in the molecule, CH2ClF, SeCl2, O3 (CO2, NH3, PO43-), the central atom is the first atom in the chemical formula.

Except when the first atom in the chemical formula is Hydrogen (H) or fluorine (F). In which case the central atom is the second atom in the chemical formula.

Find the central atom for the following:

1) H2Oa) Hb) O2) PCl3a) Pb) Cl

3) SO3a) Sb) O4) CO32-a) Cb) O

5) BeH2a) Beb) H6) IO3- a) Ib) O


Bonding

RULES FOR DRAWING LEWIS STRUCTURES

  • Count the no. of valence electrons.

  • If the species has a –n charge, add n to the electrons

  • If the species has a +n charge, subtract n from the electrons

  • Draw a skeletal structure.

  • If C is present, place C at the centre.

  • If C is not present, place the LEAST electronegative atom at the centre.

    Note: H is never the in the center.

  • Complete the octets of the outer atoms (except for H) by adding lone pairs of electrons (including the 2 electrons shared with the central atom)

  • If there are any electrons left over, place them on the central atom as lone pairs.

  • If the central atom does not have a complete octet, rearrange lone pairs on the outer atoms to form double bonds between the central and outer atoms. Continue doing until the central atom’s octet is satisfied.

  • If the species is charged, place it inside brackets and write the charge outside the brackets.


Bonding

Lewis structures

Bond pair _____

Lone pair . .

Draw Lewis structures of the following molecules:

(a) H2O , NH3 , CO2 , OCl2 , PCl4+

(b) SO2 , NO+ , OCN- , COF2 , CO32- , NO2- , O3

(c) BeCl2 , BH3 , PCl5

In which of the above obey the octet rule?


Exceptions to octet rule

Exceptions to octet rule

  • Molecules with an odd number of electrons

  • Molecules in which an atom has less than an octet

  • 3.Molecules in which an atom has more than an octet


Bonding

1. Odd Number of Electrons

NO

Number of valence electrons = 11

Resonance occurs when more than one valid Lewis structure can be written for a particular molecule (i.e. rearrange electrons)

NO2

Number of valence electrons = 17

Molecules and atoms which are neutral (contain no formal charge) and with an unpaired electron are called Radicals

O2


Bonding

2. Less than an Octet

Includes Lewis acids such as halides of B, Al and compounds of Be

BCl3

Group 3A atom only has six electrons around it

However, Lewis acids “accept” a pair of electrons readily from Lewis bases to establish a stable octet


Bonding

AlX3

Aluminium chloride is an ionic solid in which Al3+ is surrounded by six Cl-.

However, it sublimes at 192°C to vapour Al2Cl6 molecules

B2H6

A Lewis structure cannot be written for diborane.

This is explained by a three-centre bond – single electron is delocalized over a B-H-B


Bonding

Octet Rule Always Applies to the Second Period = n2 ; number of orbitals

2s, 2px, 2py, 2pz ---orbitals cannot hold more than two electrons

Ne [He]; 2s2, 2px2, 2py2, 2pz2

n = 2

n = 3


Third period n 2 32 9 orbitals

Third Period ; n2 = 32 = 9 orbitals

Ar [Ne]; 3s2, 3px2, 3py2, 3pz23d0 3d0 3d0 3d0 3d0

n = 3


Bonding

3. More than an Octet

Elements from the third Period and beyond, have ns, np and unfilled nd orbitals which can be used in bonding

PCl5

P : (Ne) 3s2 3p3 3d0

Number of valence electrons = 5 + (5 x 7) = 40

10 electrons around the phosphorus

SF4

S : (Ne) 3s2 3p4 3d0

Number of valence electrons = 6 + (4 x 7) = 34

The Larger the central atom, the more atoms you can bond to it – usually small atoms such as F, Cl and O allow central atoms such as P and S to expand their valency.


Multiple bonds

Multiple bonds

  • Bond Strength

    Triple bonds > Double bonds > Single bonds

Page 93


Bonding

The attraction between the 2 nuclei for 3 electron pairs in a triple bond is > that for 2 electron pairs in a double bond which is > than that for 1 electron pair in a single bond.

(2) Triple bonds are shorter due to greater attraction between the

bonding electrons and the nuclei with more electrons in the bond.


Bonding

Bond strength and length of covalent bonds

  • Strength

    Triple bonds > Double bonds > Single bonds

  • Length

    Single bonds > Double bonds > Triple bonds


Bonding

Coordinate (dative) bonding

  • In some molecules and polyatomic ions, both electrons to be shared come from the same atom forming the coordinate or dative bond.

  • Carbon monoxide (CO) can be viewed as containing one coordinate bond and two "normal" covalent bonds between the C atom and the O atom.

How do you draw the Lewis structure?

Page 94


Dissolving hydrogen chloride gas in water

Dissolving hydrogen chloride gas in water

  • Something similar happens. A hydrogen ion (H+) is transferred from the chlorine to one of the lone pairs on the oxygen atom.

  • The H3O+ ion is variously called the hydroxonium ion.


Bonding

Other examples:

  • The reaction between ammonia and boron trifluoride, BF3

In BF3, there are only 6 electrons in the outer shell of boron.

There is space for the B to accept a pair of electrons.


Bond polarity

Bond Polarity

  • Due to difference in electronegativity value between the 2 atoms in the bond.

  • Unequal distribution of electron density results in small charges on the atoms

    ( δ+ and δ- )

    Example

A dipole is established when two electrical charge of opposite sign are separated by a small distance.

Dipole moment


Non polar molecule

Non-polar molecule

  • A molecule can possess polar bonds and still non-polar.

  • Check the Geometry of the molecule:

  • The polar bonds are arranged symmetrically so as to give zero net direction of charge.

    i.e. Overall dipoles cancel so that there is no overall dipole.


Non polar polar covalent bonds

Non-polar & polar covalent bonds

Non-polar Covalent bond

  • No difference in electronegativity value – bond consists of

    2 ____________ atoms.

  • _______ net charge.

    Examples :

Polar Covalent bond

  • Due to the difference in electronegativity value – bond consists of

    2 ____________ atoms.

  • _______ net charge.

    Examples:


Polar molecule

polar molecule

  • In the water molecule,

  • O-H bonds are significantly polar

  • The bent structure makes the distribution of those polar bonds asymmetrical.


Bonding

  • Some molecules have very low polarity - so low as to be regarded as non-polar,


Bonding

For CO2 each C-O bond is polar since O is more

electronegative than C.

Why is the molecule non-polar?


Vsepr theory

VSEPR theory

  • The shapes of simple molecules and ions can be determined by using the Valence Shell Electron Repulsion (VSEPR) theory.

  • Electron pairs around the central atom repel each other

  • Bonding pairs and lone pairs arrange themselves to be as far apart as possible


Principles of vsepr theory

Principles of VSEPR theory

  • Find the number of electron pairs / charge centres in the valence shell of the central atom.

  • Electron pairs / charge centres repel each other to the positions of minimum energy in order to gain maximum stability.

  • Pairs forming a double or triple bond act as a single bond

  • Non-bonding pairs repel more than bonding pairs.


A few vsepr shapes

A Few VSEPR Shapes


Nine possible molecular shapes

Nine possible molecular shapes


Practice

Practice

  • Methane (CH4) – tetrahedral

  • Ammonia (NH3) – pyramidal

  • Water (H2O) – bent

  • Carbon Dioxide (CO2) - linear


Methane ammonia and water

Methane, Ammonia and Water


Bonding

Valence Shell Electron-Pair Repulsion Theory (VSEPR)

Procedure

  • Sum the total Number of Valence Electrons

    Drawing the Lewis Structure

    2.The atom usually written first in the chemical formula is the Central atom in the Lewis structure

  • Complete the octet bonded to the Central atom. However, elements in the third row have empty d-orbitals which can be used for bonding.

  • If there are not enough electrons to give the central atom an octet try multiple bonds.

    Predicting the Shape of the Molecule

  • Sum the Number of Electron Domains around the Central Atom in the Lewis Structure; Single = Double = Triple Bonds = Non-Bonding Lone Pair of Electrons = One Electron Domain

  • From the Total Number of Electron Domains, Predict the Geometry and Bond Angle(s); 2 (Linear = 180º); 3 (Trigonal Planar = 120º); 4 (Tetrahedral = 109.5º); 5 (TrigonalBipyramidal = 120º and 90º); 6 (Octahedral = 90º)

  • Lone Pair Electron Domains exert a greater repulsive force than Bonding Domains. Electron Domains of Multiple Bonds exert a greater repulsive force than Single Bonds. Thus they tend to compress the bond angle.


Lone pairs and bond angles

Lone pairs and bond angles

Lone pairs are held closer to the nucleus than the bonding pairs.

The distance between the lone pair electrons and the bonding pairs of electrons is shorter than the distance between the bonding pairs to each other.

Repulsion due to lone pairs causes the bond angles to become smaller


Bonding

Order of repulsion :

lone pair-lone pair > lone pair- bonding pair > bonding pair – bonding pair

Methane, CH4

Bond angle is 109.50

Ammonia, NH3

Greater repulsion by lone

pair of electrons.

Bond angle is smaller

than 109.50(1050)

Water, H2O

Even greater repulsion by

two lone pair of electrons.

Bond angle is even smaller

(1050)


Predicting the shape of ions

Predicting the shape of ions

Consider

NH4+ , H3O+ , NO2-

NH4+

As the 4 negative charge centres repel each other and take up positions in space to be as far apart as possible, the electron pairs are distributed in tetrahedral arrangement.

H3O+

As the 4 negative charge centres repel each other and take up positions in space to be as far apart as possible, the electron pairs are distributed in tetrahedral arrangement.

With one lone pair of electrons, the actual structure is trigonal pyramid with a bond angle of 1070 for H-O-H bond.

Read page 104


Bonding

NO2-

As the 3 negative charge centres repel each other and take up positions in space to be as far apart as possible, the electron pairs are distributed in trigonal planar arrangement.

With one lone pair of electrons, the actual structure of the ion is bent with a bond angle of about 1170for O-N-O bond.


Bonding

Molecules with more than one central atom

Consider

N2H4 , C2H2


Intermolecular forces

Intermolecular forces

  • Forces between molecules.

  • Does not exist in giant structure (ionic cpds, metals & giant covalent structure)

http://chemtools.chem.soton.ac.uk/projects/emalaria/index.php?page=13


Bonding

Intermolecular Forces: are generally much weaker than covalent or ionic bonds. Less energy is thus required to vaporize a liquid or melt a solid. Boiling points can be used to reflect the strengths of intermolecular forces (the higher the Bpt, the stronger the forces)

Hydrogen Bonding : the attractive force between hydrogen in a polar bond (particularly H-F, H-O, H-N bond) and an unshared electron pair on a nearby small electronegative atom or ion

Very polar bond in H-F.

The other hydrogen halides don’t form hydrogen bonds, since H-X bond is less polar. As well as that, their lone pairs are at higher energy levels. That makes the lone pairs bigger, and so they don't carry such an intensely concentrated negative charge for the hydrogens to be attracted to.


Bonding

Hydrogen Bonding & Water


Van der waals forces1

Van der Waals forces

  • Electrons can at any moment be unevenly spread producing a temporary instantaneous (fluctuating) dipole.

  • An instantaneous dipole can induce another dipole in a neighbouring particle resulting in a weak attraction between the two particles.

  • The forces of attraction between temporary or induced dipoles are known as Van der Waals’ forces (London Dispersion Forces).

  • Van der Waals’ forces increases with increasing mass.


Bonding

London Dispersion Forces –

significant only when molecules are close to each other

Due to electron repulsion, a temporary dipole on one atom can induce a similar dipole on a neighboring atom

Prof. Fritz London


Bonding

  • The ease with which an external electric field can induce a dipole (alter the electron distribution) with a molecule is referred to as the "polarizability" of that molecule

  • The greater the polarizability of a molecule the easier it is to induce a momentary dipole and the stronger the dispersion forces

  • Larger molecules tend to have greater polarizability

    • Their electrons are further away from the nucleus (any asymmetric distribution produces a larger dipole due to larger charge separation)

    • The number of electrons is greater (higher probability of asymmetric distribution)

  • thus, dispersion forces tend to increase with increasing molecular mass

    • Dispersion forces are also present between polar/non-polar and polar/polar molecules (i.e. between all molecules)


Bonding

  • The strength of the intermolecular forces determines how easily the molecules will separate and hence the melting and boiling points.

Why is there an increasing boiling points of the noble gases as

you go down the group?

The boiling points of the noble gases are

Helium-269°C

Neon-246°C

Argon-186°C

Krypton-152°C

Xenon-108°C

Radon-62°C


Bonding

  • Because electrons are always moving around very quickly, the charges switch around all the time.

  • the more electronsin a molecule / atom, the stronger these Van der Waals or London forces are.

  • This is seen in the increasing boiling points of the noble gases as you go down the group.

    The boiling points of the noble gases are

    Helium-269°C

    Neon-246°C

    Argon-186°C

    Krypton-152°C

    Xenon-108°C

    Radon-62°C


Group 7

Group 7

Cl2 : gas

I2 : solid

Iodine molecule is made up of larger atoms with more electrons compared to chlorine.

With more electrons moving around, the temporary dipole will be larger.

The larger atoms in the molecule means that the valence electrons are less strongly held,

Hence the induced dipoles will be larger.


Permanent dipoles

Permanent Dipoles

  • van der Waals forces are present between covalent molecules with no H atom attached to N, O or F.

  • E.g. van der Waals forces are present between HCl molecules.

  • There’s also other intermolecular forces beween the molecules :

    permanent dipole - permanent dipole


Bonding

  • These intermolecular forces between polar molecules are stronger than between non-polar molecules. (all things being equal)

  • For polar substances with similar RMM, the higher the dipole moment, the stronger the dipole-dipole attractions and the higher the boiling points.


Compare molecules with same rmm

Compare molecules with same RMM

  • Propane (C3H8) and ethanal (CH3CHO) both with RMM = 44

  • Ethanal has a higher bp.

    Ethanal

    - is a polar molecule

    - has stronger intermolecular forces (van der waals & dipole-dipole interactions) between the molecules of ethanal than between the propane molecules.

It is not true that polar molecules have stronger

intermolecular forces and hence higher bp than

non-polar molecules.

Non-polar molecules with higher RMM might have higher bp.

Read further for a few exceptions


Bonding

Boiling points increase for polar molecules of similar mass, but increasing dipole:


Hydrogen bonding

Hydrogen bonding

  • A hydrogen bond is a weak type of force that forms a special type of dipole-dipole attraction which occurs when a hydrogen atom bonded to a strongly electronegative atom exists in the vicinity of another electronegative atom with a lone pair of electrons. These bonds are generally stronger than ordinary dipole-dipole and dispersion forces, but weaker than true covalent and ionic bonds.

  • Hydrogen bonding is present between covalent molecules with H atoms attached to O, N and F


2 factors affecting the bond

2 factors affecting the bond

  • Strength of H bond

    The larger the electronegativity of H and the other atom (N, O or F), the stronger the H bond.

    Strength F > O > N

  • Number of them that can be formed between neighbouring molecules


Hydrogen bonding1

Hydrogen bonding

Although the strength is such F > O > N, HF can only form 1 H bond to 1 neighbour.

H2O can form 2, thus promoting more intermolecular interactions .

The collective strength of the H bonds in water is greater than the strength of the H bonds in HF because each O atom (with 2 lone pairs) in the water molecule can form 2 H bonds with 2 other water molecules, whereas each F atom in HF molecule can only form 1 H bond with another HF molecule.


Bonding

Ammonia molecule has 3 N-H bonds. N is larger and < electronegative than F, has far weaker H bonds due to lower electron density on the N atom (only 1 lone pair) compared to O and F.


Bonding

  • When the RMM is large, we expect the boiling point to be high because larger molecules have more space for electron distribution and more possibilities for instantaneous dipole moment.

  • However,

Greater intermolecular

force because H2O, HF,

NH3 all exhibit hydrogen

bonding .

tend to have higher viscosity

than those that do not have

H bond.

Substances which have multiple

H bonds exhibit even higher

Viscosity.


Factors affecting h bonding

Factors affecting H bonding

  • Electronegativity

    Cannot occur without significant electronegativity difference between H and the atom it is boded to.

    E.g. Both PH3 and NH3

    have trigonal pyramidal shape but only NH3 has H bonding.

  • Atomic size

    When the radii of the 2 atoms differ greatly, their nuclei cannot achieve close proximity when they interact resulting in weak interaction.


Bonding

Is there any hydrogen bonding between the molecules if CH3F?

H H

H C F H C F

H H

H is not joined directly to F in each molecule,

hence no hydrogen bonding between the molecules.


Bonding

Is there any hydrogen bonding between the molecules of ethanol?


Effects of h bonding on physical properties

Effects of H bonding on physical properties

Hydrogen bonding affects

  • the boiling points of water, ammonia, hydorgen fluoride and other molecules

  • the solubility of simple covalent molecules such as ammonia, methanol and ethanoic acid in water

  • the density of water and ice.

  • the viscosity of liquids, e.g. the alcohols.


Bonding

Van der Waals forces are made of dipole-dipole and London dispersion forces

Group 4A hydrides

Groups 4, 5, 6A hydrides


Comparisons of the physical properties of ionic compounds vs covalent substances

Comparisons of the Physical Properties of Ionic Compounds VS Covalent Substances


Bonding

Recall:

Inionic compounds, the ions are held together by strong ionic bonds in a giant ionic lattice.

In simple covalent molecules, the attractive forces between the molecules are known as intermolecular forces or van der Waal’s forces, which is weaker than the ionic bonds.

In giant covalent molecules, the atoms are held together by strong covalent bonds in a giant covalent lattice.


Bonding

Types of Covalent Substances

Covalent substances can be divided into 2 categories as shown in the table below:


Giant covalent lattice

Giant covalent lattice

  • The atoms can be either same like silicon and carbon (graphite and diamond) or of 2 different elements such as silicon dioxide.

  • Allotropes are two (or more) crystalline forms of the same element, in which the atoms ( or molecules) are bonded differently.


Bonding

C60 Fullerene

  • Molecular structure – consists of individual C60 molecules

    with covalent bonds within the molecule.

  • 60 C atoms are arranged in

    hexagons and pentagons.

  • Van der Waal’s forces between molecules.


Bonding

Properties


Bonding

Properties


Bonding

Properties


Bonding

4Metallic Bonds

Metals consist of positive ions surrounded by a 'sea of moving electrons'.

The negative 'sea of electrons' attracts all the positive ions and cements everything together.

Metallic bonds are the results of the strong forces of attraction between the negative electrons and the positive ions.

Hence, metals have high melting points and high boiling points.


Bonding

Physical Properties of Metals


Bonding

Explanation

  • The valence electrons do not belong to any particular atom, hence, if sufficient force is applied to the metal, 1 layer of metals can slide over another without disrupting the metallic bonding.

  • The metallic bonding in metal is strong and flexible and so metals can be hammered into thin sheets (malleability) or drawn into lonng wires (ductility) without breaking.

    If atoms of other elements are added by alloying, the layers of ions will not slide over each other so readily. The alloy is thus less malleable and ductile and consequently harder and stronger.


Solubility

solubility

  • ‘Like tends to dissolve like’. Polar substances tend to dissolve in polar solvents, such as water, whereas non-polar substances tend to dissolve in non-polar solvents, such as heptane or tetrachloromethane.

  • Organic molecules often contain a polar head and a non-polar carbon chain tail. As the non-polar carbon chain length increases in an homologous series the molecules become less soluble in water.

  • Ethanol is a good solvent for other substances as it contains both polar and non-polar ends.


Bonding

  • Water will mix with polar liquids such as ethanol. The oppositely charged ends of the different molecules attract one another forming hydrogen bonds.


Bonding

  • Gases are generally slightly soluble in water.

  • A small number of gases are highly soluble because they react with water to release ions.

    Example,

    SO2(g) + H2O(g) H+(aq) + HSO3-(aq)

    This solution is known as sulfurous acid , a major component of acid rain


Summary

Summary


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