Slide 1 : IONIC
BONDING
Key Questions: : Key Questions: What evidence is there for the existence of ions in ionic compounds?
How do ions form?
What is the definition of an ionic crystal?
What is the definition of ionic bonding?
What does the model of ionic bonding tell us about the properties of ionic compounds?
Slide 3 : THE IONIC BOND Ionic bonds tend to be formed between elements whose atoms need to “lose” electrons to gain the nearest noble gas electronic configuration (n.g.e.c.) and those which need to gain electrons. The electrons are transferred from one atom to the other.
Sodium Chloride
Na ——> Na+ + e¯ and Cl + e¯ ——> Cl¯
1s2 2s2 2p6 3s1 1s2 2s2 2p6 1s2 2s2 2p6 3s2 3p5 1s2 2s2 2p6 3s2 3p6
or 2,8,1 2,8 2,8,7 2,8,8
An electron is transferred from the 3s orbital of sodium to the 3p orbital of chlorine; both species end up with the electronic configuration of the nearest noble gas the resulting ions are held together in a crystal lattice by electrostatic attraction.
Slide 4 : ELECTRON
TRANSFER Mg ——> Mg2+ + 2e¯ and 2Cl + 2e¯ ——> 2 Cl¯ Mg Cl Cl e¯ e¯ THE IONIC BOND FORMATION OF MAGNESIUM CHLORIDE
Slide 5 : EVIDENCE FOR IONS - ELECTRON DENSITY MAPS X-ray diffraction measurements of crystals enable chemists to build up electron density maps. Electron density maps look like the contours on a geographical map. Each contour on an electron density map joins points with the same electron density and this allows atoms and ions to be identified.
Figure 3 shows an electron density map for sodium chloride.
Figure 3?
The electron density map for sodium chloride. The labels on the contours are in electrons per 10–30 m3.
Slide 6 : GIANT IONIC CRYSTAL LATTICE Cl-
Chloride ion
Na+
Sodium ion Oppositely charged ions held in a regular
3-dimensional lattice by electrostatic attraction
The arrangement of ions in a crystal lattice depends on the relative sizes of the ions The Na+ ion is small enough relative to a Cl¯ ion to fit in the spaces so that both ions occur in every plane.
Slide 7 : GIANT IONIC CRYSTAL LATTICE Each Na+ is surrounded by 6 Cl¯ (co-ordination number = 6)
and each Cl¯ is surrounded by 6 Na+ (co-ordination number = 6). Oppositely charged ions held in a regular
3-dimensional lattice by electrostatic attraction
The arrangement of ions in a crystal lattice depends on the relative sizes of the ions
Slide 8 : GIANT IONIC CRYSTAL LATTICE Each Na+ is surrounded by 6 Cl¯ (co-ordination number = 6)
and each Cl¯ is surrounded by 6 Na+ (co-ordination number = 6). Oppositely charged ions held in a regular
3-dimensional lattice by electrostatic attraction
The arrangement of ions in a crystal lattice depends on the relative sizes of the ions
Slide 9 : Positive ions
also known as cations; they are smaller than the original atom.
formed when electrons are removed from atoms.
the energy associated with the process is known as the ionisation energy
1st IONISATION ENERGY (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¯
Other points
Successive IE’s get larger as the proton:electron ratio increases.
Large jumps in value occur when electrons are removed from shells nearer the nucleus because there is less shielding and more energy is required to overcome the attraction. If the I.E. values are very high, covalent bonding will be favoured (e.g. beryllium). THE FORMATION OF IONS
Slide 10 : Negative ions
known as anions
are 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 AFFINITY
The energy change when one mole of gaseous atoms acquires one mole of electrons (from infinity) to form one mole of gaseous negative ion
e.g. Cl(g) + e¯ ——> Cl¯(g) and O(g) + e¯ ——> O¯(g)
The greater the effective nuclear charge (E.N.C.) the easier an electron is pulled in. THE FORMATION OF IONS
Slide 11 : COVALENT
BONDING
Slide 12 : Definition consists of a shared pair of electrons with one electron being
supplied by each atom either side of the bond.
compare this with dative covalent bonding
atoms are held together
because their nuclei which
have an overall positive charge
are attracted to the shared electrons COVALENT BONDING + +
Slide 13 : Definition consists of a shared pair of electrons with one electron being
supplied by each atom either side of the bond.
compare this with dative covalent bonding
atoms are held together
because their nuclei which
have an overall positive charge
are attracted to the shared electrons
Formation between atoms of the same element N2, O2, diamond,
graphite
between atoms of different elements CO2, SO2
on the RHS of the table;
when one of the elements is in the CCl4, SiCl4
middle of the table;
with head-of-the-group elements BeCl2
with high ionisation energies; COVALENT BONDING + +
Slide 14 : • atoms share electrons to get the nearest noble gas electronic configuration
• some don’t achieve an “octet” as they haven’t got enough electrons
eg Al in AlCl3
• others share only some - if they share all they will exceed their “octet”
eg NH3 and H2O
• atoms of elements in the 3rd period onwards can exceed their “octet” if
they wish as they are not restricted to eight electrons in their “outer shell”
eg PCl5 and SF6 COVALENT BONDING
Slide 15 : Orbital theory
Covalent bonds are formed when orbitals, each containing one electron, overlap. This forms a region in space where an electron pair can be found; new molecular orbitals are formed. SIMPLE MOLECULES The greater the overlap the stronger the bond. orbital containing 1 electron orbital containing 1 electron overlap of orbitals provides a region in space which can contain a pair of electrons
Slide 16 : HYDROGEN Another hydrogen atom also needs one electron to complete its outer shell Hydrogen atom needs one electron to complete its outer shell atoms share a pair of electrons to form a single covalent bond
A hydrogen MOLECULE is formed H WAYS TO REPRESENT THE MOLECULE PRESSING THE SPACE BAR WILL ACTIVATE EACH STEP OF THE ANIMATION
Slide 17 : HYDROGEN CHLORIDE Cl Hydrogen atom also needs one electron to complete its outer shell Chlorine atom needs one electron to complete its outer shell atoms share a pair of electrons to form a single covalent bond WAYS TO REPRESENT THE MOLECULE PRESSING THE SPACE BAR WILL ACTIVATE EACH STEP OF THE ANIMATION
Slide 18 : METHANE C Each hydrogen atom needs 1 electron to complete its outer shell A carbon atom needs 4 electrons to complete its outer shell Carbon shares all 4 of its electrons to form 4 single covalent bonds WAYS TO REPRESENT
THE MOLECULE PRESSING THE SPACE BAR WILL ACTIVATE EACH STEP OF THE ANIMATION
Slide 19 : AMMONIA N Each hydrogen atom needs one electron to complete its outer shell Nitrogen atom needs 3 electrons to complete its outer shell Nitrogen can only share 3 of its 5 electrons otherwise it will exceed the maximum of 8
A LONE PAIR REMAINS WAYS TO REPRESENT
THE MOLECULE PRESSING THE SPACE BAR WILL ACTIVATE EACH STEP OF THE ANIMATION
Slide 20 : WATER O Each hydrogen atom needs one electron to complete its outer shell Oxygen atom needs 2 electrons to complete its outer shell Oxygen can only share 2 of its 6 electrons otherwise it will exceed the maximum of 8
2 LONE PAIRS REMAIN WAYS TO REPRESENT
THE MOLECULE PRESSING THE SPACE BAR WILL ACTIVATE EACH STEP OF THE ANIMATION
Slide 21 : HYDROGEN H H H H H H H H both atoms need one electron to complete their outer shell atoms share a pair of electrons to form a single covalent bond DOT AND CROSS DIAGRAM
Slide 22 : METHANE C H H H H C H H H H H C H H H each atom needs one electron to complete its outer shell atom needs four electrons to complete its outer shell Carbon shares all 4 of its electrons to form 4 single covalent bonds DOT AND CROSS DIAGRAM
Slide 23 : AMMONIA N H H H N H H H H N H H each atom needs one electron to complete its outer shell atom needs three electrons to complete its outer shell Nitrogen can only share 3 of its 5 electrons otherwise it will exceed the maximum of 8
A LONE PAIR REMAINS
Slide 24 : WATER O H H O H H each atom needs one electron to complete its outer shell atom needs two electrons to complete its outer shell Oxygen can only share 2 of its 6 electrons otherwise it will exceed the maximum of 8
TWO LONE PAIRS REMAIN H O H
Slide 25 : OXYGEN O each atom needs two electrons to complete its outer shell each oxygen shares 2 of its electrons to form a
DOUBLE COVALENT BOND O O O
Slide 26 : Bonding Atoms are joined together within the molecule by covalent bonds.
Electrical Don’t conduct electricity as they have no mobile ions or electrons
Solubility Tend to be more soluble in organic solvents than in water;
some are hydrolysed
Boiling point Low - intermolecular forces (van der Waals’ forces) are weak;
they increase as molecules get a larger surface area
e.g. CH4 -161°C C2H6 - 88°C C3H8 -42°C
as the intermolecular forces are weak, little energy is required to
to separate molecules from each other so boiling points are low
some boiling points are higher than expected for a given mass
because you can get additional forces of attraction SIMPLE COVALENT MOLECULES
Slide 27 : Although the bonding within molecules is strong, that between molecules is weak. Molecules and monatomic noble gases are subject to weak attractive forces.
Instantaneous dipole-induced dipole forces
Because electrons move quickly in orbitals, their position is
constantly changing; at any given instant they could be anywhere
in an atom. The possibility will exist that one side will have more
electrons than the other. This will give rise to a dipole... VAN DER WAALS’ FORCES
INSTANTANEOUS DIPOLE-INDUCED DIPOLE FORCES
Slide 28 : Although the bonding within molecules is strong, that between molecules is weak. Molecules and monatomic noble gases are subject to weak attractive forces.
Instantaneous dipole-induced dipole forces
Because electrons move quickly in orbitals, their position is
constantly changing; at any given instant they could be anywhere
in an atom. The possibility will exist that one side will have more
electrons than the other. This will give rise to a dipole...
The dipole on one atom induces dipoles on nearby atoms
Atoms are now attracted to each other by a weak forces
The greater the number of electrons, the stronger the attraction
and the greater the energy needed to separate the particles. VAN DER WAALS’ FORCES
INSTANTANEOUS DIPOLE-INDUCED DIPOLE FORCES
Slide 29 : Although the bonding within molecules is strong, between molecules it is weak. Molecules and monatomic gases are subject to weak attractive forces.
Instantaneous dipole-induced dipole forces
Electrons move quickly in orbitals, so their position is
constantly changing; at any given time they could be
Anywhere in an atom. The possibility exists that one side has
More electrons than the other. This will give rise to a dipole...
The dipole on one atom induces dipoles on others
Atoms are now attracted to each other by a weak forces
The greater the number of electrons, the stronger the attraction
and the greater the energy needed to separate the particles.
NOBLE GASES ALKANES
Electrons B pt. Electrons B pt.
He 2 -269°C CH4 10 -161°C
Ne 10 -246°C C2H6 18 - 88°C
Ar 18 -186°C C3H8 26 - 42°C
Kr 36 -152°C VAN DER WAALS’ FORCES
INSTANTANEOUS DIPOLE-INDUCED DIPOLE FORCES
Slide 30 : ‘The ability of an atom to attract the electron pair in a covalent bond to itself’
Non-polar bond similar atoms have the same electronegativity
they will both pull on the electrons to the same extent
the electrons will be equally shared
Polar bond different atoms have different electronegativities
one will pull the electron pair closer to its end
it will be slightly more negative than average, d-
the other will be slightly less negative, or more positive, d+
a dipole is formed and the bond is said to be polar
greater electronegativity difference = greater polarity
Pauling Scale a scale for measuring electronegativity ELECTRONEGATIVITY
Slide 31 : ‘The ability of an atom to attract the electron pair 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
H
2.1
Li Be B C N O F
1.0 1.5 2.0 2.5 3.0 3.5 4.0
Na Mg Al Si P S Cl
0.9 1.2 1.5 1.8 2.1 2.5 3.0
K Br
0.8 2.8 ELECTRONEGATIVITY INCREASE INCREASE
Slide 32 : Occurrence occurs between molecules containing polar bonds
acts in addition to the basic van der Waals’ forces
the extra attraction between dipoles means that
more energy must be put in to separate molecules
get higher boiling points than expected for a given mass DIPOLE-DIPOLE INTERACTION Mr °C
CH4 16 -161
SiH4 32 -117
GeH4 77 -90
SnH4 123 -50
NH3 17 -33
PH3 34 -90
AsH3 78 -55
SbH3 125 -17 Mr °C
H2O 18 +100
H2S 34 -61
H2Se 81 -40
H2Te 130 -2
HF 20 +20
HCl 36.5 -85
HBr 81 -69
HI 128 -35 Boiling points
of hydrides
Slide 33 : Occurrence not all molecules containing polar bonds are polar overall
if bond dipoles ‘cancel each other’ the molecule isn’t polar
if there is a ‘net dipole’ the molecule will be polar
HYDROGEN CHLORIDE TETRACHLOROMETHANE WATER POLAR MOLECULES NET DIPOLE - POLAR NON-POLAR NET DIPOLE - POLAR
Slide 34 : Evidence place a liquid in a burette
allow it to run out
place a charged rod alongside the stream of liquid
polar molecules are attracted by electrostatic attraction
non-polar molecules will be unaffected POLAR MOLECULES NET DIPOLE - POLAR NON-POLAR
Slide 35 : BOILING POINTS OF HYDRIDES Mr °C
CH4 16 -161
SiH4 32 -117
GeH4 77 -90
SnH4 123 -50
NH3 17 -33
PH3 34 -90
AsH3 78 -55
SbH3 125 -17 Mr °C
H2O 18 +100
H2S 34 -61
H2Se 81 -40
H2Te 130 -2
HF 20 +20
HCl 36.5 -85
HBr 81 -69
HI 128 -35 GROUP IV GROUP V GROUP VI GROUP VII The values of certain hydrides are not
typical of the trend you would expect
Slide 36 : BOILING POINTS OF HYDRIDES The boiling points of the hydrides increase with molecular mass. CH4 has the lowest boiling point as it is the smallest molecule. CH4 SiH4 GeH4 PbH4 Larger molecules have greater intermolecular forces and therefore higher boiling points GROUP IV
Slide 37 : BOILING POINTS OF HYDRIDES NH3 has a higher boiling point than expected for its molecular mass. There must be an additional intermolecular force. NH3 GROUP V
Slide 38 : BOILING POINTS OF HYDRIDES H2O has a very much higher boiling point for its molecular mass. There must be an additional intermolecular force. H2O GROUP VI
Slide 39 : BOILING POINTS OF HYDRIDES HF has a higher boiling point than expected for its molecular mass. There must be an additional intermolecular force. HF GROUP VII
Slide 40 : BOILING POINTS OF HYDRIDES GROUP IV
GROUP V
GROUP VI
GROUP VII H2O HF NH3 The higher than expected boiling points of NH3, H2O and HF are due to intermolecular HYDROGEN BONDING
Slide 41 : BOILING POINTS OF HYDRIDES GROUP IV
GROUP V
GROUP VI
GROUP VII
Slide 42 : an extension of dipole-dipole interaction
gives rise to even higher boiling points
bonds between H and the three most electronegative elements,
F, O and N are extremely polar
because of the small sizes of H, F, N and O the partial charges are
concentrated in a small volume thus leading to a high charge density
makes the intermolecular attractions greater and leads
to even higher boiling points HYDROGEN BONDING
Slide 43 : HYDROGEN BONDING - ICE each water molecule is hydrogen-bonded to 4
others in a tetrahedral formation
ice has a “diamond-like” structure
volume is larger than the liquid making it
when ice melts, the structure collapses
slightly and the molecules come closer; they
then move a little further apart as they get
more energy as they warm up
this is why…
water has a maximum density at 4°C
ice floats. hydrogen bonding
lone pair
Slide 44 : HYDROGEN BONDING - ICE hydrogen bonding
Slide 45 : HYDROGEN BONDING - HF Hydrogen fluoride has a much higher boiling point than one would expect for a molecule with a relative molecular mass of 20
Fluorine has the highest electronegativity of all and is a small atom so the bonding with hydrogen is extremely polar F H F H H F H F d + d ¯ d + d ¯ d + d ¯ d + d ¯ hydrogen bonding
Slide 46 : A dative covalent bond differs from 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.
Lewis base a lone pair donor
Lewis acid a lone pair acceptor DATIVE COVALENT (CO-ORDINATE) BONDING 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.
Slide 47 : Boron trifluoride-ammonia NH3BF3
Boron 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.
Slide 48 : MOLECULAR
SOLIDS
Slide 49 : IODINE
At room temperature and pressure, iodine is a greyish solid. However it doesn’t need to be warmed much in order to produce a purple vapour. This is because iodine is composed of diatomic molecules (I2) which exist in an ordered molecular crystal in the solid state. Each molecule is independent of the others, only being attracted by van der Waals’ forces. Therefore, little energy is required to separate the iodine molecules. MOLECULAR SOLIDS
Slide 50 : COVALENT NETWORKS
GIANT MOLECULES
MACROMOLECULES
They all mean the same!
Slide 51 : DIAMOND, GRAPHITE and SILICA
Many atoms joined together in a regular array
by a large number of covalent bonds
GENERAL PROPERTIES
MELTING POINT Very high
structure is made up of a large number of covalent bonds,
all of which need to be broken if atoms are to be separated
ELECTRICAL Don’t conduct electricity - have no mobile ions or electrons
but... Graphite conducts electricity
STRENGTH Hard - exists in a rigid tetrahedral structure
Diamond and silica (SiO2)... but
Graphite is soft GIANT (MACRO) MOLECULES
Slide 52 : GIANT (MACRO) MOLECULES DIAMOND
MELTING POINT VERY HIGH
many covalent bonds must be broken to separate atoms
STRENGTH STRONG
each carbon is joined to four others in a rigid structure
Coordination Number = 4
ELECTRICAL NON-CONDUCTOR
No free electrons - all 4 carbon electrons used for bonding
Slide 53 : GIANT (MACRO) MOLECULES GRAPHITE
MELTING POINT VERY HIGH
many covalent bonds must be broken to separate atoms
STRENGTH SOFT
each carbon is joined to three others in a layered structure
Coordination Number = 3
layers are held by weak van der Waals’ forces
can slide over each other
ELECTRICAL CONDUCTOR
Only three carbon electrons are used for bonding which
leaves the fourth to move freely along layers
layers can slide over each other
used as a lubricant and in pencils
Slide 54 : GIANT (MACRO) MOLECULES DIAMOND GRAPHITE
Slide 55 : GIANT (MACRO) MOLECULES SILICA
MELTING POINT VERY HIGH
many covalent bonds must be broken to separate atoms
STRENGTH STRONG
each silicon atom is joined to four oxygens - C No. = 4
each oxygen atom are joined to two silicons - C No = 2
ELECTRICAL NON-CONDUCTOR - no mobile electrons
Slide 56 : METALLIC
BONDING
Slide 57 : METALLIC BONDING Involves a lattice of positive ions surrounded by delocalised electrons
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.
Slide 58 : METALLIC BONDING Involves a lattice of positive ions surrounded by delocalised electrons
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 in regular close packed 3-dimensional crystal lattices.
Slide 59 : METALLIC BONDING Involves a lattice of positive ions surrounded by delocalised electrons
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 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.
Slide 60 : METALLIC BOND STRENGTH Depends on the number of outer electrons donated
to the cloud and the size of the metal atom/ion. The strength of the metallic bonding in sodium is relatively weak because each atom donates one electron to the cloud. Na
Slide 61 : METALLIC BOND STRENGTH Depends on the number of outer electrons donated
to the cloud and the size of the metal atom/ion. The strength of the metallic bonding in sodium is relatively weak because each atom donates one electron to the cloud.
The metallic bonding in potassium is weaker than in sodium because the resulting ion is larger and the electron cloud has a bigger volume to cover so is less effective at holding the ions together. Na K
Slide 62 : METALLIC BOND STRENGTH Depends on the number of outer electrons donated
to the cloud and the size of the metal atom/ion. The strength of the metallic bonding in sodium is relatively weak because each atom donates one electron to the cloud.
The metallic bonding in potassium is weaker than in sodium because the resulting ion is larger and the electron cloud has a bigger volume to cover so is less effective at holding the ions together.
The metallic bonding in magnesium is stronger than in sodium because each atom has donated two electrons to the cloud. The greater the electron density holds the ions together more strongly. Na Mg K
Slide 63 : METALLIC PROPERTIES MOBILE ELECTRON CLOUD ALLOWS THE CONDUCTION OF ELECTRICITY For a substance to conduct electricity it must have mobile ions or electrons.
Because the ELECTRON CLOUD IS MOBILE, electrons are free to move throughout its structure. Electrons attracted to the positive end are replaced by those entering from the negative end. Metals are excellent conductors of electricity
Slide 64 : MALLEABLE CAN BE HAMMERED INTO SHEETS
DUCTILE CAN BE DRAWN INTO RODS AND WIRES
As the metal is beaten into another shape the delocalised electron cloud continues to bind the “ions” together.
Some metals, such as gold, can be hammered into sheets thin enough to be translucent. METALLIC PROPERTIES Metals can have their shapes changed relatively easily
Slide 65 : HIGH MELTING POINTS
Melting point is a measure of how easy it is to separate individual particles. In metals it is a measure of how strong the electron cloud holds the + ions.
The ease of separation of ions depends on the...
ELECTRON DENSITY OF THE CLOUD
IONIC / ATOMIC SIZE
PERIODS Na (2,8,1) < Mg (2,8,2) < Al (2,8,3)
m.pt 98°C 650°C 659°C
b.pt 890°C 1110°C 2470°C METALLIC PROPERTIES Na+ Al3+ Mg2+ MELTING POINT INCREASES ACROSS THE PERIOD
THE ELECTRON CLOUD DENSITY INCREASES DUE TO THE GREATER NUMBER OF ELECTRONS DONATED PER ATOM. AS A RESULT THE IONS ARE HELD MORE STRONGLY.
Slide 66 : HIGH MELTING POINTS
Melting point is a measure of how easy it is to separate individual particles. In metals it is a measure of how strong the electron cloud holds the + ions.
The ease of separation of ions depends on the...
ELECTRON DENSITY OF THE CLOUD
IONIC / ATOMIC SIZE
GROUPS Li (2,1) < Na (2,8,1) < K (2,8,8,1)
m.pt 181°C 98°C 63°C
b.pt 1313°C 890°C 774°C METALLIC PROPERTIES MELTING POINT INCREASES DOWN A GROUP
IONIC RADIUS INCREASES DOWN THE GROUP. AS THE IONS GET BIGGER THE ELECTRON CLOUD BECOMES LESS EFFECTIVE HOLDING THEM TOGETHER SO THEY ARE EASIER TO SEPARATE. Na+ K+ Li+