PERIODICITY (p. 318)
- Repeating chemical properties of the
chemical elements that occurs when layed
out in order of atomic mass is called
periodicity
- Periodic table first devised by Mendeleev
in 1870s (based on atomic weights)
- order of elements
in periodic table
anomalies
e.g. K comes
before Ar
s-block
p-block
f-block
d-block
-Most elements are metals ⇒ basic, ionic oxides e.g. BaO, Sc2O3
-Non metals ⇒ covalent, acidic oxides e.g. SO2, NO2
- Diagonal band from B to At are metalloids – semi-conductors or
some metallic character ⇒ neutral or amphoteric oxides e.g. SiO2
-Elements divided into blocks depending on the valence electron
configuration (sub-shell that is partially filled in Aufbau principle)
-E.g. p-block elements have p orbitals partly filled
The period number is the principle quantum number n of the valence
orbital
-Elements in same valence electron configuration and so have similar
chemistry
-Mendeleev accurately predicted properties of Ga (group III) and Ge
(group IV) before they were discovered
1
- Groups I to VIII (s- and p- block) are known as main group or
representative elements
- d-block called transition metals
-Incomplete d shell – can form a variety of oxidation states e.g.
Fe salts can be either Fe2+ (4s0 3d6) or Fe3+ (4s0 3d5)
- Zn, Cd, Hg different as d-shell remains filled (s2 d10) – only s
electrons are lost during ionisation to M2+
Lanthanide Metals
Actinide Metals
TRANSITION
METALS
Zn
Cd
Hg
-f-block metals e.g. La, U, Pu have valence electrons in f orbitals
(quantum number l = 3)
Ion formation for Main Group Elements (p. 321)
When main group elements forming ions, electrons are either added to
the partially filled shell to complete octet….
e.g.
N
[He]2s22p3
N3[Ne]
S
[Ne]2s22p3
S2[Ar]
Or electrons are removed to reveal the stable, noble gas core e.g.
Na
[Ne]3s1
Na+
[Ne]
Ca
[Ar]4s2
Ca2+
[Ar]
Species with the same valence electron configuration are isoelectronic
e.g. N3-, Ne and Mg2+ are all 1s2 2s2 2p6 (10 electrons)
Likewise, Se2-, Kr and Rb+ are isoelectronic (36 electrons)
2
Explaining Periodicity
-physical properties of atoms show striking periodicity:
⇒ affects ease with which atoms give up or accept electrons
⇒ need to explain physical trends to explain chemistry of the elements
Can explain most trends in periodic table by:
-increasing nuclear charge across periods (left to right)
-increasing valence shell number down groups (top to bottom)
Effective Nuclear Charge (p. 322)
- Electrons in inner orbitals shield valence electrons very effectively
but:
- Electrons in same orbitals shield valence electrons less effectively
-As electrons are added within one valence shell (i.e. across a period)
shielding stays approximately constant
BUT:
Z increases (more protons are added)
⇒ effective nuclear charge increases
⇒ electrons held more tightly on RHS of period
⇒ affects:
-ionisation energy
-atomic size
-electronegativity
⇒ reactivity
3
Atomic radius (p. 323)
-Decreases across periods due to increasing effective nuclear charge
(pulls electrons closer to nucleus, lowers energies of valence orbitals)
Alkali metal atoms largest - single s electron loosely held so strays far
-Increases down groups due to successive shells being
occupied (progressively further from nucleus)
p
Ex 8 ractise
.2, Q
s
8.40 8.38,
Ionic Radius (p. 325)
greater positive charge ⇒ smaller cations
(exposed noble gas core poorly shielded)
⇒ Cations are smaller than parent atoms
greater negative charge ⇒ larger anions
(due to mutual repulsion of extra electrons
cf. radii of isoelectronic F- and Mg2+)
⇒ Anions are larger than parent atoms
4
Ionic Radius (p. 325)
Diagonal
relationships mirror
similar chemistry:
e.g. Li+ and Mg2+
both have ionic
radius 60-65 pm
also Na+ and Ca 2+
Atomic & Ionic Radius
Q: For each pair of species which has the larger radius?
h
Ex 8 ave a g
.3, Q o
s
8.46 8.44,
i) K or K+
K is large as single 3s electron is well shielded
K+ is a well-shielded argon core
K 220 pm
K+ 133 pm
ii) Mg2+ or Al3+
both ions are isoelectronic (same shielding)
Ne noble gas core remains
Al3+ feels greater nuclear charge
Mg2+ 65 pm
Al3+ 50 pm
iii) N3- or O22
3
4
Ne
Mg
Al
5
Atomic & Ionic Radius
Q: For each pair of species which has the larger radius?
i) K or K+
K is large as single 3s electron is well shielded
K+ is a well-shielded argon core
ii) Mg2+ or Al3+
both ions are isoelectronic (same shielding)
Ne noble gas core remains
Al3+ feels greater nuclear charge
iii) N3- or O2both anions are isoelectronic (same shielding)
Ne noble gas configuration
O2- feels greater nuclear charge
h
Ex 8 ave a g
.3, Q o
s
8.46 8.44,
K 220 pm
K+
N O133 pm
Ne 2
3
4
Mg2+ 65 pm
Al3+ 50 pm
N3O2-
171 pm
140 pm
Ionisation Energy (p. 329)
-Minimum energy required to remove an electron from an atom in the
gas phase:
X
X+ + e-
∆Hº = 1st ionisation energy
- I.E. determined by effective nuclear charge:
- I.E. increases across period (more protons added but shielding stays
same; electrons pulled in closer
by greater nuclear charge)
- I.E. decreases down group
(valence electrons become
more distant from nucleus)
-I.E. is highest for stable noble
gas configurations
6
Ionisation Energy (p. 329)
Q: Put these elements in order of increasing first ionisation energy:
i) Ca Be Ba
A: metals are in same group but different periods
⇒ all have 2s2 electrons shielded by noble gas core
⇒ effective nuclear charge depends only on distance from nucleus
⇒ effective nuclear charge depends on n (period no. quantum number)
⇒ larger Z ⇒ lower I.E.
Ba < Ca < Be
1st I.E. kJ/mol
500
600
900
Ionisation Energy (p. 329)
Q: Put these elements in order of increasing first ionisation energy:
ii) He Li H
A: H and He are in same period
⇒ same number of electron shells
⇒ He greater nuclear charge
⇒ both He electrons in same orbital (poor shielding)
⇒ electrons held tighter in He than H
Ex 8 read
.4,
8.54 Qs 8.5
2,
&8
.56
Li has single 2s electron (further from nucleus)
⇒ well shielded by inner 1s electrons
⇒ 2s electron more loosely held than single H 1s electron
Li < H < He
1st I.E. kJ/mol
519
1310
2370
7
Electronegativity: the tendency of an atom to attract
electron density towards itself in a chemical bond
(p.369)
electron-poor
electron-rich
H
F
F is much more
electronegative than H
⇒ polar bond
greatest for elements with high (endothermic) ionisation energy
and high (exothermic) electron affinity
e.g. F, O, N
lowest for elements with low I.E. and low electron affinity e.g. K, Ba
amongst elements, electronegativity:
- increases across periods
– due to decreasing atomic size (greater nuclear
charge outweighs minimal extra electron shielding
- decreases down groups
– due to increasing atomic size (electrons not
attracted so strongly by spread out charges as pointlike charges)
8
⇒ F is most electronegative element
Ex 9 try
.2, Q
Cs is most electropositive element
9.36
between elements of similar electronegativity
e.g. C-H, N-Cl or I-I or C-Se there is no polarity in bond
if elements of very different electronegativity (> 2) ⇒ complete
electron transfer ⇒ ionic solid e.g. K+Br-, Ba2+O2-
Oxidation State (p. 134-135)
- the charge an atom would have if the valence electrons in a molecule
or complex were completely transferred to the most electronegative
atoms
e.g. in H-O-H
oxidation state of H is +1
oxidation state of O is -2
e.g. iodine trifluoride (IF3)
oxidation state of F is always -1
oxidation state of I is +3
e.g. in manganate ion, MnO4oxidation state of O is -2 (total of -8 for all oxygens)
oxidation state of Mn is +7 can write as MnVIIO4-
e.g. in [Co(H2O)2Cl4]2-
??
[CoII(H2O)2Cl4]2-
e.g. in H2SeO4
??
H2SeVIO4
9
Oxidation State (p. 134-135)
Q: What are the oxidation states of the underlined atoms in the
following molecules/ions?
Nitrate ion: NO3assign ox. no. for most electronegative atoms first:
O is most electronegative - needs 2 electrons to complete octet
⇒ O has oxidation state of -2
⇒ O provides six negative charges
-2
⇒ ion is -1 charged overall
+5
⇒ N must have +5 oxidation state
O
O
-2
N
-2
O
Oxidation State (p. 134-135)
Q: What are the oxidation states of the underlined atoms in the
following molecules/ions?
Arsenous acid: H3AsO2
assign ox. no. for most electronegative atoms first:
O is most electronegative - needs 2 electrons to complete octet
⇒ O has oxidation state of -2
⇒ O provides four negative charges
⇒ H is in +1 oxidation state
⇒ H provides three positive charges
p
⇒ molecule is neutral overall
Ex 4 ractise
.4, Q
s 4.4
4.48
⇒ As must have +1 oxidation state
6,
&
4.50
10
CHEMICAL BONDING (Ch. 9 & 10)
‘Bonding’ is general term referring to forces that hold together any
type of chemical species: molecules, groups of molecules, atoms or
ions
Intermolecular bonding acts between molecules e.g.
Van der Waals’ forces
Dipole-dipole forces
Ion dipole forces
Hydrogen bonding
-weak (2-25 KJ/mol) but responsible for bulk, physical properties
of matter
Intramolecular bonding holds together atoms e.g.
Ionic bonding
Metallic bonding
Covalent bonding
-strong (150-400 KJ/mol) ⇒ responsible for chemical properties
INTRAMOLECULAR BONDING
What is a chemical bond?
-an attractive force holding together atom(s) or ions(s) that makes
them function as a unit.
-bond forms if it makes the system is lower in energy than when
atoms are apart.
Energy input required to break bond = bond strength or bond
energy
e.g. the Cl-Cl bond has a strength of 243 kJ/mol
the O=O bond has a strength of 499 kJ/mol
11
The way atoms are bonded together shapes physical and chemical
properties of a compound.
e.g. graphite is grey, soft, conductor of electricity
diamond is a transparent insulator, the hardest substance known
both are carbon – allotropes of the same element
e.g. C and Si have similar chemistry but:
SiO2 is a brittle unreactive crystalline solid
CO2 is a gas
Types of Chemical Bond
Atoms can bond to each other by three different principal means:
Ionic bonding (e.g. NaCl)
Covalent bonding (e.g. H2O, CO2)
Metallic bonding (e.g. Cu, Fe, K)
Ionic Bonding (p.359)
when K atom and Cl atom brought
together – energetically favourable
for electron to be transferred from
K to Cl:
- both K+ and Cl- have stable [Ar]
electron configuration with an octet
of electrons i.e. 3s23p6
K
+
Cl
K+
_
Cl
(Lewis formula of KCl)
12
KCl is an ionic solid:
-conducts electricity when in aqueous solution or
molten (when its ions are free to move in liquid)
- ionic bonds are strong!
- ionic solids have high melting points (750 °C for KCl)
(3000 °C for MgO)
Covalent Bonding (p. 366)
-When two identical atoms approach
– no strong electrostatic forces
-Small forces e.g. two H atoms:
-Repulsion:
-Attraction:
- two negative electron
clouds
- two positive nuclei
(raises energy)
protons-electrons
(lowers energy)
At optimum distance apart electron
density resides between nuclei (electrons
feel extra attraction of both nuclei
13
Energy of system reaches a minimum (-458 KJmol-1 for H2) at
separation corresponding to bond length (74 pm in H2)
At smaller internuclear distances proton-proton repulsion
dominates, raising energy
A H-H covalent
bond has been
formed. The bond
strength (energy
needed to pull
atoms apart again)
is -458 KJmol-1
can represent H2 as
H:H
– each H has full valence shell (n = 1) and the same electron
configuration as He
H2 is stable
H:H is the Lewis structure of H2
Lewis structure: a two-dimensional representation of the valence
electrons contained in a molecule using dots as electrons
- Lewis structures useful way of counting valence electrons - the
electrons in outer shell that are potentially available for bonding
14
Other homonuclear diatomic molecules form by making covalent
bond
e.g. F has seven valence electrons so shares one electron:
2
The Lewis structure of F2
each F obtains octet of electrons (four orbitals in n = 2 shell
remember!) and electron configuration [Ne]
note F2 has six non-bonded or lone pairs of electrons in the valence
shell (n = 2)
Octet rule: atoms proceed as far as possible toward completing
their octets by sharing electron pairs in covalent bonds
Electronegativity: the tendency of an atom to attract
electron density towards itself in a chemical bond
(p.369)
15
The greater the difference in electronegativity between
two elements, the greater the polarity of the bond
extra ionic character causes stronger bonds –
electronegativity is measure of H-X bond strength
compared to H-H and X-X
r
Ex 9 ead
.
try 2 and
Q9
.40
Polar Covalent Bonds (p. 369)
In H2 and other homonuclear diatomic
molecules e.g. N2, Cl2, both atoms identical
so electron density is arranged symmetrically
H
F
Most bonds are polar
covalent: electrons not
transferred between atoms
but is unequally shared
due to slightly different
electronegativity
Ionic bond – other extreme – very different
atoms so electrons completely transferred
very different electronegativity
16
Molecule Polarity (p. 409)
- bond polarity can cause polyatomic molecules to become
polar e.g. H2O, H2C=O
- some molecules have polar bonds but no dipole moment due
to symmetry e.g.
F
O
C
B
O
F
F
Cl
C
Cl
Cl
Cl
CHCl3 has a dipole moment however due to imbalance of electron
density:
Cl
Non polar
C
Cl
CCl4 Cl
Cl
H
C
Cl
Cl
Cl
Polar
CHCl3
Importantly for life, water is
a polar molecule:
+
polar molecules spontaneously align
themselves in an electric or magnetic field:
17
polar molecules are affected by electrical field
e.g. presence of a statically-charged rod:
Non-polar liquid
e.g. CCl4
Polar liquid
e.g. H2O
Lewis structures of polyatomic molecules (p.372)
- octet rule can be describe electron arrangement in many-atom molecules
- Lewis structures of molecules are a 2-D representation of molecule – do
not describe its shape
-a shared electron pair (covalent bond) can also be drawn as a line
between atoms
easier when two or more electron pairs are shared
e.g. for O2:
O O
two bonded pairs of electrons
four lone (non-bonded) pairs
18
rules for drawing Lewis structures:
-count total number of valence electrons
-take first element of formula as central atom e.g. C in CO2 or P in PCl5
-arrange substituent atoms around central atom so each atom has noble
gas configuration
observations from Lewis structures:
- In NH3 N’s complete octet includes lone pair
- In C2H4 two electron pairs between carbons – C=C double bond
- BF3 has only six valence electrons – octet incomplete – BF3 is an electron
deficient compound
Q: Write the Lewis structure of hypochlorous acid, HClO
A: first find number of valence electrons on each atom (see periodic table):
H: 1
O: 6
Cl: 7
total of 14 valence electrons (7 electron pairs)
In acids, H atom always attached to O (e.g. CH3COOH, HONO)
so atomic arrangement must be HOCl
-form single bonds between atoms H:O:Cl or H-O-Cl
-five electron pairs remain
H O Cl
pr
Ex 9 actise
Q9. .3-9.5 &
44,
Q9
.46
-three more to complete Cl octet
-two more to complete O octet
(H valence shell only holds 2 electrons)
19
Resonance Structures (p. 377)
For molecules with multiple bonds we can sometimes write more than
one Lewis structure e.g. the cyanate (NCO-) ion
_
_
N C O
N C O
electrons can be rearranged so formal negative charge lies on either N
or O
resonance structures differ only in location of one bond
real structure is an ‘average’ of resonance forms – a resonance hybrid
e.g. carboxylate anion has two resonance forms:
_ 1.5 bonds for each
O
_
CH3 C
O
C-O (bond order =
1.5)
O
-each O carries half
a negative charge
CH3 C
O
se e
Ex 9
.8
-)
Nitrate ion (NO3 has has three resonance forms:
__
O
_
O
N
+
O
O
_
_
O
_
N
+
O
_
O
O
N
+
O
all three N-O bonds are identical - each has 2/3 of a negative charge
20
Q: Suggest two resonance forms for sulphur dioxide, SO2. Hint: the
S atom lies between the two oxygens and both SO bond lengths are
the same.
A: First draw basic Lewis structure. All atoms are Gp VI so 18
valence electrons (9 pairs)
Use 2 e pairs for single bonds (O-S-O) so 7 electron pairs remain
use one pair to form double bond
(O–S=O) so 6 remain
lone pairs required to complete octet 3 1 2
adding lone pairs gives:
remember formal charges
O
_O S
+
pra
ct
Ex 9 ise
.6
Interchanging bonds (an picture by shifting two electron
pairs) gives the other resonance structure:
_O S O
+
O S O_
+
two S-O bond lengths and strengths are equal
(bond order = 1.5 for each S-O)
Q9. try
52,
Q
9.56
21
Incomplete Octets (p.381)
-Some compounds have fewer valence electrons than full octet
F
BF3
F B
F
6 valence electrons (3 electron pairs)
BF3 is an electron deficient compound
read
Ex 9
.9
Expanding the Octet (p.381)
-When drawing Lewis structure for SF6 we must place S as central
atom and first satisfy octet rule for F:
F
F S
F
F
F
SCl2
Cl
F
SF6
Octet
expanded
S
Cl
Octet rule obeyed
48 valence electrons (6 from S, 7 from each F)
S has 12 electrons around it – exceeds the octet rule by four
electrons. How?
22
Unlike 2nd period elements (B, C, N, O, F) 3rd period (and below)
non-metals have accessible empty d orbitals:
empty 3d orbitals not much higher in energy than 3p orbital- can
hold extra electrons
- in SF6 there are six bonded pairs around central atom:
two electron pairs are housed in a sulphur 3d orbital
In PCl5 only one d orbital is required:
rea
Ex 9 d
& 9 .10
.12
23
VSEPR – Predicting the shapes of molecules (p. 400-403)
Lewis structure is a 2-D scheme– gives no idea about 3-D
molecular shape
number of electron pairs around central atom can be used to predict
shape of molecule using VSEPR model
Valence Shell Electron Pair Repulsion
Main ideas:
- electron pairs around central atom lie as far apart as possible in
order to minimise repulsions between them
- non-bonded (lone) pairs require slightly more angular space than
bonded pairs as they lie closer to surface of central atom
Consider electron deficient BeCl2 – only two bonding electron
pairs
– 180° apart is lowest energy arrangement (minimum repulsion)
so BeCl2 is a linear molecule
Cl Be Cl
In BF3 there are three electron pairs on boron - 120° apart is
lowest energy arrangement (minimum repulsion)
F
F B
120º
F
so BF3 is a trigonal molecule
– although B-F bond is polar, molecule
has no dipole moment
24
In CH4 are four electron pairs on carbon – in 3-D space H atoms
furthest apart if C-C-H angles are all 109.5° - tetrahedral geometry
H
109.5°
C
H
H
H
so methane is a tetrahedral molecule
Lewis structure of ammonia suggests 4 electron pairs also in same
valence orbitals – NH3 is isoelectronic with CH4
N
H
H
H
NH3 has four tetrahedrally-arranged electron
pairs - three H atoms form base of a pyramid –
NH3 structure is a trigonal pyramid
- might expect H-N-H angle to be tetrahedral 109.5°
- but one of ammonia’s pair is non-bonding so requires more angular
space - hydrogens pushed together - H-N-H angle is only 107°
25
In water (isoelectronic again with ammonia and methane) there are
two lone pairs – even more distortion of tetrahedral bond angle
⇒ H-O-H angle only 104.5°
O
H
H
104.5°
- makes water even more polar molecule than would be with exact
tetrahedral geometry
Se
fig.1 e
0.1
Geometries of valency expanded central atoms
Five electron pairs – one 3d orbital
occupied
trigonal bipyramid – three
electron pairs arranged trigonally
(120º apart) in same plane and
two electron pairs perpendicular to
plane
e.g. PCl5
26
Geometries of valency expanded central atoms
Six electron pairs –
octahedral – six
electron pairs arranged
90 apart from each other
e.g. SF6
VSEPR geometries summary (p. 401)
27
Q: Predict the structure of first noble gas compound to be
synthesised, XeF4. Is it a polar molecule?
A: Xe central atom (8 valence electrons) so 36 valence electrons
total
F always completes its octet with single bond – so Lewis structure
is:
Six e pairs around Xe – octahedral geometry
But two lone pairs so two structures possible:
(lone pairs adjacent, 90° apart)
OR
(lone pairs 180° apart)
(lone pairs adjacent, 90° apart)
(lone pairs 180° apart)
- lone pairs repel more than bonded pairs must be placed 180° apart–
XeF4 is square planar
- All F atoms lie n same plane so no dipole moment
28
Q: Predict the structure of the PCl4+ cation
A: P central atom (4 valence electrons due to charge) so 32 valence
electrons total
P has stable octet (no expansion of octet required)
Lewis structure is:
Cl
Cl
P Cl
+
Cl
Four electron pairs around P so Cl arranged tetrahedrally:
Cl
+P
Cl
Cl
Cl
p
Ex 1 ractise
0.1
&1
Qs 1
0.
0
10.1 .8, 10.1 2,
2&
0
10.1 ,
4
- in VSEPR multiple bonds are treated as containing one electron
pair – so formaldehyde (H2C=O, three substituent atoms on central
carbon) is trigonal – all atoms lie in same plane
Q: Predict the structure of SO2. Is it non-polar like CO2?
A: S central atom (6 valence electrons) so 18 valence electrons
total
18 valence electrons, S is central atom. There are two Lewis
resonance structures:
_O S O
+
O S O_
+
VSEPR can be applied to any one of the resonance structures in
determining molecular geometry
29
-in VSEPR double bonds do not influence molecule shape
-therefore S surrounded by 3 electron pairs – one in single bond
-one in double bond
-one lone pair
-electron pairs arranged trigonally around S making SO2 a V-shaped
molecule:
+
S O
_
O
unlike CO2, SO2 is a polar molecule. Lone pairs cause no distortion in
trigonal arrangement as electron pairs sufficiently far apart – O-S-O
bond angle is full 120 degrees
Similarly SO42- (containing two double bonds, no S lone pairs) is
tetrahedral
Q: Use VSEPR to decide whether SF4 is a polar
molecule.
A: S central atom (6 valence electrons) so 34
valence electrons total
F
F S
F
F
Lewis structure indicates one lone pair on S
total 5 electron pairs around S – trigonal
bipyramid
trigonal bipyramid has two types environment –
axial (2 sites) and equatorial (3 sites in same
plane)
two possible structures: lone pair can be placed
in axial or equatorial positions
30
equatorial always favoured – more space:
SF4 - perpendicular S-F bonds distorted away
from lone pair
Structure if
exactly trig
bipyramid:
(lone pairs always require more angular space
than bonded pairs)
VSEPR structure
allowing for
distortion caused by
S lone pair
31