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Chemical Bonds: Formation of Compounds
Chapter 11
1st shell
2nd shell
s
d subshell
p
d
f
f subshell can be placed in another loop.
Periodic System consists of periods
(n= 1, 2,…) called shells, Each
shell contains subshells (s, p, d,
f…). Elements are classified in
groups having similar chemical
characteristics. Noble gases have
filled s and p subshell, effectively
having 8 e- in the valence shell.
Transition metals populate 4s
before 3d. Before we continue with
filling 4p, 10 e- must go into 3d. We
put them in a loop since we are
.
stalled filling 4th shell.
This Periodic table has 8 groups.
Atomic properties
Ionization energy
… is the energy
required to remove
an e- from the atom.
and increases
bottom to top,
up a group.
It increases left to
right across a
period…
The behavior
exactly opposite to
that of atomic radii!
Ionization Energy vs. Atomic Number
Atomic radii
increase
right to left
across the
period,
and top to
bottom
down the
group.
Chemical Bonding
A molecule is a collection of atoms bound together. It is considered as an element if all
atoms are of the same type (e.g. H2), or a compound if it is made of different atoms.
A bond between two atoms of hydrogen will occur spontaneously if the atoms are within
a close proximity, i.e. when attractive forces between the nucleus of one atom and the eof the other overcome repulsive forces among their e-, thus pulling the atoms together.
74 pm
In a covalent bond between two hydrogen
atoms, each atom apparently has 2 electrons in
its shell. That gives each H atom the electronic
configuration of the closest noble gas, He.
Electrons are shared between two atoms.
Covalent bond
H:H
can be
..H
H
presented as
H–H
any of the
following ways:
Remember that
the line stands
for a pair of e-!
1 pm = 10-12 m
The bond that is
formed is called
covalent bond.
(Co-means partner,
valent refers to
valence electrons).
It always releases
energy.
Filled-shell e- (or core e-) are almost
never involved in the bond as they
are too close to their own nucleus.
2 He atoms
will never
form a bond
because
Energy of
He2 > 2 He.
Molecules and the Octet Rule
Elements ‘want’ to have the electronic configuration identical to that of noble gas (8e-).
When form a molecule, atoms achieve the octet (8e-) by sharing e- with other atoms.
Hydrogen is an exception, as it only needs
one more electron to fill its 1s orbital.
Electron
from C
atom
How many electrons an element needs to
satisfy the octet rule could be found in two
ways: 1) From the position of the element
in periodic table with respect to noble gas;
2) From short electron configuration:
1) O is two places left of Ne –
needs 2 e2) O: [He] 2s2 2p4
has 6 e-, needs two.
C: [He] 2s2 2p2
Needs 4 e-
Ca: [Ar] 4s2
Needs 6 e-
Four H atoms form
covalent bonds
with the C atom
All atoms in
the molecule
have the
electronic
configuration
identical to
that of the
closest noble
gas.
Gaining six electrons is difficult. For Ca to
obtain octet, it is easier to loose two e-.
Electron
from H
atom
C: 8eH: 2e-
Core electrons
Valence
electrons
Nonmetals usually gain electrons, metals loose them.
O: 8eH: 2e-
Step 1: Find the total number of
Drawing Dot Diagrams
valence e- for each atom by adding
e- in the short electron configuration.
CO2
CO
For ions, adjust e- count accordingly [He] 2s2 2p2
[He] 2s2 2p2 [He] 2s2 p4
(subtract e- for cation, add for anion).
[He] 2s2 2p4
4e- + 2 x 6eStep 2: Assume that the first non4 e- + 6 eTotal: 16ehydrogen atom in the formula of
Total: 10e
the group is the “central” (less
O–C–O
electronegative) atom. Connect
C–O
16e – 4e = 12eperipheral atoms with the central
Technically, no ‘central’ atom
in lone pairs
atom with single bonds.
here, but the rule applies as C
..
..
The central atom must form multiple bonds,
is less electronegative atom.
:O
– C – .O. :
hence H can never be the central atom.
.
.
or
10e – 2e = 8eStep 3: Subtract 2e- for each bond
..
..
in
lone
pairs
from total #e to get #e in lone pairs.
:O
..
. . – C – .O. :
Step 4: Put in the remaining
or
: C – .O. :
..
..
electrons, two at a time, as lone
:O
pairs. Satisfy octet to the terminal
. . – C – .O. :
atoms first; if there are any e- pairs
left, put them on the central atom.
Step 5: Check that each atom has
octet satisfied (doublet for H). If not,
move e- pair(s) from the adjacent
atom to form multiple bonds.
Practice with CO32-, SO3, etc.
..
..
O
..=C=O
..
..
: .O. – C Ξ O :
:CΞO:
..
:OΞC–O
. .:
Equivalent structures, or resonance forms.
..
: C – .O. :
C2H4
CO324e + (3 x 6e) + 2e = 24eO
O
C
total
in lone pairs
..
:O
..
C
..
O:
..
:.O. :
Octet rule not
satisfied for C
..
:O
..
Which C is the central atom?
2x4e + 2x1e + 2x7e = 24e-
BOTH ! 2x4e+4x1e=12e- total
total
H 12e – (5x2e) = 2e
H
O
24e – (3 x 2e) = 18e-
C
C
H
in lone pairs
H
H
H
C
C
H
H
:.O. :
O:
..
H
:C
C
H
Octet rule not
satisfied for C
H has no e- to give
in for double bonds
Two more resonance
structures for CO32ion are possible.
SO3 is identical to this; try it yourself.
H
H
2C
C2H2Cl2
C cannot
pull
electrons
from Cl
H
H
C
C
: Cl
. . : : Cl
. .:
H
..
: Cl :
H
C
C
C
C
H
:Cl
..:
:Cl
. .:
H
..
:Cl :
These are NOT
resonance structures !
The bond between metals and non-metals is usually ionic. Metals give away
their e- and become positively charged (cations). Nonmetals accept them
and become anions. The ionic bond is formed as a result of attraction
between oppositely charged ions. The compound is called ionic compound,
and the three-dimensional ordered network of the ions is called ionic lattice.
Electronegativity and the
Polar Covalent Bond
The difference in EN defines the bond.
∆EN = 0, covalent; ∆EN = 1.0, polar
covalent (23% ionic); ∆EN = 1.9, polar
covalent (60% ionic); ∆EN > 1.9, ionic.
Electrons are rarely
shared equally
between atoms.
Electronegativity (EN) is numerical rating
of an atoms ability to attract to itself the
shared electrons in a covalent bond.
Generally, electronegativity of metals is low,
and that of nonmetals is high.
The least electronegative atom (except
H!) is the central atom in dot structures.
Polar covalent bond is a covalent bond in
which e- are shared unequally (large ∆EN).
A partial negative charge (δ-)
occurs on the more EN atom.
A partial positive charge (δ+)
occurs on the less EN atom.
Ionic and covalent are two extremes at
the ends of a continuum bonding types.
The Shape of the Molecules
Valence Shell Electron Pair Repulsion (VSEPR) theory is the
model mostly used to predict molecular shape.
Electron pairs on the central atom repel one another.
The two dimensional dot structure of methane, CH4. gives the angles
between electron pairs of 90o. But the dot structure angles are arbitrary.
Molecules are three dimensional, and the electron pairs would be further
away if the third dimension is considered. In fact, the shape of methane
molecule is tetrahedral; the bond angles between electron pairs is 109.5o.
H
|
H–C–H
|
H
Four electron pairs around an atom assume tetrahedral arrangement.
When there are not enough electrons for single bonds the molecule
forms multiple bonds and the structure differs. VSEPR theory treats
each multiple bond as a single electron group, because it occupies
roughly the same region of space. The number of electron groups
around an atom is called the atom’s steric number (SN).
Dot structures of formaldehyde
and acetylene are arbitrarily
shown with angles of 90o.
Their true geometry has
bond angles of 120o and
180o, respectively.
:O:
||
H–C–H
H H
|
|
CΞC
formaldehyde
acetylene
Dot structures
:O:
||
C
H
H–CΞC–H
H
True geometry
If the central atom has a
lone pair of electrons, that
electron pair is included
in the molecular shape.
The dot diagram of
ammonia presents the
atom in a plane.
The steric number on
nitrogen is 4 (3 bonding
pairs and a lone pair).
The e- pairs on the N
assume tetrahedral
arrangement.
..
H–N–H
|
Electrons in the lone
H
pair occupy more
VSEPR arrangements of electron groups around an atom
having no lone pair electrons
O=C=O
We describe the shape
of the molecule, not that
of its electrons. Lone
pair(s) of electrons are
therefore ignored.
The molecule of NH3
has a pyramidal shape.
space than the
bonding pairs. They
squeeze H-N-H bond
angle to approx. 107o.
Rule of thumb: each lone pair of e- on a
period-2 atom compresses the remaining
bond angles around that atom by ~2o.
Using VSEPR
1. Draw a dot diagram.
2. Count the number of epairs around the central
atom, including lone pairs
(i.e. the steric number, SN).
A multiple bond counts as a
single e- group.
3. Find the best
arrangement of the
electrons using SN.
4. Pretending the lone epairs are invisible,
describe the resulting
shape of the molecule.
Practice on H2O and O3.
H-O-H
bond angle
~105o.
bent
H
|
: O:
O-O-O
bond angle
120o.
SN=4
bent
H
O
||
O:
O
SN=3
Chapter 13
Water and Properties of Liquids
Liquids have intermediate properties between solids
and gases. Liquids are almost incompressible, have
definite volume and assume the shape of the container.
Densities of
liquids are
usually lower
than that of their
solids. Water is
an exception.
Evaporation or vaporization is the escape of
molecules from liquid into gaseous state.
During evaporation, liquid that stays behind is
cooler. The opposite process is condensation.
Sublimation is the escape of molecules directly
from solid into gas, bypassing liquid state.
Vapor pressure is the pressure exerted by a gas at
evaporation
equilibrium with its liquid, so that:
liquid
gas
Vapor pressure depends only on
condensation
temperature, not on the amount of
liquid.
Open container
completely evaporates.
Closed container
reaches equilibrium
between liquid and gas.
Vapor Pressure Measurement
1 atm = 760 torr
20 oC
20
oC
a.
b.
a. The system is evacuated.
Manometer attached to the
flask shows equal pressure
in both legs.
b. Water is added.
Liquid evaporates.
Manometer shows
increase in pressure.
20 oC
30 oC
c.
d.
c. Equilibrium established.
Manometer shows constant
pressure difference, 17.5 torr.
d. Temperature raised to 30 oC.
Equilibrium reestablished.
Manometer shows constant
pressure difference of 31.8 torr.
Vapor pressure
and temperature
1 atm = 760 torr
Vapor pressure of any
gas at the boiling point
is equal to the
atmospheric pressure.
Vapor pressure of
ethyl ether is the
highest at any temp.
TBP
TBP
Vapor pressure:
Ether > Alc. > Water.
Rate of evaporation:
Ether > Alc. > Water.
proportional to vapor
pressure.
TBP
Volatility
Boiling point:
Ether < Alc. < Water
Substances that readily evaporate are volatile.
Vapor pressure of ethyl ether at 20 oC: 442.2 torr
Volatile
Vapor pressure of water at 20 oC: 17.5 torr
Vapor pressure of mercury at 20 oC: 0.0012 torr
Moderately volatile
Nonvolatile
Boiling Point Curves
Normal Boiling Point
Boiling point at standard pressure
(1 atm, or 760 torr).
Each point on the curve represents a
vapor-liquid equilibrium at a
particular temperature and pressure.
At 500 torr, ethyl ether boils at
~22 oC, alcohol at ~68 oC, and
water at 89 oC.
Freezing or Melting Point
The temperature at which the solid
and liquid are in equilibrium.
Changes of State
Majority of substances change phases
upon heating: solid liquid gas.
1 atmosphere
pressure
TBP ethyl ether
TBP alcohol
TBP water
34.6oC
78.4oC
100.0oC
Heating curve for a pure
substance
CO2 is an exception (dry ice sublimes).
A – B: solid state
B – C: melting
C – D: liquid state D – E: evaporation
E – F: vapor state
Temperature is constant during melting
and boiling – all heat used to break
solid (at boiling point) or liquid forces.
liquid
solid
evaporation
condensation
melting
freezing
gas
liquid
Heat of Fusion and Heat of Vaporization
We learned before that amount of heat
Qheating = (mass) (spec.heat) (temp.change)
depends on mass and temp. change.
Energy (heat) needed to change 1 g
Energy (heat) needed to change 1 g
of a solid at its melting point into Constant
of a liquid at its boiling point into
liquid is heat of fusion.
temperature! vapor is heat of vaporization.
Qfusion = (mass) (spec.heat of fusion)
Qvaporization = (mass) (spec.heat of vaporization)
Example 1: How many joules is
needed to change 20.0 g of ice at
0 oC to steam at 100. oC?
Qheating = (mass) (spec.heat) (temp.change)
Qtot = Qfusion + Qheating + Qvaporization
Qfusion = (20.0 g) x (335 J/g)
Qheating = (20.0 g) x (4.184 J/goC) x (100. oC)
Qvaporization = (20.0 g) x (2260 J/g)
Hydrogen Bond
}
Qtot = 60.3 kJ
produces unusually high melting & boiling point
Hydrogen Bonding (cont.)
H bonding exists between H directly bonded
to one of the three most electronegative
elements (Fluorine, Oxygen, and Nitrogen),
and F, O or N of another molecule.
. . H bond . .
H – O :. . . H – O :
|
|
H
H
H
bonded
to O
No H bond
H
H
| .. |
H–C–O–C-H
| .. |
H
H
Ethyl ether
Surface Tension and Capillary Action
A droplet of liquid
falling forms a
sphere due to
attractions to other
liquid molecules –
surface tension.
Cohesive forces within
Spontaneous rise mercury liquid (left) are
of liquid in a
stronger than adhesive
forces between Hg and
narrow tube –
walls of the container.
capillary action.
Opposite is true for H O.
2
No H bonded
to F, O, or N
H bonds are
intermolecular
forces.
Hydrates
Some ionic solutions retain water upon evaporation. It becomes
the part of the crystalline compound – water of crystallization.
The formula is written as: ionic compound, dot , # water molecules…
CuSO4 5 H2O and name them by adding # (Latin) hydrate.
.
Copper(II) sulfate pentahydrate.
Hydrates are true compounds and the water is an integral part of it.
Formula mass CuSO4 5 H2O: 63.55+32.07+64.00+5x18.02 = 249.7
Percent composition of water is (5x18.02 / 249.7) x 100 = 36.08%
dry CuSO4 – white
Hydrate = blue
.
Water can be removed by intense heat: CuSO4
The reaction is reversed when water is added.
Water, a Unique Liquid
Water indicator
. 5 H O(s) CuSO (s) + 5 H O(g)
2
4
2
δO
H
H
Water covers ~75% of Earth. 97% of water is in the oceans. Only
3% is fresh water, of which 2/3 is locked up in ice polar caps.
δ+
Solid form (ice) has lower density than liquid water.
Water is very stable molecule, can stand temperatures up to 2000 oC. It does not
conduct electricity when pure, but decomposes into H2 and O2 in solutions of ions.
2 H2 + O2 --> 2 H2O + 484 kJ
Water can be formed by
2 C2H2(g) + 5 O2 4 CO2 + 2 H2O(l) + 1212 kJ
Combustion,
Neutralization,
HCl(aq) + NaOH(aq) --> NaCl(aq) + 2 H2O
Metabolic reaction
C6H12O6(aq) + 6 O2 6 CO2(g) + 6 H2O(l) + 2519 kJ
Water reactions with metals:
Cold water reacts with Na, K, Ca:
Steam reacts with Zn, Al and Fe:
Reactions of water
Na + H2O H2 + NaOH
Fe + H2O(g) --> H2 + Fe3O4
Remind yourself of the activity series: the above six metals are the most active. Another
three metals are more active than H: Pb, Sn, and Ni and react with acids only; Cu, Ag,
Hg and Au are below H in the series and do not react with acids or H2O.
Water also reacts with certain nonmetals.
Anhydride means:
without water.
Most reactive:
2 F2 + 2 H2O(l) --> 4 HF(aq) + O2
Less reactive:
Cl2(g) + H2O(l) HCl(aq) + HOCl(aq) To test whether a metal or
nonmetal is an anhydride, try
Least reactive: C(s) + H2O(g) CO(g) + H2(g)
Water reacts with oxides of metals and non-metals:
Basic anhydride: CaO + H2O(l) Ca(OH)2(aq)
Acid anhydride:
CO2(g) + H2O(l) H2CO3(l)
Water Purification
Screening, flocculation and sedimentation,
sand filtration, aeration, disinfection.
Hard water contains Mg2+ and Ca2+ ions
Additional water purification is done by
distillation, Ca2+, Mg2+ precipitation, ion
exchange and demineralization.
to remove H2O until all
hydrogen is removed.
Ca
OH ∆
OH
CaO + H2O
H2SO4 ∆ SO3 + H2O
HW, Chapter 11 (p.241) : 3, 5, 15, 31, 41
3. Which one in each pair has the larger
radius? Ca atom or Ca2+ ion; Cl atom
or Cl- ion; Mg2+ ion or Al3+ ion; Na
atom or Si atom; K+ ion or Br - ion.
Explain.
5. Using the table of electronegativity
values (Table 11. 5, p.227) indicate
which element is more positive and
which is more negative in these
compounds:
H2O; RbCl; NH3; PbS; PF3; CH4.
15. How many electrons must be gained
or lost for the following to achieve a
noble gas electron configuration?
K atom; Al ion; Br atom; Se atom.
31. Draw Lewis structures for the
following: NCl3; H2CO3; C2H6; NaNO3.
41. Use VSEPR theory to predict the
shape of these molecules:
SiH4; PH3; SeF2.
.
HW, Chapter 13 (p.302): 5, 11 13, 21
5. In which of the following substances
would you expect to find hydrogne
bonding? C3H7OH; H2O2; CHCl3; PH3;
HF.
11.Name these hydrates: BaBr2 · 2 H2O;
AlCl3 · 6 H2O; FePO4 · 4 H2O.
13.How many moles of compound are in
25.0 g of Na2CO3 · 10 H2O?
21. How many joules of energy are
needed to change 275 g of H2O from
15 oC to steam at 100 oC?
.