Download Lecture 24 (Slides) October 18

Reducing Abilities of Group 1 and 2
Metals
2 K(s) + 2 H2O(l) → 2 K+ + 2 OH- + H2(g)
I1 = 419 kJ
I1 = 590 kJ
I2 = 1145 kJ
Ca(s) + 2 H2O(l) → Ca2+ + 2 OH- + H2(g)
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General Chemistry: Chapter 9
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Oxidizing Abilities of the Halogen
Elements
(Group 17)
2 Na + Cl2 → 2 NaCl
Cl2 + 2 I- → 2 Cl- + I2
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General Chemistry: Chapter 9
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Acid-Base Nature of Element Oxides
Basic oxides or base anhydrides:
Li2O(s) + H2O(l) → 2 Li+(aq) + 2 OH-(aq)
Acidic oxides or acid anhydrides:
SO2 (g) + H2O(l) → H2SO3(aq)
Na2O and MgO yield basic solutions
Cl2O, SO2 and P4O10 yield acidic solutions
SiO2 dissolves in strong base, acidic oxide.
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General Chemistry: Chapter 9
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Class Examples:
• 1. Which of the following atoms and ions are
paramagnetic (i.e. have unpaired electrons).
Note: An even number of electrons does not
indicate that all electrons are paired. (a) He
atom, (b) F atom, (c) As atom, (d) F- ion (e)
Al3+ ion and (f) Fe atom.
• 2. Arrange the following in order of increasing
atomic radius: (a) Mg, Ba, Be, Sr (b) Rb+, Se2-,
Br- and Sr2+ (c] Ca, Rb, F, S (d) Fe, Fe3+, Fe2+.
Class Examples – cont’d:
• 3. Write balanced chemical equations to
represent the reactions of the following oxides
with water: (a) SO3(g), (b) P4O10(s), (c) BaO(s)
and (d) Li2O(s).
• 4. Arrange the following atoms in order of
increasing first ionization energy: (a) Fr, He,
K, Br (b) P, As, N, Sb and (c) Sr, F, Si, Cl.
• 5. Why are transition metal atoms and ions so
often paramagnetic?
Class Examples
• 6. What information does the term “degenerate
orbitals” convey?
• 7. How do a ground state and an excited state
electron configuration differ?
• 8. How many electrons are described using the
notation 4p6? How many orbitals does this
notation include? What is the shape of the
orbitals described using the 4p6 notation?
Class Examples
• 9. Are all (neutral) atoms having an odd atomic
number paramagnetic? Are all atoms having an
even atomic number necessarily diamagnetic?
Explain.
• 10. The following electron configurations do
not correspond to the ground electronic state of
any atom. Why? (a) 1s22s22p64s1, (b)
1s22s22p63s23p63d2 (c)
1s22s22p63s23p63d84s24p2.
Chemical Bonding
• We’ll begin with a few familiar ideas from
high school with particular reference to Main
Group Elements.
• Observation: Noble Gases (He, Ne, Ar, Kr..)
are stable in monatomic form in all phases
(gas, liquid and solid). Other Main Group
elements/atoms are not similarly stable in
monatomic form. These atoms are held
together by covalent and/or ionic bonds.
Chemical Reactions/Bond Formations
are Associated with Energy Changes
• Energy is released when two or more isolated
atoms form chemical bonds. The simplest
examples involve the formation of
homonuclear diatomic molecules. For
example:
• H(g) + H(g) → H2(g)
ΔH0 = -436 kJ
• N(g) + N(g) → N2(g)
ΔH0 = -946kJ
• Cl(g) + Cl(g) → Cl2(g)
ΔH0 = -242kJ
Chemical Reactions/Bond Formations
are Associated with Energy Changes
• The energy released when chemical bonds are
formed accounts for the stability of chemical
compounds. We could reverse the
thermochemical equations given on the
previous slide and see, for example:
• H2(g) → H(g) + H(g)
ΔH0 = +436 kJ
• We see that 436kJ of energy must be supplied
to break the bonds in one mole of H2
molecules.
Ionic and Covalent Bonds
• When Main Group elements react, electrons
can be transferred (usually from a metal to a
nonmetal) to form ionic bonds. In other cases,
pairs of electrons can be shared (usually
between nonmetal atoms) to form covalent
bonds. In both cases valence electrons are
somehow “rearranged” when new chemical
bonds are formed. Bond “strengths” vary
widely from one compound (or element) to the
next.
Ionic Bonds – Main Group Elements
• In the simplest cases ionic bonds are formed
when a metal and a nonmetal react.
• Metal +
Nonmetal → Binary Ionic
↑
↑
• Metal atoms have low ionization energies, low
electron affinities and low electronegativity.
• Nonmetal atoms have high ionization energies,
high electron affinities and high
electronegativity.
Ionic Bonds – Main Group Elements
• Metal atom loses one or more electrons to
form a positive ion. The positive ion has a
noble gas electron configuration.
• Nonmetal atom gains one or more electrons to
form a negative ion. The negative ion usually
has a noble gas electron configuration.
• In the simplest cases we have two monatomic
ions. Complications such as polyatomic ions
will arise frequently.
Formation of LiF from Li and F Atoms
• Li (1s22s1) → Li+ (1s2) + e• F ( 1s22s22p5) + e- → F- (1s22s22p6)
• In this example the Li+ and F- ions have,
respectively, the electron configurations of He
and Ne. Ionic compounds formed at room
temperature are usually solids. Why? The
compound formed must be electrically neutral.
Why?
Formation of LiF from Li and F Atoms
• In practice, lithium fluoride would be formed
from the reaction of lithium metal and fluorine
gas (F2 molecules). The corresponding
balanced chemical equation is:
• 2 Li(s) + F2 (g) → 2 LiF(s)
• Ionic compounds are solids at anything close
to room temperature. The oppositely charged
ions are held together (bonded) by very strong
electrostatic forces.
Main Group Binary Ionic Compounds
Compound
Cation
Anion
Chemical
Formula
Potassium iodide
K+
I-
KI
Calcium selenide
Ca2+
Se2-
CaSe
Lithium oxide
Li+
O2-
Li2O
Aluminum oxide
Al3+
O2-
Al2O3
Magnesium
nitride
Mg2+
N3-
Mg3N2
Francium astatide
Fr+
At-
FrAt
Covalent Chemical Bonds
• The electronic structures of atoms and
molecules have many features in common.
Individual atoms often possess unpaired
electrons. These atoms are usually chemically
unstable. Two such atoms can come together
to form a molecule with no unpaired electrons.
This process can involve the formation of
covalent chemical bonds and is highly
exothermic.
Covalent Chemical Bonds
• In isolated atoms most of the electrons will be
found in pairs in a number of different atomic
orbitals. When atoms combine to form
molecules, valence shell electrons are
rearranged. Two electrons from different
atoms can “pair up” to form a single covalent
bond where the bonding molecular orbital is
associated with more than one atom.
Covalent and Ionic Bonds
• Before moving on to Lewis structures lets
practice (review) looking at chemical formulas
and identifying whether covalent bonds, ionic
bonds or both are important for the compounds
under consideration.
Covalent and Ionic Bonds
Substance (phase)
Chemical Formula
Chlorine (gas)
Cl2(g)
Potassium chloride (solid)
KCl(s)
Iron (III) hydroxide (solid)
Fe(OH)3(s)
Hydrogen fluoride (gas)
HF(g)
Sucrose (solid)
C12H22O11(s)
Ammonium nitrate (solid)
NH4NO3(s)
Ethanol (liquid)
CH3CH2OH(l)
Cobalt (III) fluoride (solid)
CoF3(s)
Copper (II) sulfate
pentahydrate (solid)
CuSO4∙5H2O(s)
Ionic and/or Covalent
Bonds?
Covalent Chemical Bonds
• We will consider molecular orbitals in more
detail shortly. Before doing so we will use
Lewis electron dot structures to represent
bonding in both covalent and ionic
compounds. In these structures valence shell
electrons are represented by dots. The rest of
the atom – the nucleus and the core (nonvalence) electrons are represented using the
chemical symbol for the element.
Lewis Theory: An Overview
1. Valence e- play a
fundamental role in
chemical bonding.
2. e- transfer leads to
ionic bonds.
3. Sharing of e- leads to a
covalent bond.
4. e- are transferred or
shared to give each atom
a noble gas configuration,
the octet.
Gilbert Newton Lewis (1875-1946)
General Chemistry: Chapter 10
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Lewis Symbols and Lewis Structures
• A chemical symbol represents the nucleus
and the core e-.
• Dots around the symbol represent valence e.
•
• Si •
•
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••
I •
••
General Chemistry: Chapter 10
••
Ar
••
••
••
• Se
•
•
••
Bi
•
•
•
••
Sb
•
•
•
••
• Al •
•
••
As
•
•
•
••
••
P
• •
•
••
N
• •
•
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Lewis structures
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General Chemistry: Chapter 10
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Covalent Bonding: An Introduction
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General Chemistry: Chapter 10
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Coordinate Covalent Bonds
+
H
Cl
H
H
N
H
•• Cl
••
••
H
••
N
••
H
H
H
FIGURE 10-2
•Formation of the ammonium ion, NH4+
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General Chemistry: Chapter 10
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Multiple Covalent Bonds
• In most cases Main Group elements are
surrounded by eight valence shell electrons
when molecules are formed – octet rule (H is
an impt exception!). Such molecules can
contain single bonds (one pair of bonding
electrons), double bonds (two pairs of bonding
electrons and triple bonds as well as lone pairs
of electrons.
Lewis Structures
• Class Examples: Construct Lewis structures
for H2, F2, O2, N2, CO and CN-. In which
species are multiple bonds required to satisfy
the octet rule? What type(s) of experimental
evidence would suggest that the Lewis
structures for these species have physical
meaning?