Transcript Ch 6.1 Chemical Bonding

Some observations from the final
exam…
Review from Chapter 4 & 5
• Can atoms of two different elements have the
same ground state electron configuration?
• NO!!!!
• The number of protons is unique for
each element, therefore an atom of that
element has a unique number of
electrons (atoms are neutral, p+ = e-)
• The electron configurations are different
b/c different atoms have different
number of electrons.
Review from Chapter 4 & 5
• Write the electron configuration of the
following:
•
•
•
•
Argon:
Calcium:
Iron (Fe):
Bromine (Br):
1s22s22p63s23p6
1s22s22p63s23p64s2
1s22s22p63s23p64s23d6
1s22s22p63s23p64s23d104p5
Core electrons [Ar]
i.e., noble gas core
Outer electrons – outside
of the noble gas core
Ions
• Ions are atoms or groups of bonded
atoms that have a positive or negative
charge
• Ions are formed by losing or gaining
electrons!
• an atom gains electron(s) and becomes a
negative ion (anion)
• an atom loses electron(s) and becomes a
positive ion (cation)
Review from Chapter 4 & 5
Ionization energy decreases (takes less
energy to remove an electron)
• Metals vs. non-metals
Electron affinity increases
Review from Chapter 4 & 5
• Let’s look at electron configurations of
ions…
1s22s22p63s23p6
• Argon:
• Argon is a noble gas. It’s valence shell
has a stable octet (8 electrons). It will
not form an ion because it already has
an extremely stable electron
configuration.
Review from Chapter 4 & 5
• What kind of ion would the following atoms
form? Write the electron configuration of
the ion.
• Argon:
1s22s22p63s23p6
• Calcium:
1s22s22p63s23p64s2 Ca2+
• Calcium will be more stable if it can
achieve an electron configuration
similar to Ar. To do so, it will form a
cation, Ca2+, by getting rid of its two
valence electrons
Review from Chapter 4 & 5
• What kind of ion would the following atoms
form? Write the electron configuration of the ion.
22s22p63s23p64s23d104p6
1s
• Krypton:
22s22p63s23p64s23d104p56 Br1s
• Bromine (Br):
• Bromine has a high electron affinity. It is
very close to having an octet in its valence
shell. It will gain an electron to fill its
valence shell and become an anion, Br-, and
its electron configuration is the same as Kr.
Review from Chapter 4 & 5
• Transition metals and ions…
22s22p63s23p64s23d6
1s
• Iron (Fe):
• Transition metals have 2 valence electrons.
Many transition metals can form cations
with different charges. When transition
metals form ions, they will lose electrons
from highest s sublevel first, then they may
lose electrons from the d sublevel as well.
We will talk more about transition metals
later…
Chapter 6 - Chemical Bonding
Section 6.1 – Introduction to
Chemical Bonding
Combinations
• Few elements exist
as independent
particles, most
substances are
made up of
combinations of
atoms held together
by chemical bonds.
Chemical Bond
A mutual electrical
attraction between
the nuclei and
valence electrons of
different atoms that
binds the atoms
together.
Why Atoms Bond to Other Atoms
Most atoms are less stable existing by
themselves (they are at a relatively high
potential energy). Nature favors
arrangements in which potential energy is
minimized. Bonding creates more stable
arrangements of matter in lower potential
energy states. (Mama Nature doesn’t like to
sweat!).
Bonding
• Valence electrons
are redistributed.
Three main types:
Covalent
Ionic
Metallic
Not male bonding!
Types of Chemical Bonding
Ionic
vs.
Covalent
• Electrical attraction between
cations and anions.
• Sharing of electron
• Due to gain or loss of
pairs between atoms
electrons (also called
• H2
electron transfer) between
atoms.
• H2O
Unique Covalent Bonding of C and Si
• Can form 4 covalent
bonds. Can also bond
with itself to form
long-chain molecules
of different sizes and
shapes.
Metallic Bonding (not in text – yet)
• The attraction
between metal atoms
and the surrounding
sea of electrons.
Bonding usually falls between ionic and covalent.
• This depends on how strongly the atoms of each
element attract electrons. So we look at the
electronegativity difference (atom x - atom y;
F – Cs = 4.0 – 0.7 = 3.3).
• 0……..0.3…………….1.7………………3.3
• nonpolar polar
ionic
covalent covalent
• ex. O2
H2O
NaCl
• (POLAR = uneven sharing of electrons.)
Non-polar Covalent
• A covalent bond in which the bonding
electrons are shared equally by the bonded
atoms, resulting in a balanced distribution
of electrical charge.
Example: Two hydrogen atoms combine to form H2 gas.
Polar Covalent Bonds
δ+
δ−
Example: Hydrogen chloride
HCL
δ (Greek lower case delta)
means “partial”
• A covalent bond in
which the bonded
atoms have an
unequal attraction
for the shared
electrons. Such
bonds are polar, in
that they have an
unequal distribution
of charge.
Rules for determining which type
• Electronegativity
differences, or percentages
of ionic character
determine the bond type:
• (Electronegativity values
for elements are on p.151)
difference
%
Ionic
> 1.7
50
Polar C. 1.7-0.3
50-5
Non-polar <0.3
<5
covalent
Assignments – Due Wednesday
BOP
• 6.1 Worksheet
• 6.1 Review Problems