Download chemical bond

Ch 6
Sec 6.1
Atoms seldom exist as independent particles in nature. Almost
everything is a combination of atoms that are held together by a
chemical bond.
A chemical bond is a mutually electrical attraction between the
nuclei and valence e- of different atoms.
3 types of chemical bonds:
1.
2.
3.
Ionic
Covalent
Metallic
Ionic - electrical attraction between large #’s of cations and anions.
Transfer of e- from a metal to a non-metal.
Covalent – the two atoms share the e- .
Normally between non-metal and non-metals.
Metallic – the chemical bonding that results from the attraction
between metal atoms and the surrounding sea of e- .
(Show overheads)
Electronegativity chart – pg 151
Difference in Electronegativity = Polarity
(overhead)
Non-polar bond – the e- is shared equally.
Polar bond – the e- is shared unequally. One atom attracts the estronger. We say the atom that attracts the e- more is the more
negative atom.
Ex. Classify the bonds of each of the following elements with Sulfur.
Which atom is more negative?
Bonding between
S and:
Electronegativity
Difference
Bond Type
Na
H
Cs
Cl
S
For most general purposes:
metal + non-metal = ionic bond
non-metal + non-metal = covalent bond
More neg. atom
ex.
Compound
H2
NaCl
H2O
Electronegativity
Difference
Bond Type
More Negative
atom
Sec. 6-2
A molecule is a neutral group of atoms that are held together by
covalent bonds (sharing). i.e. must be a non-metal and a non-metal
Molecular compound – a chemical compound whose simplest units
are molecules.
ex. Water, sugar, etc.
Molecular formula – shows the types and # of atoms combined in a
single molecule of a molecular compound.
ex. H2O has 2 H and 1 O
Diatomic molecule – can form a molecule when bonded with itself.
It must contain 2 atoms.
H, N, O, F, Cl, Br, I can form diatomic molecules.
(They form a 7 and there are 7 of them)
By sharing e- in overlapping orbital's, each H atom in the H2 molecule
experiences the effect of a stable 1s2 electron configuration.
(show orbital’s)
Ex. F2
Ex. HF
Electron dot notation – electron configuration that only uses
valence e- and uses dots to represent e- .
# of valence e-
Electron dot notation
examples
Write the electron dot notation for :
1.
P
2.
Si
3.
S
4.
Cl
5.
Xe
Lewis Structures – electron dot notation can also be used to
represent molecules.
Ex. H2
Ex. F2
Ex. Cl2
Structural Formula – just shows the bond, not the unshared e- .
ex. H2O
Draw the Lewis Structures and Structural Formula for:
Molecular
Formula
NH3
H2S
CH3I
Lewis Structure
Structural
Formula
Double and Triple bonds – needed when there is not enough
valence electrons to go around.
Molecular Form
Structural Formula
C2H4
CH2O
CO2
N2
O3
O3 show resonance – which shows multiple Lewis dot structures.
Sec 6-3
Most of the rocks and minerals that make up the Earth’s crust consists
of (+) and (-) ions held together by ionic bonding.
An ionic compound transfers electron so that the compound becomes
neutral. Metal + non-metal = ionic bonding
(Overhead)
Ionic bond will shatter when hammered. As a layer is shifted toward
the next, like charges are beside each other causing them to repel
and shatter.
Bond energy – energy needed to break 1 bond between 2 atoms.
** Overall, ionic bonds are stronger than covalent bonds.
Polyatomic Ions – a charged group of covalently bonded atoms.
Know this list!!
Sec 6-4
Metallic bonding – the chemical bonding that results from the
attraction between metal atoms and the surrounding sea of e- .
(Overhead)
1.
The luster of metals is caused by the emission of photos when the
excited electron returns to ground state.
2.
This sea of electrons allows for the conductivity of heat and
electricity.
3.
Metals are malleable and ductile because the metallic bonding is the
same in all directions throughout the solid. One plane of atoms can
slide past another with out resistance.
Metallic
1. Type of elements
involved.
2. What happens to
the electrons?
3. Conductivity
4. Physical State at
room temperature.
5. Bond Strength
6. Other
Characteristics.
Ionic
Covalent
Sec 6-5
Board notes