a. the basics
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Chapter 1: Chemical Bonding. 1.1 Forming and Representing Compounds. A. The Basics. scientists have studied the way in order to. elements and compounds appear in nature. - PowerPoint PPT PresentationTRANSCRIPT
A. The Basics scientists have studied the way
in order to
elements and compounds appear in nature categorize
most are combined with in nature (called ) and are
1.1 Forming and Representing Compounds
Chapter 1: Chemical Bonding
chemical bonding
metals non-metalsores solids
a few are found in theirmetals pure form…precious metals
metals (except ) in pure form are
Hgsolids
combine with one another to form
the only elements that are found in in nature are the
non-metals
solids, liquids or gases
nevercombined form noble gases
atoms in such a way that they create a
the can only hold and therefore satisfies the octet rule when it has
gain, lose or share electrons full outer energy
level… octet rule
first energy level
two e
two e in it
octet rule is a guideline …not all elements follow it at all times
called the
are the electrons in the
they are the only electrons involved in
for representative elements ( ) group number (ignore the “1” in front of groups above 10) tells you the number of
period number tells you the number of
valence e outermost
chemical bonding
groups 1,2 and 13-18 valence electrons
energy levels occupied by electrons
energy level of an atom
for many , the number of valence electrons
transition metals is not as predictable… it depends on the environment
eg)iron can be Fe3+ or Fe2+
the can be used to determine ion charge number of valence e
eg)Fe3+ had 3 valence electrons
around the ion
B. Electron Dot Diagrams you can’t see atoms and electrons,
therefore it is convenient to to show the structure and formation of
an is one such model
consists of the with
chemical bonds
electron dot diagram
symbol for the element dots
when drawing the diagrams, look up the , then place number of valence e dots around the symbol clockwise
draw models
representing thevalence e
for a maximum offour dots
if you have more electrons to place, go back to the start pairing up the e
Na
Al
Ca
Si
P
O
Cl
Ar
topof the symbol and
a orbital is called a and is (at this level)
a orbital contains a
the of an atom is the maximum number of that it can form (equals the number of )
full lone pairnot involved in bonding
half full bonding electron
bonding capacitysingle covalent bonds
bonding e
O
lone pairs
bonding e
Try These
H
He
Be
C
F
Draw the Lewis diagram (electron dot diagram) for each of the following:
S
Mg
K
Br
P
C. Ionic Bonding an is the
most have
they tend to these electrons and become
electrostatic attraction between oppositely charged ions
losepositive ions (cations)
ionic bond
metals three or fewer valence e
Na Na+
Crash course Lewis dot diagram
most have non-metals more than four valence e
they tend to and become
gain electrons negative ions (anions)
O2-
O
after ions form, the attraction between the
when drawing the electron dot diagrams for ionic compounds:
ions
the number of electrons by the must the number of electrons by the
the on the compound must be
you may have to have of the to balance out the
positive chargeandnegative chargedraws thetogether, forming anionic bond
lost metalequal gainednon-metal
net chargezero
more than onemetal and/or non-metal
charges
Examples
NaCl
MgO
Na Cl
OMg
Examples
CaF2
K2S
K
KS
F
F
Ca
O
O
O
Mg
Mg
Fe
Fe
N
MgN
Fe2O3
Mg3N2
notice the following about the diagrams:
the metal has (since they them)
the non-metal has the valence level
both ions have and the
charges = charges
no valence electronslose
filled
square bracketscharge
positive negative
D. Covalent Bonding a is formed when
compounds containing are called
two non-metals share a pair of electrons
covalent bond
covalent bonds molecular compounds
electron dot diagrams used to show molecular compounds are called Lewis
structures
ions are not formed!!!
Covalent Bond video
instead of transferring electrons, valence electrons are now to satisfy the
the electrons that are shared are called
sharing two or three pairs of electrons between two atoms results in a
shared
a bonding pair
double or triple bond,
octet rule
respectively
to draw the structures:
place the atom with the in the
arrange all other atoms around it as as possible
to make sure that all atoms have the (remember that hydrogen only needs electrons to be satisfied)
most bonding electronscentre
symmetrically
share electronsoctet rule satisfied
two
Covalent bonding video
P
eg) PH3
H
H
H
P
HH
H
Try These
1. HCl
2. CH4
3. F2
4. NBr3
Draw the Lewis diagram (electron dot diagram) for each of the following:
6. N2
5. C2H4
H Cl
F F
N Br BrBr
N N
C HHH
HC
H HC
H H
a is another way of drawing molecules
to draw them, figure out the Lewis structure then replace all
E. Structural Formulas
eg) PH3
Lewis Diagram Structural Diagram
P HH
H
P HH
H
structural formula (diagram)
shared pairs of e with a line and leave off the lone pairs
Try These
1. HCl
2. CH4
3. F2
4. NBr3
Draw the structural formula for each of the following:
6. N2
5. C2H4
H Cl
C HH
H
H
F F
N Br Br
Br
CH H
CH H
N N
most metals are at room temperature which means that there must be
metals form or with
in all the atoms
F. Metallic Bonding
solids
strong attractive forces
metallic bonding share all the valence e
the valence electrons are , which means they are from one atom to another
holding the atoms of a pure metal together
DO NOT covalent ionic bondsother metal atoms
delocalizedfree to move
metallic bonds are made up of a network of positive metal
ions a is the
this theory helps explain the of metals
eg) good conductors of electricity and heat, ductility, malleability
in a “sea” of electrons
metallic bond electrostatic force of attraction between the positive metal ions and the negative sea of electrons
properties
metal cations
“sea” of delocalized electrons
Metallic Bond Model
the of an element is the relative measure of the ability of an atom to
there is an attraction between the of an atom and the
A. Electronegativity electronegativity
valence e in an adjacent atom
1.2 The Nature of Chemical Bonds
attract electrons in a chemical bond
nucleus (protons)
nucleuselectrons
each element is designated a number to represent how strong it’s nucleus is at attracting another atom’s valence e
trend on periodic table – electronegativity anddecreases down group
electronegativity means
since do not readily react with other substances, electronegativities have been assigned to them
higher greater attraction(affinity)
increases across period
noble gasesnot
understanding electronegativity has contributed to the knowledge of bonding in ionic and molecular compounds
decreases
increases
Electronegativity and the Periodic Table
as you move from left to right across a period, both the and
B. Size & Electronegativity
electronegativity, however size of the atom
atomic number increasedecreases
here is why size decreases across a period:
the size of an atom depends on the of the containing the
in any given period, the valence e of each atom occupy the
as you move across the period, the and thus the in the nucleus
there is a between the and when there are more , therefore the atom is
radius energy level valence e
same energy level
atomic number increases number of protons increases
greater amount of attractionnucleus e
protonssmaller
3 p+
1 valence e
Li
6 p+
4 valence e
C
9 p+
7 valence e
F
Period 2 Elements
so, the next question is “why does electronegativity increase when atomic size across a period decreases?”
the strength of the attraction (and therefore electronegativity) between oppositely charged particles depends on two factors:
the between the charges – the attractive force between opposite charges with the between them
the of the charges – the attractive force is to the
distance
decreases square of the distance
magnitudedirectly proportional
amount of charge
this means that an atom that is and has lots of (like fluorine) will have a amount of electrostatic attraction (electronegativity) for the of another atom
big atoms have but they are by the therefore have a amount of attraction (electronegativity) for the of another atom
smallprotons
very largee
lots of protonsshielded inner levels of e
smalle
Cs SiF Si
nucleus of cesium
valence electrons of silicon
nucleus of fluorine
valence electrons of silicon
distance between nucleus of cesium and valence electrons of silicon
distance between nucleus of fluorine and valence electrons of silicon
electronegativities can be related to bond types:
ionic bonds occur between
C. Bond Type & Electronegativity
metals have electronegativities and will while non-metals have and will
the two ions that are formed will
metals and non-metals
lowlose e high
gain e
attract each other
electronegativities
and form achemical bond
covalent bonds occurs between
if you look at two atoms that have the electronegativity, like in H2(g), the two nuclei of the atoms will attract the
the electrons are
non-metallic atoms
same
electronswith exactly thesame strength
shared equally between the two atoms
when two non-metals that have electronegativities share electrons, the sharing is
the element with the electronegativity pulls the
different
no longer equal
higher
e closer to itself
this results in one end of the bond having a and the other end of the bond having a
+
slightly negative charge()slightly positive charge
(+)
bonds that have are called
unequal sharing of electronspolar covalent bonds
also called since the bonds have
bond dipoles
oppositely charged ends
H – F
“arrow” points towards element with higher electronegativity (-)
“+” at the end that is +
Bond Dipole Arrows
+ -
Try These:
Draw the bond dipole arrow, label the + and ends, and state the bond type (polar, nonpolar, ionic)
1. H – H
2. N – H
3. B – F
4. S – O
5. P – H
6. C – H
7. Cl – Cl
8. Si – Cl
9. O – H
10. Na – Cl
nonpolar polar
polar
polar
polar
nonpolar
nonpolar
polar
polar
ionic
+
+
+
+
+
+
-
-
-
-
-
-
0.8
2.0
0.8
0.4
1.3
1.2
you can use the difference in electronegativity between two atoms to determine
bond character
mostly ionic polar covalent slightly polar covalent
non-polar covalent
3.3 1.7 0.5 0
Difference in Electronegativity
as you can see,
bonding is considered a and there is
continuum
bond classification is not simple
no clear distinctionbetween ionic and covalent bonding
ionic compounds have
they form so that are as as possible this is called a
crystal structure
3-D array of alternating positive and negative ions crystal lattice
A. Ionic Crystals
2.1 Three Dimensional StructuresChapter 2: Chemical Bonding
oppositely charged ionsclose together
since all the in the lattice are the
attractive forces same, you cannot call them
each positive ion is attracted to of the negative ions around it
the chemical formula is the lowest whole number ratio
molecules…all
(and vice versa)
for that type of crystal
sodium chloride
eg) NaCl has a 1:1 ratio of Na ions to Cl ions
shape also depends on the relative size of the ions and the charges on the ions
there are many different and they all depend on the way the ions
crystal shapespack together
B. Structure of Molecules in molecular compounds, covalent bonds exist between
these compounds exist as that have a and therefore are not necessarily written with the
specific pairs of atoms
molecules given number of atoms lowest whole number ratio
C. VSEPR & Structure of Molecular Compounds
electron pair repulsion is
the states that molecules adjust their shapes so that valence e-
shape is determined around the
valence shell electron pair repulsion (VSEPR) theory are as far away from each other as possible
not always equal… it is greatest between twoless between a and aand lowest between two
central atom
lone pairs (LP),LP bonding pair (BP),
BP’s
VSEPR
shapes can be classified into five categories:
1. linear – central atom is bonded to and has lone pairs, or there is only two atoms in the molecule eg) CO2(g), HCN(g), HCl(g)
C OO
C NH
ClH
two other atoms zero
2. trigonal planar – central atom is bonded to and has lone pairs eg) CH2O(l)
O
HH
C
3. tetrahedral – central atom is bonded to and has lone pairs eg) CH4(g)
H
H H
H
C
three other atoms zero
four other atomszero
4. pyramidal – central atom is bonded to and has lone pair
eg) NH3(g)
5. bent – central atom is bonded to and has either lone pairs eg) H2O(l), HNO(g)
H
H
HN
HH
O
three other atomsone
OH
N
two other atoms one or two
the code has two numbers:
we can use a to determine the shape of a molecule around the
code
1. the number of attached to the central atom 2. the number of on the central atom
central atom
atoms
lone pairs
eg) NH3(g)
H
H
HN
CH4
H
H H
H
C
3 - 1
pyramidal
4 - 0
tetrahedral
Code Shape Example
4 – 0
3 – 0
3 – 1
2 – 1
2 – 2
tetrahedral
trigonal planar
pyramidal
bent
bent
CH4
CH2O
NH3
HNO
H2O
***all other codes arelinear
VSEPR shapes
D. Polar Bonds & Polar Molecules a molecule that contains
can be overall
the individual are that can be to each other
if the bond dipoles are and , they each other out resulting in a
this canceling happens in symmetrical molecules
if the bond dipoles , the entire molecule will have a
polar covalent bondsnonpolar
bond dipoles vectorsadded
equal in strengthopposite in direction cancel
nonpolar molecule
do not cancel slightly positive and slightly negative end…called dipoles
general rules: tetrahedral: if all atoms
attached have the same pull (in or out), if different atoms attached
trigonal planar: if all atoms attached have the same pull (in or out), if different atoms attached
pyramidal: as long as there is a difference in electronegativity between the atoms
bent:
linear: …look at electronegativity difference
nonpolarpolar
nonpolarpolar
polar
polar
polar or nonpolar
Examples
1. H2O 2. HCl
4. C2HI3. C2H2
O
H H
H Cl
HH C C IH C Cnp np
polarpolar
nonpolarpolar
Try These
1. HF 2. CH4
4. PI3
H F
polar
polar
nonpolar
3. N2
5. C2H6
N Nnp
PI
II
CH
HH
H
nonpolar
nonpolar
HH C Cnp
H
H H
H
A. Types of Forces
responsible for state, melting point, boiling point etc.
are the forces of attraction intermolecular forces between molecules
are the forces of attraction intramolecular forces within molecules
(eg. ionic or covalent bonding)
2.2 Intermolecular Forces
they are theweakest of all forces
there are three types of intermolecular forces that we will look at:
1. Dipole-Dipole Forces electrostatic force of attraction between the
attract the in other molecules and vice versa
dipoles of polar molecules
+
+
+
+
+-
- -
-
-
slightly negative “poles” slightly positive “poles”
2. Hydrogen Bonding this is a special type of dipole-dipole
interaction that is
hydrogen bonding is the attraction between a
when hydrogen is bonded to a such as O, F or N, the electrons are pulled from it
very strong
hydrogen on one molecule which is bonded to O, F or N, to the O, F or N of an adjacent molecule
highly electronegative element
far away
since hydrogen doesn’t have any other its is basically exposed
this proton is then able to be attracted not only to the
electrons,
pole
OH
H
OH H
O
H
H
O
HH
proton
but also to the lone pairs
3. London (Dispersion) Forces the attractive force that occurs between
molecules is called
the result of the electrostatic attraction of
electrons in atoms and molecules are always in
for brief instances, the distribution of electrons becomes which produces very weak
induced dipoles
constant rapid motion
distorted
allLondon Dispersion force
dipoles
this induces a dipole in the when
this process throughout the substance, causing that
even though this force lasts and is , the overall effect in a substance is
temporary dipole like charges repel each other
“disperses”flickering dipoles
attract each other
only a momentvery weak
significant
adjacent molecule
LD forces are affected by two factors:
1.size of the atoms – means higher probability of creating
2.shape of the molecule –
the more between molecules, the the force of attraction
more e
temporary dipoles
contacthigher
Youtube videos
Intramolecular vs intermolecular
http://www.youtube.com/watch?v=S8QsLUO_tgQ&feature=related
Scale of Forces
very low
very high
Intermolecular Forces
(between)
Intramolecular Forces
(within) London Dispersion Dipole – Dipole Hydrogen Bonding
metallic ** wide range ionic
network covalent
eg) diamond, SiC, SiO2
covalent
LDDD
HB ionic covalent
networkcovalent
A. States of Matter – Read p. 72 – 74!!! the of a substance depends on the
between its particles
2.3 Relating Structures and Properties
solids have the forces of attraction, liquids have the and gases have intermolecular attractions between the particles
state strength of the attractive forces
greatestnext greatest
very few if any
B. Melting & Boiling Points both melting and boiling points are indicators
as to the
when melt or boil, must be broken
strength of attractions within or between molecules
metals metallic bonds
when melt or boil, must be broken
ionic compounds ionic bonds
when melt or boil, must be broken
these forces are much therefore it takes a lot less
as you the number of intermolecular forces, the melting and boiling points
molecular compoundsonly intermolecular forces
(covalent bonds DO NOT break)
weakerenergy to separate the molecules
increaseincrease
Order of bp’s
Using the scale of forces you can order compounds based on their relative bp’s
Ex. From Highest to Lowest Network covalent compound (ex. SiO2) Ionic compound Molecular compound with HB, DD, LD Molecular compound with DD, LD Molecular compound with LD (if 2 molecular
compounds have LD only then bigger molecule or molecule with more electrons has higher bp)
C. Mechanical Properties of Solids the mechanical properties of solids are
determined by the types of bonds in the substance delocalized e cause bonds to be non-directional, which allows a solid to be malleable and ductile
solids that do not have delocalized e have directional bonds, which causes them to be brittle and hard
eg) metals
eg) ionic compounds
D. Conductivity electric current is the directional flow of
electrons or ions
metals are good conductors of electricity because the delocalized valence e are free to move
solid ionic compounds have valence electrons that are held solidly in place therefore they cannot conduct electricity
when ionic compounds melt or dissolve in water, the ions are able to move past one another which allows them to carry an electric current
in most molecular compounds, valence electrons are not free to move through the molecule therefore they are not able to conduct electricity
when molecular compounds melt or dissolve in water, they do not form ions and therefore they do not carry an electric current