Loading presentation...

Present Remotely

Send the link below via email or IM

Copy

Present to your audience

Start remote presentation

  • Invited audience members will follow you as you navigate and present
  • People invited to a presentation do not need a Prezi account
  • This link expires 10 minutes after you close the presentation
  • A maximum of 30 users can follow your presentation
  • Learn more about this feature in our knowledge base article

Do you really want to delete this prezi?

Neither you, nor the coeditors you shared it with will be able to recover it again.

DeleteCancel

Make your likes visible on Facebook?

Connect your Facebook account to Prezi and let your likes appear on your timeline.
You can change this under Settings & Account at any time.

No, thanks

Chemical Bonding

No description
by

Madison D.

on 3 February 2015

Comments (0)

Please log in to add your comment.

Report abuse

Transcript of Chemical Bonding

Chemical Bonding
Electronegativity
The higher the electronegativity of an atom, the greater the attraction for bonding electrons.

is a property of an atom which increases tendency to attract the electrons of a bond.
Trends; It decreases as the atomic number increases, because of the increased distance between valence electrons and nucleus (greater radius)
Why is this important for the formation of bonds?
Think of it as a tug of war. The atoms are the people and the rope is electrons. Which ever person is stronger wins because they had enough strength to pull more of the rope to their side then the opponent.
This is the same for electronegativity, which ever atom is more electronegative ends up having more of a pull, resulting in a higher density of electrons at the end of a reaction.
what happens if the people are equally strong?

If the atoms are equally electronegative, both have the same tendency to attract the bonding pair of electrons, resulting in the electrons to be half way between each atom.
increases
Decreases
electrons
Ionization energy
is the minimum amount of energy required to remove an electron from a gaseous atom or ion
An element with a low ionization energy is more likely to lose electrons and form cations.
An element with a high ionization energy is more likely to steal an electron and form anions.
Ionization energy generally goes off of atomic radius, the further away an electron is from the nucleus, the easier it is for it to be taken. The larger the radius, the smaller the amount of energy is required to take the electron from the outer most orbital.
So, if you think of our tug of war analogy and put low ionization energy against high electronegative who would win the electron?
low ionization
high electronegative
High electronegativity implies that an element strongly wants to take electrons in a reaction.
If ionization energy is low it implies that it really wants to give up a electrons to a reaction
vs.
High electronegative wins!

Ionic bonds
Important key ideas
Ionic bonds are formed when a metal and a non-metal exchange electrons
Involves electrostatic forces between oppositely charged ions
Ions that have lost one or more electrons are called cations (metals)
Ions that have gained one or more electrons are called anions (non-metals)
Covalent Bonds
Ways in which atoms form molecules
In this example the sodium atom is giving one of its valence electrons to the chlorine atom. The sodium becomes a cation and the becomes chlorine the anion.
Important key ideas
Covalent bonds are formed by the sharing of electrons between atoms
Can occur between atoms of the same element, elements close to each other on the periodic table, and between metals and non-metals
Primarily occurs between non-metals
In this example Phosphorous is sharing three of its unpaired valence electrons with the Chlorine atoms. In the end all of the molecules are satisfied with eight electrons.
You can also predict the types of bonds based on their electronegativites!
Calculation example: (use periodic table on slide 2)
0 = non-polar covalent ex)N2
0.4 - 1.7 = polar covalent ex) H2O
>1.7 = ionic ex) MgS
AlCl3
3.0 - 1.5 = 1.5
therefore polar covalent !
The force of attraction in a covalent bond is between a pair of electrons and two adjacent positive nuclei, rather than between a cation and an anion as in an ionic bond
Metallic bonds
The electromagnetic interaction between delocalized electrons known as conduction electrons
Electrons are passed from one atom to another
Conduction electrons gather in an "electron sea" and the metal nuclei within metals
In this example Sodium atoms have come together and their one valence electron shares space with eight other corresponding electrons on the other Sodium atoms. The electrons can move freely within the "electron sea" and each electron becomes detached from its parent atom. The metal is held together by the force of attraction between the nuclei and delocalized electrons.
Covalent Compounds
Soft, tend to be gases, liquids or soft solids
Poor conductors of heat and electricity
Brittle
Flexible
Properties
Most covalent compounds only have a few atoms and the forces between these molecules are weak therefore most covalent compounds have low melting, boiling points and low enthalpy of vaporization.
Examples of covalent compounds

H2O - water
HCl - hydrogen chloride
CH4 - methane
CO2 - carbon dioxide
Ionic compounds
Properties
Hard
Conduct electricity when they are dissolved in water
Brittle
Ionic solids are good insulators
Ionic compounds have high melting and high boiling points. Recall that ionic bonds are formed between negative and positive ions therefore their force of attraction is high. To break these positive and negative charges apart, it takes a large amount of energy. Therefore, a lot of energy is required to melt ionic compounds or cause them to boil.
Examples of Ionic compounds

NaCl - sodium chloride
CaCl2 - calcium chloride
K2O - potassium oxide
MgO - magnesium oxide
Metallic compounds
Properties
Hard
Dense
Malleable
Ductile
Conduct heat and electricity
Metallic compounds have high melting and boiling points because of the strong forces of attraction between the delocalized electrons and the nuclei of the atoms, because of this strong force it requires a lot of energy to melt or boil.

Examples of Metallic compounds

Ni3Al - Nickel aluminide
FeS - Ferric Sulphate
How to physically visualize your molecules
....
The introduction on how to draw
Lewis Dot Structures
Lewis Dot Structures

Valence electrons are the electrons in the outer most orbitals. A quick and easy way to determine them are by using the periodic table. On the periodic table the vertical columns (Groups) are numbered 1-18. Group 1 has 1 ve- , group 2 has 2ve-, group 13 has 3ve- etc.. (except for He 2 ve-)

The transitional elements
(Groups 3-12)

Group 3: 3 ve-
Group 4: 2 to 4 ve-
Group 5: 2 to 5 ve-
Group 6: 2 to 6 ve-
Group 7: 2 to 7 ve-
Group 8: 2 or 3 ve-
Group 9: 2 or 3 ve-
Group 10: 2 or 3 ve-
Group 11: 1 or 2 ve-
Group 12: 2 ve-
Step 1: Find total number of valence electrons
Example: NCl3

N is in group 15 therefore 5ve-
Cl is in group 17 therefore 7ve- but because there is 3 atoms multiply by three 7x(3)

Total number of valence electrons = 26
Continued..
For atoms to be "happy" it must follow the octet rule, meaning it needs 8 valence electrons around it to be satisfied.
There are some exceptions that we will see later.
Step 2:
Make atoms "happy"
Step 3: Determine the number of bonds in the molecule

To determine this you will subtract your answer from step 2 by step 1 and divide that answer by 2
Example: NCl3
There is 4 atoms and each atom needs 8 valence electrons therefore 4 x 8 = 32
Example: (26ve-) - 32 = 6 6/2 = 3
Total bonds = 3

Step 4: Choose a central atom
The central atom of a molecule is
usually
the least electronegative atom.
In this case nitrogen will be our central atom because there is 3 atoms of chlorine
Step 6: Draw a "skeleton" structure
N
Cl
Cl
Cl
Step 7: Place valence electrons around atoms and place remaining electrons around
the central atom (depending on how "full" the central atom is (octet rule))
Complete the octets around each of the outer atoms. If there are not enough electrons to complete the octets, the skeletal structure from step 5 is most likely incorrect. Try a different arrangement.
Complete the octet for the central atom with the remaining electrons. If there are any bonds left over from Step 3, create double bonds with lone pairs on outside atoms.
Your complete Lewis dot structure should look like this.
Bonds account for 2 valence electrons.
You can check by counting valence electrons and making sure that answer is the same as in step 1
Exceptions for the Octet Rule
Molecules with an odd number of electrons
The best way to approach this is by trial and error
Ex) NO2
Just draw
Molecules that contain double bonds next to single bonds (Resonance Structures)
Ex) CO3
This ion has a total of: 4 + (3x6) +2 =
24 valence electrons
(Note that the extra 2 electrons is due to its charge)
The two electrons in the double bond have been shared equally between the central atom of carbon and each of the three oxygen atoms
Continued ..
Molecules with more than an Octet
To expand Octets, elements that are the central atom must be
in or greater than the third period or have an atomic number greater than 10
Ex) PCl5
How to predict polarity based on their molecular shape
VSEPR Theory
Valence Shell Electron Pair Repulsion
Electron Pair Geometry
: The arrangement in 3D space of all electron pairs surrounding the central atom
Molecular Geometry
: The arrangement in 3D of all ligands surrounding the central atom. Molecular geometry is based off of electron pair geometry. Multiple bonds are treated as single bonds while determining molecular geometry
Consider this example of Hydrogen Fluoride, which is a covalent compound with a single bond. Recall our tug of war analogy of electronegativity and since Fluorine has the "bigger pull" the electron density is shifted more towards Fluorine because it's more electronegative. This shift of electrons is indicated by an arrow.
The + sign at the beginning of the arrow simply means that the H end of the molecule becomes partially positive.
Shapes of molecules
Based on ligands and electrons
General Formula
Central atom - "A"
Ligands - "B"
Lone Pairs - "E"
Ex) NH3's general formula is
AB3E
Non-dipole moment
Ex) CH4
This is a non dipole moment because each dipole cancels each other.
Non-polar molecules
do not
have dipole moments and polar molecules
do
have dipole moments. Thus the dipole moment is a tool to predict the polarity of a molecule and determine its shape.
We can also use our tug of war analogy in this situation. If two people who were each equally strong pulling on either ends of the rope who would win? Nobody. Its almost as if their strengths cancel each other out. Somewhat the same situation here in our non-dipole moment. Each dipole has canceled each other out resulting in a non-dipole moment.
If the molecule is symmetrical it's non-polar
If it is asymmetrical it's polar
Asymmetrical therefore polar
Symmetrical therefore non-polar
Intermolecular Forces
Intramolecular forces (bonding forces) exist within molecules and influence the chemical properties. Intermolecular forces exist between molecules and influence the physical properties.
Types of intermolecular forces
Dispersion forces
Attractive forces that arise as a result of temporary dipoles induced atoms or molecules. These forces are present in
all
molecules, whether they are polar or nonpolar. The tendency of an electron cloud to distort in this way is called polarizability.
Polarizability
is the ease with which the electron distribution in the atom or molecule can be distorted. Molecules with less energy and larger radii have higher polarizability. Higher polarizability means stronger attraction.
The positive end of an atom is attracted to the negative end of the other and vice-versa.
Dipole-Dipole forces
Hydrogen Bonding
Polar molecules that contain hydrogen atoms bonded directly to fluorine, oxygen, or nitrogen.

Exists between an ion and partial charge at one end of a polar molecule
Ion-Dipole Force
Intermolecular forces affecting properties
How properties of liquids are affected by intermolecular forces
Surface tension; amount of energy required to stretch or increase surface of a liquid
strong intermolecular forces means high surface tension
results from the net inward force experienced by the molecules on the surface of a liquid
Cohesion:
the intermolecular bonding of a substance where its attractiveness forces them to hold a certain shape of the liquid. (between like molecules)
Capillary Action: the ability of liquids to rise in tubes through adhesive forces
Adhesive
: When forces of attraction between unlike molecules occur. (between unlike molecules
Water has formed a concave meniscus
Mercury has formed a convex meniscus
The different shapes are determined by the adhesive forces of the molecules of the liquids with the walls of the tub and the cohesive forces between molecules of water.
continued..
is a measure of how well substances flow

Viscosity:
Stronger intermolecular forces means higher viscosity
It is related to the ease with which molecules can move past each other
Melting Point:
Stronger intermolecular forces means higher melting point

Vapor pressure:
temperature at which a solid changes into a liquid
Boiling Point:
temperature at which the substance's vapour pressure equals the surrounding air pressure
Stronger intermolecular forces means higher boiling point
is the equilibrium pressure of a vapor above its liquid
Stronger intermolecular forces means lower vapor pressure
•In this example the molecules are escaping the liquid and becoming a vapor, the pressure is increasing and some of the gas molecules will strike the liquid molecules and condense back into it. Once the rate of condensation of the gas equals the rate of evaporation of the liquid an equilibrium state has been reached.
Practice questions
Question: Identify the compound that has the highest boiling point explain in terms of intermolecular forces. He and Ar
Things to think about...
1. Which intermolecular forces are in play here ? Keep in mind which forces are the strongest
2. Which compound has more electrons (larger) more polarizable
Strongest to least: Ion-ion>hydrogen bonding>dipole-dipole>dispersion
Answer:
Both Helium and Argon only have dispersion forces but since Argon is larger and therefore has more electrons, it is more polarizable than Helium. Therefore Argon experiences more dispersion forces and a greater amount of energy would be required to break the stronger force of attraction. In a final coonclusion the boiling point of Argon is higher.
Question: Explain why ethane C2H6 melts at -183 degrees Celsius and nonane C9H20 melts at -54 degrees Celsius
Answer:
C2H6 and C9H20 are both non-polar and only have dispersion forces. C9H20 has a greater molar mass because it has more electrons, making it more polarizable than C2H6. Therefore, C9H20 has greater dispersion forces. Because of its stronger forces it then requires more energy to melt then C2H6 does.
Kinetics
Collision Theory
reaction rate depends on number of collisions per unit time.
the fraction of these collisions that succeed in producing products.
Factors Affecting Rates of Reaction
By increasing the concentration you are increasing the amount of particles in a given space and therefore collide more often. As concentration, the rate of reaction also increases.
An increased temperature will increase energy which will lead to particles striking one another more frequently and more forcefully. As temperature increases, the rate of reaction will increase
By increasing the partial pressure of a gas it increases the number of collisions and the opportunity for particles to collide. However it has no effect on reactions involving reactants in liquid or solid phases.


Pressure:
Temperature:
Concentration:
Potential Energy Diagrams
Activation energy:
The minimum kinetic energy that the reacting species must have in order to react.

Enthalpy:
Potential energy that must be given off or absorbed as heat

Activated Complex:
An intermediate state that is formed during the conversion of reactants into products

Catalyst
: A substance that increases the rates of chemical reactions without being used up
Graphical representation of the energy changes that take place during a chemical reaction
The geometric orientation of particles must also be favourable for a successful collision
Exothermic: a reaction that releases heat
Endothermic: a reaction that absorbs energy
These are basic potential energy diagrams.
The exothermic combination reaction's profile is simply the reverse of that for the endothermic decomposition reaction
Thermochemical and H notation equations
Thermochemical equations include the change of energy in the equation
A + B
AB
+ 100kj/mol
rxn
H notation requires the change of energy to be written separate form the equation
A + B
AB
H = - 100kj/mol
rxn
Activation energy and Catalysis
A Catalyst provides an additional reaction mechanism with a lower activation energy, which results in an increased reaction rate. The catalyst is not consumed during the reaction.
You can think of a catalyst as a short cut when you're walking home from school. Its faster and takes less energy to take the short cut than it would to take your normal route
Short cut (Catalyst)
Activation Energy w/no catalyst
Average rate reaction:
change in a measurable quantity of a chemical species
change in time
How to measure the rate of a reaction
Rate laws look like this:
(k) is the rate constant for the reaction
Note that the exponents have nothing to do with the coefficients in the balanced equation. They must be determined experimentally.

How to mathematically calculate exponents
Steps
1. Choose two rates in which something is constant
2. In this case divide rate 2 over rate 1
3. Cancel what you have the same of on the top and bottom
4. Divide your rate 2 by 1 and do the same with initial rate
5. Divide your answer from initial rate by your answer from rates

How to mathematically calculate the constant (k)
Steps
1. Divide initial rate of any of the trails by the the molarity of F2 and ClO2

Exothermic: Synthesis reaction example
Energy released
Endothermic: Decomposition reaction example
100kj/mol + AB
A + B
AB
A + B
H = +100kj/mol
Energy absorbed
This is a potential energy exothermic diagram. You can see the difference with the activation energy when a catalyst has been added to the reaction. The activation complex has dropped a sufficient amount.
BE SURE TO WATCH OUT FOR YOUR UNITS
CH3Br + OH- → CH3OH + Br-

Rate = k[CH3Br][OH-]

A rate that express how the rate depends on concentration is a
rate law
An expresses how the concentration depends on time
Types of rate laws:
integrated rate law
Integrated Rate Law
First order:
Integrated first-order rate law:
ln[A] = -kt + ln[A]
0
Important things to note
The concentration of A is depend on time, if the initial concentration and the rate constant (k) is known. The concentration of A can be calculated.
The equation above is in the form of y = mx + b. Where a plot of y versus x is a straight line. Thus, plotting the natural logarithm of concentration versus time ALWAYS gives a straight line. This fact can be used to determine if the order is first order.
y = ln[A] x = t m = -k b = ln[A]
0
Second order:
Integrate second-order rate law:
Important things to note
A plot of 1/[A] versus (t) will produce a straight line with a slope equal to the rate constant (k).
The equation above shows how [A] depends on time and can be used to calculate [A] at any time t. Provided that the values of k and [A] are known.
0
Positive slope
Negative slope
Zero order:
Integrated zero-order rate law: [A] = -kt + [A]
0
Half-Life
The time required for a reactant to reach half its original concentration
First-order half-life:
You can calculate t1/2 if k is known and k if t1/2 is known.

Note that half-life is not depended on concentration.
Second-order half-life:
Zero-order half-life:
t1/2 is dependent on [A]
SUMMARY
There are several methods for converting qualitative comparisons to actual numerical values of electronegativities of the elements. One widely used scale is
Linus Pauling's scale:

Dipole Moments
Polar or non-polar?
Crystalline solids
Molecular solids:
A solid with regular arrangement of its components
When a molecular compound crystallizes, the intermolecular forces of the compound hold the crystal together. Melting the compound disrupts these intermolecular forces, but does not disturb the covalent bonds (intrmolecular forces) within each molecule.
A solid made up of molecules and held together by intermolecular forces. Discrete, relatively small molecules at the lattice points
Ionic solids:
A solid containing cations and anions that dissolve in water to give a solution containing the separated ions, which are mobile and free to conduct electric currents. Have ions at lattice points
When solid sodium chloride is placed in water and dissolves the ions are distributed throughout. The resulting solution has free ions and thus conduct electric currents.
Atomic solids:
A solid that contains atoms at its lattice points
Atomic solid: Diamond
Ionic solid: Sodium chlorine
Molecular solid: Ice
2 electrons for 1 bond
Cohesive > Adhesive forms water drops

Cohesive < Adhesive forms film

Cohesive = Adhesive forms a pool
2-
BrF3, SF4, XeF5+, ICl3, XeOF4, XeF4, SbF5
F has exceeded the Octet and it is
not
an element that can exceed the Octet Rule (not in or greater than third period)
Other examples that expand octet:
TIP: Draw the lewis dot structures
Full transcript