Transcript Bonding An Introduction to Chemical Reactions

Bonding
An Introduction to Chemical
Reactions
Pg. 156-215
Chemical Bonds
• Properties of many materials can be
understood in terms of their microscopic
properties:
– connectivity between atoms,
– three dimensional shape of the molecule.
• When atoms are strongly attracted to one
another = chemical bond
• What causes this attraction between atoms?
Electrostatics
• Electrostatics- attraction and repulsion
determines bonding between atoms and
forces of attraction that can exist between
molecules.
Coulomb’s Law
q1 & q2 are charges on particles 1 & 2
D is the distance between the particles
ke = constant
• As distance between charges increases, the
electrostatic force __________.
• As the charge on the particles increases, the
electrostatic force __________.
Questions to Consider
• Where do these charges exist in an atom?
• How does the organization of the atom’s
electrons affect this electrostatic force?
• Make a connection between reactivity of
atoms in the periodic table and the
organization of electrons using this concept
of electrostatic force.
Bond Types
• 1. Ionic Bonds___________________________________
_____ These oppositely charged ions are
attracted to each other through electrostatic
forces.
• 2. Covalent Bonds• 3. Metallic Bonds-
Metallic Bonds
“Positive ions in a sea of mobile electrons.”
Delocalized Valence Electrons
Metallic Bonds
• Form between two or more metals
• Atoms of metals achieve stability by sharing their
valence electrons. Delocalized valance electrons.
• Metallic bonds are the attractive forces between
fixed positive ions and the moving valence
electrons of the metal.
Composition of Selected Alloys
Stainless Steel
Coinage Silver
74% Fe, 18% Cr, 8% Ni,
0.18% c
90% Ag, 10% Cu
Plumber’s Soder
67% Pb, 33% Sn
Brass
67% Cu, 33% Zn
18 Carat Gold
75% Au, 10-20% Ag, 515% Cu
60% Ni, 40% Cr
Nichrome
Ionic Bonds
• Static electricity and the clothes dryer
• Static electricity is the basis for ionic bonds.
• Octet Rule dictates that some substances gain
electrons- __________, while others lose
electrons- ___________.
• Positive and negative ions are attracted to one
another.
Ionic Bonds
Characteristics of Substances
with Ionic Bonds
1.
2.
3.
4.
5.
Composed of _______
Have ________ melting
points
Solids at room temperature,
many soluble in water
________________________
________________________
Tend to be ______________
Covalent Bonds
• Formed by a shared pair of electrons
between two atoms.
• Molecule =
Glycine- AA
Types of Formulas
• Molecular formulaindicates the number of
atoms that are in a single
molecule of a compound.
C6H12O6
• Empirical formulaindicates the lowest whole
number ratio of atoms in a
molecule. CH2O
• Structural Formulasspecifies which atoms are
bonded to each other in a
molecule
Structural Formulas- Lewis
Structures
• Valence electrons are indicated around the
symbol for the element
Oxygen has 6 valence
electrons
Nitrogen has 5
valence electrons
Drawing Lewis Structures
• Imagine each side (top, bottom, left, right)
of the symbol of the element can hold 2
electrons for a total of 8 electrons.
• Each side will hold one electron first, then
will double up.
• In covalent bonding the number of single
electron sides (unpaired electrons)
indicates the number of covalent bonds the
atom must have to satisfy its octet.
• Oxygen has 6 valence
electrons.
• Two unpaired
electrons means that
oxygen must form two
bonds to satisfy its
octet.
• Draw the Lewis
structure for the
following:
– Chlorine
– Phosphorus
– Carbon
Lewis Structures
• Atoms share electrons
to fill their octets.
• A solid line indicates a
shared pair of
electrons.
• Dots are used to
indicate unshared pairs
of electrons.
Formation of a single covalent bond
Double and Triple Bonds
• A unique characteristic of covalent compounds
is their ability to form multiple bonds between
two atoms.
• Refer back to the Lewis Structures for nitrogen
and oxygen.
• Nitrogen needs to share three electrons
• Oxygen needs to share two electrons.
Technique for Drawing
Lewis Structures
1. Determine the number of valence electrons in
each atom making up the molecule
2. Add the valence electrons and divide by two
3. Draw the “skeleton.” If carbon is present, place
it at the center of the molecule.
4. Distribute the pairs of electrons around the
skeleton to satisfy each atoms octet. (Remember:
Hydrogen only needs two electrons to fill its
octet.)
Practice
• Draw Lewis Structures for the following
compounds:
–
–
–
–
Ammonia
Ethyne- C2H2
Carbon Dioxide
HCN
Exceptions to the Octet Rule
• Atoms with more than an octet
– SF4
• Molecules with an odd number of electrons
– NO
– Generally short lived, unstable molecules
Properties of Molecular
Compounds
• Composed of 2 or more ___________
• ___________ electrons in bond formation
• Can be solids, liquids, or gases at room
temperature.
• Some are soluble in water, others are not.
• Tend to be ___________________
conductive.
• Generally have _________ melting points.
Questions to Consider for
Lewis Structures
• What does it mean to “share” electrons in the
formation of a bond.
• In your experience, is “sharing” always equal?
• Pick a bond in your Lewis structure and decide if
the sharing of electrons is equal or unequal. Why
is it so?
• How might this “sharing” affect the physical and
chemical characteristics of the molecule?
Covalent Bonds- Are the Atoms
Really “Sharing” Electrons?
Chlorine
Hydrogen
Covalent Bond Types
• Polar Covalent Bondselectrons in bond are
________________.
• Nonpolar Covalent
Bonds- electrons in
bond
__________________
_________________.
Polar Covalent Bonds
When a bond is classified as polar covalent (H-O), the
atom with the higher electronegativity has the greater
attraction for the shared electrons
·As a result, a charge unbalance is produced in the
molecule + by H and – by O
Dipole = charge unbalance
d+ H –O d• The “positive” and “negative” ends of the dipole are
not real charges (such as positive and negative ions)
because no electrons have actually been transferred
between the atoms. The dipole represents only an
unbalanced charge distribution along the bond.
Nonpolar Covalent Molecules
BrINCl HOF Elements
• Diatomics- elements that can combine with
themselves in a nonpolar covalent molecule
to form a stable compound.
• Memorize!
Electronegativity
Bond Type by Electronegativity
Electronegativity
Difference
<= 0.4
Between 0.4 and 2.0
>= 2.0
Bond Type
Water’s polarity allows it to pull
at the ions in an ionic crystal.
Metallic Vs Ionic Bonding
• Much easier to deform materials with metallic than
with ionic bonding. Why?
NaCl (s)
Ag (s)
• Sliding atom planes over each other (deformation) very
unfavorable energetically in ionic solids!
•  metals are ductile & ceramics (ionic) are brittle
Intermolecular Vs Intramolecular Forces
Intermolecular forces are forces ___________ molecules.
Arises from interaction between dipoles. Bond Polarity
Intramolecular forces _________________________________
Intermolecular vs Intramolecular
•
41 kJ to vaporize 1 mole of water (inter)
•
930 kJ to break all O-H bonds in 1 mole of water (intra)
Generally, intermolecular forces are
much weaker than intramolecular
forces.
Types of Intermolecular Forces
1. Dipole-Dipole Forces
2. Hydrogen Bonding Forces
3. London Dispersion Forces
Tend to be less than 15% as strong as covalent or
ionic bonds.
“Measure” of intermolecular force
boiling point
DHvap
melting point
DHfus
DHsub
Intermolecular Forces
1. Dipole-Dipole Forces:
solid
liquid
Intermolecular Forces
2. Hydrogen Bond: a special dipole-dipole interaction between
the hydrogen atom in a polar N-H, O-H, or F-H bond and an
electronegative O, N, or F atom.
A
H…B
or
A
A & B are N, O, or F
H…A
Intermolecular Forces
3. London Dispersion Forces:
London Dispersion Forces among
nonpolar molecules
separated
Cl2
molecules
instantaneous
dipoles
Chemical Reactions
• A process in which one or more substances
are converted into new substances with
different physical and chemical properties.
• Reactant- a substance that enters into a
chemical reaction.
• Product- a substance that is produced by a
chemical reaction.
The Reason for Reactions
• During a chemical reaction, new substances
are produced as existing bonds are broken,
atoms are rearranged, and new bonds are
formed.
• Substances undergo chemical reactions with
other substances _____________________
Chemical Equations
• Describes what happens in a chemical reactionsimilar to mathematic equations.
• Word Equations- give the names of the reactants
and the products.
Calcium + oxygen yields calcium oxide
• Formula Equations-chemical symbols replace the
names of the reactants and products.
Ca + O2  CaO
Law of Conservation of Mass and
Balancing Chemical Equations
• Matter is neither created nor destroyed
during a chemical reaction. Therefore, all
the atoms that were present at the start of
the reaction must be present at the end of
the reaction.
Balanced?
Ca + O2  CaO
• Coefficients are used in chemical equations
to balance an equation.
• Subscripts cannot be changed once the
compound is written. Changing the
subscript would change the compound!
Ca + O2  CaO
A coefficient of 2 is placed in front of calcium
and calcium oxide to balance the equation.
2Ca + O2  2CaO
Steps to Balance Chemical
Equations
1. Write the formula equation with the correct
symbols and formulas.
Na + Cl2  NaCl
2. Count the number of atoms of each element on
each side of the arrow.
3. Balance atoms by using coefficients.
2Na + Cl2  2NaCl
4. Check your work by counting atoms of each
element.
Edible Equations
• 1. Gather several thin pretzel sticks and a package
of M&Ms.
• 2. Use the pretzels and M&Ms to make models of
the following chemical reactions:
2KClO3  2KCl + 3O2
U+ 3F2  UF6
Cd + HCl  CdCl2 + H2
Cs2 + O2  CO2 + SO2
• 3. How do your models illustrate the Law of
Conservation of Matter?
Practice
• Sodium phosphate is used to cut grease.
Write a balanced equation for the reaction in
which iron (II) chloride reacts with sodium
phosphate to produce sodium chloride and
iron (II) phosphate.
• Chlorine reacts with lithium bromide to
produce lithium chloride and bromine.
Classifying Chemical Reactions
• Types:
1. Direction Combination Reactions
(Synthesis)- two or more reactants come
together to form a single product
A + B  AB
2 Na (s) + Cl2 (g)  2 NaCl (s)
4Fe (s) + 3O2 (g)  2 Fe2O3 (s)
Direct Combination (Synthesis)
2. Decomposition Reactions (Analysis)
• A reaction in which a single compound is
broken down into two or more smaller
compounds or elements.
AB  A + B
2H2O (l)  2H2 (g) + O2 (g)
Decomposition Reactions
(Analysis)
• 3. Single Replacement Reaction (REDOX)an uncombined element displaces an
element that is part of a compound.
A + BX  AX + B
BX and AX are generally ionic compounds and A
and B are elements.
Mg (s) + CuSO4 (aq)  MgSO4 (aq) + Cu (s)
Fe (s) + CuSO4 (aq)  FeSO4 (aq) + Cu (s)
Single Replacement Reaction
(REDOX)
Single Replacement Reactions
• A more active element will replace a less
active element.
• Table J- Activity Series
• Substances higher in the table will replace
substances lower in the table
Pb(s) + CuSO4 (aq) 
Single Replacement Reactions
•
•
•
•
Mg (s) + 2HCl (aq) 
2Al (s) + 3ZnCl2 (aq) 
Al (s) + NaCl (aq) 
NaCl (aq) + H2(g) 
4. Double Replacement Reactions• Atoms or ions from two different
compounds replace each other. An
identifying characteristic of a double
replacement reaction is the presence of two
compounds as reactants and two compounds
as products. “Switch Partners”
AX + BY  AY + BX
CaCO3 + 2HCl  CaCl2 + H2CO3
Double Replacement Reactions
• Double replacement reactions do not occur unless:
• The reactants are dissolved in water so that the
compounds can separate into ions.
• And one of the following:
1.
2.
3.
• Table F, which shows the solubilities of various
ionic substances in water,can be used to help us to
determine if a precipitate is formed.
Predict if the following Double Replacement
reactions will occur and indicate why the
reaction does or does not occur.
AgNO3 + NaCl  AgCl + NaNO3
KOH + Al(NO3)3  KNO3 + Al(OH)3
NaOH + HCl  NaCl + H2O
KBr + NaNO3  KNO3 + NaBr