Transcript Types of Changes in Matter

2.1 Science and Technology in Society
2.2 Changes in Matter
2.3 Balancing Chemical Reactions
2.4 Chemical Amount
2.5 Classifying Chemical Reactions
2.6 Chemical Reactions in Solution
Chemical Reactions
2.2 Changes in Matter
I
II III IV V
Types of Changes in Matter

Physical Changes
 Change in State
 No new substance
Types of Changes in Matter

Nuclear Changes
 Changes within the nucleus
 Fusion and Fission (applets)
Types of Changes in Matter

Chemical Changes
 A change in chemical bonds
 Hoffman Demo
Signs of a Chemical Change
Production of heat and light
 Formation of a gas
 Formation of a precipitate
 Colour change

Why Do Reactions Occur?


The kinetic molecular theory states
that matter is made up of tiny
particles in continuous random
motion.
The average kinetic energy (the
energy of motion) of the particles
depends on the temperature of the
particles.
Solid
Liquid
Gas
Types of Particle Motion
 Translational
motion: the
motion of a particle in a
straight line
 Rotational
movement: a
spinning or turning of a
molecule.
Motion Continued:
 Vibrational
movement: the
back and forth motion within a
molecule
 These
types of particle
movement can be discussed
with the three states of matter
KMT Continued:


The particles this theory refers to
may be atoms (Na), ions (Na+) or
molecules (CO2).
As the particles move, they collide
with each other and objects in
their path.
Collision Theory

Reaction rate depends on the
collisions between reacting particles.

Successful collisions occur if the
particles...
 collide with each other
 have the correct orientation
 have enough kinetic energy to
break bonds
Collision Theory

Activation Energy (Ea)
 minimum energy required for a
reaction to occur
Activation
Energy
Collision Theory

Activation Energy
 depends on reactants
 low Ea = fast rxn rate
Ea
Factors Affecting Rxn Rate

Surface Area
 high SA = fast rxn rate
 more opportunities for collisions
 Increase surface area by…
- using smaller particles
- dissolving in water
Factors Affecting Rxn Rate

Concentration
 high conc = fast rxn rate
 more opportunities for collisions
Factors Affecting Rxn Rate

Temperature
 high temp = fast rxn rate
 high KE
- fast-moving particles
- more likely to reach activation
energy
Factors Affecting Rxn Rate

Temperature
Analogy: 2-car collision
5 mph “fender bender”
50 mph “high-speed crash”
Factors Affecting Rxn Rate

Catalyst
 substance that increases rxn rate
without being consumed in the rxn
 lowers the activation energy
 Burn a sugar cube with a catalyst
Exothermic Reaction
reaction that
releases
energy
 products have
lower energy
than reactants

energy
released
2H2(l) + O2(l)  2H2O(g) + energy
Endothermic Reaction
reaction that
absorbs
energy
 reactants have
lower energy
than products

energy
absorbed
2Al2O3 + energy  4Al + 3O2
Law of Conservation of Mass

mass is neither created nor destroyed
in a chemical reaction
total mass stays the same
 atoms can only rearrange

4H
36 g
2O
4H
2O
4g
32 g
Chemical Reactions
2.3 Balancing Chemical
Reaction Equations
I
II III IV V
Chemical Equations
 Chemical
reactions are
represented by chemical
equations
A
chemical equation lists all of
the compounds that participate
in the reaction.
Chemical Equations
A+B  C+D
REACTANTS
PRODUCTS
Chemical Equations
Writing Equations
2H2(g) + O2(g)  2H2O(g)

Identify the substances involved.

Use symbols to show:
 How many? - coefficient
 Of what? - chemical formula
 In what state? - physical state
Writing Equations
Two atoms of aluminum react with
three units of aqueous copper(II)
chloride to produce three atoms of
copper and two units of aqueous
aluminum chloride.
• How many?
• Of what?
• In what state?
2Al(s) + 3CuCl2(aq)  3Cu(s) + 2AlCl3(aq)
Describing Equations
Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g)
• How many?
• Of what?
• In what state?
One mole of solid zinc reacts with
two moles of aqueous
hydrochloric acid to produce one mole
of aqueous zinc chloride and one
mole of hydrogen gas.
Balancing Steps
1. Write the unbalanced equation.
2. Count atoms on each side.
3. Add coefficients to make #s equal.
Coefficient  subscript = # of atoms
4. Reduce coefficients to lowest
possible ratio, if necessary.
5. Double check atom balance!!!
Helpful Tips
Balance one element at a time.
 Update ALL atom counts after adding
a coefficient.
 If an element appears more than
once per side, balance it last.
 Balance polyatomic ions as single
units.
 “1 SO4” instead of “1 S” and “4 O”

Balancing Example
Aluminum and copper(II) chloride react
to form copper and aluminum chloride.
2 Al + 3 CuCl2  3 Cu + 2 AlCl3
2 1
Al
1 2
3 1
Cu
1 3
6 2
Cl
3 6
The Mole
What
is
the
Mole?
 A counting number (like a dozen)
 Avogadro’s number (NA)
 1 mol = 6.02  1023 items
A
large amount!!!!
 A mole of a compound is an
observable quantity. It can be
weighed and used in experiments.
LEFT: Proceeding clockwise from the top, samples
containing one mole each of copper, aluminum, iron,
sulfur, iodine, and (in the center) mercury.
A. What is the Mole?

1 mole of hockey pucks would
equal the mass of the moon!

1 mole of basketballs would fill a
bag the size of the earth!
 1 mole of pennies would cover the Earth 1/4
mile deep!
 One mole = 6.02 x 1023 atoms,
molecules, or ions (Avogadro’s
number)
When coefficients are used to
balance chemical equations, they
express a mole-to-mole ratio of
the products and reactants.
These numbers do not
represent the exact
number of moles for the
reactants or products,
but rather give a ratio
we can use to compare
quantities.
Molar Mass
 Mass of 1 mole of an element or compound.
 Atomic mass tells the...
 atomic mass units per atom (amu)
 grams per mole (g/mol)
 Round to 2 decimal places
Molar Mass Examples
 carbon
12.01 g/mol
 aluminum
26.98 g/mol
 zinc
65.39 g/mol
Molar Mass Examples
 water
 H2O
 2(1.01) + 16.00 = 18.02 g/mol
 sodium chloride
 NaCl
 22.99 + 35.45 = 58.44 g/mol
Molar Mass Examples
 sodium bicarbonate
 NaHCO3
 22.99 + 1.01 + 12.01 + 3(16.00)
 sucrose
= 84.01 g/mol
 C12H22O11
 12(12.01) + 22(1.01) + 11(16.00)
= 342.34 g/mol
Molar Conversions
molar
mass
6.02  1023
MASS
NUMBER
MOLES
IN
GRAMS
OF
PARTICLES
(g/mol)
(particles/mol)
Molar Conversions
n = m/M
n = number of moles
m = mass (g)
M = Molar Mass (g/mol)
Molar Conversions
p = n * NA
n = number of moles
p = particles
NA = Avogadro's number
(6.02 x 1023 particles per mole)
Molar Conversion Examples
 How many moles of carbon are in
26 g of carbon?
26 g C 1 mol C
12.01 g C
= 2.2 mol C
Molar Conversion Examples
 How many molecules are in 2.50
moles of C12H22O11?
6.02  1023
2.50 mol molecules
1 mol
= 1.51  1024
molecules
C12H22O11
Molar Conversion Examples
 Find the mass of 2.1  1024
molecules of NaHCO3.
2.1  1024
molecules
1 mol
84.01 g
6.02  1023 1 mol
molecules
= 290 g NaHCO3
Chemical Reactions
2.5 Classifying
Chemical
Reactions
I
II III IV V
Combustion

the burning of any substance in O2 to
produce heat
A + O2  B
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
Combustion

Products:
 contain oxygen
 hydrocarbons form CO2 + H2O
4 Na(s)+ O2(g)  2 Na2O(s)
C3H8(g)+ 5 O2(g)  3 CO2(g)+ 4 H2O(g)
Synthesis/Formation

the combination of 2 or more
substances to form a compound

only one product
A + B  AB
Synthesis
H2(g) + Cl2(g)  2 HCl(g)
Synthesis

Products:
 ionic - cancel charges
 covalent - hard to tell
2 Al(s)+ 3 Cl2(g)  2 AlCl3(s)
Decomposition

a compound breaks down into 2 or
more simpler substances

only one reactant
AB  A + B
Decomposition
2 H2O(l)  2 H2(g) + O2(g)
Decomposition

Products:
 Binary ionic - break into elements
 others - hard to tell
2 KBr(l)  2 K(s) +
Br2(l)
Single Replacement

one element replaces another in a
compound
 metal replaces metal (+)
 nonmetal replaces nonmetal (-)
A + BC  B + AC
Single Replacement
Cu(s) + 2AgNO3(aq)  Cu(NO3)2(aq) + 2Ag(s)
Double Replacement
ions in two compounds “change
partners”
 cation of one compound combines
with anion of the other

AB + CD  AD + CB
Double Replacement
Pb(NO3)2(aq) + K2CrO4(aq)  PbCrO4(s) + 2KNO3(aq)
Double Replacement

Products:
 switch negative ions
 one product must be insoluble
(check solubility table)
Pb(NO3)2(aq)+ 2KI(aq)  PbI2(s)+2KNO3(aq)
NaNO3(aq)+ KI(aq)  N.R.
Chemical Reactions in
Solution
Section 2.6
Why in Solution?




Often reactions are carried out in solution in
order to speed up the rate of reaction.
Dissolving something makes the particles
smaller (increase surface area)
We can add heat or stir
May be a way to dilute volatile substances to
make the reaction safer.
Solubility
Refers to an ionic compound’s ability to
dissolve in water
 If a compound is soluble (completely
dissolves in water) it is considered to be
aqueous.
 Aqueous solutions are labeled with the
letters (aq) after the chemical formula
 Example: NaCl(aq)

If an ionic compound is not soluble in
water, it will remain as a solid (it will form
a precipitate).
 Solids are labeled with a (s) after the
compound formula.
 Example: Ag2SO4(s)

How to use the Solubility
Table


The solubility table will tell us if an
ionic compound is soluble in water.
Steps to follow:
1) Determine the non-metal of your
compound and find it on the top
of the chart.
2)
Look beneath the non-metal on the
chart. If the metal it is matched
with is in the section labeled “high
solubility”, the compound is soluble
in water.
- the compound is then labeled
with a (aq)
3)
If the metal is located in the
section labeled “low solubility”,
your compound is not soluble in
water.
-the compound is then labeled with a
(s)

SEE PAGE 8 IN YOUR DATA
BOOKLET
Examples: Are these ionic
compounds soluble in water?
Ag+ and SO42-
K+
and NO3
-
Ca2+ and S2-
Li+ and OH-