Intro to Reactions

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Intro to Reactions
Hey! Screw you, pal!
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Intro to Reactions
Hey! Screw you, pal!
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Intro to Reactions

• Chapter 20 Chemical Reactions

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Signs of a Chemical Reaction

• Evolution of heat and light
• Formation of a gas
• Formation of a precipitate
• Color change

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Law of Conservation of Mass

• mass is neither created nor destroyed in a
  chemical reaction

• total mass stays the same

• atoms can only rearrange

4 H 2 O
4 H 2 O
36 g
4 g
32 g
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Chemical Equations

• AB ? CD

REACTANTS
PRODUCTS
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Chemical Equations
p. 246
, Yields
(aq)
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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

• Remember the diatomic elements.

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Describing Equations

• Describing Coefficients
• individual atom atom
• covalent substance molecule
• ionic substance unit

3 molecules of carbon dioxide 2 atoms of
magnesium 4 units of magnesium oxide
3CO2 ? 2Mg ? 4MgO ?
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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?

2
(aq)
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Describing Equations
Zn(s) 2HCl(aq) ? ZnCl2(aq) H2(g)

• How many?
• Of what?
• In what state?

One atom of solid zinc
reacts with
two molecules of aqueous hydrochloric acid

• to produce

one unit
and one
of aqueous zinc chloride
molecule of hydrogen gas.
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Starting a Reaction

• Always takes a little energy
• Energy goes into breaking bonds in the reactants
• Can use different forms of energy
• Heat
• Electricity
• Light

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Forming Bonds Makes Energy

• Releases energy
• Energy is conserved
• Chemical Energy- energy stored in the bonds of
  the chemicals.
• Reactions have an energy change
• Soak up energy (sweating)
• Release energy (fire)

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Exothermic Reactions

• If breaking bonds takes less energy than making
  them- it releases energy
• Exo- outside
• therm- heat
• Exothermic reactions release energy
• Get hot
• Give off light
• Or release electricity

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Change is down
Chemical Energy
Energy Released
Reactants
Products

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Endothermic Reactions

• If breaking bonds takes more energy than making
  them- it absorbs energy
• Endo- inside
• therm- heat
• Endothermic reactions absorb energy
• Get cold
• Require heat or energy or they stop

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Change is up
Heat is released
Chemical Energy
Reactants
Products

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Balanced Equation

• Atoms cant be created or destroyed
• All the atoms we start with we must end up with
• A balanced equation has the same number of each
  element on both sides of the equation.

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O

C
C
O
O
O

• C O2 CO2
• This equation is already balanced
• What if it isnt already?

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O

C
C
O
O

• C O2 CO
• We need one more oxygen in the products.
• Cant change the formula, because it describes
  what is

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C
C
O

O

O
C
O
C

• Must have started with two C
• 2 C O2 2 CO

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Rules for balancing

• Write the correct formulas for all the reactants
  and products
• Count the number of atoms of each type appearing
  on both sides
• Balance the elements one at a time by adding
  coefficients (the numbers in front)
• Check to make sure it is balanced.

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Never

• Change a subscript to balance an equation.
• If you change the formula you are describing a
  different reaction.
• H2O is a different compound than H2O2
• Never put a coefficient in the middle of a
  formula
• 2 NaCl is okay, Na2Cl is not.

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Example
H2
H2O
O2

Make a table to keep track of where you are at
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Example
H2
H2O
O2

R
P
H
2
2
O
2
1
Need twice as much O in the product
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Example
H2
H2O
O2

2
R
P
H
2
2
O
2
1
Changes the O
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Example
H2
H2O
O2

2
R
P
H
2
2
O
2
1
2
Also changes the H
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Example
H2
H2O
O2

2
R
P
H
2
2
4
O
2
1
2
Need twice as much H in the reactant
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Example
H2
H2O
O2

2
2
R
P
H
2
2
4
O
2
1
2
Recount
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Example
H2
H2O
O2

2
2
R
P
H
2
2
4
4
O
2
1
2
The equation is balanced, has the same number of
each kind of atom on both sides
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Example
H2
H2O
O2

2
2
R
P
H
2
2
4
4
O
2
1
2
This is the answer
Not this
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Examples

• AgNO3 Cu Cu(NO3)2 Ag
• Mg N2 Mg3N2
• P O2 P4O10
• Na H2O H2 NaOH
• CH4 O2 CO2 H2O

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Examples of Balancing Equations

• a) Pb(NO3)2 K2CrO4 ? PbCrO4 KNO3
• b) MnO2 HCl ? MnCl2 H2O Cl2
• c) C3H6 O2 ?CO2 H2O
• d) Zn(OH)2 H3PO4 ? Zn3(PO4)2 H2O

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Answers

• a) Pb(NO3)2 K2CrO4 ? PbCrO4 2KNO3
• b) MnO2 4HCl ? MnCl2 2H2O Cl2
• c) 2C3H6 9O2 ?6CO2 6H2O
• d) 3Zn(OH)2 2H3PO4 ? Zn3(PO4)2 6H2O

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More Examples

• e) CO Fe2O3 ?Fe CO2
• f) CS2 Cl2 ?CCl4 S2Cl2
• g) CH4 Br2 ? CH3Br HBr
• h) Ba(CN)2 H2SO4 ? BaSO4 HCN

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More Answers

• e) 3CO Fe2O3 ?2Fe 3CO2
• f) CS2 3Cl2 ?CCl4 S2Cl2
• g) CH4 Br2 ? CH3Br HBr BALANCED
• h) Ba(CN)2 H2SO4 ? BaSO4 2HCN

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4 Types of Reactions

• Synthesis (combination)
• Decomposition
• Single Replacement
• Double Replacement
• Combustion

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Synthesis (Combination) Reactions
Two or more substances combine to form a new
compound.
A X ? AX

• Reaction of elements with oxygen and sulfur
• Reactions of metals with Halogens
• Synthesis Reactions with Oxides
• There are others not covered here!

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Decomposition Reactions
A single compound undergoes a reaction that
produces two or more simpler substances
AX ? A X

• Decomposition of
• Binary compounds H2O(l ) ? 2H2(g) O2(g)
• Metal carbonates CaCO3(s) ? CaO(s) CO2(g)
• Metal hydroxides Ca(OH)2(s) ? CaO(s) H2O(g)
• Metal chlorates 2KClO3(s) ? 2KCl(s) 3O2(g)
• Oxyacids H2CO3(aq) ? CO2(g) H2O(l
  )

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Single Replacement Reactions
A BX ? AX B
BX Y ? BY X
Replacement of

• Metals by another metal
• Hydrogen in water by a metal
• Hydrogen in an acid by a metal
• Halogens by more active halogens

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The Activity Series of the Metals

• Lithium
• Potassium
• Calcium
• Sodium
• Magnesium
• Aluminum
• Zinc
• Chromium
• Iron
• Nickel
• Lead
• Hydrogen
• Bismuth
• Copper
• Mercury
• Silver
• Platinum
• Gold

HCl Cu ? No reaction 2 HCl 2 Zn ? 2 ZnCl H2
Metals can replace other metals provided that
they are above the metal that they are trying to
replace.
Metals above hydrogen can replace hydrogen in
acids.
Metals from sodium upward can replace hydrogen in
water
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The Activity Series of the Halogens

• Fluorine
• Chlorine
• Bromine
• Iodine

Halogens can replace other halogens in
compounds, provided that they are above the
halogen that they are trying to replace.
2NaCl(s) F2(g) ?
2NaF(s) Cl2(g)
???
MgCl2(s) Br2(g) ?
???
No Reaction
MgFl2 Cl2
MgCl2(s) Fl2(g) ?
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Double Replacement Reactions
The ions of two compounds exchange places in
an aqueous solution to form two new compounds.
AX BY ? AY BX
One of the compounds formed is usually a
precipitate, an insoluble gas that bubbles out
of solution, or a molecular compound, usually
water.
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Combustion Reactions
A substance combines with oxygen, releasing a
large amount of energy in the form of light and
heat.

• Reactive elements combine with oxygen

P4(s) 5O2(g) ? P4O10(s)
(This is also a synthesis reaction)

• The burning of natural gas, wood, gasoline

C3H8(g) 5O2(g) ? 3CO2(g) 4H2O(g)
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Collision Theory

• In order to react molecules and atoms must touch
  each other.
• They must hit each other hard enough to react.
• Anything that increase these things will make the
  reaction faster.

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Things that Affect Rate

• Temperature
• Higher temperature faster particles.
• More and harder collisions.
• Faster Reactions.
• Concentration
• More concentrated closer together the molecules.
• Collide more often.
• Faster reaction.

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Things that Affect Rate

• Particle size
• Molecules can only collide at the surface.
• Smaller particles bigger surface area.
• Smaller particles faster reaction.
• Smallest possible is molecules or ions.
• Dissolving speeds up reactions.
• Getting two solids to react with each other is
  slow.

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Things that Affect Rate

• Catalysts- substances that speed up a reaction
  without being used up.
• Inhibitor- a substance that blocks a catalyst,
  slowing the reaction down
• Enzymes are biological catalysts- made by plants
  and animals to control reactions
• Heat destroys most catalysts