Download Chemical Reactions - Northside Middle School

Cool Chemistry Show!
Unit 2:
Chapter 6
Chapter Vocabulary
1. Physical Change
2. Solution
3. Homogeneous
4. Solute
5. Solvent
6. Chemical Change
7. Product
8. Reactant
9. Chemical Reaction
10.Precipitate
11.Concentration
12.Molarity
13.Mole
14.Saturated Solution
15.Supersaturated
Solution
16.Polymer
17.Chemical Test
18.Acid-Base Indicator
19.Rate of Reaction
20.Surface Area
21.Catalyst
22.Endothermic
Change/Rxn
23.Exothermic Change/Rxn
24.Activation Energy
25.Kinetic Energy
26.Law of Conservation of
Energy
27.Gravitational Potential
Energy
28.Heat
29.Temperature
Chapter Vocabulary
30.Compound
31.Ion
32.Anion
33.Cation
34.Ionic Compound
35.Polyatomic Ion
36.Molecular
Compound
37.Covalent Bond
38.Oxidation Number
39.Synthesis Reaction
40.Decomposition
Reaction
41.Single-Replacement
Reaction
42.Double Replacement
Reaction
43.Salts
44.Acid
45.Base
46.Neutralization
47.Titration
48.Endpoint
49.Buffer
50.pH
51.Oxidation
52.Reduction
53.Redox Reaction
Chemical and Physical
Changes
Act 1
Properties of Matter
Properties
• Words that describe matter
(adjectives)
• Physical Properties- a property
that can be observed and
measured without changing the
composition.
• Examples: color, hardness, m.p., b.p.
• Chemical Properties- a property
that can only be observed by
changing the composition of the
material.
• Ex: Iron rusting
Classification of Matter
Matter
Mixtures
Pure Substance
Mixtures
• Combination of 2 or more
pure substances
• Each substance retains its
individual chemical
properties
• Composition variable
• Most everyday matter
Mixture examples
• Sand and water
• Salt and water
Classification of Matter
Matter
Mixtures
Homogeneous
Pure Substance
Heterogeneous
Mixtures
• Heterogeneous:
• Mixture is not uniform in
composition
• EX: Chocolate chip cookie,
gravel, soil.
• Homogeneous:
• Same composition throughout;
called “solutions”
• Ex: Kool-aid, air, salt water
Solutions
• Homogeneous mixture
• Mixed molecule by molecule
• Can occur between any state of
matter:
• gas in gas
• liquid in gas
• gas in liquid
• solid in liquid
• solid in solid (alloys), etc.
Solutions: Parts
• Solute:
• Substance being dissolved
• Salt, sugar
• Solvent:
• Substance that does the
dissolving
• Water, alcohol
Saturated & Supersaturated
Solutions
• Concentration: ratio of solute to
solvent
• 1.5M (Molarity) : number of moles
of solute dissolved in 1 liter of
solution
• Dilute: less solute
• Concentrated: more solute
• Saturated: no more solute will
dissolve
• Supersaturated: contains more
solute than normal
Pure Substance
• Same fixed composition and
properties
• Ex: sand, water, oxygen,
sodium
1. Elements
2. Compounds
Classification of Matter
Matter
Mixtures
Homogenous
Pure Substance
Heterogeneous
Elements
Compounds
Elements
•
•
•
•
•
Simplest kind of matter
All one kind of atom.
On Earth 91 occur naturally
Unique name and symbol
Symbol 1, 2 or 3 letters
• 1st: always capitalized
• 2nd and 3rd: lower case
Compounds
• Made of two or more atoms, chemically
combined
• Can be broken down only by chemical
methods
• When broken down, the pieces have
completely different properties than
the original compound.
• NaCl: Table salt
• Na: Flammable metal
Cl: poisonous
gas
• Formula: combination of chem symbols
to show what makes up the compound
Matter
Mixtures
Heterogeneous
Dirt, blood, milk
Physical Changes
Pure Substances
Homogenous
Elements
Compounds
Lemonade, gas, steel
Oxygen, gold, iron
Salt baking
soda, sugar
Chemical Changes
Compound or Mixture
Compound
Mixture
Made of one kind
of material
Made of more than
one kind of material
Made by a
chemical change
Made by a
physical change
Definite
composition
Variable
composition
Which is it?
Mixture
Element
Compound
Physical Changes
• A change that changes
appearances, without
changing the composition.
• Ex. Boil, melt, cut, bend,
split, crack
Chemical Changes
• A change where a new form of
matter is formed.
• Ex. Rust, burn, decompose,
ferment
• Reactants: starting materials
• Products: ending materials
Evidence of a Chemical
Reaction
• Temp change
• Release: light/heat
• Absorb
•
•
•
•
Color change
Odor change
Production of gas bubbles
Formation of a precipitate
(solid)
Polymers
• Giant molecules of
repeating groups of
atoms (monomers)
• Linked through
intermolecular forces
• Can be molded into useful
objects
Polymers
• Cross-linking
• Short bridges b/w long polymer
chains
• Gives polymer new properties
monomer
polymer
Cross-linking
Polymers
• Polyester: clothing
• Polyethylene: blow-molded
beverage bottles, auto gas
tanks, and pipe
• Epoxy adhesives
Biological Polymers
• Carbohydrates: sugars
• Lipids: fatty acids
• Proteins: amino acids
• DNA/RNA: nucleotides
Check your own
Portfolio
• Chapter Vocabulary
• Act 1
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What do you think?
Summary
Checking up Q’s
What do you think now?
Chemical Essential Q’s
Act 1 Quiz!
1. Is the property below chemical or physical ?
Mercury is a silvery liquid.
2. Which of the following is NOT evidence of a
chemical change?
a) formation of a precipitate b) formation of a gas
c) melting of a solid
d) color change
3. The materials forming a solution are called the
____ and _____.
4. When no more solute will dissolve in the solvent
the solution is __________.
a) dilute b) concentrated c) saturated d) supersaturated
5. For each determine if the material is an element,
compound, heterogeneous or homogeneous.
a) pepperoni pizza
b) gold c) steel d) soda pop
Ion
Name
Ion
NH4+
ammonium
PO43-
phosphate
NO2-
nitrite
CO32-
Carbonate
NO3-
nitrate
SO32-
Sulfite
OH-
hydroxide
SO42-
Sulfate
CN-
cyanide
BrO3-
bromate
MnO4- permanganate IO3-
Name
iodate
More Chemical Changes
Act 2
Chemical Test for Gases
• Chemical test: procedure or
chemical rxn used to ID a
substance
• Oxygen
• Glowing splint; burst into flames
http://www.youtube.com/watch?v=NGEc6wMedDE
• Hydrogen
• Burning splint; loud pop
http://www.youtube.com/watch?v=7s2dfXcoOyI&feature=related
• Carbon Dioxide
• Limewater: Ca(OH)2 Forms a ppt
http://www.youtube.com/watch?v=21aBA0jUJXU
Indicators for Acids and
Bases
• Indicators used to determine
acids an bases
• Substances that change color
when react with acid or base
• Phenol red
• Turns yellow in presence of an
acid
• Litmus Paper
Check your own
Portfolio
• Chapter Vocabulary
• Act 1
•
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What do you think?
Summary
Checking up Q’s
What do you think now?
Chemical Essential Q’s
• Act 2
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What do you think?
Summary
Checking up Q’s
What do you think now?
Chemical Essential Q’s
Act 2 Quiz!
Act 2 Quiz
1. Which of the following is a test for carbon dioxide?
a) Burning splint causes a pop
into flame
c) Fmn of white ppt in limewater
b) glowing splint burst
d) global warming
2. What is the acid-base indicator used in lab?
a) Thymol Blue
c) Methyl yellow
b) Methyl red
d) Phenol red
3. The indicator in lab turned (color)
of
(acid or base) .
4. Write the formula.
Sodium sulfate
5. Write the name.
Ra(NO3)2
in the presence
1. Which of the following is NOT a physical process?
a) Sharpening a pencil
c) An explosion of dynamite
b) melting ice
d) popping a balloon
2. Which of the following is evidence of a chemical
reaction?
a) Formation of gas
b) disappearance of liquid
c) Formation of smaller pieces
d) formation of crystals
3. Which of the following is a test for carbon dioxide?
a) Burning splint causes a pop
into flame
c) Fmn of white ppt in limewater
b) glowing splint burst
d) global warming
4. What is the acid-base indicator used in lab?
a) Thymol Blue
c) Methyl yellow
b) Methyl red
d) Phenol red
5. The indicator in lab turned (color)
of
(acid or base) .
in the presence
Investigate
Water
Baking
soda
Baking
Powder
Antiacid
Tablet
Heated
Vinegar
Ammonia
No
More
Intense No
bubbles bubbles bubbles bubbles
CO2
Slow
Intense Intense No
bubbles bubbles bubbles bubbles
CO2
CO2
Fast
Fast
Intense No
bubbles bubbles bubbles bubbles
CO2
CO2
Chemical Names and
formulas
Act 3
Ions
• Form by losing or gaining
electrons
• Anions
• Typically nonmetals
• Gain electrons
• Negatively charged
• Cations
• Typically metals
• Lose electrons
• Positively charged
Ionic compound
• Bond b/w cation and anion
• Electrostatic force
• Conducts electricity when liquid
or dissolved in water
Naming ions and ionic
compounds
• Name cation first and the anion second
• Monatomic cations use the element name
• Monatomic anions take their name from the
root of the element name plus the suffix –ide
• Oxidation numbers of transition metals are
written as roman numerals in parentheses
• If the compound contains a polyatomic ion,
simply name to ion
• Examples
Examples
• Write the formula for the pairs of
elements
•
•
•
•
Barium and oxygen
Strontium and iodine
Radium and chlorine
Lithium and chlorine
• Write the formula for each :
•
•
•
•
Lithium oxide
Calcium bromide
Sodium oxide
Aluminum sulfide
Examples
• Write the formula for each :
•
•
•
•
Ammonium and sulfite ions (SO32-)
Barium and nitrate ions (NO3-)
Magnesium hydroxide (OH-)
Sodium phosphate (PO43-)
Examples
• Write the formula for each :
•
•
•
•
Copper (I) sulfide
Tin (IV) fluoride
Gold (III) cyanide
Lead (II) sulfide
• Write the names:
• Mn2O3
• HgF2
Molecular Compounds
•
•
•
•
Covalent bonds
Sharing electrons
Nonmetals only
Typically not conductive
Molecular Nomenclature
Prefixes
PREFIX
monoditritetrapentahexaheptaoctanonadeca-
NUMBER
1
2
3
4
5
6
7
8
9
10
Molecular Nomenclature:
Examples
• CCl4
• carbon tetrachloride
• N2O
• dinitrogen monoxide
• SF6
• sulfur hexafluoride
More Molecular Examples
• arsenic trichloride
• AsCl3
• dinitrogen pentoxide
• N2O5
• tetraphosphorus decoxide
• P4O10
Diatomic Molecules
•
•
•
•
•
•
•
•
•
H
I
Br
O
Cl
F
N
Hi Bronclif!
aka... HI BrONClF
Check your own Portfolio
• Chapter Vocabulary
• Act 1
•
•
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•
What do you think?
Summary
Checking up Q’s
What do you think now?
Chemical Essential Q’s
• Act 2
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What do you think?
Summary
Checking up Q’s
What do you think now?
Chemical Essential Q’s
• Act 3
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What do you think?
Summary
Checking up Q’s
What do you think now?
Chemical Essential Q’s
Act 3 Quiz! (1pt each)
Name
Formula
Cation
Anion
Sodium Nitrate
Br-
MgBr2
Li +
O2-
K2S
Calcium hydroxide
Sr2+
Strontium bromate
Fe2(CO3)3
Potassium sulfate
(NH4)3PO4
Fe3+
CO32-
Act 3 Quiz! (1.5pt each)
Name
Formula
Cation
Br-
MgBr2
Li +
K2S
Calcium nitride
Iron III carbonate
Fe2(CO3)3
Ammonium
Phosphate
(NH4)3PO4
Anion
NH4+
O2-
Chemical Equations
Act 4
Representing Chemical
Reactions
• Have two parts:
• Reactants - the substances
you start with
• Products- the substances you
end up with
• Reactants  Products
In a chemical reaction
• The way atoms are joined is
changed
• Atoms aren’t created or destroyed.
• Can be described several ways:
1. Word equation
Copper + chlorine  copper (II) chloride
2. Skeleton equation
Cu(s) + Cl2(g)  CuCl2(l)
Skeleton Equation
• Uses formulas and symbols
to describe a reaction
• All chemical equations are
sentences that describe
reactions.
Symbols in equations
• Arrow separates the
reactants from the products
• Read “reacts to form”
• The plus sign = “and”
• (s) after the formula = solid
• (g) after the formula = gas
• (l) after the formula = liquid
Symbols used in
equations
• (aq) after the formula dissolved in water, an
aqueous solution
Convert these to
equations
• Solid iron (III) sulfide reacts
with gaseous hydrogen chloride
to form iron (III) chloride and
hydrogen sulfide gas.
• Nitric acid (HNO3) dissolved in
water reacts with solid sodium
carbonate to form liquid water
and carbon dioxide gas and
sodium nitrate dissolved in
Now, read these:
• Fe(s) + O2(g)  Fe2O3(s)
• Cu(s) + AgNO3(aq) 
Ag(s) + Cu(NO3)2(aq)
Pt
• NO2 (g)  

N2(g) + O2(g)
Balanced Equation
• Atoms can’t 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.
C
+
O
O

O C
• C + O2  CO2
• This equation is already
balanced
• What if it isn’t?
O
C
+
O
O

C
O
• C + O2  CO
• We need one more oxygen in the
products.
• Can’t change the formula,
because it describes what it is
(carbon monoxide in this
example)
C
+
O
O

C
O
C
O
• Must be used to make
another CO
• But where did the other C
come from?
C
+
C
O
O

C
O
C
O
• Must have started with two
C
• 2 C + O2  2 CO
Rules for balancing:
 Assemble, 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) - save H and O until LAST!
 Check to make sure it is balanced.
• 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.
Example
H2 + O2  H2O
Make a table to keep track of where you
are at
Example
H2 + O2  H2O
R
P
2 H 2
2 O 1
Need twice as much O in the product
Example
H2 + O2 
R
P
2 H 2
2 O 1
Changes the O
2 H2O
Example
H2 + O2 
2 H2O
R
P
2 H 2
2 O 1 2
Also changes the H
Example
H2 + O2 
2 H2O
R
P
2 H 2 4
2 O 1 2
Need twice as much H in the reactant
Example
2 H2 + O2 
2 H2O
R
P
2 H 2 4
2 O 1 2
Recount
Example
2 H2 + O2 
2 H2O
R
P
4 2 H 2 4
2 O 1 2
The equation is balanced, has the same
number of each kind of atom on both sides
Example
2 H2 + O2 
2 H2O
R
P
4 2 H 2 4
2 O 1 2
This is the answer
Not this
Balancing Examples
• _AgNO3 + _Cu  _Cu(NO3)2 + _Ag
• _Mg + _N2  _Mg3N2
• _P + _O2  _P4O10
• _Na + _H2O  _H2 + _NaOH
• _CH4 + _O2  _CO2 + _H2O
Balancing Examples
• 2AgNO3 + Cu  Cu(NO3)2 + 2Ag
• 3Mg + N2  Mg3N2
• 4P + 5O2  P4O10
• 2Na + 2H2O  H2 + 2NaOH
• 2Na + 2H2O  H2 + 2NaOH
• CH4 + 2O2  CO2 + 2H2O
Act 4 Quiz Part I
•
Write the skeleton equation for:
1. Solid potassium reacts with iron (II) sulfate
dissolved in water to form potassium sulfate
dissolved in water and solid iron.
2. Solid phosphorous and oxygen gas reacts
to form solid diphosphorus trioxide.
• Balance the following equations:
3. Li + AlCl3  LiCl + Al
4. Ca + Au(NO3)3  Ca(NO3)2 + Au
• Write a sentence for the equation (remember roman
numerals with transition metals)
5. AgNO3(aq) + BaCl2(aq)  Ba(NO3)2 (aq) + AgCl
Types of Chemical
Reactions
• OBJECTIVES:
• Identify a reaction as
combination, decomposition,
single-replacement, doublereplacement, or combustion
• Predict the products of
combination, decomposition,
single-replacement, doublereplacement, and combustion
reactions.
Types of Reactions
•
•
•
•
•
There are millions of reactions.
Can’t remember them all
Fall into several categories.
We will learn 5 major types.
Will be able to predict the
products.
• For some, we will be able to
predict whether they will
happen at all.
#1 - Combination Reactions
(Synthesis)
• Combine - put together
• 2 substances combine to make
one compound.
• Ca +O2 CaO
• SO3 + H2O  H2SO4
• We can predict the products if
they are two elements.
• Mg + N2 
#1 - Combination Reactions
(Synthesis)
• Combine - put together
• 2 substances combine to make
one compound.
• Ca +O2 CaO
• SO3 + H2O  H2SO4
• We can predict the products if
they are two elements.
• 3Mg + N2 Mg3N2
Write and balance
• Ca + Cl2 
• Fe + O2  iron (II) oxide
• Al + O2 
• Remember that the first step
is to write the correct
formulas
• Then balance by using
coefficients only
#2 - Combustion
• Means “add oxygen”
• A compound composed of only
C, H, and maybe O is reacted
with oxygen
• If the combustion is complete,
the products will be CO2 and
H2O.
• If the combustion is incomplete,
the products will be CO
(possibly just C) and H2O.
Examples
• C4H10 + O2 
• C6H12O6 + O2 
Examples
• 2C4H10 + 13O2  8CO2 +
10H2O
• C6H12O6 + 6O2  6CO2 + 6H2O
#3 - Decomposition
Reactions
• decompose = fall apart
• one reactant falls apart into two or
more elements or compounds.
electricity
• NaCl 

 Na + Cl2
• CaCO3  CaO + CO2
• Note that energy is usually required
to decompose
#3 - Decomposition
Reactions
• Can predict the products if
it is a binary compound
• Made up of only two
elements
• Falls apart into its elements
electricity
• H2O   

• HgO  
#3 - Decomposition
Reactions
• Can predict the products if
it is a binary compound
• Made up of only two
elements
• Falls apart into its elements
• 2H2O electricity
2H2 + O2
 
• 2HgO  
2Hg + O2
#4 - Single Replacement
• One element replaces another
• Reactants must be an element
and a compound.
• Products will be a different
element and a different
compound
• Na + KCl  K + NaCl
• F2 + LiCl  LiF + Cl2
#4 Single Replacement
• Metals replace other metals
(and hydrogen)
• K + AlN 
• Zn + HCl 
• Think of water as HOH
• Metals replace one of the H,
combine with hydroxide.
• Na + HOH 
#4 Single Replacement
• Metals replace other metals
(and hydrogen)
• 3K + AlN K3N + Al
• Zn + 2HCl  ZnCl2 + H2
• Think of water as HOH
• Metals replace one of the H,
combine with hydroxide.
• Na + HOH  NaOH + H2
#4 Single Replacement
• We can tell whether a reaction
will happen
• Some chemicals are more
“active” than others
• More active replaces less active
• There is a list on page 462 called the Activity Series of
Metals
• Higher on the list replaces lower.
#4 Single Replacement
• Note the
* concerning Hydrogen
• H can be replaced in acids by
everything higher
• Li, K, Ba, Ca, & Na replace H from
acids and water
Challenge!
•
•
•
•
Fe + CuSO4 
MgCl2 (aq) + Zn (s) 
Ag + Cu(NO3)2 
F2 + NaBr 
• Br2 + NaF 
#4 Single Replacement
•
•
•
•
Fe + CuSO4 FeSO4 + Cu
MgCl2 + Zn no
Ag + Cu(NO3)2 no
F2 + 2NaBr 2NaF + Br2
• Br2 + 2NaF no
#5 - Double
Replacement
• Two things replace each other.
• Reactants must be two ionic
compounds or acids.
• Usually in aqueous solution
• NaOH + FeCl3 
• The positive ions change place.
• NaOH + FeCl3 Fe+3 OH- + Na+1
Cl-1
• NaOH + FeCl3 Fe(OH)3 + NaCl
#5 - Double
Replacement
• Has certain “driving forces”
• Will only happen if one of the
products:
• doesn’t dissolve in water and
forms a solid (a “precipitate”), or
• is a gas that bubbles out, or
• is a covalent compound (usually
water).
Solubility rules
1. All salts of alkali metals and ammonium
are soluble
2. All chlorides, bromides and iodides are
soluble with silver, lead, and mercury
3. All nitrates are soluble
4. All sulfates are soluble except with
calcium, barium, strontium, and lead
5. All carbonates, phosphates, hydroxides,
and sulfides are insoluble except with
alkali metal or ammonium.
Complete and balance
• Assume all of the following
reactions take place:
• CaCl2 + NaOH 
• CuCl2 + K2S 
• KOH + Fe(NO3)3 
• (NH4)2SO4 + BaF2 
Complete and balance
• Assume all of the following
reactions take place:
• CaCl2 + 2NaOH  2NaCl + Ca(OH)2
• CuCl2 + K2S  2KCl + CuS
• 3KOH + Fe(NO3)3  Fe(OH)3 + 3K(NO3)
• (NH4)2SO4 + BaF2  BaSO4 + 2NH4 F
Act 4 Quiz Part 2
• For each equation
a.
b.
c.
d.
1.
2.
3.
4.
5.
ID (1pt each equation)
Predict the products (2pt each equation)
Balance (1pt each equation)
Write 2 of the 5 as sentences (2pt extra
credit)
K + B2O3
Al + S8
C10H22 + O2
H3PO4 + Ca(OH)2
Mg3P2
Act 4 ReQuiz Part 2
• For each equation
a.
b.
c.
d.
1.
2.
3.
4.
5.
ID (1pt each equation)
Predict the products (2pt each equation)
Balance (1pt each equation)
Write 2 of the 5 as sentences (2pt extra
credit)
Na + O2
Ca + Al2S3
C10H22 + O2
KCl
LiNO3 + Mg(OH)2
Chemical Energy
Act 5
Endothermic and
Exothermic
• Endothermic
• Process that absorbs heat energy
• Internal energy of system
increases
• Example from lab?
• Cold pack made with NH4NO3
• 2H2O + energy
2H2 + O2
Endothermic and
Exothermic
• Exothermic:
• Process that releases heat energy
• Internal energy of system
decreases
• Example from lab?
• Hot pack made with Na2CO3
• 2H2 + O2
2H2O + energy
Importance of energy
• In order for chem rxn to take place,
particles from the reactants must
collide
• Particles must have enough kinetic
energy to break existing bonds
• Activation energy: minimum amount
of energy required for a chem rxn to
start
• Bond breaking
• Endothermic; add energy
• Bond making
• Exothermic; release energy
Exothermic
• Bond forming >
bond breaking
• Products have
less energy than
the reactants
• ∆H: heat of
reaction
• Energy difference
b/w reactants and
products
Endothermic
• Bond forming <
bond breaking
• Reactants have
less energy than
the products
• Ea: activation
energy
Conservation of Energy
• Law of conservation of Energy
• Energy absorbed from or released
to surroundings is equal to the
change in the energy system
• Total energy in a closed system
remains the same.
Conservation of Energy
• Law of conservation of Energy
• Energy is not created or
destroyed, only transferred
• Some types of energy:
• Light
• Heat
• Sound
• Nuclear
• Kinetic
• Chemical
Heat and Temperature
• Related but not the same
• Heat is a type of energy; can be
transferred
• Temperature related to kinetic
energy of the molecules; what
is measured
1.
Act 5 Quiz
The minimum energy required for a chemical
reaction to begin is _________.
2.
__________ is the measurement of the kinetic
energy of particles.
3.
What is the Law of conservation of Energy?
4.
The process of breaking bonds is ________ and
the process of making bonds is ________.
5.
Give an example of an endothermic reaction
from lab.
EC. Fill in the missing slots AND ID the
graph as endothermic or exothermic
A
D
C
B
Reaction Rates
Act 6
Factors Affecting Rate of
Reaction
• Rate of reaction
• How fast a rxn takes place
• Decrease of concentration of reactants over
time
• Increase of concentration of products over
time
• Factors (Increase particles colliding)
•
•
•
•
Surface area
Concentration of reactants
Temperature
Catalyst
Factors Affecting Rate
of Reaction
•
Surface Area
•
When the surface area of reactants
increases, the reaction rate increases.
•
•
•
•
•
•
•
Chemical reactions occur when reactants
collide at the surface of other reactants.
Smaller particle size
Greater surface area
Greater chance for collisions to occur
Chemical reaction will proceed faster
Lab example?
Antacid tablets: crushed vs. whole
Factors Affecting Rate
of Reaction
• Concentration
• When reactants are more concentrated,
the rate of a chemical reaction will
increase.
• There is a greater chance that reactant
particles will collide when they are more
concentrated.
• More collisions mean a faster reaction rate.
• Lab example?
• Vinegar: 3 different concentrations
• HCl: 2 different concentrations
Factors Affecting Rate
of Reaction
Temperature




•
•

Temperature increases, the rate of a
chemical reaction increases.
Average kinetic energy of molecules of
reactants increases as temperatures
increase.
More reactant particles with enough
energy produces more successful
collisions
The reaction will proceed faster
Lab example?
Tea and antacid tablets
•
Hot vs. Cold water
Factors Affecting Rate
of Reaction
•
Catalyst
•
The presence of a catalyst will speed
up a chemical reaction.
•
•
•
•
•
•
A catalyst lowers the amount of energy
needed to start a reaction (activation
energy).
Energy needed for successful collisions is
less
There will be more successful collisions
The chemical reaction will proceed faster
Lab example?
MnO2 and H2O2
Act 6 Quiz
1. List the four factors that influence the rate of a
reaction.
2. For each reaction, predict it will proceed faster (F) or
slower (S) as a result.
a. The reactants are made into smaller pieces.
b. The reactants are diluted.
c. The reactants are heated.
d. The reactants are kept in large pieces.
3. Pick one condition from above and explain why it
effects the reaction like it does.
4. A catalyst speeds up a reaction by ______.
5. Pick one of the 4 factors. Give the lab example and
explain why the rate increased.
Acids and Bases
Act 7
Red and Blue Litmus Paper
Phenolthalien
Magnesium
• Reacts with acid
Bromothymol Blue
Acidic
Neutral
Basic
Some Definitions
• Arrhenius
• Acid: when dissolved in water,
increases the concentration of
hydrogen ions.
• Base: when dissolved in water,
increases the concentration of
hydroxide ions.
Some Definitions
• Brønsted–Lowry
• Acid: Proton donor
• Base: Proton acceptor
What Happens When an
Acid Dissolves in Water?
• Water acts as a
Brønsted–Lowry
base and pulls a
proton (H+) from
the acid.
• As a result, the
conjugate base of
the acid and a
hydronium ion are
formed.
Conjugate Acids and
Bases:
• From the Latin word conjugare,
meaning “to join together.”
• Reactions between acids and bases
always yield their conjugate bases
and acids.
Properties of Acids
• Reacts with metals to produce
hydrogen gas
• The effect of acids on indicators
(from lab)
• Neutralization of bases
• Sour taste
• Have a pH less than 7
Acid and Base Strength
• Strong acids are
completely
dissociated in water.
• Their conjugate bases
are quite weak.
• Weak acids only
dissociate partially
in water.
• Their conjugate bases
are weak bases.
Strong Acids
• Recognize and write ionization
equations for the following five
strong acids
• HI
HI (aq)
H+ + I• HBr
• HCl
• HNO3
• H2SO4
Properties of Bases
•
•
•
•
Bases are electrolytes
The effect of bases on indicators
Neutralization of acids
Water solutions of bases taste
bitter and feel slippery
• Have a pH greater than 7
Acid and Base Strength
• Substances with
negligible acidity do
not dissociate in
water.
• Their conjugate bases
are exceedingly
strong.
Strong Bases
• Recognize and write ionization
equations for the following four
strong bases
• NaOH
Na+ + OH• Ca(OH)2
• Ba(OH)2
• KOH
Neutralization
• When acids and bases react hydrogen ion
and hydroxide ion form water
• Remaining ions form a salt
• HCl(aq) + NaOH(aq)
HOH(l)+ NaCl(aq)
• Resulting solution is neutral
• Buffer:
• A solution that resist change in pH when acid
or base is added.
pH
• pH is defined as the negative
base-10 logarithm (ten fold) of
the hydrogen ion concentration.
pH = −log [H+]
pOH
• pOH is defined as the negative
base-10 logarithm (ten fold) of
the hydroxide ion concentration.
pOH = −log [OH-]
Act 7 Quiz
1. Label as an acid or base and write the ionization
reactions and balance for each
a) Ca(OH)2
b)H2S
c)HCl
d)LiOH
2. Write the equation and ID the acid, base, conjugate
acid and conjugate base.
H2SO4(aq) + H2O(aq)
H3O+ + HSO4-(aq)
3. What does pH stand for?
4. When acids and bases react ______ and a ______
are formed.
5. Phenolthalein is color in the presence of an acid.
6. Red litmus paper is color in the presence of a base.
Color Reactions that
Involve the Transfer of
Electrons
Act 8
Redox reactions
• Oxidation
• When substance loses electrons
• Lab example: Iron lost electrons to form Fe2+ that
dissolve in solution
• Reduction
• When substances gains electrons
• Lab example: Cu2+ gained electrons from the iron to
form copper atoms that plate out as solid
• Redox rxn
• Process of oxidation and reduction together
• Remember!
• Lose Electrons Oxidation; Gain Electrons Reduction
Redox Reactions
• Rusting
• Water and oxygen corrode iron metal to
rust
• Iron loses electrons to b/c Fe3+
• O2 gain electrons to b/c O2• Prevention by protective covering, i.e.
painting
Lab
1. Did a rxn take place? How do you know?
•
•
•
Zn + CuSO4
What type of rxn?
What is oxidized? Why? What is reduced?
Why?
2. Did a rxn take place? How do you know?
•
•
•
Al + CuCl2
What type of rxn?
What is oxidized? Reduced?
Lab
3. Did a rxn take place? How do you know?
•
•
•
Mg + CuCl2
What type of rxn?
What is oxidized? Reduced?
4. Did a rxn take place? How do you know?
•
•
•
Cu + Al(NO3)2
What type of rxn?
What is oxidized? Reduced?
Act 8 Quiz
1. Na + Cu(NO3)2
a. What type of rxn (of the 5 types)?
b. Predict the products.
c. What is oxidized? Why? What is reduced?
Why?
2. What is needed for iron to rust? What can
be done to protect iron from rusting?