Intro to Reactions (p. 282

1
Intro to Reactions (p. 282 285)

• Ch. 9 Chemical Reactions

2
Chemical Reaction

• Chemical change
• Atoms of one or more substances (reactants) are
  rearranged into new substances (products)
• Signs of a chemical reaction
• Change of temperature or light
• Formation of a gas (bubbles)
• Formation of a precipitate
• Odor
• Color change
• Change in volume

3
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
4
Chemical Equations

• AB ? CD

REACTANTS
PRODUCTS
Starting substances Substances formed
5
Chemical Equations
p. 283
6
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!

7
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 ?
8
Writing Equations

• Two atoms of solid aluminum react with three
  units of aqueous copper(II) chloride to produce
  three atoms of solid copper and two units of
  aqueous aluminum chloride.

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

2 Al (s)
3 CuCl2 (aq)
3 Cu (s)
?
2 AlCl3 (aq)
9
Describing Equations
Zn(s) 2HCl(aq) ? ZnCl2(aq) H2(g)

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

One atom of solid zinc
react with
two molecules
of aqueous hydrochloric acid
to produce one unit
of aqueous zinc chloride
and one molecule of
hydrogen gas.
10
II. Balancing Equations

• Ch. 9 Chemical Reactions

11
Why is there a 2 after the oxygen?
__________________
Oxygen is diatomic
1
2
10
4
2
8
4
12
12
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!!!

13
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

14
Counting Atoms
H 3 PO4 1

• H3PO4
• 2KC2H3O2
• Mg3(PO4)2
• 3Al2(SO4)3

K 2 C2H3O2 2
Mg 3 PO4 2
Al 6 SO4 9
15
Practice Balancing!

• All the atoms on the reactants side must equal
  all the atoms on the products side
• First, you need to count all the atoms
• H2 O2 ? H2O

H
2
2
O
2
1
16
How to Balance

• Are they equal?
• Add coefficients to change the number of atoms
• H2 O2 ? H2O

Balanced!! Good job ?
2
2
H
2
4
2
4
O
2
1
2
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Balancing Example

• Aluminum and copper(II) chloride react to form
  copper and aluminum chloride.

3
3
2
2
1 1 1 1 2
3
2 ?
? 2 ? 6
Good Job!
3 ? 6 ?
? 3
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Practice
Great!
2

• Sn HF ? SnF2 H2
• Cu AgNO3 ? Cu(NO3)2 Ag

2 2
1 2 2
1 1 1
Sn H F
2
2
2 2
1 1 1
1 1 2
Cu Ag NO3
2
Good Job!
19
III. Types of Chemical Reactions

• Ch. 9 Chemical Reactions

20
Combustion

• The burning of any substance in O2 to produce
  heat, usually a hydrocarbon
• Products are always carbon dioxide and water vapor

CXHY O2 ? CO2 H2O
CH4(g) 2O2(g) ? CO2(g) 2H2O(g)
21
Synthesis (composition)

• the combination of 2 or more elements or
  compounds to form a more complex compound
• Basic Form A X ? AX

22
Examples of Synthesis Reactions

• 1) Metal    oxygen  ?   metal oxide
• EX. 2 Mg(s)    O2(g)  ?    2 MgO(s)
• 2) Nonmetal    oxygen  ?   nonmetallic oxide
• EX. C(s)    O2(g)  ?    CO2(g)
• 3) Metal nonmetal  ?    salt
• EX. 2 Na(s)    Cl2(g)  ?    2 NaCl(s)
• 4) A few nonmetals combine with each other.
• EX. 2 P(s)    3 Cl2(g)  ?    2 PCl3(g)

23
Decomposition

• A single compound breaks down into its component
  parts of simpler compounds
• Basic form AX ? A X

24
Examples of Decomposition Rxns

• 1) Some oxides, when heated, decompose.
• EX. 2 HgO(s)  ?    2 Hg(l)    O2(g)
• 2) Some are produced by electricity.
• EX. 2 H2O(l)  ?    2 H2(g)    O2(g)
  EX. 2 NaCl(l)  ?    2 Na(s)    Cl2(g)

25
Single Replacement

• a more active element takes the place of another
  element in a compound and sets the less active
  one free
• Basic Form A BX ? AX B or
  AX Y ? AY X
• Examples of replacement reactions
• 1) Replacement of a metal in a compound by a more
  active metal.
• Fe(s)    CuSO4(aq) ?  FeSO4(aq)    Cu(s)

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Examples of Single Replacement Rxns

• 2) Replacement of hydrogen in water by an active
  metal.
• Note it is helpful to think of water as HOH
  (HOH-)
• 2 Na(s)  2 H2O(l)  ? 2 NaOH(aq)    H2(g) or
  2 Na(s) 2 HOH(l)
  ? 2 NaOH(aq)    H2(g)
• Mg(s) 2 H2O(g)  ?   Mg(OH)2 (s)    H2(g)
  or Mg(s) 2 HOH(g)  ?
   Mg(OH)2 (s)    H2(g)

27
Examples of Single Replacement Rxns

• 3) Replacement of hydrogen in acids by active
  metals.
• EX. Zn(s)    2 HCl(aq)  ?    ZnCl2(aq)    H2(g)
  
• 4) Replacement of nonmetals by more active
  nonmetals.
• EX. Cl2(g)    2 NaBr(aq)  ?    2 NaCl(aq)  
   Br2(l)
• If
• EX. Br2(g)    2 NaCl(aq)  ?    No Reaction!

28
Single Replacement

• NOTE Refer to the activity series for metals and
  nonmetals to predict products of replacement
  reactions.
• If the free element is above the element to be
  replaced in the compound, then the reaction will
  occur.
• If it is below, then no reaction (NR) occurs.

29
Double Replacement (Ionic)

• ions in two compounds change partners cation
  of one compound combines with anion of the other
• Occurs between ions in aqueous solution. A
  reaction will occur when a pair of ions come
  together to produce at least one of the
  following
• a precipitate (s)
• water (l)
• a gas (g)
• Basic form AX     BY   ?    AY    BX
• Complex form ABX CDY ?
  ADY CBX

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Examples of Double Replacement Rxns

• 1) Formation of precipitate (s)
• NaCl(aq)  AgNO3(aq) ? NaNO3(aq) AgCl(s)
• BaCl2(aq)  Na2SO4(aq) ? 2 NaCl(aq) BaSO4(s)
• 2) Formation of water (l). (If the reaction is
  between an acid and a base it is called a
  neutralization reaction.)
• HCl(aq)    NaOH(aq)  ?  NaCl(aq)    HOH(l)
• 3) Formation of a gas (g).
• HCl(aq)    FeS(s)  ?    FeCl2(aq)    H2S(g)

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4) Formation of a product which decomposes.

• Ex CaCO3(aq) 2 HCl(aq) ? CaCl2(aq)
  H2CO3 (aq) ?    CaCl2 (aq)    CO2
  (g)    H2O(l)
• Ex 2 NaOH(aq) (NH4)2S (aq)  ? Na2S (aq) 2
  NH4OH (aq)   ? Na2S (aq)   2 NH3 (g)
   H2O(l) 

32
Examples of unstable products (you need to know
these three)

• A) H2CO3(aq)  ?   CO2(g) H2O(l)
• Carbonic acid, as in soft drinks, spontaneously
  decomposes when it is formed.
• B) NH4OH(aq)  ?   NH3(g) H2O(l)
• Ammonium hydroxide produces the odor of ammonia
  gas because whenever it is formed it
  spontaneously decomposes into ammonia water.
• C) H2SO3(aq)  ?   SO2(g) H2O(l)
• Sulfurous acid, as a bleaching agent and
  disinfectant, spontaneously decomposes when it is
  formed.

33
Double Replacement

• NOTE Use the solubility rules to decide whether
  a product of an ionic reaction is insoluble in
  water and will thus form a precipitate. If a
  compound is soluble in water then it should be
  shown as being in aqueous solution, or left as
  separate ions.
• If it products are both in solution (aqueous)
  then
• No Net Ionic Reaction occurs!

34
Aqueous Rxns Double Replacement

• Ex in the reaction involving the ionic compounds
  silver nitrate and potassium chloride we have
• AgNO3 (aq) KCl (aq) ? AgCl (s) KNO3 (aq)
• The driving force for double replacement
  reactions is the removal of ions from solution

35
Precipitation Reactions

• Double replacement rxns that result in an
  insoluble precipitate
• Solubility amount of a substance that can be
  dissolved in a given quantity of water
• The solubility of an ionic compound determines
  whether it will precipitate or not.
• The reaction of KI and Pb(NO3)2
• 2KI(aq) Pb(NO3)2(aq) ? PbI2(s) 2KNO3 (aq)
• The PbI2 is an insoluble ionic compound that will
  precipitate out of the solution

36
Precipitation Reactions

• Can we predict whether an ionic compound will be
  soluble or not?
• If an ionic compound is insoluble it means that
  neighboring ions have an attraction for each
  other that is greater than the attraction of
  water for the ions
• Unfortunately, there are no clear rules for
  solubility based on physical properties of ions.
• General behaviors of certain ions are observed by
  consulting A Solubility Table!

37
Neutral Molecular Compounds

• Even though the neutral molecular compound may be
  soluble in aqueous solution, its formation is
  essentially irreversible.
• Thus, ions are effectively removed from solution
  by this irreversible process
• The neutralization reaction of HCl and NaOH
• HCl (aq) NaOH(aq) ? H2O(l) NaCl(aq)
• The formation of the covalent compound (H2O) from
  the proton and hydroxide ions is essentially
  irreversible and drives the double replacement
  reaction (even though we would consider H2O to be
  "highly soluble" in H2O).

38
Gas Formation in Double Replacement

• When a double replacement reaction involves the
  formation of a gas ( the gas is not soluble in
  H2O) the loss of the gas can drive the double
  replacement reaction (i.e. the gas is lost -
  therefore, it is an irreversible process) 
• Gasses that can form from ionic compounds
  include
• CO2 (carbon dioxide)
• H2S (hydrogen sulfide - smells like rotten eggs)
• NH3 (ammonia).

39
Formation of CO2 from carbonic acid

• The bicarbonate ion (HCO3-) can combine with a
  proton to produce carbonic acid (H2CO3)
• HCl (aq) NaHCO3 (aq) ? NaCl (aq) H2CO3 (aq)
• Carbonic acid in water is unstable
  spontaneously decomposes to form water and carbon
  dioxide gas
• H2CO3 (aq) ? H2O (l) CO2 (g)
• The carbon dioxide gas is lost, thus the
  formation of carbonic acid is irreversible and
  drives the double replacement reaction. The
  reaction is thus
• HCl (aq) NaHCO3 (aq) ? NaCl (aq) H2O (l)
  CO2 (g)

40
Gases produced in Spontaneous Rxns

• The three typical gasses produced take place with
  the following spontaneous reactions (only when
  they appear on the product side)
• KNOW THEM
• H2CO3(aq)  ?   CO2(g) H2O(l)
  
• NH4OH(aq)  ?   NH3(g) H2O(l)
• H2SO3(aq)  ?   SO2(g) H2O(l)

41
Net Ionic Reactions - DR

• Complete ionic equation equation that shows all
  dissolved ionic compounds as dissociated free
  ions
• Net ionic equation Equation for a reaction in
  solution that shows only those particles that are
  directly involved in the chemical change
• Spectator Ion an ion that appears on both sides
  of an equation and is not directly involved in
  the reaction

42
Neutralization of Nitric Acid KOH

• The molecular equation would be
•  HNO3 (aq) KOH (aq) ? KNO3 (aq) H2O (l)
• The complete ionic equation would be
•   H1 (aq) NO3 -1 (aq) K1 (aq) OH -1 (aq)
  ?
• K1 (aq) NO3 -1 (aq) H2O (l)
• The net ionic equation would therefore be
•   H1 (aq) OH -1 (aq) ? H2O (l)
• Spectator Ions NO3 -1 (aq) K1 (aq)

43
Net Ionic Reactions! - DR
Molecular (or Full) Equation

• AgNO3(aq) KCl(aq) ? AgCl(s) KNO3(aq)

CIE
Ag(aq) NO3 (aq) K(aq) Cl(aq) ? AgCl(s)
K(aq) NO3(aq)
Ag(aq) Cl(aq) ? AgCl(s)
NIE
K(aq) and NO3(aq)
SI
44
Net Ionic Reactions! - DR
Molecular (or Full) Equation

• Pb(NO3) 2(aq) 2LiCl(aq) ? PbCl2 (s)
  2LiNO3(aq)

CIE
Pb2(aq)2NO3(aq)2 Li(aq)2Cl(aq)?PbCl2(s)2Li
(aq)2NO3(aq)
Pb2(aq) 2Cl(aq) ? PbCl2 (s)
NIE
Li(aq) and NO3(aq)
SI