Quiz
0. How many ways are commonly used for denoting a molecule/ion?
1. Write the molecular formula of the following compounds:
Sodium Chloride
Calcium Phosphate
Copper (II) Chloride
Nitric Oxide
Hydrocyanic acid
Potassium Permagnate
Sodium Borate Decahydrate
2. Write the name of the following compounds:
CuSO4．5H2O
Na2ClO2
CuSO
3
3. Write the meaning of the following prefixes:
Mono-, Tetra-, Penta-, Hexa-, Octa-, Deca-
Answer
0. How many ways are commonly used for denoting a molecule/ion?
(1) Molecular structural formula, (2) Ball-and-stick model, (3) Tube structure, (4) Space-filling representation.
1. Write the molecular formula of the following compounds:
Sodium Chloride NaCl
Calcium Phosphate Ca3(PO4)2
Copper (II) Chloride CuCl2
Nitric Oxide
HNO2
Hydrocyanic acid HCN
Potassium Permagnate KMnO4
Sodium Borate Decahydrate NaBO3.10H2O
2. Write the name of the following compounds:
CuSO3
CuSO4．5H2O
Na2ClO2
Copper (II) sulfate
Sodium chlorite
Copper (II) sulfate pentahydrate
3. Write the meaning of the following prefixes:
Mono-1, Tetra-4, Penta-5, Hexa-6, Octa-8, Deca-10
Chapter 3
Chemical Reactions
The process that brings about a chemical change
The carbon cycle is evident in fossils
like this one, which are found in
limestone, a form of calcium
carbonate. The carbon atoms in
limestone were once part of carbon
dioxide molecules in the
atmosphere. They were then taken
up in the shells of marine organisms.
When the organisms died, the shells
settled to the bottom of the ocean
and became compacted into
limestone. Millions of years later, we
dig up the limestone and use it to
construct buildings. Some of the
limestone is also heated to make
quicklime in a process that releases
the carbon atoms once again to the
atmosphere as carbon dioxide.
Skeletal Equation
ReactantsProducts
Staring materials
Substances formed in a
chemical reaction
A reagent is a reactant only when it is being used in a
particular reaction.
sodium+watersodium hydroxide+hydrogen
Na+H2ONaOH+H2
Skeletal equation
Chemical Equations
Law of Conservation of Mass: Atoms are neither created nor
destroyed in a chemical reaction.
Na+H2ONaOH+H2
2Na+2H2O2NaOH+H2
Balanced expression of chemical
reaction=chemical equation
Stoichiometric coefficients which
give the molar ratios of the reactants
and products
Molecules
Reaction Conditions
States: gas(g), liquid(l), aqueous(aq), solid(s)
2Na(s)+2H2O(l)2NaOH(aq)+H2(g)
Temperature
CaCO3(s)
CaO(s)+CO2(g)
High temperature
Other conditions: pressure, reaction time, catalysts …
Balancing Chemical Equations
H2+O2H2O
!
Danger!
H2+O22H2O
2H2+O22H2O
2H2(g)+O2(g)2H2O(l)
H2+O2H2O2
Change the
stoichiometric
coefficients only!
2H+OH2O
H2+1/2O2H2O
Balancing A Chemical Reaction
C4H10+O2CO2+H2O
C4H10+O24CO2+5H2O
C4H10+(13/2)O24CO2+5H2O
2C4H10+13O28CO2+10H2O
2C4H10(g)+13O2(g)8CO2(g)+10H2O(l)
!
Figure 3.6
When solutions of silver nitrate and potassium chromate are mixed, a
precipitate of red silver chromate, Ag2CrO4, forms.
Precipitation reaction
Insoluble substance
potassium chromate+silver nitrate
Silver chromate+potassium nitrate
K2CrO4(aq)+2AgNO3(aq)Ag2CrO4(s)+2KNO3(aq)
Soluble substance
Figure 3.7
These two beakers contain solutions with different concentrations of
the same solute. The lighter color of the solution on the left (a) shows
that the solute is less concentrated than in the solution on the right
(b). In the molecular-level view, we see that there are more solute
particles in a given volume of the more concentrated solution.
solvent
Dissolve
Solution=solvent+solute
Concentration
The amount of solute molecules in a given
volume of solution
C
C

solute amount
solution volume
(g/ml, mol/l, kg/m ...)
105.0 g NaOH
25.0 ml aqueous solution
4.2 g / ml
42 g/mol
3
 4.2(g/ml)
 0.1 (mol/ml)  10 (mol/l)
2
Figure 3.8
Sodium chloride consists of sodium ions and chloride ions. When it
is added to water (left), the ions separate and spread throughout the
solvent (right). The solution consists of water molecules, sodium
ions, and chloride ions. There are no NaCl molecules present at any
stage. The overlays show only the solute.
Hydration of ions
Electrolyte: a substance that dissolves to give a solution that contains ions.
Strong electrolytes: mostly ions. Weak electrolytes: mostly molecules
Nonelectrolytes: no ions
Nonelectrolyte
Nonelectrolyte
Figure 3.9 Pure water is a poor conductor of electricity, as shown by
the almost imperceptible glow of the bulb in the circuit (a). However,
when ions are present, as in an electrolyte solution, the solution does
conduct. The ability of the solution to conduct is low if it is a weak
electrolyte (b) but significant if it is a strong electrolyte (c), even if the
solute concentration is the same in each case.
Figure 3.10
In a nonelectrolyte solution, the solute remains as molecules and
does not break up into ions. Methanol, CH3OH, is a nonelectrolyte
and is present as molecules when it is dissolved in water.
Figure 3.11 The formation of a silver chloride precipitate occurs
immediately as silver nitrate solution is added to a solution of sodium
chloride.
Strong electrolyte
AgNO3(aq)+NaCl(aq)AgCl(s)+NaNO3(aq)
Figure 3.12 A series of scenes in a solution of sodium chloride. A
sodium ion and a chloride ion move together, linger near each other
for a time because of the attraction of their opposite charges, and
then move apart. The loose, transient association of oppositely
charged ions is called an ion pair. The solution is shown both with
solvent molecules, for realism, and without, for clarity.
Figure 3.13 In water, ions are hydrated; that is, they are surrounded
by a cluster of water molecules bound loosely to the ion. Note that a
hydrated cation (a) is surrounded by water molecules oriented so that
the O atom is closest to the ion, whereas a hydrated anion (b) has
water molecules attached through their hydrogen atoms. The number
of hydrating molecules depends on the size of the ion, but for most
ions it is approximately six.
Hydration of ions
Figure 3.14
In this precipitation reaction, yellow lead(II) chromate is formed when
lead(II) nitrate solution is added to a solution of potassium chromate.
Pb(NO3)2(aq)+K2CrO4(aq)PbCrO4(s)+2KNO3(aq)
Quiz
• Explain the following concepts:
(1) Electrolyte (2) Hydration
• What is the real meaning of “aq” in a
chemical equation?
Net Ionic Equations
AgNO3(aq)+NaCl(aq)AgCl(s)+NaNO3(aq)
Complete ionic equation






Ag (aq)  NO (aq)  Na (aq)  Cl (aq)  AgCl(s)  NO (aq)  Na (aq)
3
3
Net ionic equation
spectator ions


Ag (aq)  Cl (aq)  AgCl(s)
Figure 3.15
Two depictions of a precipitation reaction that results when the ions
in two electrolyte solutions are mixed (left beakers). The top right
beakers show the fate of all four types of ions. By imagining the ionic
reaction without the spectator ions (bottom right beakers), we can
focus on the essential process described by the net ionic equation.
Figure 3.16
How to write a net ionic equation. Write the balanced overall equation
(top), Then show all ionic solutes as separate ions in the complete
ionic equation (second line), and delete the spectator ions. The result
is the net ionic equation (bottom).
Figure 3.17 Another example of a precipitation reaction. This time, a
solution of mercury(I) nitrate is being added to a solution of
potassium iodide, and the insoluble product, mercury(I) iodide, is
precipitated. Notice that a yellow color forms first. Mercury(I) iodide
has two solid forms. The yellow form precipitates first but is soon
converted to the more stable orange form.


Hg (aq)  I (aq)  HgI(s)
Hg (aq)  NO(aq)  K (aq)  I (aq)  HgI(s)  NO (aq)  K (aq)
3
3
Figure 3.18
The shape of this shell is a result of a precipitation reaction in which
the shellfish secreted calcium ions at certain points on its surface.
The calcium ions reacted with carbonate ions in the surrounding
water. The colors of the shell are due to iron impurities that were
captured in the solid as it formed.
Ca2  (aq)  CO2-(aq)  CaCO (s)
3
3
The Reactions of Acids and Bases
• HCl (hydrochloric acid)
• CH3COOH (Acetic acid)
H   H2O(l)  H3O (aq)
• NaOH (sodium hydroxide)
• NH4OH(ammonium hydroxide)

Arrhenius acid

HCl(g)  H2O(l)  H3O (aq)  Cl (aq)
Arrhenius base
CH3COOH(aq)  H2O(l)  H3O  (aq)  CH3COO (aq)


NaOH(aq)  H2O(l)  Na (aq)  OH  H 2O(l)
Strong/Weak Acids
HCl (hydrochloric acid)
(Almost completely ionized in aqueous solution)
HCl(g)  H2O(l)  H3O (aq)  Cl (aq)
CH3COOH (Acetic acid)
(incomplete ionized in aqueous solution)


CH3COOH(aq)  H2O(l)  H3O (aq)  CH3COO (aq)
Strong/Weak Bases
• NaOH (sodium hydroxide)
(Almost completely ionized in aqueous solution)


NaOH(aq)  H2O(l)  Na (aq)  OH  H 2O(l)
•NH4OH(ammonium hydroxide)
(Incompletely ionized in aqueous solution)


NH 4 OH(aq)  H2O(l)  NH 4 (aq)  OH  H 2 O(l)
The strong acids and bases in water
• HBr(aq), HCl(aq), HI(aq), HNO3(aq),
HClO4(aq), HClO3(aq), H2SO4(aq)
• Group 1 hydroxides, Alkaline earth metal
hydroxides
Neutralization
• Base+Acid Salt +Water +(Others)
• HCl(aq)+NaOH(aq)NaCl(aq)+H2O(l)
• 2HNO3(aq)+Mg(OH)2Mg(NO3)2(aq)+2H2O(l)


H  OH  H 2O

H  NH 3  NH

4
Neutralization=proton transfer
Gas-Formation Reactions
!!!
• 2NaCl(s)+H2SO4(l)Na2SO4(s)+2HCl(g)
• FeS(s)+2HCl(aq)FeCl2(aq)+H2S(g)
• CaCO3(s)+H2SO4(aq)CaSO4(s)+H2CO3(
aq)H2O+CO2(g)
The reaction of acids with salts is a proton
transfer reaction that may produce gas or a
compound that decomposes into a gas.
Redox Reactions
6CO2(g)+6H2O(l)C6H12O6(s)+6O2(g)
(photosynthesis reaction)
CH4(g)+2O2(g)CO2(g)+2H2O(l)
(Natural gas reaction)
2Mg(s)+O2(g)2MgO(s)
Mg(s)+Cl2(g)MgCl(s)
Zn(s)+2HCl(aq)ZnCl2(aq)+H2(g)
Anything in common?
Figure 3.26
An example of an oxidation reaction: magnesium burning brightly in
air. Magnesium also burns brightly in water and in carbon dioxide;
consequently, magnesium fires are very difficult to extinguish.
e

2Mg(s)+O2(g)2MgO(s)
Oxidized (reducing agent)
Reduced (Oxidizing agent)
Figure 3.27
When bromine is poured on red phosphorus, a vigorous reaction
takes place. In the reaction phosphorus is oxidized and bromine is
reduced.
P(s)+5Br(s)PBr5(s)
e

Oxidized (reducing agent)
Reduced (Oxidizing agent)
They Are All Redox Reactions
6CO2(g)+6H2O(l)C6H12O6(s)+6O2(g)
(photosynthesis reaction)
CH4(g)+2O2(g)CO2(g)+2H2O(l)
(Natural gas reaction)
2Mg(s)+O2(g)2MgO(s)
Mg(s)+Cl2(g)MgCl(s)
Zn(s)+2HCl(aq)ZnCl2(aq)+H2(g)
Figure 3.28
The common oxidation numbers of main-group elements. Notice the
tendency of elements in the same group to assume the same
oxidation number.
Figure 3.29 : How to determine an oxidation number. Each atom is imagined to be a
separate ion. Certain ions are assigned charges by using the rules in Toolbox 3.3, and the
charge on the central atom is determined by considering the overall charge on the
species. (a) Oxide “ions” in an oxoanion are given the charge of 2; because there are
four oxygen atoms and the overall charge on the anion is 2, the charge on the central
atom must be 6. (b) This molecule has three chlorine atoms with oxidation numbers of 
1, an oxygen atom (2), and a hydrogen atom (1). The sum of these oxidation numbers is
4 and the overall charge on the molecule is 0. Thus, the central atom must have an
oxidation number of 4.
Determine Oxidation Number
SO2
X+2(-2)=0x=4
SO
2
4
x+4(-2)=-2x=6
Figure 3.30 When a strip of zinc is placed in a solution that contains
Cu2 ions, the blue solution slowly becomes colorless and copper
metal is deposited on the zinc. In this redox reaction, the zinc metal is
reducing the Cu2 ions to copper and the Cu2 ions are oxidizing the
zinc metal to Zn2 ions. (a) The reaction. (b) A visualization of the
process.
2
2
Zn(s)  Cu (aq)  Zn (aq)  Cu(s)
Figure 3.31 (a) Copper reacts slowly with dilute nitric acid to give blue
Cu2 ions and the colorless gas nitric oxide, NO. (b) When copper reacts
with concentrated nitric acid, nitrogen dioxide, NO2, is produced instead
of NO. The blue solution is turned green by this brown gas.
Cu(s)  8H (aq)  2NO3- (aq,dilute)  Cu 2 (aq)  2NO(g)  4H2O(l)
Cu(s)  4H (aq)  2NO3- (aq,concent.)  Cu 2 (aq)  2NO2 (g)  2H2O(l)
Figure 3.32
Aluminum reacts vigorously with hydrochloric acid to form soluble
aluminum chloride and water.
2Al(s)  6HCl(aq)  2AlCl 3 (aq)  3H 2 (g)
Al(s)  H  (aq)  Al 3 (aq)  H 2 (g)
Mass balance :
Al(s)  2H  (aq)  Al 3 (aq)  H 2 (g)
Charge balance :
2Al(s)  6H  (aq)  2Al 3 (aq)  3H 2 (g)
Case Study 3
Astronauts on the space shuttle must change the canisters of lithium hydroxide
daily. Here, Sidney Gutierrez carries out the task. Two canisters are used, and one
is changed every 12 hours so that the capacity to remove carbon dioxide remains
stable.
A Better Solution:
4KO2(s)+2CO2(g)K2CO3(s)+3O2(g)
CO2(g)+2H2(g)C(s)+2H2O(l)
2H2O(l)2H2(g)+O2(g)
(Each element can be recovered and
reused!)
CO2(g)+2LiOHLi2CO3(s)+H2O(l)
Figure 3.33 The three main types of chemical reactions discussed in
this chapter can be distinguished by the type of change taking place.
(a) In a precipitation reaction, ions mix and one combination of ions
is insoluble. (b) In a neutralization reaction, hydrogen ions are
transferred from an acid to a base. (c) In a redox reaction, electrons
are transferred from a reducing agent to an oxidizing agent.
Figure 3.34 : We can predict the products of a reaction by examining
the reactants. (a) Two soluble salts may form a precipitate. (b) An acid
and a hydroxide react to form a salt and water. (c) When two elements
react, a redox reaction generally occurs. A metal and nonmetal react
to form an ionic compound and two nonmetals react to form a
molecular compound. (d) In combustion reactions, organic
compounds react with oxygen to form carbon dioxide and water.
Three Most Important Types of Reactions
• Precipitation
(Soluble salts exchange ionsionic solids
• Proton transfer
(Neutralization, Gas Formation)
• Electron transfer
(Redox Reaction)
Assignment for Chapter 3
17,25,33,37,43,51,62