CCST 4124

Chemistry for Life Science

Chapter 3
Chemical Reactions
Dr. T.H. YUI 1
Chemical Formula
Chemical Formula
➢ Represents the composition of elements inside the molecule.

➢ Usually arranged starting from low e-ve to high e-ve elements.

➢ Example: H2O (There are two Hs and one O in a H2O molecule.)
➢ e-ve: H = 2.1, O = 3.5
➢ Not necessary
➢ Example: C2H5OH

➢ For neutral ionic compounds, the charges of ions are usually NOT
labelled.
➢ Example: NaOH
➢ Composition: Na+ and OH-
➢ Example: H2SO4
➢ Composition: two H+ and one SO42- 2
Chemical Equation
Chemical Equation
➢ An equation shows the formulae of reactants and products.
➢ The state of matter is labeled at the bottom right corner of the species.
i. l: liquid
ii. g: gas
iii. s:solid
iv. aq: aqueous solution (for something that is dissolved in water)
➢ A chemical equation should be balanced in order to show the relative
number of substances involved in a chemical reaction.
Na(s) + Cl2(g) → NaCl(s) (X)
2Na(s) + Cl2(g) → 2NaCl(s) (O)
➢ If the reaction is unidirectional(irreversible), species in left-hand
side are reactants and in right-hand side are products.
➢ Chemical reaction could occur in both direction(reversible, will
appear in future chapter Equilibria), which means that reactants
and products are depending on the direction of reaction 3
Acid – Base Chemistry
Acid & Base – Common Definition
➢ pH value: 0-14
➢ pH < 7 = Acid pH = 7 = Neutral pH > 7 = Base

Depend on what you reacting

Acid & Base in Chemistry
➢ Relative terms Eg, water not must be neutral

➢ Many Definitions
➢ Arrhenius Definition
➢ BrØ nsted-Lowry Definition
➢ Lux-Flood Definition
➢ Lewis Definition
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Arrhenius Definition of Acid
Arrhenius Definition – Acid
➢ A species produces H+ in water
➢ e.g. HCl(aq), H2SO4(aq)…etc

H+ vs H3O+
➢ Recall: H+ → Small size → High charge density → Unstable →
Not exist alone in water
➢ H+ associated with H2O to form hydronium ion (H3O+)
H+(aq) + H2O(l) → H3O+(aq)
➢ Therefore, the equation for ionization of hydrochloric acid in
water should be written as:
HCl(aq) + H2O(l) → H3O+(aq) + Cl-(aq)
NOT
HCl(aq) → H+(aq) + Cl-(aq) 5
Arrhenius Definition of Base
Arrhenius Definition – Base
➢ A species produces OH- in water
➢ e.g. NaOH(aq), NH4OH(aq)…etc

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BrØ nsted-Lowry Definition of Acid –
Base
BrØ nsted-Lowry Definition
➢ An acid-base reaction involves transfer of a proton (hydrogen ion)
from an acid to a base.

BrØ nsted-Lowry Definition – Acid

➢ Proton (Hydrogen ion) Donor

BrØ nsted-Lowry Definition – Base

➢ Proton (Hydrogen ion) Acceptor

Examples:
HCl(aq) + H2O(l) → H3O+(aq) + Cl-(aq)
CO32-(aq) + H2O(l) → HCO3-(aq) + OH-(aq)
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Strong Acid
Strong Acids Strong proton donor

Acids that dissociate completely in water

HA(aq) + H2O(l) → H3O+(aq) + A-(aq)

Common strong acids:

Hydrochloric acid HCl
Hydrobromic acid HBr
Nitric acid HNO3
Sulphuric acid H2SO4
Perchloric acid HClO4
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Strong Base Strong proton accepter

Strong Bases
Base that ionize completely in water
MB(aq) → M+(aq) + B-(aq)

Common strong bases:

Group I Metal Hydroxides NaOH, KOH…etc
*Group I Metal Oxides Na2O, K2O…etc
Calcium Hydroxide Ca(OH)2

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Acid-Base Reaction and Neutralization
Acid-Base Reaction
➢ Broad definition
➢ Many acid-base theories
➢ Reaction between acid and base
Neutralization
➢ Specific acid-base reaction
➢ Acid reacts with equivalent base to create water and a salt

Acid + Base → Salt + Water

Example:
HCl + NaOH → NaCl + H2O
Net ionic equation:
H+ + OH- → H2O 10
Exercise
Predict the products and write down balanced chemical equations
for each of the following reactions:

1. HNO3(aq) + Ca(OH)2(aq) →

2. NaHSO4(aq) + NaOH(aq) →

3. Al(OH)3(s) + H3PO4(aq) →

4. K2O(s) + HCl(aq) →

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Formation of Ionic Compounds
Ionic Bonding
Electrostatic attraction between oppositely charged ions
(i.e.: Cations and Anions)
Compounds formed by ionic bonding are called ionic compound.

Recall Chapter 2 p.10…

Group 1 to Group 3 Metals (except Beryllium, Be) tend to form cations by
losing their outermost electrons in order to obtain stable duplet/octet
structure.

Group 5 to Group 7 Elements tend to form anions by gaining electrons in

order to obtain stable duplet/octet structure.
➢Usually formed between metals and non-metals (NOT necessary)

Examples:
2Na(s) + Cl2(g) → 2NaCl(s)
2Mg(s) + O2(g) → 2MgO(s) 12
Exercise
Predict the products and write down balanced chemical equations
for each of the following reactions:

1. K(s) + I2(s) →

2. Mg(s) + N2(g) →

3. In(s) + O2(g) →

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Formation of Covalent Compounds
Covalent Bonding
Two atoms share one or more pairs of electrons
Each shared pairs of electrons forms one covalent bond

Group 4 to Group 7 Elements tend to form covalent compounds by

sharing electrons in order to obtain stable duplet/octet structure.

Hydrogen (Group 1), Beryllium (Group 2) and Boron (Group 3)

are the three exceptions in their corresponding group. They form
covalent compounds.

➢Usually formed between non-metals and non-metals (again, NOT necessary)

Examples:
C(s) + O2(g) → CO2(g)
2B(s) + 3Cl2(g) → 2BCl3(s) 14
Exercise
Predict the products and write down balanced chemical equations
for each of the following reactions:

1. C(s) + S8(s) →

2. Si(s) + O2(g) →

3. Se(s) + Br2(g) →

4. I2(s) + Cl2(g) →

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Redox Reactions
Redox – REDuction & OXidation
Reduction
➢ Gaining electrons
Oxidation
➢ Losing electrons

Reducing agent
➢ A species losing electrons
➢ Being oxidized
Oxidizing agent
➢ A species gaining electrons
➢ Being reduced

Redox Reaction must involve electron transfer!! 16

Exercise
Label the type of the following reactions:

1. Na → Na+ + e- Oxidation

2. Zn2+ + 2e- → Zn Reduction

3. MnO4- + 8H+ + 5e- → Mn2+ + 4H2O Reduction

Oxidation
4. 2S2O32- → S4O62- + 2e-

5. H2SO4 + Ba(OH)2 → BaSO4 + 2H2O

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Exercise
Label the role of red species in the following reactions:

1. Na → Na+ + e- Reducing agent

2. Zn2+ + 2e- → Zn
Oxidizing agent

Oxidizing agent
3. MnO4- + 8H+ + 5e- → Mn2+ + 4H2O

4. 2S2O32- → S4O62- + 2e- Reducing agent

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Oxidation Number
Oxidation Number
➢ The hypothetical charge of an atom if all of its bonds to different atoms
were fully ionic.
+ve = A Species lost electrons
-ve = A Species gained electrons
0 = Element
➢ Sum of O.N. of a species = Overall charges of compounds/ion

Example:
Mg → Mg2+ + 2e-
➢ O.N. of Mg = 0
➢ O.N. of Mg2+ = +2

Redox Reaction must involve change in oxidation number!!

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Oxidation Number
Rules:
1. For monoatomic ions, O.N. = Charge.
➢ Na+ in Na2O = +1
➢ O2- in Na2O = -2

2. For covalent compounds, higher e-ve species will be more negative one.
➢ C in CO2: +4
➢ O in CO2: -2

3. For polyatomic ions, sum of O.N. of all atoms equals to the charges of ion.
➢ H in OH-: +1
➢ O in OH-: -2

4. One element may have more than one oxidation states.

➢ O in H2O: -2 Same element bond together =0 eg. O =-2 -> -1

➢ O in H2O2: -1 20
Oxidation Number

Transition metals: Various

(Out of syllabus)

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Exercise
Label the oxidation number of all elements in each of the following
species:

1. H2SO4

2. F2O

3. S4O62-

4. NH4+

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Exercise
Which of the followings is/are redox reaction?

1. C(s) + O2(g) → CO2(g)

2. HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

3. I2(aq) + 2S2O32-(aq) → S4O62-(aq) + 2I-(aq)

4. NaCl(aq) → Na+(aq) + Cl-(aq)

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