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General Chemistry
CUES102
Course Synopsis
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Fundamentals, Stoichiometry;
Atoms-molecules-ions,
Mass relationships in chemical reactions;
Reactions in Aqueous Solution;
Gases;
Intermolecular forces: liquids-solids;
Quantum theory and the electronic structure of atoms;
Periodic relationships among the elements;
Chemical bonding: Basic concepts, Molecular geometry and hybridisation of
atomic orbitals;
Thermochemistry;
Rates of Reactions;
Chemical Equilibrium;
Acids and Bases; Acid-Base and Solubility Equilibria;
Atmospheric chemistry;
Entropy, free energy & equilibrium;
Electrochemistry;
Organic Chemistry, synthetic and natural organic polymers.
Coursework
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Tutorial Quiz
Assignments
Tests
Practicals
Textbooks
• Chemistry, 9th edition, Raymond Chang, 2007,
McGraw-Hill Company
Study of Chemistry
• Specialised vocabulary
– Electronic, quantum leap, equilibrium, catalyst, chain reaction, critical mass,...
• A cook – a practising Chemist
– Leavening of bread using baking powder
– Squeeze lemon juice over sliced pears to prevent them from browning or over
fish to minimize odour
– Adding vinegar to water in which you are to poach eggs
– Adding meat tenderizer to a pot roast
– Using a pressure cook
• Chemist – study of matter and the changes it undergoes
– looks at the macroscopic world and visualize the particles and events at the
microscopic/atomic world (nanotechnology)
– e.g. Rust
• Modern tools used to probe the microscopic world
Scientific Method
• A systematic approach to research
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Define the problem
Perform experiments
Make careful observations
Record information or data about the system
Interpretation (to try & explain the observed phenomenon)
Hypothesis formulation based on the data obtained (tentative explanation for
a set of observations)
– Further experiments devised to test the hypothesis
• Data obtained may be:
– Qualitative: general observation about the about the system
– Qualitative: comprising number obtained by repeated measurements of the
system
• Information is summarised in a concise way through a law: a concise
verbal or mathematical statement of a relationship between phenomena
that is always the same under the same conditions
• Theory: a unifying principle that explains a body of facts and/or those laws
that are based on them
Scientific Method
Scientific Method
Classification of matter
Three states of matter
States of matter
Physical and chemical properties of matter
Measurement
SI Base Units
Prefixes used with SI Units
Mass & Weight
• Amount of matter in an object Vs the pull of the
force of gravity on an object
• Mass is independent of location while weight is
dependent
• Mass units are kilograms, kg, but grams commonly
used in the lab.
Mass and Weight
Volume
• Volume has SI-derived units of cubic meter (m3),
• Again cubic centimeters (cm3 or mL) or cubic decimeter (1
dm3= 1 L) are commonly used in chemistry
1 cm3 = (1x10-2 m)3 = 1x10-6 m3
1 dm3 = (1x10-1 m)3 = 1x10-3 m3
1 L = 1000 mL
= 1000 cm3
= 1 dm3
Density
• Mass per unit volume (d = m/V)
• Density is kg/m3 in derived SI units but g/cm3 is
commonly used:
1 g/cm3 = 1 g/mL = 1000 kg/m3
1 g/L = 0.001 g/mL
1. A piece of platinum metal with density of 21.5 g/cm3 has a volume of 4.49 cm3. What
is its mass?
2. The density of sulfuric acid in a certain car battery is 1.41 g/mL. Calculate the mass of
2.42 x 10-6 m3 of the liquid.
Temperature
Temperature
• The melting point of lead is 327.5 °C. What is the
melting point in the Farheinheit scale?
• The boiling point of ethanol measured by a
student in America was found to be 172.9 °F.
Calculate the boiling point in the degrees Celsius
scale.
• Convert 77 K (the boiling point of liquid nitrogen)
to degrees Celsius
Handling numbers: Scientific Notation
• Use of the scientific notation: N x 10n
• Numbers in chemistry tend to be much larger, and much smaller,
than numbers in “every day life”
• Express the following in scientific notation:
(a) 602214179000000000000000 (number of atoms in 12 g of
carbon)
(b) 0.000000000139 m (Length of the bond between two carbon
atoms in Benzene)
Scientific Notation
• Find the answers to the following:
(a) (7.4 x 103) + (2.1 x 103) =
(b) (4.31 x 104) + (3.9 x 103) =
(c) (2.22 x 10-2) - (4.10 x 10-3) =
(d) (4.0 x 104) x (5.0 x 102) =
(e) (4.0 x 10-5) x (7.0 x 103) =
(f) (6.9 x 107) ÷ (3.0 x 10-5) =
Significant figures
• Guidelines
– Non-zero integers always count as significant figures e.g. 845 cm, 134.2 kg
– Zeros between non-zero digits are significant, e.g. 101 g/mL; 0.201 g; 30807
kg.
– Zeros to the left of the first non-zero digit are not significant, e.g. 0.000325
mol
– If a number is greater than 1, then all zeros written to the right of the decimal
point count as significant, e.g. 3.00 A; 10.0 joules; 0.3002 mL; 10.507 g; 4.60
W.
– For number without decimal points, trailing zeros may or may not be
significant, e.g. 40 m, 400 cm, (40), (100), 4.0 x 102 nm, 6.00 x 10-5 mol.
Significant figures
• Guidelines
– In addition and subtraction, the answer cannot have more digits to the
right of the decimal point than either of the original numbers
– In multiplication and division, the number of significant figures in the
final product or quotient is determined by the original number that
has the least number of significant figures
• Example: calculate the volume of a box that is 34.49cm long, 23.0cm wide, and
15 cm high
• Example: calculate the density of a piece of wood with a mass of 25.6 g and a
volume of 20.265 cm3.
– Exact numbers obtained from definitions or counting indivisible
things, e.g. 5 cars , 3 molecules, 35 atoms have an infinite number of
significant figures:
• Exact definitions e.g. 1 L = 1000 mL = 1 x 103 cm3 ; 1 in = 2.54 cm
• If we have 9 objects and each weighs 5.0 kg, then the total mass of the objects
is 45 kg
Significant figures
• Guidelines
– Types of zeros
• Leading zeros e.g. 0.0025
• Captive zeros e.g. 1.0085
• Trailing zeros e.g. 100.0
• Practice Exercise:
– Carry out the following arithmetic operations and round off the
answers to the appropriate number of significant figures:
(a) 26.5862 L + 0.27 L,
(b) 9.1 kg - 4.682 g,
(c) 7.1 x 104 dm X 2.2654 x 102 dm,
(d) 6.54 g ÷ 86.5542 mL,
(e) (7.55 x 104 m) – (8.62 x 103 m)
Dimensional Analysis
• Procedure for conversion between units
– To convert from one unit to another use the
equivalence statement that relates two units.
• e.g. 1 cm3 = 1 mL, 1 kg = 1000 g, 1 m = 100 cm
– Derive the appropriate unit factor by looking at
the direction of the required change (to cancel the
unwanted units).
– Multiply the quantity to be converted by the unit
factor to give the quantity with the desired units
Conversion between units
• Example 1: Convert 57.8 meters to centimeters
1. To convert from one unit to another use the equivalence statement
that relates two units.
• ? cm = 57.8 m
• 1 cm = 1x10-2 m
2. Derive the appropriate unit factor by looking at the direction of the
required change (to cancel the unwanted units).
3. Multiply the quantity to be converted by the unit factor to give the
quantity with the desired units.
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Conversion between units
• Example:
– Calculate the number of cubic centimeters in 6.2
m3.
? cm3 = 6.2 m3
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Therefore, ? cm3 =
= 6.2 x 106 cm3
The answer is conveniently expressed in scientific notation
Conversion between units
• Example
– The density of gold is 19.3 g/cm3. Convert the
density to units of kg/m3.
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Therefore, ? kg/m3 =
= 1.93 x 104 kg/m3
Summary
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Study of chemistry involves three basic steps: observation, representation and
interpretation
– Observation: measurements in the macroscopic world
– Representation: use of shorthand symbols and equations for communication
– Interpretation: based on atoms and molecules (microscopic world)
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Classification and properties of matter (Substances have unique physical and
chemical properties)
– The simplest substances in chemistry are elements. Compounds are formed by the
combination of atoms of different elements
– Elements are arranged in a periodic table according to their related properties
– All substances can in principle exist in 3 states: solid, liquid and gas. Interconversion effected
by temperature change
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Measurement
– Length, volume, temperature and density
– SI units are used to express all physical quantities in all sciences
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Handling numbers
– Scientific notation (addition and subtraction, multiplication and division)
– Significant figures
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Dimensional analysis
– Interconversion between units of measurement