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CHEMISTRY SEMESTER ONE
COURSE OUTLINE
CHEMISTRY: SEMESTER ONE: COURSE OUTLINE
Welcome to Introduction to Chemistry (Semester 1)! This Outline includes information about the
learning outcomes, delivery, scheduling, assessments and resources required for this course.
COURSE DESCRIPTION
This course is a study of the fundamental principles of chemistry with particular reference to chemical
reactions and equations; energy in chemical systems; atomic and molecular structure; stoichiometry and
chemical equilibrium; properties of gases, liquids, solids and solutions. This theory component is
accompanied by a laboratory program which introduces basic techniques of quantitative and qualitative
chemistry.
This course is designed for students seeking a degree or diploma in a field of science or technology. It is
also suitable as an elective course for General Interest or Arts students.
PREREQUISITES

A previous Chemistry course at the Grade 12 level (Senior-year high school Chemistry (US),
Chemistry 12 or ABE Provincial Level Chemistry (C+ minimum) (Canada) and

University-preparatory Mathematics (Principles of Math 12 or ABE Provincial Level Math (C
minimum) Canada)

A Physics course at the Grade 11 level (Chemistry 11 or ABE Advanced Physics (Canada); Junioryear high school Physics (US) is also recommended.
RESOURCES
Textbook selection varies by instructor and may change from year to year. When this course was
developed (2011) the following textbooks were in use:
Petrucci, Harwood, Herring and Madura. General Chemistry: Principles & Modern Applications.
9th ed. Prentice Hall
Kotz, John C. and Paul M. Treichel (2003). Chemistry & Chemical Reactivity, 5th edition.
Florence, KY: Thomson/Brooks-Cole.
Reading assignments in the Units for this course are based on Petrucci et al. (9th ed.)
A scientific calculator is required, but programmable calculators are not allowed on exams.
A lab kit is required if you are studying the lab component of this course via distance delivery.
An internet connection (preferably high-speed) is required to participate in this course.
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CHEMISTRY SEMESTER ONE
COURSE OUTLINE
Textbooks, calculators, and lab kits will be available through your College’s bookstore. Syllabus, Student
Manual, and all other resources will be available through the online learning environment set up for this
course. Please see instructor’s Syllabus for a complete list of the current required textbooks and other
resources.
ABOUT THE LABS
The theory section of this course is accompanied by a Laboratory component consisting of 8 required
labs. Labs may be available in one (or more) of several different delivery formats:



Entirely by distance: utilizing lab kits and Remote Web-Based Science Laboratory (RWSL)
In one or more concentrated lab weeks/weekends on campus
On a weekly basis in conjunction with on-campus Chemistry course lab schedule
Check with your College to find out which delivery option(s) may be available at your institution. More
information about the lab component of this course is available in the Student Manual for Introduction
to Chemistry (Online): Semester One LABS.
COMPLETION TIME
This course is designed to be completed over a scheduled 15-week period. Students with a weak
background in Chemistry may choose to spend several weeks before the course begins to review units 1
to 4. A time range for each unit is suggested below: your instructor may modify the delivery to
accommodate student strengths and weaknesses.
SUGGESTED STUDENT SCHEDULE (PROGRESS CHART – TO BE COMPLETED BY INSTRUCTOR)
Title
Approx.
weeks
Introduction and
1
Lab
correlation
Assignment
Introduction forum
Review or Unit 1: Basic Concepts of Chemistry
Unit 2: Atoms, Molecules and Ions
0-1
Unit 3: Chemical Reactions
0-1
Unit 4: Stoichiometry
0-1
Unit 5: Energy and Chemical Reactions
1-2
Unit 6: Structure of atoms
1
Unit 7: Structure of atoms and periodic trends
1
Unit 8: Bonding and molecular structure
1-2
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CHEMISTRY SEMESTER ONE
COURSE OUTLINE
Unit 9: Orbitals and Theories of Chemical Bonding
1-2
Unit 10: Gases
1
Unit 11: Intermolecular forces and Liquids
1-2
Unit 12: Solutions and Their Behaviour
1
ASSESSMENT
Exact weighting of course components will vary from institution to institution. The following assessment
profile is suggested:
Assessment item
Theory
% Of total Grade
Assignments and/or quizzes
5-20%
Midterm Test(s)
20-40%
Final Examination
30-40%
Theory total:
80%
Laboratory
Laboratory Reports
15%
Laboratory Test
5%
Laboratory total:
20%
Course Total
100%
CONTACTING YOUR INSTRUCTOR
Your instructor can be contacted by a variety of digital means including email, discussion boards/forums
or messaging within your online classroom. Your instructor will also be available by phone. Online office
hours may be available via live chat and/or conferencing applications (e.g. Adobe Connect, Skype,
Elluminate, etc.).
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COURSE OUTLINE
CHEMISTRY SEMESTER ONE: COURSE UNITS
This course has been organized into 12 units. Each unit includes learning outcomes, course notes
written by your instructor, a reading assignment, and suggested problems for completion and/or for
submission to your instructor. To begin each unit, visit your online classroom and review the list of unit
tasks to be completed.
Information about material to be covered in each unit is detailed below:
UNIT 1: BASIC CONCEPTS OF CHEMISTRY
Overview:
 Chemistry and its methods – laws, hypotheses and theories - goals of science and the ethical and
moral dimension

Classifying matter

Elements and molecules

Compounds and molecules

Physical properties

Physical and chemical change

The tools of quantitative chemistry
Learning Outcomes:
 Understanding differences between hypotheses, laws and theories.

Understanding connection between the kinetic-molecular theory and states of matter

Understanding difference between pure substances and mixtures and the difference between
homogeneous and heterogeneous mixtures

Understanding the importance of and differences between macroscopic and microscopic (atomic)
descriptions of matter

Identifying names and symbols of elements and correct usage of the terms atom, element, molecule
and compound

Knowing commonly used physical properties and understanding difference between intensive and
extensive properties
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
Understanding the difference between physical and chemical change

The tools of quantitative chemistry
o
Understanding the International System of Units

Base quantities and the SI base units, their names and abbreviations

Scientific prefixes
o
Understanding difference between precision and accuracy
o
Understanding experimental error and standard deviation
o
Mathematics of chemistry

Understanding exponential notation

IMPORTANT – entering numbers into calculator using exponential notation

Understanding correct application of significant figures

Using dimensional analysis to solve problems

Graphing and correct labeling of axes

Application of problem solving strategies
UNIT 2: ATOMS, MOLECULES AND IONS
Overview:
 Atomic structure

o
Numbers of proton and neutron within the nucleus and connection to “atomic number”
and “mass number”
o
Isotopes and connection to number of neutrons in the nucleus
o
Atomic masses
o
“Relative atomic mass” otherwise called atomic weight being average atomic mass of
representative sample of the element. It depends on isotopic composition of sample
and atomic masses of each isotope.
The Periodic Table
o
Elements ordered by atomic number
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
o
General features and concepts of groups containing elements with similar properties
o
Main group and transition metals
o
Special place occupied by lanthanides and actinides
o
Division between metals and non-metals
o
Brief overview of periodic properties of elements
Molecules, Compounds and Formulae
o



Formulae and models
Ionic compounds
o
Cations and anions
o
Monatomic ions and the periodic table
o
Polyatomic ions
o
Formulae of ionic compounds- note: compounds are neutral
o
Naming of molecular compounds
Atoms, molecules and the mole
o
“Amount of substance (n)” is an SI base quantity – the SI unit for amount of substance is
the mole abbreviated to mol. The actual number of particles in the mole is an
experimental quantity, Avogadro’s number (NA = 6.022x1023 mol-1).
o
The atomic or molar mass is mass of one mole of atoms or molecules
Describing compound formulae
o
connection between formula of compound and mass of the elements therein
o
connection between formula of compound and weight percent of the elements therein
o
empirical formula from weight percent composition – connection to molecular formula
o
mass spectrometry and molecular formulae
o
hydrated compounds
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Learning Outcomes:
 Understanding that

o
the structure of the atom consists of an incredibly dense, positively charged nucleus
surrounded by negatively charged electrons which occupy most of the space.
o
the nucleus is made up of protons and neutrons each with very similar masses.
o
the chemical properties of the atom are determined by the number of protons in the
nucleus, called the “atomic number”.
o
the “mass number” of an atom is the sum of the neutrons and protons in the nucleus.
o
isotopes of an element contain the same number of protons in the nucleus but different
numbers of neutrons
o
the “atomic mass” is the mass of the atom in question divided by the mass of a carbon12 atom multiplied by 12.
o
that mass spectrometer can be used to determine the atomic mass and isotopic
abundance.
o
atomic masses are very close to integer values but not exactly. This is due to very slight
differences between the masses of a proton and a neutron. The atomic mass is however
slightly less than the sum of the masses of the protons and neutrons due to a “mass
defect” which will be discussed later in the course.
o
“relative atomic mass” (called in the text “atomic weight”) is the average mass of a
representative sample of an element taking into account the atomic masses of its
various isotopes and the relative isotopic abundances.
Be able to perform calculations:
o
Connecting an element to the numbers of protons and neutrons in its nucleus
o
connecting isotopic abundances and atomic masses to relative atomic masses

Understand general features of periodic table

Understand in general terms the properties of the various elements of the table

Understand the concept of ions (cations and anions) monatomic and polyatomic

Understand connection between ionic charge of monatomic ions and the periodic table

Know names, formulae and charges of common polyatomic and monatomic ions
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
Understand connection between ionic solids and the strong electrostatic interactions in the crystal
lattice

Understand the naming of simple molecular compounds according to their formulae

Understand the concept of the mole and its relation to molar mass

Be able to convert between masses, numbers of moles and molar masses

Be able to calculate weight percentages of the various elements in a compound to its molecular
formula.

Be able to determine the empirical formula of a compound from the weight percentage of its
constituent elements and understanding the connection of empirical formula to molecular formula

Understand the use of mass spectrometry to obtain molecular formulae

Understand the concept of hydrated ionic compounds

Be able to calculate numbers of water molecules in an ionic compound from experimental
dehydration data
UNIT 3: CHEMICAL REACTIONS
Overview:
 Balancing chemical reactions

Equilibria and the concept of reactions not going to completion

Terminology of solutions

Dissolution of ionic solids in solvents such as water to give electrically conducting solutions

Dissolution of molecular solids in water to give non-conducting solutions

Concept of weak and strong acids according to the extent of ionization

Concept of net ionic reactions as applied to precipitation and gas-forming reactions

Oxidation-reduction reactions and the concept of oxidation numbers
Learning Outcomes:
 Understand information conveyed in a balanced chemical reaction

Be able to balance chemical reactions – remember must know reactants and products to do this
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
Understand that not all reactions go to completion but result in a “dynamic equilibrium”

Understand nature of chemical reactions in solution and the terms solution, solute and solvent

Understand that when an ionic solid dissolves in water the solid is fully dissociated into separate ions
and that such solutions are electrical conductors.

Understand that solutions of compounds that when dissolved in water do not conduct electricity are
called non-electrolytes.

Understand that some molecular compounds when dissolved in water break apart into ions. In some
instances they break apart completely (e.g. HCl to produce H+ + Cl-). These are called strong
electrolytes. In other cases, e.g. acetic acid (CH3COOH) we develop an equilibrium between the
neutral CH3COOH and the ions H+ and CH3COO-. This is an example of a weak electrolyte.

Know the electrolyte solubility guidelines.

Be able to predict the outcome of precipitation reactions using the solubility guidelines

Be able to write precipitation reactions as net ionic reactions and understand the concept of
spectator ions.

Know the common strong and weak acids and bases.

Understand the Arrhenius definition of acids and bases

Understand the Brönsted-Lowry definition of acids and bases

Understand the idea that a strong acid dissociates 100% into H3O+ in water whereas a weak acid only
partially dissociates.

Know how to write a net ionic reaction for the reaction of a strong acid (e.g. HCl) and a strong base
(e.g. NaOH)

Understanding how non-metal and metal oxides act as acids and bases respectively

Be able to write net ionic reactions for gas forming reactions

Be able to identify oxidation-reduction reactions and to recognize oxidizing and reducing agents

Understanding the connection between oxidation-reduction reactions and electron transfer

Know the rules for determining the oxidation number for an element in a compound and be able to
use them to determine oxidation numbers

Understand the role of oxidation numbers in identifying an oxidation-reduction reaction.
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
Be able to balance oxidation-reduction reactions
UNIT 4: STOICHIOMETRY
Overview:
 Meaning of a balanced reaction

Concepts of limiting reagent, percentage yield and use of chemical analysis to determine ,

Combustion

Solution stoichiometry and molarity

pH and titrations

Oxidation-reduction titrations

Use of spectroscopy in quantitative analytical chemistry
Learning Outcomes:
 Understand the:

o
mass relationships in a balanced chemical reaction
o
use of dimensional analysis and the stoichiometry factor to determine mass
relationships between reactants and products.
Be able to
o
identify the limiting reagent for reactions in which one reactant is in limited supply, and
thence to determine how much product is formed and how much of the other reactant
is unused.
o
calculate the percentage yield of a chemical reaction
o
use chemical analysis data to determine amount of substance in a sample
o
determine the formula of an organic compound using combustion data

Understand the concept of the concentration scale “molarity”

Be able to calculate the molarity of a solution from solute mass data

Understand how to prepare a dilute solution quantitatively from a concentrated one

Understand the concept of pH and be able to convert H3O+ concentrations to pH and vice versa.
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
Be able to perform stoichiometry calculations for reactions in solution.

Understand the concept of a titration and standardization of a solution by titration

Be able to perform calculations involving titrations including determination of molar mass

Understand the basis of oxidation-reduction titrations and be able to perform calculations involving
such titrations

Understand the basis of quantitative analysis by spectroscopy including choice of appropriate
wavelength and generation of a calibration curve.

Be able to manipulate data from quantitative analysis by spectroscopy.
UNIT 5: ENERGY AND CHEMICAL REACTIONS
Overview:
 Forms of energy and conservation of energy

Heat and temperature

First law of thermodynamics

Concept of enthalpy and state functions

Enthalpies of reaction and standard state

Calorimetry

Hess’s Law

Standard Enthalpy of Formation
Learning Outcomes:
 Be able to describe the various forms of energy and understand concept of conservation of energy

Be able to use the common energy unit, the joule

Understand the connection between temperature and thermal energy (heat)

Understand the thermodynamic concepts, system, surroundings, directionality and extent of heat
flow, and thermal equilibrium

Be able to use the idea of specific heat capacity and temperature change to calculate extent of
thermal energy transfer
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
Understand:
o
the idea of thermal energy and changes of state and be able to perform calculations
relating to the amount of heat required in such processes
o
first law of thermodynamics and concept of work noting in particular the algebraic sign
conventions used
o
the concept of “enthalpy” and that at constant pressure enthalpy equals heat
transferred in the process
o
the concept of “state functions”
o
enthalpy changes for a chemical reaction
o
the concept of and definition of “standard state”
o
the idea of constant pressure calorimetry and be able to calculate enthalpies of reaction
from temperature changes in simple “coffee-cup” type calorimeter
o
the idea of constant volume (bomb) calorimeter and be able to calculate heats of
reaction from experimental data

Be able to use Hess’s Law to calculate enthalpies of reaction by suitable manipulation of chemical
actions whose enthalpies are known

Understand concept of “standard enthalpy of formation” and be able to use them to calculate
enthalpies of reaction
UNIT 6: STRUCTURE OF ATOMS
Overview:
Electromagnetic radiation and quantization
Atomic line spectrum of hydrogen, Balmer series and Bohr atom
Particle-wave duality
Uncertainty principle
Quantum numbers, shells and sub-shells
Shapes of atomic orbitals
Paired and unpaired electrons
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Learning Outcomes:
 Understand
o
electromagnetic (EM) radiation and be able to use the equation which connects frequency,
wavelength and the speed of light
o
the electromagnetic spectrum in particular the UV, visible and infrared regions
o
the quantization of light and be able to use Plank’s equation to convert between energy of
a photon and its frequency.
o
the photoelectric effect

Be able to calculation the energy of a mole of photons of a given frequency or wavelength

Understand
o
the idea of atomic line spectra, the success of the Balmer equation and the Bohr model for
understanding the structure of the hydrogen atom

Be able to calculate ground and excited state energies for the hydrogen atom

Understand
o
the particle-wave duality and de Broglie’s equation and be able to use this equation to
interconvert between particle velocity and wavelength
o
the concept of the “uncertainty principle”
o
the concept of and connection between the quantum numbers, n, l, and ml and the idea of
shells and sub-shells.
o
the shapes of s orbitals and the connection to electron density
o
the shapes of p and d orbitals
o
the property of electron spin and the quantum number ms
o
the idea of diamagnetism, paramagnetism and ferromagnetism and their connection to
paired and unpaired electrons
UNIT 7: STRUCTURE OF ATOMS AND PERIODIC TRENDS
Overview:
 Pauli exclusion principle, shells and sub-shells
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CHEMISTRY SEMESTER ONE
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
Sub-shell ordering and effective nuclear charge

Electron configurations o elements of periodic table

Connection between position in periodic table and chemical and physical properties
Learning Outcomes:
 Understand
o
the Pauli exclusion principle and how it relates to the number of electrons that can be
accommodated in electron shells
o
the ordering of subshells
o
the concept of “effective nuclear charge, Z*
o
the electron configurations of the main groups elements in periods 1 and 2.
o
the electron configurations of the transition metals
o
the electronic configurations of ions
o
how the atomic properties, atomic size, ionization energy and electron affinity relate to
position in the periodic table
o
how chemical properties relate to periodic trends
UNIT 8: BONDING AND MOLECULAR STRUCTURE
Overview:
 Ionic and covalent bonding

Lewis dot diagrams and octet rule

Resonance diagrams

Exceptions to octet rule

VSEPR theory

Bond polarity, electronegativity and formal charge

Molecular polarity
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
Bond order, bond length and bond strength

Enthalpies of bond formation
Learning Outcomes:
 Understand
o
the two different bonding types: concepts, ionic and covalent
o
the concept of valence electrons and using Lewis dot symbols for main group elements
o
Lewis electron dot structures for molecules and the octet rule

Be able to draw Lewis dot structures for molecules containing main group elements

Understand

the concept of “isoelectronic” species

the concept of atom formal charges and be able to calculate them

the concept of resonance and be able to identify and draw resonance structures

exceptions to the octet rule, including species with less than and those with more than eight.

the idea of molecules with unpaired electrons (free radicals)

VSEPR theory and be able to use it to predict molecular

the concept of “bond polarity” and “electronegativity”

the concept of “formal charge” and its relation to electronegativity and be able to calculate the
formal charges on atoms in a molecule.

the connection between bond polarity and the experimentally measurable quantity, molecular
polarity and be able to predict whether or not a molecule possesses a permanent dipole

the concept of bond order and its connection to bond length and bond strength

the concept of bond energy and be able to use it to estimate enthalpies of reaction
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UNIT 9: ORBITALS AND THEORIES OF CHEMICAL BONDING
Overview:
 Valence bond theory and orbital overlap

Hybridization

Multiple bonding and delocalized systems

Molecular orbital theory

Bond order

Molecular orbital diagrams for simple diatomic molecules

Application of Molecular orbital theory to unsaturated and aromatic systems
Learning Outcomes:
Understand
o
the concept of valence bond theory and orbital overlap model
o
the concept of hybridization to accommodate varying bond angles
o
application of valence bond theory to multiple bonding
o
application of valence bond theory to benzene and other delocalized systems
o
the three principles of molecular orbital theory – note one should add a fourth namely
the extent of bonding depends crucially on the extent to which the orbitals can overlap.
o
The concept of bond order and be able to calculate bond order for chemical bonds in
simple molecules
o
the reason for the non-existence of certain diatomic molecules such as Be2 because of
their zero bond order

Be able to construct molecular orbital diagrams for simple diatomic molecules

Understand
o
application of molecular orbital theory to simple heteronuclear diatomic molecules
o
application of molecular orbital theory to resonance structures and delocalized orbitals
noting that from n atomic orbitals we generate n molecular orbitals
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o
application of molecular orbital theory to benzene
UNIT 10: GASES
Overview:
 Pressure concept, units

Boyles Law

Charles Law And idea of absolute temperature scale

Combined gas law and ideal gas law

Use of ideal gas law to determine numbers of moles of sample and molar mass

Gas mixtures

Kinetic theory of gases and distribution of gas molecule velocities

Effusion, diffusion and molar mass determination
Learning Outcomes:
 Understand the concept of gas pressure being due to gas phase molecules impinging on a surface

Understand the concept of pressure being force per unit area

Understand the proper units for pressure

Be able to convert between different pressure units

Understand Boyles Law and be able to use it in calculations

Understand Charle’s Law and the idea of extrapolating to hypothetical zero volume at a certain
temperature (-273.15C)

Understand the definition of Kelvin temperature as T(K) = T(C) + 273.15

Be able to use Charles Law in calculations

Be ale to use the combined gas law in calculations

Understand Avogadro’s Hypothesis and the development of the ideal gas law

Understand that the above law allows determination of the number of moles of a gas and be able to
use such ideas in calculations
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
Be able to the calculate molar mass of a gas from its mass, temperature, volume and pressure

Be able to calculate molar masses from gas density

Be able to perform calculations relating to the gas laws, chemical reactions and stoichiometry

Understand the concept of gas mixtures and partial pressures

Be able to perform calculations relating to the pressures of gas mixtures

Understand the concept of kinetic theory of gases and the idea of a distribution of gas molecule
velocities

Understand the idea mean square velocity (mean square speed) and its connection to kinetic energy.

Be able to perform calculations which connect mean squared velocity, molar mass and temperature

Understand diffusion and effusion and its connection to molar mass

Be able to perform calculations which connection effusion to molar mass

Appreciate the fact that the ideal gas law is only valid as the pressure of a gas gets very low and
deviations can be considerable at high pressure and low temperature.
UNIT 11: INTERMOLECULAR FORCES AND LIQUIDS
Overview:
 Understand basis of intermolecular forces

Types of intermolecular forces

Special place of “hydrogen bonding” and its importance throughout the physical and biological world

Condensation of gases

Vapour pressure, boiling and enthalpies of vaporization

Critical behaviour and supercritical state

Surface tension and capillarity

Viscosity
Learning Outcomes
 Understand that it is forces between its molecules that cause a gas to condense into a liquid or for a
liquid to solidify. These are called “intermolecular forces”
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CHEMISTRY SEMESTER ONE
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
Understand the various intermolecular forces (in order of decreasing strength): 1) ion-ion; 2) ionpermanent dipole molecule; 3) permanent dipole molecule-permanent dipole molecule; 4)
permanent dipole molecule-induced dipole molecule; 5) induced dipole molecule-induced dipole
molecule

Understand how these forces affect the physical properties of matter

Understand the special place of “hydrogen bonding”, the interaction X-H·····X where X is typically O,
N F or Cl. Interaction is strong since X is very electronegative and H is very small allowing close
approach.

Understand how the strong hydrogen bonding in water gives this “liquid of life “ its unique
properties

Be able to explain the unique properties of water

Understand the connection between kinetic theory of matter and condensation.

Understand the idea of enthalpies of vaporization and be able to use them in calculation

Understand the concept of vapour pressure and the fact that liquids boil when their vapor pressure
equals atmospheric pressure

Be able to use the Clausius-Clapeyron equation to connect vapour pressure, temperature and
enthalpy of vaporization

Understand the idea of critical temperature and pressure and the idea of a supercritical fluid

Understand the concept of surface tension and capillary rise

Understand the concept of viscosity and its connection to strong intermolecular forces
UNIT 12: SOLUTIONS AND THEIR BEHAVIOUR
Overview:
Concentration scales
Solubility process for solid/liquid, liquid/liquid and gas/liquid equilibria
Enthalpies of solution
Colligative properties
Nature of and properties of different types of colloids
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CHEMISTRY SEMESTER ONE
COURSE OUTLINE
Learning Outcomes
 Be able to convert between the various concentration scales: weight percent, molality, mole fraction
and molarity

Understand the dissolution process as a dynamic equilibrium

Understand the concept of liquids dissolving in liquids and the terms miscibility and immiscibility.

Understand the process of ionic solids dissolving in a liquid and the concept of like dissolving like

Understand the origin of enthalpy of solution and be able to use such data in enthalpy calculations

Understand and be able to use Henry’s law in gas solubilities

Understand the effect of temperature on solubilities and Le Chatelier’s Principle

Understand the basis of the colligative properties, vapour pressure depression, freezing point
depression, boiling point elevation and osmotic pressure and to know the limitations of the methods

Be able to use data on colligative properties to determine molar masses of substances dissolved in a
liquid

Understand the basis of colloids and their importance in the world
NANSLO Chemistry Core Units and Laboratory Experiments
by the North American Network of Science Labs Online,
a collaboration between WICHE, CCCS, and BCcampus
is licensed under a Creative Commons Attribution 3.0 Unported License;
based on a work at rwsl.nic.bc.ca.
Funded by a grant from EDUCAUSE through the Next Generation Learning Challenges.
Creative Commons Attribution 3.0 Unported License
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