CHM110H5 Chemical Principles 1
The Complete Exam Bible · Fall 2026CHM110H5 Overview
- 0.5 credit
- Fall 2026
- University of Toronto Mississauga
- In person
- Five laboratory experiments
What Chemical Principles 1 covers
CHM110H5 is the first-year chemistry course at the University of Toronto Mississauga in which matter and its transformations are studied at the macroscopic level. The Fall 2026 syllabus names five topic families: stoichiometry, phases of matter, equilibria, thermodynamics and electrochemistry.
- Final exam is the largest piece The Final Exam carries 34%; two 45 minute midterms carry 14% each.
- Two pass conditions Attend at least four of the five labs, and average at least 40% on the midterms and final exam; missing either caps the final mark at 40%.
- Units decide marks Concentration, gas and heat problems turn on choosing the right unit and sign.
- Three units, one thread Stoichiometry feeds equilibrium, and free energy links equilibrium constants to cell potentials.
How CHM110H5 is assessed
| Component | Weight | Format |
|---|---|---|
| Short Lab Reports | 12% | Experiments 1, 3 and 5, 4% each |
| Long Lab Reports | 16% | Experiments 2 and 4, 8% each |
| Tutorial Quizzes | 10% | Five quizzes, 2% each, in tutorial sections |
| Midterm 1 | 14% | 45 minute test in your lecture section, October 5, 2026 |
| Midterm 2 | 14% | 45 minute test in your lecture section, November 9, 2026 |
| Final Exam | 34% | Written in the final exam period |
| Laboratory attendance · hurdle | Attend at least four of the five labs to pass the course | |
| Midterms and Final Exam minimum · hurdle | Weighted average of at least 40% across the two midterms and the final exam; below that the final mark is capped at 40% |
The six weighted components total 100%. Two pass conditions carry no weight of their own. Laboratory attendance: missing more than one lab caps the final mark at 40%. Midterms and final exam: your weighted average on the two midterms and the final exam must reach at least 40%, or the final mark is capped at 40% whatever the lab reports and quizzes add. A missed midterm with granted consideration moves its weight to the final exam.
Assessment structure
Current CHM110H5 dates
| Date | Item | Control |
|---|---|---|
| October 5, 2026 | Midterm 1 | 45 minutes, in your own lecture section. |
| November 9, 2026 | Midterm 2 | 45 minutes, in your own lecture section. |
Dates are as published in the Fall 2026 course syllabus. Confirm exact deadlines and submission settings in the live LMS.
What CHM110H5 covers
Nine chapters follow the three syllabus units and the five laboratory experiments, from measurement and stoichiometry to equilibrium, calorimetry and electrochemical cells.
Measurement, Units and Significant Figures
Uncertainty, precision and accuracy, significant figure rules and dimensional analysisWhy measurement opens the course
CHM110H5 begins with a block of background skills that every later topic leans on: SI and common non-SI units (grams, metres, litres, moles, atmospheres, bar and mmHg), converting between them, dimensional analysis, uncertainty, significant figures, and the difference between accuracy and precision.
- SI and non-SI units
- Dimensional analysis chains
- Counting significant figures
- Rounding sums versus products
- Exact numbers
- Accuracy versus precision
Writing and Balancing Chemical Equations
Reaction types, atom balance, mole ratios and limiting reactantsWhy balancing comes first in CHM110H5
Writing and balancing chemical equations sits in the background toolkit for Unit 1 Stoichiometry, alongside units, notation and nomenclature, and the course outcomes ask you to classify reactions by type and balance them.
- Conservation of atoms in equations
- Coefficients versus subscripts
- Balancing by inspection
- Clearing fractional coefficients
- Combination, decomposition and displacement
- Neutralisation, precipitation and redox
Solution Concentration Units
Molarity, molality, normality, mole fraction and percent by mass, with conversions through densityFive ways to say how much is dissolved
CHM110H5 expects you to define and know when to use five concentration measures: molarity (M), molality (m), normality (N), mole fraction (X) and percent. Each is a ratio of solute to some size of the mixture.
- Molarity, molality and normality definitions
- Mole fraction and percent by mass
- Temperature dependence of volume-based units
- Density as the mass to volume bridge
- Choosing a calculation basis
- Why molality usually exceeds molarity
Ideal Gas Law and Real Gas Behaviour
Boyle, Charles and Avogadro relations, PV = nRT, molar mass from gas data and van der Waals correctionsWhy gases get their own chapter
In CHM110H5 the gas laws sit inside the first unit alongside stoichiometry and solutions, and the learning outcome is specific: you should be able to describe the experiments that led to the ideal gas law, decide when it is safe to use, and explain the factors that make real gases stray from it.
- Boyle, Charles and Avogadro laws
- Deriving PV = nRT and choosing R
- Hydrogen collected over water
- Dumas molar mass of a vapour
- Van der Waals constants a and b
- High pressure and low temperature deviations
Laboratory Skills and the Iodine Clock
Volumetric glassware, dilution, timing a clock reaction and safe laboratory practiceWhat the first lab is really testing
The opening experiment in CHM110H5 is described in the course manual as a low-stakes skills lab.
- Three competing reactions in the iodine clock
- Why starch turns dark blue only after ascorbic acid runs out
- Preparing a standard solution in a volumetric flask
- Dilution with C1V1 = C2V2 into a 30.0 mL mixture
- Glassware tolerances from the course manual
- WHMIS, pre-lab and clean-up rules
Equilibrium and Le Chatelier's Principle
Equilibrium constants, reaction quotient, ICE tables, stress and shift, and carbon dioxide solubilityWhat equilibrium means in this course
Unit 2 of CHM110H5 is built around chemical equilibrium.
- Dynamic equilibrium and the meaning of K
- Reaction quotient Q versus K
- Kp, Kc and pure solids or liquids
- ICE tables and the 5 percent test
- Le Chatelier shifts: concentration, pressure, temperature, catalyst
- Henry's law and dissolved carbon dioxide
Acid Dissociation and Spectrophotometry
Beer-Lambert law, acid and base forms of an indicator and finding pKa from absorbanceWhat Experiment 4 measures
This chapter covers the CHM110H5 lab where you find the dissociation constant of bromophenol blue with a spectrophotometer.
- Bromophenol blue acid and base forms
- Indicator dissociation constant K_In and pK_In
- Beer-Lambert law A = εbc
- Choosing the two absorption maxima
- Absorbance ratio method for [HIn]/[In⁻]
- Blank, cuvette and wavelength errors
Calorimetry, Enthalpy and Hess's Law
Heat capacity, calorimeter calibration, enthalpy of reaction and combining reactions by Hess's lawEnthalpy as heat at constant pressure
Most reactions you run in the CHM110H5 lab happen in open containers, so pressure stays fixed while the system is free to expand or contract.
- Enthalpy as constant-pressure heat
- The sign of q from q = m c ΔT
- Calorimeter calibration with warm and cold water
- Extrapolating T-final from a cooling curve
- Reaction enthalpy per mole of limiting reagent
- Hess cycle for MgO formation
Entropy, Free Energy and Electrochemical Cells
Entropy, Gibbs free energy, the link to K and cell potential, and galvanic versus electrolytic cellsWhy entropy joins enthalpy
By this point in CHM110H5 you can already calculate enthalpy changes with heat capacities and Hess's law.
- Entropy and the second law
- Gibbs free energy: ΔG = ΔH − TΔS
- Sign cases and crossover temperature
- ΔG° and the equilibrium constant K
- Cell potential, ΔG° and K
- Labelling a galvanic cell
It is a 0.5 credit course with lectures, tutorials and a laboratory program, and Grade 12 Chemistry and Grade 12 Advanced Functions are prerequisites.
The three units
Unit 1, stoichiometry, starts with the atom, states of matter and the difference between pure substances and mixtures, then asks you to choose among the concentration units M, m, N, X and percent, to classify and balance reactions, and to use the ideal gas law while explaining why real gases deviate from it.
Unit 2, chemical equilibrium, defines the equilibrium state, asks how temperature, pressure, concentration and solvent change the extent of reaction, and has you calculate equilibrium constants and concentrations, including the conditions under which simplified calculations are valid.
Unit 3, thermodynamics and electrochemistry, covers state functions, the first law, heat capacities, Hess's law, entropy and free energy, the links between the standard free energy change, the equilibrium constant and the standard cell potential, and the parts of galvanic and electrolytic cells.
The laboratory program
Five experiments run beside the lectures: an iodine clock skills lab, the behaviour of gases, Le Chatelier's principle and the solubility of carbon dioxide, finding an indicator's dissociation constant by spectrophotometry, and calorimetry with Hess's law.
Each has a pre-lab and a report. The syllabus expects you to relate lecture theory to laboratory practice, identify sources of experimental error and write organised scientific reports. Attendance is a pass condition: you must personally attend at least four of the five labs, and online safety training must be completed before the first experiment.
A second pass condition sits on the test side: across the two midterms and the final exam your weighted average must be 40% or higher, otherwise the final mark cannot exceed 40%.
How this guide is organised
Nine chapters follow the units and the experiments in the order you meet them. Each chapter works a full problem with its own numbers, shows where marks are usually lost, and ends with practice you can check.
A mixed practice set and an exam-morning summary close the guide. Use the chapters in step with the recommended Silberberg reading, and redo each worked example from a blank page before the matching tutorial quiz or midterm.
Worked example · free
Gas volume from a decomposition reaction
- 2Moles of KClO3 = 2.45 g divided by 122.55 g/mol = 0.01999 mol.
- 2Mole ratio 3 O2 to 2 KClO3 gives 0.01999 times 1.5 = 0.02999 mol O2.
- 3V = nRT/P = 0.02999 times 0.08206 times 298 divided by 1.00 = 0.7333 L.
- 1Report three significant figures from 2.45 g and 1.00 atm: 0.733 L.
Key terms
- Significant figures
- The digits in a measured value that are known reliably plus one estimated digit; they set how precisely a calculated result may be reported.
- Dimensional analysis
- A method that multiplies a quantity by conversion factors equal to one so that units cancel and the answer arrives in the required unit.
- Molarity
- Concentration expressed as moles of solute per litre of solution, written M; it changes slightly with temperature because volume does.
- Molality
- Concentration expressed as moles of solute per kilogram of solvent, written m; it does not depend on temperature because it uses masses.
- Mole fraction
- The moles of one component divided by the total moles of all components; it has no units and the fractions of a mixture sum to one.
- Ideal gas law
- The relation PV = nRT linking pressure, volume, amount and absolute temperature for a gas whose particles have no volume and do not attract.
- Equilibrium constant
- The ratio of product to reactant activities at equilibrium, each raised to its coefficient; only a change of temperature changes its value.
- Le Chatelier's principle
- The rule that a system at equilibrium responds to an imposed change by shifting in the direction that partly counteracts it.
- Hess's law
- The statement that the enthalpy change of an overall reaction equals the sum of the enthalpy changes of steps that add up to it.
- Gibbs free energy
- The state function G = H minus TS; a negative change at constant temperature and pressure marks a spontaneous process.
CHM110H5 FAQ
How is the final mark made up?
Short lab reports carry 12% and long lab reports 16%, five tutorial quizzes carry 10% in total, each of the two midterm tests carries 14%, and the final exam carries 34%, which together make 100%. Two conditions sit on top of those weights: attending at least four of the five labs, and a weighted test and exam average of at least 40%. Failing either one limits the final mark to a maximum of 40%.
When are the midterm tests and how long do they last?
Midterm 1 is written on October 5, 2026 and Midterm 2 on November 9, 2026. Each test lasts 45 minutes and must be written during your own regularly scheduled lecture section; students may not sit a term test in a different lecture section.
What happens if I miss a lab?
Attendance at a minimum of four of the five labs is required to pass. A lab missed without documentation scores zero with no reweighting, and missing more than one lab caps the final mark at 40%.
What is the difference between molarity and molality?
Molarity divides moles of solute by litres of solution, while molality divides by kilograms of solvent. For dilute aqueous solutions they are close; as concentration rises the solution volume and solvent mass diverge and so do the two numbers.
Which calculator and textbook does the course use?
The course materials list Silberberg and Amateis, the Chemistry textbook, in any edition, together with any non-programmable scientific calculator, plus a lab coat, vented goggles and gloves for the laboratory.
How are late lab reports penalised?
Late term work such as a lab report loses 20% of its mark per day for up to five days; after five days late it receives zero. A late penalty may be waived when the policy conditions in the syllabus are met.
Questions students ask in each CHM110H5 chapter
- Why do I round to decimal places when adding but to significant figures when multiplying?
- Addition combines absolute uncertainties, so the measurement with the coarsest decimal place controls how far the sum can be trusted. Multiplication scales relative uncertainty, so the factor with the fewest significant figures dominates the error in the product or quotient. Chapter 01 · Measurement, Units and Significant Figures
- Why can't I just change H2O to H2O2 when oxygen does not balance?
- Because the subscript is part of what the substance is. H2O2 is hydrogen peroxide, a completely different compound from water, so the equation would describe another reaction. Only the coefficient in front of a formula can be adjusted, which multiplies the whole unit. Chapter 02 · Writing and Balancing Chemical Equations
- Why is molarity usually smaller than molality for a water-based solution?
- Molarity divided by molality equals the kilograms of solvent in one litre of solution. Dissolved solute occupies part of each litre, so less than about 1 kg of water remains, and the litre-based number comes out lower. Chapter 03 · Solution Concentration Units
- Why do I need to subtract water vapour when collecting hydrogen over water?
- The burette holds a mixture of hydrogen and water vapour, and the measured total pressure includes both. Only the hydrogen came from the metal, so its partial pressure, found by removing the temperature-dependent vapour pressure of water, is what goes into PV = nRT. Chapter 04 · Ideal Gas Law and Real Gas Behaviour
- Why does the iodine clock stay colourless for a while and then change suddenly?
- Ascorbic acid destroys iodine faster than hydrogen peroxide makes it, so free iodine stays near zero. When the ascorbic acid is exhausted, iodine accumulates, joins iodide as triiodide and binds starch, producing the dark blue colour within moments. Chapter 05 · Laboratory Skills and the Iodine Clock
- Does adding a catalyst change the equilibrium constant?
- No. A catalyst lowers the activation barrier for the forward and reverse reactions by the same factor, so equilibrium is reached sooner but the final composition and the value of K are identical to the uncatalysed system. Chapter 06 · Equilibrium and Le Chatelier's Principle
- Why is the acid-only solution treated as 100 percent HIn?
- Adding 2 mL of 0.20 M HCl raises the hydronium concentration far above K_In, so Le Chatelier's principle drives the dissociation back almost completely. The small leftover In⁻ is negligible compared with measurement noise, which makes SA3 a practical reference for the pure yellow form. Chapter 07 · Acid Dissociation and Spectrophotometry
- Why can't I just use the highest temperature the thermometer showed as the final temperature?
- The cup starts losing heat before the water is fully mixed, so the true peak is never displayed. Extending the straight cooling section back to the instant of mixing recovers a higher, loss-corrected value, and using the lower raw maximum makes the calorimeter constant come out too large. Chapter 08 · Calorimetry, Enthalpy and Hess's Law
- If ΔG° is positive, does the reaction not happen at all?
- Not quite. A positive ΔG° means K is less than one, so only a small amount of product forms at equilibrium. The actual ΔG = ΔG° + RT ln Q can still be negative when product concentrations are very low, so some forward reaction occurs. Chapter 09 · Entropy, Free Energy and Electrochemical Cells
How to study for the exam
Work in step with the units. After each lecture week, redo the matching worked example from a blank page and attempt the recommended Silberberg problems for that chapter. Before every lab, finish the pre-lab and write the calculation you will need with symbols, so that only numbers remain on the day. Keep an error log of unit, sign and significant figure slips and read it before each tutorial quiz.
In the last week before each midterm, time yourself on mixed sets of 45 minutes.
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