Bridge to university · Sciences

Bridge: chemistry skills for first-year science

Moles, molar mass, significant figures, unit conversions, balancing, naming compounds, gas laws and dilution, reviewed before first-year chemistry.

37 cards6 topicsAbout 15 minutes a roundPreparation for a typical first-year general chemistry course

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About this deck

General chemistry in first year assumes a toolkit that many students only half remember from Grade 11 and 12. This deck rebuilds it. You will review the mole and molar mass, mole ratios, limiting reagents and percent yield, significant figures and scientific notation, dimensional analysis for unit conversions, balancing equations and recognizing reaction types, naming ionic and molecular compounds and acids, the ideal gas law and the simple gas laws, and concentration and dilution. Short worked examples show the arithmetic so you can check your own method.

The deck suits anyone heading into a science, health science, engineering or kinesiology program that includes a chemistry course. Go through it with a periodic table and a calculator beside you, and redo each example before you flip the card. Pay attention to units on every line; most first-year chemistry errors are unit errors. Other frequent slips include changing subscripts instead of coefficients when balancing, using Celsius in gas law equations, treating coefficients as mass ratios instead of mole ratios, and rounding too early in a multi-step calculation. Keep extra digits until the final answer, then round to the correct number of significant figures.

Before you start: Grade 11 chemistry and comfortable algebra; Grade 12 chemistry helps but is not essential for this review.

All 37 cards

Cover the definitions with a sheet of paper and test yourself, or use the study modes above.

Moles

Mole: The amount of substance containing 6.022 x 10^23 particles; it links counts of atoms or molecules to grams.Example: 1 mol of carbon atoms is 6.022 x 10^23 atoms
Avogadro's number: 6.022 x 10^23 per mole: the number of particles in one mole of anything.Example: 0.50 mol of water holds about 3.01 x 10^23 molecules
Molar mass: Mass in grams of one mole of a substance, found by adding atomic masses from the periodic table; unit g/mol.Example: H2O: 2(1.01) + 16.00 = 18.02 g/mol
Grams to moles: Divide the mass by the molar mass: n = m/M.Example: 36.0 g of water / 18.02 g/mol = 2.00 mol
Mole ratio: Coefficients in a balanced equation give the proportions in which substances react and form, counted in moles, never in grams.Example: 2H2 + O2 -> 2H2O: 2 mol of hydrogen per 1 mol of oxygen
Limiting reagent: The reactant that runs out first and so fixes the maximum product. Compare moles available with moles needed.Example: 3 mol H2 with 1 mol O2: oxygen runs out first
Percent yield: (actual amount obtained / theoretical amount) x 100%.Example: Expected 12.0 g, obtained 10.2 g: 85.0%
Empirical formula: The simplest whole-number ratio of atoms in a compound.Example: C6H12O6 reduces to CH2O

Measurement

Significant figures: Digits that carry meaning in a measurement: all nonzero digits, zeros between them, and trailing zeros after a decimal point.Example: 0.00450 has three
Leading zeros: Zeros before the first nonzero digit only place the decimal point and never count as significant.Example: 0.0072 has two significant figures
Sig figs in multiplication and division: Round the answer to the smallest number of significant figures found among the inputs.Example: 2.5 x 3.14 = 7.85, reported as 7.9
Sig figs in addition and subtraction: Round the answer to the fewest decimal places found among the inputs.Example: 12.11 + 0.3 = 12.41, reported as 12.4
Exact numbers: Counted values and defined conversions such as 1000 mL in 1 L have unlimited precision and never limit the answer.Example: 3 beakers is exactly 3
Unit conversion (dimensional analysis): Multiply by factors equal to 1, such as 1 km / 1000 m, arranged so unwanted units cancel and the wanted ones remain.Example: 72 km/h x 1000 m/km x 1 h/3600 s = 20 m/s
Scientific notation: Writing a number as a x 10^n with 1 <= a < 10. Every digit in a is significant.Example: 0.000340 = 3.40 x 10^-4
Accuracy versus precision: Closeness to the true value, versus closeness of repeated measurements to one another.Example: Tight cluster far from the target: precise but not accurate

Reactions

Balancing an equation: Adjust coefficients, never subscripts, until every element has the same number of atoms on both sides.Example: C3H8 + 5O2 -> 3CO2 + 4H2O
Law of conservation of mass: Matter is neither created nor destroyed in a chemical change, which is why equations must balance.Example: Mass of reactants equals mass of products in a closed system
Combustion of a hydrocarbon: Reaction with oxygen that produces carbon dioxide and water when oxygen is plentiful.Example: CH4 + 2O2 -> CO2 + 2H2O
Single displacement: One element replaces another in a compound: A + BC -> AC + B.Example: Zn + CuSO4 -> ZnSO4 + Cu
Double displacement: Two compounds swap ions: AB + CD -> AD + CB, often forming a precipitate, water or a gas.Example: AgNO3 + NaCl -> AgCl + NaNO3
Synthesis and decomposition: Building one product from simpler substances, and breaking one compound into simpler ones.Example: 2H2 + O2 -> 2H2O builds; 2H2O -> 2H2 + O2 breaks apart

Naming

Naming ionic compounds: Name the cation first, then the anion with an ide ending for simple anions. No prefixes are used for counts.Example: MgCl2 is magnesium chloride
Stock system (Roman numerals): Shows the charge of a metal that can form more than one ion, written in brackets right after the metal name.Example: FeCl3 is iron(III) chloride
Naming molecular compounds: Two nonmetals take prefixes (mono, di, tri, tetra, penta) to show counts; mono is dropped on the first element.Example: N2O4 is dinitrogen tetroxide; CO is carbon monoxide
Polyatomic ions: Charged groups of atoms that stay together, such as sulfate SO4^2-, nitrate NO3^-, hydroxide OH^- and ammonium NH4^+.Example: Na2SO4 is sodium sulfate
Naming acids: Binary acids use hydro...ic (HCl, hydrochloric acid). For oxyacids, ate becomes ic and ite becomes ous.Example: H2SO4 is sulfuric acid; HNO2 is nitrous acid
Writing a formula from charges: Combine ions so the total charge is zero, using the smallest whole-number ratio.Example: Al^3+ and O^2- give Al2O3

Gases

Ideal gas law: PV = nRT, with R = 8.314 L kPa/(mol K) when pressure is in kPa and volume in litres.Example: 1.00 mol at 273.15 K and 101.325 kPa occupies about 22.4 L
Kelvin temperature: The absolute scale used in gas calculations: add 273.15 to the Celsius value. Celsius in PV = nRT gives nonsense.Example: 25 degrees C = 298.15 K
Boyle's law: At constant temperature and amount, pressure and volume are inversely related: P1V1 = P2V2.Example: 2.0 L at 100 kPa squeezed to 1.0 L reaches 200 kPa
Charles's law: At constant pressure, volume is proportional to Kelvin temperature: V1/T1 = V2/T2.Example: 1.0 L at 300 K heated to 600 K becomes 2.0 L
STP: Standard temperature and pressure. IUPAC now uses 273.15 K and 100 kPa; older books use 101.325 kPa, so check your course.Example: At 273.15 K and 100 kPa, 1 mol of ideal gas fills about 22.7 L

Solutions

Molar concentration: Moles of solute per litre of solution, written mol/L or M: c = n/V.Example: 0.50 mol in 0.250 L gives 2.0 mol/L
Dilution: Adding solvent keeps the moles of solute fixed, so c1V1 = c2V2.Example: 100 mL of 2.0 M made up to 500 mL gives 0.40 M
Strong versus weak acid: One that ionizes almost completely in water, versus one that ionizes only partly.Example: HCl is strong; acetic acid in vinegar is weak
pH: Equal to -log of the hydronium ion concentration. Each step down of 1 means ten times more acidic.Example: [H3O+] = 1 x 10^-3 mol/L gives pH 3