IB Chemistry SL and HL review

1.1.1 Units for Molar mass?

grams per mole or g/mol or g mol-1 M is the symbol for molar mass

1.3.1: Balance chemical equations when all reactants and products are given. Distinguish between coefficients and subscripts.

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1.3.2: Identify the mole ratios of any two species in a balanced chemical equation. Use balanced chemical equations to obtain information about the amounts of reactants and products.

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1.3.3 Apply the state symbols (s), (l), (g) and (aq). Encourage the use of state symbols in chemical equations.

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1.4.1: Calculate stoichiometric quantities and use these to determine experimental and theoretical yields. Mass is conserved in all chemical reactions. Given a chemical equation and the mass or amount (in moles) of one species, calculate the mass or amount of another species.

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1.4.2: Determine the limiting reactant and the reactant in excess when quantities of reacting substances are given. Given a chemical equation and the initial amounts of two or more reactants: identify the limiting reactant calculate the theoretical yield of a product calculate the amount(s) of the reactant(s) in excess remaining after the reaction is complete.

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1.4.3: Apply Avogadro's law to calculate reacting volumes of gases.

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1.5.1: Define the terms solute, solvent, solution and concentration (g dm-3 and mol dm-3). Concentration in mol dm-3 is often represented by square brackets around the substance under consideration, eg [CH3COOH].

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1.5.2: Carry out calculations involving concentration, amount of solute and volume of solution.

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1.5.3: Solve solution stoichiometry problems. Given the quantity of one species in a chemical reaction in solution (in grams, moles or in terms of concentration), determine the quantity of another species.

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1.1.1: Describe the mole concept and apply it to substances.

12.00 grams of carbon-12 contains 6.02 x 10^23 atoms of carbon-12. 6.02 x 10^23 = Avogadro's Number. 6.02 x 10^23 per mole = Avogadro's constant (symbol L). The mass of one mole of any substance is known as the molar mass (symbol M). For example, 1 mole of hydrogen atoms contains 6.02 x 10^23 hydrogen atoms therefore the hydrogen atoms have 1/12th the mass of carbon-12 atom and a resulting mass of 1.01 grams. Typically elements are made up of a mixture of isotopes and have relative atomic masses.

1.1.2: Calculate the number of particles and the amount of substance (in moles).

Be able to convert between the amount of substance (in moles) and the number of atoms, molecules or formula units. 1 mole = 6.02 x 1023 formula units of that substance. Breaking down atoms within molecules: for example 1 mole of water (H2O) contains 2 moles of hydrogen atoms and 1 mole of oxygen atoms. 1 mol of H2O = 18g = 6.02 x 10^23 mol 2 mol of H atoms = 2.02g = 1.2 x 10^24 mol 1 mol of O atoms = 16.00g = 6.02 x 10^23 mol For more understanding watch the link below: http://www.youtube.com/watch?v=qLHUTTtIE7A

1.2.6 Determine the empirical formula and/or the molecular formula of a given compound. Determine the: - empirical formula from the percentage composition or from other suitable experimental data - percentage composition from the formula of a compound - molecular formula when given both the empirical formula and the molar mass

Empirical formula is determined by dividing the percentage composition of each element within a compound by its relative mass. The percentage composition is found by dividing the total mass of an element within a compound by the relative mass of that compound and multiplying by 100.

1.2.2: Molecular mass?

Is the sum of the component atoms within one molecule of the compound. This refers only to simple molecular compounds. For example the molecular mass of H20 is 2.02 + 16.00 = 18.02 g mol-1

1.2.3 Units of relative atomic/molecular/formula mass?

NO UNITS!! RELATIVE=NO UNITS

1.1.1 The definition of a mole?

The definition of a mole is an Avogadro number of particles of any substance. 1 mole = 6.02 x 10^23.

1.2.5:Empirical formula?

The empirical formula is the simplest possible ratio of atoms in a compound. The empirical formula is the literally the formula obtained by an experiment.

1.2.5: Define the terms empirical formula and molecular formula.

The empirical formula is the simplest possible ratio of atoms in a compound. The empirical formula is the literally the formula obtained by an experiment. The molecular formula is a multiple of the empirical formula.

1.2.1 Define the term molar mass (M) and calculate the mass of one mole of a species.

The mass of one mole of any substance is known as the molar mass and has the symbol M. The units are g/mol or g mol-1. Calculate the molar mass of Sulfuric acid: H2SO4 (2x1.01) + 32.06 + (4x16.00) = 98.07 g/mol

1.2.2: Atomic mass?

The mass of the atom of a specific element. It is usually expressed in atomic mass units (a.m.u.)

1.2.5: Molecular formula?

The molecular formula is a multiple of the empirical formula.

1.2.3 Relative atomic mass (Ar)?

The relative atomic mass of an element is the average of all the naturally occurring isotopes of the element RELATIVE TO CARBON-12. Ar is the symbol for relative atomic mass.

1.2.3 Relative formula mass?

The relative formula mass is used for ionic compounds. Ionic compounds are composed of metal and nonmetals. For example the relative formula mass of KCl = (39.10 + 35.45) = 74.55.

1.2.3 Relative molecular mass (Mr)?

The relative molecular mass is the average of all atomic masses of the atoms in one molecule relative to Carbon-12. For example the relative molecular mass of glucose C6H12O6 = (6 x 12.01) + (12 x 1.01) + (12 x 1.01) = 180.18. Mr is the symbol for relative molecular mass.

1.2.2: Distinguish between atomic mass, molecular mass and formula mass

The term molar mass (in g mol-1) can be used for all of these. atomic mass =mass of an atom of a specific element molecular mass =sum of the component atoms in one molecule Formula mass = sum of the component atoms in a substance that is not simple molecular

1.2.3: Define the terms relative molecular mass (Mr) and relative atomic mass (Ar).

The terms have no units. Relative atomic mass = the average

1.2.2: Formula mass?

This term is used when the substance is not simple molecular i.e silicon dioxide (macro molecular structure) The formula mass can also be applied to ionic substances (which also only exist in giant structure form). The structure can be simplified to give the simplest ratio of ions in the structure which is then taken as the formula. Typical units are g mol-1 or formula units

1.2.4: State the relationship between the amount of substance (in moles) and mass, and carry out calculations involving amount of substance, mass and molar mass.

moles=mass/ Mr; Mr=mass(grams)/moles; mass=(Mr)(moles) COMPARE IN MOLES