Chemical formulae, equations, calculations (1e)

Law of Conservation of Mass

Total mass of reactants is always equal to total mass of products

RFM Calculation

• multiply each A_r by how much there is

• add them together

Avogadro’s Number

6.023 × 10²³

Moles and Mass

n = M/M_r

Calculating Molar Mass/ no. of moles

• calculate M_r

• divide mass given by M_r

Calculating Reacting Mass

• find reacting moles (given or calculate with M / M_r)

• find ratio between given substance and substance to find

• find no. of moles

• use no. of moles to find mass

Balancing Equations using Reacting Mass

• write unbalanced equation

• write down masses

• calculate moles using mass and M_r

• use moles to find ratio

• use ratio to balance equation

Reasons for not getting 100% yield in a reaction

• some reactants left behind

• reaction may be reversible- high yield is impossible since products are continually turning back into reactants

• some products lost during purification or separation stages like filtration or distillation

• side reactions: substances reacting with gas in the air or impurity in reactant

• products lost during transfer between containers

Thoretical Yield

• amount of product that would be obtained under perfect practical and chemical conditions

• calculated from balanced equation and the reacting masses

Actual Yield

• recorded amount of product obtained

• always less than theoretical yield

Percentage Yield

• compares actual yield to theoretical

• for economic reasons, objective of every company is to have yield % as high as possible → to reduce costs & wastes and increase profits

• good way to measure how successful a chemical process is

Calculating Percentage Yield

• find actual yield (usually given)

• find theoretical yield (calculate using moles and mass)

• use equation: (actual/theoretical) x 100

Experiment (Finding Formulae of Simple Compounds): Aim

to determine formula of hydrated copper sulfate: CuSO₄.xH₂O

Experiment (Finding Formulae of Simple Compounds): Method

• measure mass of evaporating dish

• add known mass of hydrated salt

• heat over bunsen burner, gently stirring

• stop when salt turns from blue to white (all water lost)

• record mass of dish and contents

Experiment (Finding Formulae of Simple Compounds): Overheating the salt

decomposes and gives a larger change in mass

Experiment (Finding Formulae of Simple Compounds): Results

mass of white anhydrous salt

• measure mass of white anhydrous salt (mass of salt remaining)

mass of water

• subtract mass of white anhydrous salt from mass of known hydrated salt

• divide mass of the salt and water by masses (moles)

• simplify the ratio (multiply by 2 if decimal)

• find ratio as 1:water

• represent ratio as ‘salt.xH₂O)

Practical (Determining Formula of Magnesium Oxide): Aim

To determine the empirical formula of magnesium oxide by combustion of magnesium

Practical (Determining Formula of Magnesium Oxide): Diagram

Practical (Determining Formula of Magnesium Oxide): Method

• measure mass of crucible with lid

• add Mg sample to crucible and measure mass with lid

• measure

• strongly heat crucible over Bunsen burner for several minutes

• lift lid frequently to allow sufficient air into crucible for Mg to oxidise without letting MgO smoke escape

• continue heating until mass remains constant (max mass) → reaction is complete

• measure mass of crucible and contents

• calculate mass of crucible and contents by subtracting mass of empty crucible

Practical (Determining Formula of Magnesium Oxide): Results

• find mass of metal by subtracting mass of crucible from Mg and mass of empty crucible

• subtract mass of Mg used from mass of MgO

• divide each mass by A_r,

• simplify ratio (multiply by 2 if decimal)

• represent as M_xO_y

Practical (Determining Formula of Copper(II) Oxide): Aim

To determine the formula of copper(II)oxide by reduction with methane

Practical (Determining Formula of Copper(II) Oxide): Diagram

Practical (Determining Formula of Copper(II) Oxide): Method

• measure mass of empty boiling tube

• place metal oxide into a horizontal boiling tube and measure mass again

• support tube in horizontal position by clamp

• natural gas(methane) is passed over copper(II) oxide and excess gas is burned off

• copper(II) oxide is heated strongly with a Bunsen burner

• heat until metal oxide fully changes colour (all oxygen removed)

• measure mass of remaining powder in the tube and subtract mass of tube

Practical (Determining Formula of Copper(II) Oxide): Results (Empirical Formula)

• measure mass of powder to find mass of metal

• divide masses by A_r

• simplify ratio

• represent ratio as M_xO_y

Molecular Formula

Formula showing number and type of each atom in a molecule. Ex: ethanoic acid is C₂H₄O₂

Empirical Formula

Simplest whole number ratio of atoms of each element present in one molecule or formula unit of the compound. Ex: ethanoic acid is CH₂O

Ionic compounds are always _________ formulae

Empirical

Calculating Empirical Formulae

• write element

• write value given (% or mass)

• write A_r

• calculate moles by m/M_r

• calculate ratio of moles (multiply to make all values whole numbers)

• write final empirical formula

Calculating Molecular Formula

• find M_r of empirical formula (add masses of all atoms in the empirical formula

• divide M_r of molecular formula by M_r of empirical formula

• multiply each number in empirical formula by answer to find molecular formula

Calculating Concentration of Solutions in mol/dm³

number of moles (mol) / volume (dm³)

Avogadro’s Law

At the same conditions of temperature and pressure, equal amounts of gases will occupy the same volume of space

Molar Gas Volume at RTP

24dm³ or 24000 cm³

RTP

room temperature and pressure (20^oC and 1atm)

Gas Volume Equations

• volume = moles x 24 (dm³/mol)

• volume = moles x 24000 (cm³/mol)

Metals + Cold Water Reaction Speeds

K : violent

Na: quick

Li and Ca: less strong

Fe: slow rust

Mg, Zn, Cu: no reaction/very slow

Metal + Water Reaction Format

metal + water —→ hydroxide + hydrogen gas

Metals and Acids Reaction Rate

• only metals above Hydrogen

• more reactive metal = more vigourous reaction

• K and Na are very dangerous and react explosively

Metal + Acid Reaction Format

Metal + Acid → Salt + Hydrogen