Calculations Involving Masses Flashcards

All 52 cards in this deck

What is relative formula mass ($M_r$)?

The sum of the relative atomic masses of all the atoms shown in the formula of a substance.

What are the steps to calculate a relative formula mass?

e.g. $\mathrm{CaCO_3}$ ($A_r$: C = 12, O = 16, Ca = 40)

1. Multiply each $A_r$ by the number of those atoms: $40$, $12$, $3 \times 16 = 48$
2. Add them: $40 + 12 + 48 = 100$

What is the formula for the percentage by mass of an element in a compound?

$\%$ by mass $= \dfrac{A_r \times \text{number of atoms of that element}}{M_r \text{ of compound}} \times 100$

True or false? Relative formula mass has the unit grams.

False. Relative formula mass is a ratio and has no units.

What are the steps to find the percentage by mass of an element in a compound?

e.g. % of O in $\mathrm{H_2O}$ ($A_r$: H = 1, O = 16)

1. Find $M_r$: $2 + 16 = 18$
2. Find total mass of that element: $16$
3. Divide and $\times 100$: $\frac{16}{18} \times 100 = 88.9\%$

True or false? In $\mathrm{POCl_3}$ the mass of chlorine used in a percentage-by-mass calculation is $3 \times 35.5$.

True. You must multiply the $A_r$ by the number of atoms of that element in the formula.

What is an empirical formula?

The simplest whole number ratio of the atoms (or ions) of each element in a compound.

What is a molecular formula?

The actual number of atoms of each element in one molecule of a substance.

How do you deduce the empirical formula from a molecular formula?

e.g. ethene, $\mathrm{C_2H_4}$

Divide all the subscripts by their highest common factor.
$\mathrm{C_2H_4} \rightarrow \mathrm{CH_2}$

What are the steps to deduce a molecular formula from an empirical formula and $M_r$?

e.g. empirical formula $\mathrm{CH_2}$, $M_r = 84$

1. Find the mass of the empirical formula unit: $12 + 2 = 14$
2. Divide $M_r$ by it: $84 \div 14 = 6$
3. Multiply all subscripts by that number: $\mathrm{C_6H_{12}}$

True or false? The empirical and molecular formula of a compound can be the same.

True. e.g. water is $\mathrm{H_2O}$ for both, as the subscripts are already in the simplest ratio.

True or false? Formulae worked out from reacting masses or percentage composition are molecular formulae.

False. They are empirical formulae; you also need the $M_r$ to get the molecular formula.

What are the steps to calculate an empirical formula from reacting masses?

e.g. 0.92 g Na combines with 0.32 g O ($A_r$: Na = 23, O = 16)

1. Divide each mass by its $A_r$: $0.92/23 = 0.04$; $0.32/16 = 0.02$
2. Divide both by the smaller: $2 : 1$
3. Write as a formula: $\mathrm{Na_2O}$

In an empirical formula calculation, what do you do if the ratio comes out as 1 : 1.5?

Multiply both numbers by 2 to give whole numbers, 2 : 3.

In the magnesium oxide experiment, how is the mass of oxygen that combined worked out?

Mass of magnesium oxide formed minus mass of magnesium used.

Why must magnesium be heated to constant mass in the empirical formula experiment?

To make sure all the magnesium has reacted; otherwise the mass of oxide is too low.

Give one source of error in the magnesium oxide empirical formula experiment that makes the measured mass of oxide too low.

White smoke of magnesium oxide escapes when the lid is lifted.

State the law of conservation of mass.

No atoms are lost or made in a chemical reaction, so the total mass of the products equals the total mass of the reactants.

Why does the mass not change during a precipitation reaction in a closed flask?

It is a closed system, so no substances can enter or leave; the same atoms are just rearranged.

Why does the mass decrease when a metal carbonate is heated in an open crucible?

Carbon dioxide gas is given off and escapes from the non-enclosed system.

Why does the mass increase when magnesium is burned in an open crucible?

Oxygen from the air reacts with and adds to the magnesium, and the oxide stays in the crucible.

True or false? A mass decrease in an open flask shows that mass has been destroyed.

False. Mass is conserved; a gas product has escaped into the air.

True or false? If magnesium is burned in a sealed container, the total mass stays the same.

True. Nothing can enter or leave a closed system, so the total mass is unchanged.

What are the steps to calculate the mass of a product from a balanced equation, given the mass of a reactant?

e.g. mass of MgO from 2.4 g Mg in $2\mathrm{Mg} + \mathrm{O_2} \rightarrow 2\mathrm{MgO}$ ($M_r$: Mg = 24, MgO = 40)

1. Moles of known substance: $2.4 \div 24 = 0.1$
2. Use the equation ratio (2 : 2) to get moles of MgO: $0.1$
3. Mass $=$ moles $\times M_r$: $0.1 \times 40 = 4$ g

True or false? In a reacting mass calculation you can use the ratio of the balancing numbers directly as a ratio of masses.

False. The balancing numbers give a ratio of moles, which must be converted to masses using $M_r$.

How do you convert a mass in kilograms to grams before a mole calculation?

e.g. 2 kg

Multiply by 1000.
$2 \times 1000 = 2000$ g

What is the formula for the concentration of a solution in $\mathrm{g\ dm^{-3}}$?

concentration $= \dfrac{\text{mass of solute (g)}}{\text{volume of solution (dm}^3)}$

How do you convert a volume in $\mathrm{cm^3}$ to $\mathrm{dm^3}$?

e.g. $250\ \mathrm{cm^3}$

Divide by 1000.
$250 \div 1000 = 0.25\ \mathrm{dm^3}$

What are the steps to calculate the concentration of a solution in $\mathrm{g\ dm^{-3}}$?

e.g. 9 g of solute in $200\ \mathrm{cm^3}$ of solution

1. Convert the volume to $\mathrm{dm^3}$: $200 \div 1000 = 0.2$
2. Divide mass by volume: $9 \div 0.2 = 45\ \mathrm{g\ dm^{-3}}$

How do you find the mass of solute in a solution from its concentration in $\mathrm{g\ dm^{-3}}$?

e.g. $50\ \mathrm{g\ dm^{-3}}$ in $500\ \mathrm{cm^3}$

Mass $=$ concentration $\times$ volume in $\mathrm{dm^3}$.
$50 \times 0.5 = 25$ g

True or false? $1\ \mathrm{dm^3}$ is the same volume as $1000\ \mathrm{cm^3}$.

True. $1\ \mathrm{dm^3} = 1000\ \mathrm{cm^3} = 1$ litre.

True or false? Dividing a mass in grams by a volume in $\mathrm{cm^3}$ gives the concentration in $\mathrm{g\ dm^{-3}}$.

False. You must then multiply by 1000, or convert the volume to $\mathrm{dm^3}$ first.

How is one mole of particles of a substance defined in terms of number of particles?

The Avogadro constant number of particles, $6.02 \times 10^{23}$ atoms, molecules, formulae or ions.

How is one mole of a substance defined in terms of mass?

A mass in grams equal to the relative particle mass ($A_r$ or $M_r$) of that substance.

What is the value of the Avogadro constant?

$6.02 \times 10^{23}$ per mole.

What is the formula linking moles, mass and relative formula mass?

moles $= \dfrac{\text{mass (g)}}{M_r}$

How do you calculate the number of moles in a given mass of a substance?

e.g. 0.1 g of magnesium ($A_r$ = 24)

Divide the mass by the relative particle mass.
$0.1 \div 24 = 4.2 \times 10^{-3}$ mol

True or false? One mole of water molecules and one mole of carbon atoms contain the same number of particles.

True. One mole of any substance contains $6.02 \times 10^{23}$ particles, though their masses differ.

How do you calculate the number of particles in a given number of moles?

e.g. 3.5 mol of $\mathrm{O_2}$ molecules

Multiply moles by the Avogadro constant.
$3.5 \times 6.02 \times 10^{23} = 2.107 \times 10^{24}$ molecules

What are the steps to calculate the number of particles in a given mass of a substance?

e.g. molecules in 36 g of water ($M_r$ = 18)

1. Moles $=$ mass $\div M_r$: $36 \div 18 = 2$
2. Particles $=$ moles $\times 6.02 \times 10^{23}$: $1.204 \times 10^{24}$ molecules

How do you find the number of atoms in a given number of moles of a compound?

e.g. 2 mol of $\mathrm{H_2O}$

Multiply moles by $6.02 \times 10^{23}$, then by the number of atoms in one formula unit (3 for $\mathrm{H_2O}$).
$2 \times 6.02 \times 10^{23} \times 3 = 3.61 \times 10^{24}$ atoms

What is the limiting reactant in a reaction?

The reactant that is completely used up first, and which is not in excess.

Why is the mass of product formed controlled by the limiting reactant?

Once it is used up the reaction stops, so no more product can be made however much of the other reactant is left.

What does it mean to say a reactant is 'in excess'?

More of it is present than is needed to react with all of the other reactant, so some is left over at the end.

True or false? The limiting reactant is always the one with the smaller mass.

False. It is the one that runs out first, found by comparing moles against the ratio in the balanced equation.

What are the steps to identify the limiting reactant?

e.g. 0.1 mol Mg and 0.1 mol HCl in $\mathrm{Mg} + 2\mathrm{HCl} \rightarrow \mathrm{MgCl_2} + \mathrm{H_2}$

1. Find the moles of each reactant: Mg $= 0.1$, HCl $= 0.1$
2. Compare with the equation ratio (1 Mg : 2 HCl): 0.1 mol Mg would need 0.2 mol HCl
3. The one that runs out first is limiting: HCl

Which reactant's moles must be used when calculating the maximum mass of product?

The moles of the limiting reactant.

What is the stoichiometry of a reaction?

The ratio of the numbers of moles of the reactants and products, shown by the balancing numbers in the equation.

What are the steps to deduce the stoichiometry of a reaction from masses?

e.g. 4.8 g Mg gives 8.0 g MgO ($A_r$ Mg = 24, $M_r$ MgO = 40)

1. Convert each mass to moles: $4.8/24 = 0.2$; $8.0/40 = 0.2$
2. Divide by the smallest to get the simplest whole number ratio: Mg : MgO $= 1 : 1$

When deducing stoichiometry, how do you find the mass of a reactant that was not measured directly?

Use conservation of mass: total mass of products minus the mass of the reactant that was measured.

True or false? The mole ratio deduced from experimental masses gives the balancing numbers for the equation.

True. The simplest whole number mole ratio is the stoichiometry of the balanced equation.

In hydrated copper sulfate, $\mathrm{CuSO_4.5H_2O}$, what does the formula tell you about moles?

Each mole of $\mathrm{CuSO_4.5H_2O}$ contains 5 moles of water.