Mole Concept and Molar Mass
The Mole — A Chemical Counting Unit
Atoms are extremely small particles whose individual masses are impossible to measure directly on a balance. To work with meaningful quantities, chemists use a unit called the mole. Just as a dozen means 12 items regardless of what they are, a mole is a fixed number of particles. The concept bridges the gap between the atomic scale and the macroscopic (laboratory) scale.
Gram Atom (Mole of an Element): The atomic mass of an element expressed in grams equals one gram atom, also called one mole of that element. For example, 1 mole of carbon = 12 g, 1 mole of magnesium = 24 g, and 1 mole of uranium = 238 g.
Gram Molecule (Mole of a Compound): The molecular mass of a molecular substance expressed in grams equals one gram molecule, or one mole of that substance. For example, 1 mole of water ($H_2O$) = 18 g, and 1 mole of sucrose ($C_{12}H_{22}O_{11}$) = 342 g.
Gram Formula (Mole of an Ionic Compound): Ionic compounds do not exist as discrete molecules. The sum of atomic masses of the ions in the formula unit, expressed in grams, is called one gram formula or one mole of the ionic substance. For example, 1 mole of NaCl = 58.50 g, and 1 mole of $Na_2CO_3$ = 106 g.
Gram Ion (Mole of an Ion): The ionic mass of a charged species expressed in grams is called one gram ion or one mole of ions. For example, 1 mole of $OH^-$ = 17 g, and 1 mole of $SO_4^{2-}$ = 96 g.
Mole Equivalents — Quick Reference
•
1 gram atom of hydrogen = 1.008 g of H atoms
•
1 gram atom of sodium = 23 g of Na atoms
•
1 gram atom of carbon = 12.000 g of C atoms
•
1 gram molecule of $H_2SO_4$ = 98.0 g
•
1 gram molecule of water = 18.0 g
•
1 gram formula of NaCl = 58.50 g
•
1 gram formula of $AgNO_3$ = 170 g
•
1 gram ion of $OH^-$ = 17 g
•
1 gram ion of $CO_3^{2-}$ = 60 g
The number of moles of any substance is calculated by dividing the given mass by its molar mass. This fundamental relationship applies to elements, molecular compounds, and ionic compounds alike.
$$n = \frac{m}{M}$$
Calculates the number of moles from the mass and molar mass of a substance
$n$=Number of moles (dimensionless)(mol)
$m$=Mass of the substance(g)
$M$=Molar mass (atomic, molecular, or formula mass as appropriate)(g mol$^{-1}$)
Avogadro's Number ($N_A$)
Avogadro's number ($N_A$) is the number of particles (atoms, molecules, or ions) present in one mole of any substance. Its value is $6.02 \times 10^{23}$ particles per mole. A key insight is that one mole of any substance always contains the same number of particles, even though different substances have vastly different masses per mole.
Same Count, Different Masses: 1 mole of hydrogen (1.008 g) and 1 mole of uranium (238 g) both contain exactly $6.02 \times 10^{23}$ atoms. The heavier element has heavier atoms, so you need more mass to reach the same count.
Elements: 23 g of Na = 1 mole of Na = $6.02 \times 10^{23}$ atoms of Na. Different masses of different elements contain the same number of atoms because their atoms differ in mass.
Compounds: 18 g of $H_2O$ = 1 mole = $6.02 \times 10^{23}$ molecules. 180 g of glucose ($C_6H_{12}O_6$) = 1 mole = $6.02 \times 10^{23}$ molecules. Different masses, same molecule count.
Ions: 96 g of $SO_4^{2-}$ = 1 mole = $6.02 \times 10^{23}$ ions of $SO_4^{2-}$. 62 g of $NO_3^-$ = 1 mole = $6.02 \times 10^{23}$ ions of $NO_3^-$.
Three related formulas connect mass to the number of particles (atoms, molecules, or ions) using Avogadro's number. These are derived by combining $n = m/M$ with $N = n \times N_A$.
$$N = \frac{m \times N_A}{M}$$
Calculates the number of particles from mass, Avogadro's number, and the mass of one particle (atomic, molecular, or ionic mass)
$N$=Number of particles (atoms, molecules, or ions)(particles (dimensionless count))
$m$=Given mass of the substance(g)
$N_A$=Avogadro's number ($6.02 \times 10^{23}$)(mol$^{-1}$)
$M$=Atomic mass, molecular mass, or ionic mass(g mol$^{-1}$)
$M$ = atomic mass
→$N$ gives the number of atoms of an element
$M$ = molecular mass
→$N$ gives the number of molecules of a compound
$M$ = ionic mass
→$N$ gives the number of ions of an ionic species
Number of Atoms: To find atoms of an element, use atomic mass as $M$. For example, 10 g of Mg (atomic mass = 24) gives $\frac{10 \times 6.02 \times 10^{23}}{24} = 2.51 \times 10^{23}$ atoms.
Number of Molecules: To find molecules of a compound, use molecular mass as $M$. For example, 10 g of $H_2O$ (molecular mass = 18) gives $\frac{10 \times 6.02 \times 10^{23}}{18} = 3.34 \times 10^{23}$ molecules.
Atoms Within Molecules: Once you know the number of molecules, multiply by the subscript of each element. In $H_2O$, each molecule has 2 H atoms and 1 O atom, so 3.34 × 10²³ molecules contain 6.68 × 10²³ H atoms and 3.34 × 10²³ O atoms.
Molar Mass — Unified Mass Measure
Molar mass is the mass of one mole of a substance expressed in grams per mole (g mol$^{-1}$). It is numerically equal to the atomic mass for elements, the molecular mass for molecular compounds, or the formula mass for ionic compounds. The concept unifies the three types — gram atom, gram molecule, and gram formula — under one term.
For Elements (Gram Atom Mass): The molar mass of an element equals its atomic mass in g mol$^{-1}$. Carbon = 12 g mol$^{-1}$, sodium = 23 g mol$^{-1}$, magnesium = 24.305 g mol$^{-1}$.
For Molecular Compounds: The molar mass equals the sum of atomic masses of all atoms in the molecule. Water ($H_2O$) = $2(1) + 16 = 18$ g mol$^{-1}$. Sulphuric acid ($H_2SO_4$) = $2(1) + 32 + 4(16) = 98$ g mol$^{-1}$.
For Ionic Compounds: The molar mass equals the formula mass — the sum of atomic masses of all ions in the formula unit. NaCl = $23 + 35.5 = 58.5$ g mol$^{-1}$. $Na_2CO_3$ = $2(23) + 12 + 3(16) = 106$ g mol$^{-1}$.
Molar Mass Calculation Examples
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$H_2SO_4$: $2(1.008) + 32.06 + 4(16) = 98.08$ g mol$^{-1}$
•
Sucrose ($C_{12}H_{22}O_{11}$): $12(12) + 22(1) + 11(16) = 342$ g mol$^{-1}$
•
$MgSO_4$: $24.3 + 32.06 + 4(16) = 120.4$ g mol$^{-1}$
•
$Na_2CO_3$: $2(23) + 12 + 3(16) = 106$ g mol$^{-1}$
•
$AgNO_3$: $107.87 + 14 + 3(16) = 169.87$ g mol$^{-1}$
•
$H_3PO_4$: $3(1.008) + 30.97 + 4(16) = 98.0$ g mol$^{-1}$
The mole formula can be rearranged to find mass when the number of moles is known. This is essential when working with chemical reactions where mole ratios are given but actual masses are needed.
$$m = n \times M$$
Calculates the mass of a substance from its number of moles and molar mass
$m$=Mass of the substance(g)
$n$=Number of moles(mol)
$M$=Molar mass(g mol$^{-1}$)
Molar Volume and Gas Calculations
One mole of any ideal gas at STP (standard temperature = 0 °C or 273 K, standard pressure = 1 atm or 101.325 kPa) occupies a fixed volume of 22.414 dm³. This is called the molar volume. Different gases have vastly different molar masses, yet one mole of each occupies the same volume because in the gaseous state, the distance between molecules is roughly 300 times greater than their diameters — the masses and sizes of individual molecules do not significantly affect the total volume.
Molar Volume at STP: $V_m$ = 22.414 dm³ mol$^{-1}$ at STP. This is true only for ideal gases.
Equal Volume, Different Mass: 22.414 dm³ of $H_2$ (mass = 2.016 g) and 22.414 dm³ of $CH_4$ (mass = 16 g) both contain $6.02 \times 10^{23}$ molecules. The heavier gas has heavier molecules, but the same count in the same volume.
Unit Conversion: 1 dm³ = 1000 cm³ = 1 litre. So 22.414 dm³ = 22414 cm³ = 22.414 L.
Molar Volume Equivalents at STP
•
$H_2$: 2.016 g = 1 mol = $6.02 \times 10^{23}$ molecules = 22.414 dm³
•
$CH_4$: 16 g = 1 mol = $6.02 \times 10^{23}$ molecules = 22.414 dm³
•
$O_2$: 32 g = 1 mol = $6.02 \times 10^{23}$ molecules = 22.414 dm³
•
$CO_2$: 44 g = 1 mol = $6.02 \times 10^{23}$ molecules = 22.414 dm³
•
$N_2$: 28 g = 1 mol = $6.02 \times 10^{23}$ molecules = 22.414 dm³
The number of moles of a gas at STP can be found directly from its volume using the molar volume relationship.
$$n = \frac{V}{22.414 \text{ dm}^3 \text{ mol}^{-1}}$$
Calculates the number of moles of an ideal gas at STP from its volume
$n$=Number of moles of the gas(mol)
$V$=Volume of the gas at STP(dm³)
$22.414$=Molar volume at STP(dm³ mol$^{-1}$)
Volume to Moles: Divide the gas volume by 22.414 dm³ mol$^{-1}$. A gas occupying 500 cm³ (= 0.5 dm³) at STP contains $0.5 / 22.414 = 0.0223$ moles.
Finding Molar Mass from Volume and Mass: Once moles are found from volume, the molar mass follows from $M = m/n$ where $m$ is the measured mass of the gas sample.
Calculations with Dissociated Substances
When an ionic or molecular compound dissolves and dissociates in water, the resulting ions can be counted using the stoichiometric ratio from the balanced dissociation equation. The number of each type of ion, their individual masses, and the total charges can all be calculated from the original amount of substance.
Ion Count from Moles: From the balanced dissociation equation, multiply the number of moles of the original compound by the coefficient of each ion. For $H_3PO_4 \rightarrow 3H^+ + PO_4^{3-}$, 1 mole of $H_3PO_4$ yields 3 moles of $H^+$ and 1 mole of $PO_4^{3-}$.
Charge Calculation: Multiply the number of each ion by its charge magnitude. 1 mole of $H_3PO_4$ gives $3 \times N_A$ positive charges (from 3 moles of $H^+$, each with charge +1) and $3 \times N_A$ negative charges (from 1 mole of $PO_4^{3-}$ with charge −3). Positive and negative charges are always equal in a dissociated solution.
Mass of Each Ion: The mass of individual ions after dissociation can be found using $m_{\text{ion}} = n_{\text{ion}} \times M_{\text{ion}}$ where $n_{\text{ion}}$ is the number of moles of that ion and $M_{\text{ion}}$ is its ionic mass.
Practical Example — $H_2SO_4$: Dissolving 9.8 g of $H_2SO_4$ (0.1 moles) gives 0.2 moles of $H^+$ and 0.1 moles of $SO_4^{2-}$. The mass of $H^+$ is $0.2 \times 1.008 = 0.202$ g, and the mass of $SO_4^{2-}$ is $0.1 \times 96 = 9.6$ g. Total positive charges = $0.2 \times N_A$, and total negative charges = $0.2 \times N_A$ (each $SO_4^{2-}$ carries 2 negative charges).
Application Strategy
Worked Examples
Problem Statement
Calculate the number of gram atoms (moles) in 0.1 g of sodium.
molemolar mass
Strategic Toolkit
$$n = \frac{m}{M}$$
Step 1
Identify the given values
Mass of sodium is 0.1 g. The molar mass of sodium (from the periodic table) is 23 g mol$^{-1}$.
$$m = 0.1 \text{ g}, \quad M_{Na} = 23 \text{ g mol}^{-1}$$
Step 2
Apply the mole formula
Divide the mass by the molar mass to get the number of moles.
$$n = \frac{0.1}{23} = 0.0043 \text{ mol}$$
Step 3
Express in scientific notation
The answer is 0.0043 mol, which equals 4.3 × 10⁻³ mol.
$$n = 4.3 \times 10^{-3} \text{ mol}$$
Problem Statement
Calculate the number of gram atoms (moles) in 0.1 kg of silicon.
molemolar massunit conversion
Strategic Toolkit
$$n = \frac{m}{M}$$
Step 1
Convert mass to grams
The mass is given in kg, so first convert to grams by multiplying by 1000.
$$m = 0.1 \text{ kg} = 0.1 \times 1000 = 100 \text{ g}$$
Step 2
Identify molar mass of silicon
The molar mass of silicon is 28.086 g mol$^{-1}$.
$$M_{Si} = 28.086 \text{ g mol}^{-1}$$
Step 3
Apply the mole formula
Divide the mass in grams by the molar mass to get the number of moles.
$$n = \frac{100}{28.086} = 3.56 \text{ mol}$$
Problem Statement
Calculate the mass of $10^{-3}$ moles of $MgSO_4$.
molemolar massgram formula
Strategic Toolkit
$$m = n \times M$$
Step 1
Calculate the formula mass of MgSO₄
$MgSO_4$ is an ionic compound, so use formula mass. Add the atomic masses of Mg, S, and four O atoms.
$$M_{MgSO_4} = 24 + 32 + 4(16) = 120 \text{ g mol}^{-1}$$
Step 2
Apply the mass formula
Multiply the number of moles by the formula mass to get the mass in grams.
$$m = 10^{-3} \times 120 = 0.12 \text{ g}$$
Problem Statement
How many molecules of water are there in 10.0 g of ice? Also calculate the number of hydrogen atoms, oxygen atoms, and total atoms in the sample.
Avogadro's numbermolar massparticle count
Strategic Toolkit
$$N = \frac{m \times N_A}{M}$$
Step 1
Find the number of water molecules
Use the particle formula with mass = 10 g and molar mass of water = 18 g mol$^{-1}$.
$$N_{H_2O} = \frac{10 \times 6.02 \times 10^{23}}{18} = 3.34 \times 10^{23} \text{ molecules}$$
Step 2
Count hydrogen atoms
Each $H_2O$ molecule has 2 hydrogen atoms, so multiply the molecule count by 2.
$$N_H = 2 \times 3.34 \times 10^{23} = 6.68 \times 10^{23} \text{ atoms}$$
Step 3
Count oxygen atoms
Each $H_2O$ molecule has 1 oxygen atom, so the oxygen atom count equals the molecule count.
$$N_O = 1 \times 3.34 \times 10^{23} = 3.34 \times 10^{23} \text{ atoms}$$
Step 4
Total atoms
Add the hydrogen and oxygen atom counts together.
$$N_{\text{total}} = 6.68 \times 10^{23} + 3.34 \times 10^{23} = 1.002 \times 10^{24}$$
Problem Statement
An ideal gas is enclosed in a container of volume 500 cm³ at STP. Its mass is 0.72 g. What is the molar mass of this gas?
molar volumemolar massSTP
Strategic Toolkit
$$n = \frac{V}{22.414}$$
Step 1
Find moles from volume at STP
Convert 500 cm³ to moles using the molar volume. Since 22414 cm³ = 1 mole at STP, divide the volume by 22414.
$$n = \frac{500}{22414} = 0.0223 \text{ mol}$$
Step 2
Calculate molar mass
Use the rearranged formula $M = m/n$ with the mass 0.72 g and moles 0.0223.
$$M = \frac{0.72}{0.0223} = 32 \text{ g mol}^{-1}$$
Problem Statement
10.0 g of $H_3PO_4$ is dissolved in excess water and dissociates completely. Calculate the number of positive and negative ions produced.
Avogadro's numberdissociationion count
Strategic Toolkit
$$H_3PO_4 \rightarrow 3H^+ + PO_4^{3-}$$
Step 1
Find moles of $H_3PO_4$
Molar mass of $H_3PO_4$ = 3(1) + 31 + 4(16) = 98 g mol$^{-1}$. Divide 10 g by 98 g mol$^{-1}$.
$$n = \frac{10}{98} = 0.102 \text{ mol}$$
Step 2
Count $H^+$ ions
From the dissociation equation, 1 mole of $H_3PO_4$ produces 3 moles of $H^+$. Multiply moles by 3, then by Avogadro's number.
$$N_{H^+} = 3 \times 0.102 \times 6.02 \times 10^{23} = 1.84 \times 10^{23} \text{ ions}$$
Step 3
Count $PO_4^{3-}$ ions
From the dissociation equation, 1 mole of $H_3PO_4$ produces 1 mole of $PO_4^{3-}$. Multiply moles by Avogadro's number.
$$N_{PO_4^{3-}} = 1 \times 0.102 \times 6.02 \times 10^{23} = 6.14 \times 10^{22} \text{ ions}$$
Memorize formulas
Moles from Mass
Divide mass (g) by molar mass (g mol$^{-1}$) to get moles. Works for elements, compounds, and ions.
Formula
$$n = \frac{m}{M}$$
Particles from Mass
Multiply the mole result by Avogadro's number ($6.02 \times 10^{23}$ mol$^{-1}$) to get the count of atoms, molecules, or ions.
Formula
$$N = \frac{m \times N_A}{M}$$
Mass from Moles
Multiply the number of moles by the molar mass (g mol$^{-1}$) to get mass in grams.
Formula
$$m = n \times M$$
Gas Moles from Volume at STP
Divide gas volume (dm³) by 22.414 dm³ mol$^{-1}$ to get moles. Valid only at STP (0 °C, 1 atm) for ideal gases.
Formula
$$n = \frac{V}{22.414 \text{ dm}^3 \text{ mol}^{-1}}$$
Review
1
A mole is defined as the amount of substance that contains:
Ans
6.022 × 10²³ elementary entities
Review
Tap to reveal
Recall the definition of a mole
One mole contains a fixed number of elementary particles equal to Avogadro's constant
$$1 \text{ mol} = 6.022 \times 10^{23} \text{ entities}$$
2
Avogadro's number ($N_A$) is:
Ans
$6.022 \times 10^{23}$
Review
Tap to reveal
Recall Avogadro's constant
Avogadro's number is a fundamental constant linking moles to particle count
$$N_A = 6.022 \times 10^{23} \text{ mol}^{-1}$$
3
Which statement correctly distinguishes a gram atom from a gram molecule?
Ans
A gram atom is the molar mass of an element in grams, while a gram molecule is the molar mass of a molecular compound in grams
Review
Tap to reveal
Define gram atom
Mass of one mole of atoms of an element, numerically equal to its atomic mass in grams
$$\text{Gram atom of C} = 12 \text{ g}$$
Define gram molecule
Mass of one mole of molecules of a compound, numerically equal to its molecular mass in grams
$$\text{Gram molecule of H}_2\text{O} = 18 \text{ g}$$
4
What is the molar mass of calcium carbonate (CaCO₃)?
Ans
100 g mol⁻¹
Review
Tap to reveal
Identify the constituent atoms
CaCO₃ contains 1 Ca atom, 1 C atom, and 3 O atoms
$$\text{CaCO}_3 = \text{Ca} + \text{C} + 3\text{O}$$
Sum the atomic masses
Ca = 40, C = 12, O = 16 (each)
$$M = 40 + 12 + (3 \times 16) = 100 \text{ g mol}^{-1}$$
5
When 1 mole of Al₂(SO₄)₃ completely dissociates in water, the total number of moles of ions produced is:
Ans
5 moles
Review
Tap to reveal
Write the dissociation equation
Al₂(SO₄)₃ dissociates into its constituent ions in water
$$\text{Al}_2(\text{SO}_4)_3 \rightarrow 2\text{Al}^{3+} + 3\text{SO}_4^{2-}$$
Count total moles of ions
Add the moles of cations and anions
$$n_{\text{ions}} = 2 + 3 = 5 \text{ mol}$$
6
The molar mass of a substance is defined as:
Ans
The mass of one mole of a substance expressed in grams
Review
Tap to reveal
Molar mass definition
Molar mass links the mass of a substance to its amount in moles
$$M = \frac{m}{n} \quad \text{(units: g mol}^{-1}\text{)}$$
7
At STP, a gas occupies 2.24 dm³. How many moles of the gas are present?
Ans
0.1 mol
Review
Tap to reveal
Apply the molar volume formula
Divide the given volume by the molar volume at STP
$$n = \frac{V}{22.414 \text{ dm}^3 \text{ mol}^{-1}} = \frac{2.24}{22.414}$$
Calculate the result
Perform the division to find the number of moles
$$n \approx 0.1 \text{ mol}$$
8
A gas at STP occupies 11.2 dm³ and has a mass of 16 g. What is its molar mass?
Ans
32 g mol⁻¹
Review
Tap to reveal
Find moles from volume at STP
Convert the gas volume to moles using molar volume
$$n = \frac{11.2}{22.414} \approx 0.5 \text{ mol}$$
Calculate molar mass
Use M = m / n to find the molar mass
$$M = \frac{16 \text{ g}}{0.5 \text{ mol}} = 32 \text{ g mol}^{-1}$$
9
How many molecules are present in 9 g of water (H₂O)?
Ans
$3.011 \times 10^{23}$ molecules
Review
Tap to reveal
Calculate molar mass of H₂O
Sum atomic masses: H₂O = 2(1) + 16 = 18
$$M_{\text{H}_2\text{O}} = 18 \text{ g mol}^{-1}$$
Convert mass to moles
Apply n = m / M
$$n = \frac{9}{18} = 0.5 \text{ mol}$$
Convert moles to molecules
Multiply moles by Avogadro's number
$$N = 0.5 \times 6.022 \times 10^{23} = 3.011 \times 10^{23}$$
10
What mass of sulphuric acid (H₂SO₄) contains $3.011 \times 10^{23}$ molecules?
Ans
49 g
Review
Tap to reveal
Convert molecules to moles
Divide by Avogadro's number
$$n = \frac{3.011 \times 10^{23}}{6.022 \times 10^{23}} = 0.5 \text{ mol}$$
Calculate molar mass of H₂SO₄
Sum atomic masses: 2(1) + 32 + 4(16) = 98
$$M_{\text{H}_2\text{SO}_4} = 98 \text{ g mol}^{-1}$$
Convert moles to mass
Apply m = n × M
$$m = 0.5 \times 98 = 49 \text{ g}$$
11
The molar volume of an ideal gas at STP is:
Ans
22.414 dm³ mol⁻¹
Review
Tap to reveal
Recall molar volume at STP
One mole of any ideal gas occupies a fixed volume at standard temperature and pressure
$$V_m = 22.414 \text{ dm}^3 \text{ mol}^{-1} \text{ at STP}$$
12
What is the mass of 0.25 mol of NaOH?
Ans
10 g
Review
Tap to reveal
Calculate molar mass of NaOH
Sum atomic masses: Na = 23, O = 16, H = 1
$$M_{\text{NaOH}} = 23 + 16 + 1 = 40 \text{ g mol}^{-1}$$
Convert moles to mass
Apply m = n × M
$$m = 0.25 \times 40 = 10 \text{ g}$$
13
How many total ions are present in 0.2 mol of Al₂(SO₄)₃?
Ans
$6.022 \times 10^{23}$ ions
Review
Tap to reveal
Write the dissociation equation
Determine how many ions each formula unit produces
$$\text{Al}_2(\text{SO}_4)_3 \rightarrow 2\text{Al}^{3+} + 3\text{SO}_4^{2-}$$
Calculate total ion moles
Multiply moles of compound by the number of ions per formula unit
$$n_{\text{ions}} = 0.2 \times (2 + 3) = 1.0 \text{ mol}$$
Convert to number of ions
Multiply by Avogadro's number
$$N = 1.0 \times 6.022 \times 10^{23} = 6.022 \times 10^{23}$$
14
How many moles of magnesium atoms are present in $1.806 \times 10^{22}$ atoms of Mg?
Ans
0.03 mol
Review
Tap to reveal
Convert atoms to moles
Divide the number of atoms by Avogadro's number
$$n = \frac{1.806 \times 10^{22}}{6.022 \times 10^{23}}$$
Simplify the calculation
Cancel powers of 10 and divide coefficients
$$n = \frac{1.806}{60.22} = 0.03 \text{ mol}$$
15
A gas mixture at STP contains 11.2 dm³ of N₂ and 5.6 dm³ of O₂. What is the total mass of the mixture?
Ans
22 g
Review
Tap to reveal
Find moles of N₂
Convert volume to moles at STP
$$n_{\text{N}_2} = \frac{11.2}{22.414} \approx 0.5 \text{ mol}$$
Find moles of O₂
Convert volume to moles at STP
$$n_{\text{O}_2} = \frac{5.6}{22.414} \approx 0.25 \text{ mol}$$
Calculate total mass
Convert moles to mass for each gas and add
$$m = (0.5 \times 28) + (0.25 \times 32) = 14 + 8 = 22 \text{ g}$$
Question 1 of 15
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1
A mole is defined as the amount of substance that contains: