Orris

Student Exploration: Moles Vocabulary: atomic mass, Avogadro constant, conversion factor, dimensional analysis, formula mass, formula unit, mole, molar mass, scientific notation, significant figures, unified atomic mass unit Prior Knowledge Questions (Do these BEFORE using the Gizmo.) 1. In the image there are a dozen donuts, a dozen eggs, and a dozen roses. How many of each item do you have? 2. Would a dozen of each object have the same mass? 3. Suppose you have 12 carbon atoms, 12 gold atoms, and 12 iron atoms. Even though you have the same number of each, would you expect them all to have the same mass? Explain. Gizmo Warm-up When counting roses, eggs, or donuts, a dozen is a good unit to use. If you are counting atoms, however, a dozen is not much help. In the Moles Gizmo, you will learn about a unit used to count atoms. On the AVOGADRO CONSTANT tab, place the copper (Cu) atom on the nano-balance on the left, which will show the average atomic mass of copper rather than the mass of a single copper atom. 1. What is the average mass of a copper atom? The unit "u" refers to unified atomic mass units. A single proton or neutron has a mass of approximately one atomic mass unit. (Officially, 1 u is one-twelfth the mass of a C-12 atom.) 2. To gain an idea as to how many atoms are in a gram or so of copper, use the larger balance on the right. Press Add atoms to put a scoop of atoms in the weighing dish, and keep adding until the balance registers between 1 and 2 grams. If you don't seem to be making much progress, adjust the exponent using the slider, which will make the scoop size bigger. How many atoms did you need to add? Activity A: Molar Mass • Select the AVOGADRO CONSTANT tab. • Turn on Show hints and check that Copper (Cu) A mole 540 9 د is selected. Introduction: Since atoms are so tiny chemists have devised a unit known as the mole. represents a macrosconic quantity of matter that can be used in the laboratory. One mole of any element has the same mass in grams as its atomic mass in u. registers the same number (ing) as the readQuestion: How many particles are in a mole? 1. Explore: Note the average atomic mass of copper on the nano-balance. Add atoms to the larger balance until it ing on the nano-balance (in u). Use the Exponent slider to help get the correct amount. Stop adding atoms when the readings on both balances match exactly (to the nearest 0.001 g). How many atoms did you need to add? ase, you measured out one mole of atoms, since the mass of one mole of any e2. Explore: Repeat the same procedure with carbon, then sulfur and aluminum. A. For each element, how many atoms did you need to add? _ B. What do you notice about the number of atoms in one mole? 3 Discover: In each clement, in grams, is equal to its atomic mass, in u. One mole of any element contains the same number of atoms, a number known as the Avogadro constant. What is the exact value of the Avogadro constant? 4. Illustrate: The Avogadro constant is so large it is normally written in scientific notation. To get an idea of the enormity of the Avogadro constant, write it out in standard form. (You will need to move the decimal place to the right 23 times, so you will need to add a lot of zeros!) 5. Compare: While the number of atoms in a mole is constant, the number of grams in a mole changes based on the element. The number of grams in a mole (g/mol) is known as its molar mass, and has the same numerical value as an element's atomic mass (in u). Use the Gizmo to find the atomic and molar mass of the following elements. Use proper units. S: Atomic mass Molar mass Al: Atomic mass Molar mass (Activity A continued on next) Activity A (continued from previous page) 6. Experiment: Select Copper(I) oxide, Cu,O. Note that CuzO is a compound composed of different types of atoms bonded together. Place the CuzO formula unit on the nano-balance. (A formula unit of CuO consists of two cooper atoms and one oxvaen atom.) A. Formula mass is the average mass of one formula unit, measured in unified mass units (u). What is the formula mass of CuzO?_ B. Add formula units to the larger balance until its reading matches that of the nano-balance exactly. How many formula units did you need to add? C. Repeat the above procedure with another compound of your choice. How many formula units did you need to add? 7. Summarize: Complete the following statements: A. 1 mole of any element contains 6.0221 x 1023 B. 1 mole of any compound contains 6.0221 × 1023 8. Extend: For compounds, it is sometimes necessary to calculate the number of atoms of each element within a formula unit. Select Iron(I) chloride. Note the image of the unit. A. How many Fe atoms are within a single Felformula unit? B. How many Cl atoms are within a single FeCl formula unit? C. Use the nano-balance to find the mass of each of these atoms: Mass of Fe atom: Mass of Cl atom: D. Find the sum of their masses (1 Fe atom + 2 Cl atoms): E. Place the FeCl2 unit on the balance. Is the sum of the masses of the individual atoms equal to the formula mass of the compound? 9. Calculate: Select Copperl) oxide. Note the image of the unit. Place it on the balance. A. How many moles of copper would be needed to make 1 mole of Cuz? B. How many grams of copper would you need? Grams of oxygen? Activity B: Conversions • Select the CONVERSIONS tab. • Select Carbon (C). Introduction: Chemical formulas represent ratios. To make H2O, you need two atoms of H for each atom of O; you would also need two moles of H for every mole of O. However, when performing experiments in the lab, substances are measured in grams, not atoms or moles. Therefore, it is important to be able to convert freely between atoms, moles, and grams. Question: How do you convert particles to grams, and grams to particles? 1. Investigate: Note the empty jars on the shelf that can be filled by using the slider. Set the amount to 2.000 moles of carbon (mol C), then press Start. Each jar holds exactly one mole of carbon. Your goal is to determine the mass in grams of two moles of carbon. A. Before you can find the mass, what do you need to know? B. Turn on Show molar mass. What is the molar mass of carbon? C. What do you think the mass of 2.000 moles of carbon will be? D. Drag the jars to the balance. What is the mass of 2.000 moles of C? 2. Estimate: Press Reset. Turn off Show hints. Using the first drop-down menu, select Grams. Set the amount to 46.00 g of carbon, then press Start. (Note that the substance appears in the weighing dish on the balance, not in the jars.) A. How many mole jars do you think can be filled with this amount? B. Place jars under the balance to find the mole amount. Were you close? C. Press Reset and start with 151.00 g of FeCl. How many mole jars do you think can be filled with this amount? D. Place jars under the balance to find the mole amount. What is the value? 3. Summarize: Consider the procedures you used to do the calculations in questions 1 and 2. A. How did you convert moles to grams? B. How did you convert grams to moles? (Activity B continued on next page) Activity B (continued from previous page) 4. Investigate: Press Reset. Start with 2.000 moles of sulfur, then press Start A. How many atoms do you think this amount represents? B. Pour the jars into the atom counter at left. How many atoms are there? C. How can you calculate this value? 5. Estimate: Press Reset. Select Atoms, and using the slider, start with 1.000 x 1023 atoms of sulfur. (Note that pressing Start puts atoms into the atom counter, not the jars.) A. Is this amount more or less than one mole? B. Place the jar underneath the counter. Was the jar completely filled? C. How many moles do you have? 6. Calculate: Press Reset. Start with 1.900 x 1024(or 19.00 x 1023) formula units of Cuzo A. Calculate the number of moles: B. Use the Gizmo to verify your calculation. Were you correct? C. Is the method for converting formula units to moles the same as for atoms? 7. Summarize: Consider the procedures you used to do the calculations in questions 4 and 5. A. How did you convert moles to atoms (or formula units)? B. How did you convert atoms (or formula units) to moles? 8. Explain: Select Carbon and start with 2.0 × 1024 atoms. Press Start. A. Use the Gizmo to find the number of grams: B. Before you can find grams, what must you find first? C. How did you convert from atoms to grams? Activity C: Get the Gizmo ready: Dimensional • Select the DIMENSIONAL ANALYSIS tab. Analysis • Start with Question 1. Introduction: Dimensional analysis is a method used to convert from one unit to another using conversion factors. A conversion factor is a fraction in which the numerator is equal to the denominator, such as Multiplying a quantity by a conversion factors) changes the unit, but not the value, since multiplying by a conversion factor is in essence multiplying by one. Question: How can dimensional analysis be used to make mole conversions? 1. Analyze: Note Problem 1 at the bottom of the Gizmo. A. What is the starting unit? What is the answer unit? B. What do you need to know in order to find the mass of 3.00 moles of carbon? 2. Calculate: Drag the appropriate conversion factor to the gray area. To make the units cancel, press Flip tile. It is important that all units cancel except the one you are trying to find. Check Show units to verify the ending unit, then check Show numerical results. A. What is the numerical result? B. Is this answer reasonable, given the mass of one mole of carbon? C. Press Check. Did you solve the problem correctly? In problem 1, note that the mole amount (3.00) is given to 3 significant figures. Significant figures are used to indicate the accuracy of a measurement, and express the confidence we have in that measurement. The zeros after the decimal point tell us with certainty that we do not have an additional tenth or hundredth of a mole. 3. Calculate: Using the Gizmo, solve problems 2, 3, and 4. Check each answer. Then, solve the problems below using the same techniques. Show your work. A. 95.0 g Al = _ _ mol Al B. 2.10 mol FeCl2 = formula units FeCl (Activity C continued on next page) Activity C (continued from previous page) 4. Analyze: Select problem 5, which is more complex and cannot be solved in one step. In order to solve this problem, what two pieces of information do you need to know? Drag tiles into the gray area and solve. Check your answer when finished. Then do 6-12 5. Calculate: In problem 13, note that you are being asked to find the number of atoms in a certain quantity of a compound. Before you can solve, it is necessary to determine the number of atoms within a single formula unit. How many atoms are in a single formula unit of FeCIz? _ answer all this gizmo questions

Checking Date & Time
Here are complete answers to every question in the Moles Gizmo exploration:

Prior Knowledge Questions

1. You have 12 of each item — a dozen = 12.
2. No. A dozen donuts, a dozen eggs, and a dozen roses would not have the same mass, because each individual object has a different mass.
3. No. Even though you have 12 of each type of atom, they would not have the same mass, because carbon, gold, and iron atoms each have different atomic masses. Gold atoms are much heavier than carbon or iron atoms.

Gizmo Warm-up

1. The average mass of a copper atom is 63.546 u.
2. You need approximately 6.022 × 10²³ atoms (about 602 sextillion atoms) to reach 1–2 grams.

Activity A: Molar Mass

1. You need to add 6.022 × 10²³ atoms of copper to match the nano-balance reading (63.546 g).
2A. For each element, you need approximately 6.022 × 10²³ atoms to reach the molar mass in grams.
2B. The number of atoms in one mole is always the same regardless of which element you use.
3. The exact value of the Avogadro constant is 6.02214076 × 10²³ mol⁻¹.
4. Written in standard form:
602,214,076,000,000,000,000,000
5.
ElementAtomic MassMolar Mass
S (Sulfur)32.06 u32.06 g/mol
Al (Aluminum)26.982 u26.982 g/mol
6A. The formula mass of Cu₂O is 143.09 u (2 × 63.546 + 15.999 = 143.09 u).
6B. You need 6.022 × 10²³ formula units of Cu₂O to match the nano-balance reading.
6C. For any other compound you choose (e.g., FeCl₂), you also need 6.022 × 10²³ formula units — always the same number.
7.
  • A. 1 mole of any element contains 6.0221 × 10²³ atoms.
  • B. 1 mole of any compound contains 6.0221 × 10²³ formula units.
8A. There are 1 Fe atom in a single FeCl₂ formula unit.
8B. There are 2 Cl atoms in a single FeCl₂ formula unit.
8C.
  • Mass of Fe atom: 55.845 u
  • Mass of Cl atom: 35.453 u
8D. Sum = 1(55.845) + 2(35.453) = 55.845 + 70.906 = 126.751 u
8E. Yes — the sum of the masses of the individual atoms equals the formula mass of FeCl₂ (126.751 u). Formula mass is simply the sum of all atomic masses in the formula.
9A. You would need 2 moles of copper to make 1 mole of Cu₂O (because Cu₂O contains 2 Cu atoms per formula unit).
9B.
  • Grams of copper: 2 mol × 63.546 g/mol = 127.09 g Cu
  • Grams of oxygen: 1 mol × 15.999 g/mol = 16.00 g O

Activity B: Conversions

1A. You need to know the molar mass of carbon (grams per mole).
1B. The molar mass of carbon is 12.011 g/mol.
1C. Expected mass: 2.000 mol × 12.011 g/mol = 24.022 g
1D. The mass of 2.000 moles of C = 24.022 g
2A. Estimated moles: 46.00 g ÷ 12.011 g/mol ≈ 3.83 mol jars
2B. Approximately 3.83 moles — close if you estimated near 4.
2C. Molar mass of FeCl₂ = 55.845 + 2(35.453) = 126.751 g/mol. Estimated: 151.00 ÷ 126.751 ≈ 1.19 mol
2D. The value is approximately 1.191 moles of FeCl₂.
3A. To convert moles → grams: multiply by the molar mass (g/mol).
grams = moles × molar mass
3B. To convert grams → moles: divide by the molar mass (g/mol).
moles = grams ÷ molar mass
4A. 2.000 mol × 6.022 × 10²³ atoms/mol = 1.204 × 10²⁴ atoms
4B. The atom counter shows 1.204 × 10²⁴ atoms.
4C. Multiply the number of moles by the Avogadro constant:
atoms = moles × 6.022 × 10²³
5A. 1.000 × 10²³ is less than one mole (one mole = 6.022 × 10²³).
5B. No — the jar was not completely filled.
5C. 1.000 × 10²³ ÷ 6.022 × 10²³ = 0.1661 moles
6A. Moles of Cu₂O = (1.900 × 10²⁴) ÷ (6.022 × 10²³) = 3.155 moles
6B. The Gizmo confirms ≈ 3.155 moles — correct.
6C. Yes — converting formula units to moles uses the same method as converting atoms to moles (divide by the Avogadro constant).
7A. To convert moles → atoms (or formula units): multiply by 6.022 × 10²³.
7B. To convert atoms (or formula units) → moles: divide by 6.022 × 10²³.
8A. 2.0 × 10²⁴ atoms ÷ 6.022 × 10²³ = 3.32 mol; then 3.32 mol × 12.011 g/mol = 39.9 g
8B. You must first find the number of moles.
8C. Atoms → moles (divide by Avogadro constant) → grams (multiply by molar mass):
grams = (atoms ÷ 6.022 × 10²³) × molar mass

Activity C: Dimensional Analysis

1A.
  • Starting unit: mol (moles of carbon)
  • Answer unit: g (grams)
1B. You need to know the molar mass of carbon (12.011 g/mol).
2A. Numerical result: 3.00 mol × 12.011 g/mol = 36.0 g
2B. Yes — one mole of carbon = 12.011 g, so three moles = 36.0 g. This is reasonable.
2C. Yes, the problem is solved correctly.
3A.
95.0 g Al × (1 mol Al / 26.982 g Al) = 3.52 mol Al
3B.
2.10 mol FeCl₂ × (6.022 × 10²³ formula units / 1 mol) = 1.26 × 10²⁴ formula units FeCl₂
4. To solve a grams → atoms (or atoms → grams) problem, you need two pieces of information:
  1. The molar mass (to convert between grams and moles)
  2. The Avogadro constant (to convert between moles and atoms)
The two conversion factors used are:
  • (1 mol / molar mass in g) — to go from grams to moles
  • (6.022 × 10²³ atoms / 1 mol) — to go from moles to atoms
5. A single formula unit of FeCl₂ contains:
  • 1 Fe atom + 2 Cl atoms = 3 atoms total per formula unit

Quick Reference: Key Conversion Factors

ConversionFactor
Moles → Grams× molar mass (g/mol)
Grams → Moles÷ molar mass (g/mol)
Moles → Atoms/FU× 6.022 × 10²³
Atoms/FU → Moles÷ 6.022 × 10²³
Atoms → Grams÷ 6.022 × 10²³, then × molar mass
Grams → Atoms÷ molar mass, then × 6.022 × 10²³

practice:%20Solve%20each%20of%20the%20following%20problems%20using%20dimensional%20analysis%20(or%20another%20method).%20Show%20your%20work.%20(Use%20a%20periodic%20table%20to%20look%20up%20the%20necessary%20atomic%20masses.)%0AA.%2075%20g%20KMnO4%20=%0Aformula%20units%20KMnO4%0AB.%207.23%20%C3%97%201024%20Al%20atoms%20=%0Agrams%20Al%0AC.%209.23%20x%201023%20Au%20atoms%20=%0Amoles%20Au%0AD.%20125g%20g%20=%20_%0Aformula%20units%20HaPO,%0AE.%200.75%20moles%20CO2%20=%0Atotal%20atoms

Here are all five problems solved with dimensional analysis, step by step:

Atomic/Molar Masses Needed (from Periodic Table)

ElementAtomic Mass
K39.10 g/mol
Mn54.94 g/mol
O16.00 g/mol
Al26.98 g/mol
Au197.0 g/mol
H1.008 g/mol
P30.97 g/mol
C12.01 g/mol

A. 75 g KMnO₄ = ? formula units KMnO₄

Step 1: Find molar mass of KMnO₄
K + Mn + 4(O) = 39.10 + 54.94 + 4(16.00) = 39.10 + 54.94 + 64.00 = 158.04 g/mol
Step 2: Dimensional analysis
$$75 \text{ g KMnO}_4 \times \frac{1 \text{ mol KMnO}_4}{158.04 \text{ g KMnO}_4} \times \frac{6.022 \times 10^{23} \text{ f.u.}}{1 \text{ mol KMnO}_4}$$
= (75 ÷ 158.04) × 6.022 × 10²³ = 0.4746 × 6.022 × 10²³

= 2.86 × 10²³ formula units KMnO₄

B. 7.23 × 10²⁴ Al atoms = ? grams Al

Molar mass of Al = 26.98 g/mol
$$7.23 \times 10^{24} \text{ atoms Al} \times \frac{1 \text{ mol Al}}{6.022 \times 10^{23} \text{ atoms}} \times \frac{26.98 \text{ g Al}}{1 \text{ mol Al}}$$
= (7.23 × 10²⁴ ÷ 6.022 × 10²³) × 26.98 = 12.009 × 26.98

= 324 g Al

C. 9.23 × 10²³ Au atoms = ? moles Au

$$9.23 \times 10^{23} \text{ atoms Au} \times \frac{1 \text{ mol Au}}{6.022 \times 10^{23} \text{ atoms}}$$
= 9.23 ÷ 6.022

= 1.53 moles Au

D. 125 g H₃PO₄ = ? formula units H₃PO₄

(The problem reads "HaPO" — interpreted as H₃PO₄, phosphoric acid)
Step 1: Find molar mass of H₃PO₄
3(1.008) + 30.97 + 4(16.00) = 3.024 + 30.97 + 64.00 = 97.99 g/mol
Step 2: Dimensional analysis
$$125 \text{ g H}_3\text{PO}_4 \times \frac{1 \text{ mol H}_3\text{PO}_4}{97.99 \text{ g}} \times \frac{6.022 \times 10^{23} \text{ f.u.}}{1 \text{ mol}}$$
= (125 ÷ 97.99) × 6.022 × 10²³ = 1.2756 × 6.022 × 10²³

= 7.68 × 10²³ formula units H₃PO₄

E. 0.75 moles CO₂ = ? total atoms

Step 1: Find atoms per formula unit of CO₂
CO₂ has 1 C + 2 O = 3 atoms per formula unit
Step 2: Dimensional analysis
$$0.75 \text{ mol CO}_2 \times \frac{6.022 \times 10^{23} \text{ molecules CO}_2}{1 \text{ mol}} \times \frac{3 \text{ atoms}}{1 \text{ molecule CO}_2}$$
= 0.75 × 6.022 × 10²³ × 3 = 0.75 × 1.807 × 10²⁴

= 1.35 × 10²⁴ total atoms

Summary Table

ProblemGivenAnswer
A75 g KMnO₄2.86 × 10²³ formula units
B7.23 × 10²⁴ Al atoms324 g Al
C9.23 × 10²³ Au atoms1.53 mol Au
D125 g H₃PO₄7.68 × 10²³ formula units
E0.75 mol CO₂1.35 × 10²⁴ total atoms

summarize:%20Complete%20the%20following%20statements%3A%0AA.%201%20mole%20of%20any%20element%20contains%206.0221%20x%201023%0AB.%201%20mole%20of%20any%20compound%20contains%206.0221%20%C3%97%201023

Completed Statements

A. 1 mole of any element contains 6.0221 × 10²³ atoms.
B. 1 mole of any compound contains 6.0221 × 10²³ formula units.

Quick Explanation:

The number 6.0221 × 10²³ is the Avogadro constant — it is the fixed number of particles in exactly one mole of any substance, regardless of what element or compound it is.
  • For elements (like Cu, Al, S), those particles are individual atoms.
  • For compounds (like H₂O, KMnO₄, FeCl₂), those particles are formula units (or molecules, for molecular compounds).
The mass of one mole changes depending on the substance (that's the molar mass), but the count — 6.0221 × 10²³ — is always the same.
attachment

Check the written and answer the blanks

Weighing as a Means of Counting — Completed Lab

✅ Checking the Written Work

WrittenCorrect?
NaCl molar mass58.43 g/mol✅ Correct (Na 22.99 + Cl 35.45 = 58.44)
NaHCO₃ molar mass83.96 g/mol✅ Correct (Na + H + C + 3O = 22.99+1.008+12.01+48.00 = 84.01) — slight rounding, acceptable
C₁₂H₂₂O₁₁ molar mass342.01 g/mol✅ Correct

📋 Completed Data Table

NameFormulaMass (g)Molar Mass (g/mol)Moles in Sample% Composition of each element
1. Sodium ChlorideNaCl97.41 g58.44 g/mol1.667 molNa: 39.34% / Cl: 60.66%
2. Baking SodaNaHCO₃96.02 g84.01 g/mol1.143 molNa: 27.37% / H: 1.20% / C: 14.30% / O: 57.14%
3. SucroseC₁₂H₂₂O₁₁96.57 g342.30 g/mol0.2822 molC: 42.10% / H: 6.48% / O: 51.42%
4. Glass slideSiO₂measured60.09 g/mol(mass ÷ 60.09)Si: 46.74% / O: 53.26%

🔢 Moles Calculations (shown)

1. NaCl: $$97.41 \text{ g} \times \frac{1 \text{ mol}}{58.44 \text{ g}} = \textbf{1.667 mol NaCl}$$
2. NaHCO₃: $$96.02 \text{ g} \times \frac{1 \text{ mol}}{84.01 \text{ g}} = \textbf{1.143 mol NaHCO}_3$$
3. Sucrose (C₁₂H₂₂O₁₁): $$96.57 \text{ g} \times \frac{1 \text{ mol}}{342.30 \text{ g}} = \textbf{0.2822 mol sucrose}$$

% Composition Calculations

NaCl:
  • Na: (22.99 ÷ 58.44) × 100 = 39.34%
  • Cl: (35.45 ÷ 58.44) × 100 = 60.66%
NaHCO₃:
  • Na: (22.99 ÷ 84.01) × 100 = 27.37%
  • H: (1.008 ÷ 84.01) × 100 = 1.20%
  • C: (12.01 ÷ 84.01) × 100 = 14.30%
  • O: (48.00 ÷ 84.01) × 100 = 57.14%
C₁₂H₂₂O₁₁:
  • C: (144.12 ÷ 342.30) × 100 = 42.10%
  • H: (22.18 ÷ 342.30) × 100 = 6.48%
  • O: (176.00 ÷ 342.30) × 100 = 51.42%

Conclusions

4. Which compound has the greatest number of moles in one spoonful?
NaCl (sodium chloride) — with 1.667 mol, it has the most moles per spoonful. This is because NaCl has the smallest molar mass, so a given mass contains more moles.
5. Which compound has the greatest total number of atoms?
NaCl — 1.667 mol × 2 atoms/formula unit × 6.022×10²³ = 2.01 × 10²⁴ total atoms (NaHCO₃: 1.143 × 6 × 6.022×10²³ = 4.13×10²⁴ — actually NaHCO₃ wins here!)
Recalculating total atoms for each:
CompoundMolesAtoms per f.u.Total Atoms
NaCl1.66722.01 × 10²⁴
NaHCO₃1.14364.13 × 10²⁴ ✅ Greatest
Sucrose0.2822457.65 × 10²³
NaHCO₃ (baking soda) has the greatest total number of atoms.
6. Why can measuring mass be used as a means of counting?
Because every element has a known, fixed atomic mass, and every compound has a known molar mass. Since one mole always contains 6.022 × 10²³ particles, measuring the mass of a sample and dividing by the molar mass tells you exactly how many moles — and therefore how many atoms or molecules — are present. Mass and particle count are directly proportional, so the scale becomes a counting tool.

The mass of nacl 27.41, nahco3 26.02, c12h22o11 26.57

Recalculated with Correct Masses

Updated Data Table

NameFormulaMass (g)Molar Mass (g/mol)Moles in Sample% Composition
1. Sodium ChlorideNaCl27.41 g58.44 g/mol0.4690 molNa: 39.34% / Cl: 60.66%
2. Baking SodaNaHCO₃26.02 g84.01 g/mol0.3097 molNa: 27.37% / H: 1.20% / C: 14.30% / O: 57.14%
3. SucroseC₁₂H₂₂O₁₁26.57 g342.30 g/mol0.07762 molC: 42.10% / H: 6.48% / O: 51.42%

Moles Calculations

1. NaCl: $$27.41 \text{ g} \times \frac{1 \text{ mol}}{58.44 \text{ g}} = \textbf{0.4690 mol NaCl}$$
2. NaHCO₃: $$26.02 \text{ g} \times \frac{1 \text{ mol}}{84.01 \text{ g}} = \textbf{0.3097 mol NaHCO}_3$$
3. Sucrose: $$26.57 \text{ g} \times \frac{1 \text{ mol}}{342.30 \text{ g}} = \textbf{0.07762 mol C}{12}\text{H}{22}\text{O}_{11}$$

Updated Conclusions

4. Greatest number of moles in one spoonful?
NaCl — 0.4690 mol (smallest molar mass = most moles per gram)
5. Greatest total number of atoms?
CompoundMolesAtoms/f.u.Total Atoms
NaCl0.469025.65 × 10²³
NaHCO₃0.309761.12 × 10²⁴ ✅ Greatest
Sucrose0.07762452.10 × 10²³
NaHCO₃ (baking soda) still has the greatest total number of atoms.
6. Why can measuring mass be used as a means of counting?
Because every substance has a known molar mass. Dividing the measured mass by the molar mass gives moles, and multiplying by 6.022 × 10²³ gives the exact number of particles. Mass and particle count are directly proportional — so a balance effectively acts as a counting tool.
This is a shared conversation. Sign in to Orris to start your own chat.