Mole Concept flash cards

1. 1.What is a mole?

   A mole is the amount of substance containing as many elementary entities (atoms, molecules, ions, etc.) as there are atoms in exactly $12$ g of $^{12}C$. It equals $6.022\times10^{23}$ entities (Avogadro number).

   Hint: Defined via carbon-12.

2. 2.State the value of Avogadro's number ($N_A$) and its significance.

   $N_A = 6.022\times10^{23}\ \text{mol}^{-1}$. It is the number of elementary entities in one mole of any substance, linking the macroscopic (grams) and microscopic (atoms/molecules) scales.

   Hint: $6.022\times10^{23}$.

3. 3.Give the three key formulas to compute number of moles ($n$).

   $n = \dfrac{\text{mass}}{\text{molar mass}} = \dfrac{N}{N_A} = \dfrac{V}{22.4\ \text{L}}$ (gas at STP), where $N$ = number of particles.

   Hint: Mass, particles, or gas volume.

4. 4.What is molar mass?

   The mass of one mole of a substance, expressed in $\text{g mol}^{-1}$. Numerically equal to the atomic/molecular mass in amu.

   Hint: Mass per mole.

5. 5.Define gram atomic mass and gram molecular mass.

   Gram atomic mass = mass in grams of $1$ mole of atoms (numerically = atomic mass). Gram molecular mass = mass in grams of $1$ mole of molecules (numerically = molecular mass).

   Hint: One mole of atoms vs molecules.

6. 6.What is the atomic mass unit (amu / u)?

   $1$ amu $= \dfrac{1}{12}$ of the mass of one $^{12}C$ atom $= 1.66\times10^{-24}$ g. It is the standard unit for expressing atomic and molecular masses.

   Hint: 1/12 of carbon-12 mass.

7. 7.Distinguish average atomic mass from the mass of a single isotope.

   Average atomic mass is the weighted mean of the masses of all naturally occurring isotopes, weighted by their fractional abundances. A single isotope has one fixed mass; the periodic-table value is the average.

   Hint: Weighted by abundance.

8. 8.How do you calculate average atomic mass from isotopic data?

   $\bar{A} = \sum (f_i \times m_i)$, where $f_i$ is the fractional abundance and $m_i$ the mass of isotope $i$. E.g. for $Cl$: $0.75\times35 + 0.25\times37 = 35.5$.

   Hint: Sum of (abundance × mass).

9. 9.What is a gram equivalent?

   The mass of substance in grams equal to its equivalent weight. Number of gram equivalents $= \dfrac{\text{mass}}{\text{equivalent weight}}$.

   Hint: Mass / equivalent weight.

10. 10.Define equivalent weight.

    Equivalent weight $= \dfrac{\text{molar mass}}{n\text{-factor}}$. It is the mass that combines with or displaces $1$ mole of $H$ atoms (or $1$ mole of electrons in redox).

    Hint: Molar mass / n-factor.

11. 11.What is the n-factor for an acid, a base, and a salt?

    Acid: number of replaceable $H^+$ ions (basicity). Base: number of replaceable $OH^-$ ions (acidity). Salt: total charge of cations (or anions).

    Hint: Replaceable H+, OH-, or total charge.

12. 12.What is the n-factor in a redox reaction?

    The number of moles of electrons gained or lost per mole of the species (i.e. the total change in oxidation number per formula unit).

    Hint: Electrons transferred per mole.

13. 13.Give the equivalent weight of $H_2SO_4$ and $Ca(OH)_2$.

    $H_2SO_4$: molar mass $98$, n-factor $2$, so eq. wt. $= 49$. $Ca(OH)_2$: molar mass $74$, n-factor $2$, so eq. wt. $= 37$.

    Hint: Divide molar mass by n-factor (both = 2).

14. 14.What is the equivalent weight of $KMnO_4$ in acidic medium?

    In acidic medium $Mn$ goes from $+7$ to $+2$, so n-factor $= 5$. Eq. wt. $= \dfrac{158}{5} = 31.6$.

    Hint: n-factor = 5 (Mn: +7 → +2).

15. 15.What is the equivalent weight of $KMnO_4$ in neutral / faintly alkaline medium?

    $Mn$ goes $+7 \to +4$ (forms $MnO_2$), so n-factor $= 3$. Eq. wt. $= \dfrac{158}{3} = 52.7$.

    Hint: n-factor = 3 (Mn: +7 → +4).

16. 16.Define empirical formula.

    The simplest whole-number ratio of atoms of each element in a compound. E.g. the empirical formula of $C_6H_{12}O_6$ is $CH_2O$.

    Hint: Simplest whole-number atom ratio.

17. 17.Define molecular formula and relate it to the empirical formula.

    The molecular formula gives the actual number of atoms of each element in a molecule. Molecular formula $= (\text{empirical formula})_n$, where $n = \dfrac{\text{molecular mass}}{\text{empirical formula mass}}$.

    Hint: n × empirical formula.

18. 18.Outline the steps to determine an empirical formula from percentage composition.

    1) Take $100$ g sample so % = mass in g. 2) Divide each mass by its atomic mass to get moles. 3) Divide all mole values by the smallest. 4) Convert to nearest whole numbers (multiply if needed).

    Hint: Mass → moles → ratio → whole numbers.

19. 19.How is percentage composition of an element in a compound calculated?

    $\%\ \text{element} = \dfrac{\text{(atomic mass)}\times\text{(no. of atoms)}}{\text{molar mass of compound}}\times100$.

    Hint: Mass of element / molar mass × 100.

20. 20.What is stoichiometry?

    The quantitative study of the relationships (mass, mole, volume) between reactants and products in a balanced chemical equation, based on the mole ratios given by the coefficients.

    Hint: Quantitative mole relationships in reactions.

21. 21.What information do the coefficients in a balanced equation give?

    They give the relative number of moles (and, for gases at same T,P, the volume ratio) of reactants and products. They do NOT directly give mass ratios.

    Hint: Mole (and gas volume) ratios.

22. 22.What is a limiting reagent?

    The reactant that is completely consumed first in a reaction, thereby limiting the amount of product formed. The other reactant(s) are in excess.

    Hint: Runs out first; caps the product.

23. 23.How do you identify the limiting reagent?

    Divide the moles of each reactant by its stoichiometric coefficient. The reactant with the smallest such ratio is the limiting reagent.

    Hint: Smallest (moles ÷ coefficient) ratio.

24. 24.Define molarity ($M$) and give its formula.

    Molarity is moles of solute per litre of solution: $M = \dfrac{n_{\text{solute}}}{V_{\text{solution (L)}}}$. Unit: $\text{mol L}^{-1}$.

    Hint: Moles solute per litre solution.

25. 25.Why does molarity change with temperature but molality does not?

    Molarity depends on volume of solution, which expands/contracts with temperature; molality depends only on mass of solvent, which is temperature-independent.

    Hint: Volume varies with T; mass does not.

26. 26.Define molality ($m$) and give its formula.

    Molality is moles of solute per kilogram of solvent: $m = \dfrac{n_{\text{solute}}}{\text{mass of solvent (kg)}}$. Unit: $\text{mol kg}^{-1}$.

    Hint: Moles solute per kg solvent.

27. 27.Define normality ($N$) and relate it to molarity.

    Normality is gram equivalents of solute per litre of solution: $N = \dfrac{\text{gram equivalents}}{V(\text{L})}$. Relation: $N = M \times n\text{-factor}$.

    Hint: N = M × n-factor.

28. 28.Define mole fraction and state a key property.

    Mole fraction of a component $= \dfrac{\text{moles of that component}}{\text{total moles}}$. It is dimensionless, temperature-independent, and the mole fractions of all components sum to $1$.

    Hint: Component moles / total moles; sum = 1.

29. 29.Define mass percentage (%w/w) of a solution.

    $\%\text{w/w} = \dfrac{\text{mass of solute}}{\text{mass of solution}}\times100$. E.g. $10\%$ w/w means $10$ g solute in $100$ g solution.

    Hint: Mass solute / mass solution × 100.

30. 30.What is ppm (parts per million) as a concentration term?

    $\text{ppm} = \dfrac{\text{mass of solute}}{\text{mass of solution}}\times10^6$. Used for very dilute solutions (e.g. trace impurities in water).

    Hint: Mass ratio × 10^6.

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