Some Basic Concepts of Chemistry

Chapter overview

Chemistry deals with composition, structure, properties, and transformations of matter. Quantitative chemistry begins with atoms, molecules, formulae, moles, masses, and balanced chemical equations.

1. Matter and classification

Matter is anything that has mass and occupies space.

2. Laws of chemical combination

Law of conservation of mass

Mass is neither created nor destroyed in a chemical reaction.

$$\text{Total mass of reactants} = \text{Total mass of products}$$

Law of definite proportions

A given compound always contains the same elements in the same fixed proportion by mass.

Law of multiple proportions

When two elements combine to form more than one compound, the masses of one element that combine with a fixed mass of the other are in a simple whole-number ratio.

Gay-Lussac's law

Gases combine in simple whole-number volume ratios at the same temperature and pressure.

Avogadro's law

Equal volumes of gases at the same temperature and pressure contain equal number of molecules.

$$V\propto n$$

3. Atomic mass, molecular mass, and formula mass

Atomic mass is expressed in atomic mass unit.

$$1\ \text{u} = \frac{1}{12} \times \text{mass of one carbon-12 atom}$$

Molecular mass:

$$\text{Molecular mass} = \sum \text{atomic masses of atoms in molecule}$$

Example:

$$H_2O=2(1)+16=18\ \text{u}$$

Formula mass for ionic compounds:

$$NaCl=23+35.5=58.5\ \text{u}$$

4. Mole concept

One mole contains Avogadro number of entities.

$$N_A=6.022\times10^{23}\ \text{mol}^{-1}$$

Number of moles:

$$n=\frac{m}{M}$$

where:

• $n$ = number of moles
• $m$ = given mass
• $M$ = molar mass

Number of particles:

$$N=nN_A$$

Mass from number of particles:

$$m=\frac{N}{N_A}\times M$$

5. Molar volume of gas

At STP, one mole of an ideal gas occupies approximately:

$$22.4\ \text{L}$$

Gas mole relation:

$$n=\frac{V}{22.4} \quad \text{at STP}$$

Ideal gas equation:

$$PV=nRT$$

6. Percentage composition

Mass percentage of an element:

$$\%\text{ of element} = \frac{\text{mass of element in compound}} {\text{molar mass of compound}} \times100$$

Example for water:

$$\%\ H = \frac{2}{18}\times100 = 11.11\%$$ $$\%\ O = \frac{16}{18}\times100 = 88.89\%$$

7. Empirical formula and molecular formula

Empirical formula gives the simplest whole-number ratio of atoms.

Molecular formula gives the actual number of atoms.

Relation:

$$\text{Molecular formula} = n\times \text{Empirical formula}$$

where:

$$n= \frac{\text{molecular mass}} {\text{empirical formula mass}}$$

Example:

$$\text{Empirical formula}=CH_2O$$

If molecular mass is $180$ and empirical formula mass is $30$:

$$n=\frac{180}{30}=6$$

Therefore:

$$\text{Molecular formula}=C_6H_{12}O_6$$

8. Stoichiometry

Stoichiometry uses balanced chemical equations to calculate amounts of reactants and products.

Example:

$$2H_2+O_2\rightarrow2H_2O$$

This means:

$$2\ \text{mol }H_2 + 1\ \text{mol }O_2 \rightarrow 2\ \text{mol }H_2O$$

For methane combustion:

$$CH_4+2O_2\rightarrow CO_2+2H_2O$$

Mole ratio:

$$CH_4:O_2:CO_2:H_2O=1:2:1:2$$

9. Limiting reagent

The limiting reagent is completely consumed first and limits the amount of product formed.

Example:

$$N_2+3H_2\rightarrow2NH_3$$

If $1$ mol $N_2$ and $2$ mol $H_2$ are available, then $H_2$ is limiting because $1$ mol $N_2$ needs $3$ mol $H_2$.

10. Concentration terms

Molarity

$$M= \frac{\text{moles of solute}} {\text{volume of solution in litre}}$$

Molality

$$m= \frac{\text{moles of solute}} {\text{mass of solvent in kg}}$$

Mole fraction

$$X_A= \frac{n_A}{n_A+n_B}$$ $$X_A+X_B=1$$

Mass percentage

$$\text{Mass percent} = \frac{\text{mass of solute}} {\text{mass of solution}} \times100$$

Parts per million

$$ppm= \frac{\text{mass of solute}} {\text{mass of solution}} \times10^6$$

11. Balanced equation examples

Neutralisation:

$$HCl+NaOH\rightarrow NaCl+H_2O$$

Precipitation:

$$AgNO_3+NaCl\rightarrow AgCl\downarrow+NaNO_3$$

Combustion:

$$C_3H_8+5O_2\rightarrow3CO_2+4H_2O$$

Oxidation of magnesium:

$$2Mg+O_2\rightarrow2MgO$$

12. Optional mhchem rendering tests

If MathJax mhchem is enabled, these should render nicely:

$$\ce{2H2 + O2 -> 2H2O}$$ $$\ce{AgNO3 + NaCl -> AgCl v + NaNO3}$$ $$\ce{CH4 + 2O2 -> CO2 + 2H2O}$$

13. Solved examples

Example 1

Find moles in $18$ g of water.

$$n=\frac{m}{M} = \frac{18}{18} =1\ \text{mol}$$

Example 2

Find number of molecules in $2$ mol of $CO_2$.

$$N=nN_A = 2\times6.022\times10^{23}$$ $$N=1.2044\times10^{24}$$

Example 3

Find molarity if $0.5$ mol solute is dissolved in $250$ mL solution.

$$V=250\ \text{mL}=0.250\ \text{L}$$ $$M=\frac{0.5}{0.250}=2\ \text{M}$$

14. Common mistakes

• Using atomic mass where molecular mass is required.
• Forgetting to balance chemical equation before stoichiometry.
• Confusing molarity and molality.
• Using mL directly in molarity instead of litre.
• Ignoring limiting reagent.
• Treating empirical formula and molecular formula as always same.

15. Quick revision