Atomic Foundations of Matter Class 9 Notes Science Chapter 9

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Class 9 Science Chapter 9 Atomic Foundations of Matter Notes

Class 9 Science Exploration Chapter 9 Notes

Class 9 Science Chapter 9 Notes – Class 9 Atomic Foundations of Matter

→ Understanding Matter and Changes: Everything around us is made up of matter, and matter is composed of extremely tiny particles called atoms. Atoms combine by losing, gaining, or sharing electrons to achieve a stable electronic configuration.

→ Physical Change: A change in which no new substance is formed. Example: Dissolving salt in water.

→ The mass is conserved during a physical change.

→ Chemical Change: A change in which new substances are formed with different properties. Example: Vinegar reacting with baking soda produces carbon dioxide gas and other substances.

• In an open system, the mass may appear to decrease because the gas escapes into the surroundings.
• In a closed system, the total mass remains unchanged before and after the reaction.

→ Law of Conservation of Mass: This law was proposed by Antoine Lavoisier.

• It states that mass can neither be created nor destroyed during a chemical reaction.
• The total mass of reactants is always equal to the total mass of products.
• This law is valid for both physical and chemical changes.

→ Example: Consider a chemical reaction,
Hydrogen + Oxygen → Water
Mass of hydrogen + Mass of oxygen = Mass of water formed.

→ Law of Constant Proportions: This law was proposed by Joseph Proust. It states that a compound always contains the same elements combined in a fixed proportion by mass.

→ Example: Water always contains hydrogen and oxygen in the ratio of 1:8 by mass. If you take 9 g of water, it will always contain 1 g hydrogen and 8 g oxygen. This ratio remains the same irrespective of the source of the compound.

• This law applies only to compounds and not to mixtures.
• Mixtures can have variable composition depending on how they are prepared.

→ Dalton’s Atomic Theory: This theory was proposed by John Dalton. It provides the basis for the modern understanding of atoms and their behaviour.

→ According to this theory:

• All matter is made up of very tiny particles called atoms, which participate in chemical reactions.
• Atoms cannot be created or destroyed during a chemical reaction. They are indivisible particles.
• Atoms of the same element are identical in mass and chemical properties.
• Atoms of different elements have different masses and chemical properties.
• Atoms combine in simple whole-number ratios to form compounds.
• The relative number and kinds of atoms are constant in a given compound.

→ Significance of Dalton’s Atomic Theory:

• Dalton’s atomic theory explains the Law of Conservation of Mass and the Law of Constant Proportions.
• It states that atoms are rearranged during a chemical reaction rather than being created or destroyed. They are indivisible.
• How Atoms Combine: Atoms combine to form molecules, which are electrically neutral and can exist independently.
• A molecule shows all the properties of the substance it represents.
• Atoms combine either by sharing electrons or by transferring electrons.
• Some elements, like helium, exist as single atoms because their atoms are already stable.

→ Stability of Atoms:

• Atoms tend to complete their valence shell to achieve a stable electronic configuration (octet or duplet).
• A complete valence shell usually contains eight electrons, or two in the case of the first shell.

→ Chemical Bond: A chemical bond is the force that holds atoms together in a compound. Bond formation leads to a decrease in energy and increases the stability of the system.

→ Covalent Bond: A covalent bond is formed by the sharing of electrons between atoms. Examples: Hydrogen molecule (H₂), Oxygen molecule (O₂), Hydrogen chloride (HCl) and Water (H₂O).

→ Types of Covalent Bonds:

• Single bond: One pair of electrons is shared (one electron from each atom). It is represented by a single line between two atoms, such as H—H.
• Double bond: Two pairs of electrons are shared. It is represented by drawing two lines between two atoms, such as (0 = 0).

→ Naming Covalent Compounds: Covalent compounds are named using prefixes to indicate the number of atoms present.

• The first element retains its name, while the second element ends with “-ide”.
• Prefixes such as mono-, di-, tri-, tetra-, penta and hexa indicate the number of atoms.
• The prefix “mono-” is usually omitted for the first element but used for the second element.
• When a prefix ends with a vowel and the element name starts with a vowel, one vowel is dropped (e.g., monoxide).
• The first element retains its regular name, while the second element ends in -ide.

→ Examples: Carbon monoxide (CO), Carbon dioxide (CO₂), and Sulfur hexafluoride (SF₆).

• When hydrogen is the first element in a formula, no prefix is used before it, regardless of the number of hydrogen atoms. Example: hydrogen sulfide (H₂S).
• Some covalent compounds are known by common names, such as water (H₂O) and ammonia (NH₃).

→ Ionic Bond: The electrostatic force of attraction between oppositely charged ions forms an ionic bond. It is formed when electrons are transferred from one atom to another.

• Atoms lose or gain electrons to form ions and achieve a stable electronic configuration.
• Atoms with fewer than four valence electrons tend to lose electrons to achieve stability.
• Atoms with more than four valence electrons tend to gain electrons to complete their octet.
• Positively charged ions are called cations, while negatively charged ions are called anions. Example: Sodium (Na⁺) and chloride ions (Cl^–) combine to form sodium chloride (NaCl).

→ Structure of Ionic Compounds: Ionic compounds do not exist as single molecules but form three-dimensional crystal structures.

→ Example: In sodium chloride, each sodium ion is surrounded by six chloride ions and vice versa.

→ The regular arrangement of ions in space is called a crystal lattice. The crystal lattice helps to maintain strong forces of attraction between ions.

→ Naming Ionic Compounds: Ionic compounds are named by writing the name of the cation first, followed by the anion.

• The names of simple anions end with “-ide”.
• Metals usually form cations, while non-metals form anions.
• Some ions consist of a group of atoms and are called polyatomic ions.
• Polyatomic ions have specific names that do not end with “-ide”.
• Examples: NaCl → sodium chloride, MgO → magnesium oxide

→ Chemical Formulae of Covalent Compounds: It represents the number of atoms of each element present. The formula is written using the symbols and valencies of the elements. To write the chemical formula of a covalent compound, follow these steps:

• Write the symbols of the constituent elements of the compound.
• Write the valencies of these elements.
• Cross over the valencies of the combining atoms and write them as subscripts after the symbols of elements. If the valency is one, it is not written in the formula.

→ Example:
The formula of hydrogen sulfide-

→ Chemical Formulae of Ionic Compounds: The formula of an ionic compound is written by balancing the charges of cations and anions.

• Write the symbol of the cation first, followed by the symbol of the anion.
• Write the charges under the symbols rather than as superscripts.
• Cross over the charges (only the numbers) to obtain the formula.
• The chemical formula gives the simplest ratio of the elements in a compound. Therefore, after criss-crossing, the subscripts are divided by a common factor, if any.
• The overall compound must be electrically neutral.
• Brackets are used when more than one polyatomic ion is present in the formula.

→ Properties of Ionic and Covalent Compounds:

→ Solubility

• Ionic compounds are generally soluble in water but insoluble in organic solvents like kerosene and petrol.
• Covalent compounds are generally insoluble in water but soluble in organic solvents.

→ Electrical Conductivity

• Ionic compounds do not conduct electricity in the solid state because ions are fixed in position by strong forces.
• Ionic compounds conduct electricity in aqueous or molten state because ions are free to move.
• Covalent compounds do not conduct electricity because they do not form ions.

→ Melting and Boiling Points

• Ionic compounds have high melting and boiling points due to strong electrostatic forces.
• Covalent compounds generally have low melting and boiling points due to weaker intermolecular forces.

→ Molecular Mass of Covalent Compounds:
Molecular mass is the sum of the atomic masses of all atoms present in a molecule. It is calculated by adding the atomic masses according to the formula of the compound. Example: The molecular mass of water (H₂O) is 18 u.

Formula Unit Mass of Ionic Compounds: Ionic compounds do not exist as molecules but as formula units. A formula unit represents the simplest whole number ratio of ions in an ionic compound. Formula unit mass is the sum of the atomic masses of all atoms in the formula unit (simplest whole number ratio of ions in an ionic compound).
Example: The formula unit mass of Na₂O = (2 × 23 u)+(1 × 16 u) = 62 u

Atomic Foundations of Matter Notes Class 9

Law of Conservation of Mass

• This law was proposed by the French scientist Lavoisier in 1789. One of the most fundamental principles of chemistry is that matter is neither created nor destroyed in any chemical reaction.
• This means that the total mass of the substances present before a chemical reaction (the reactants) is always equal to the total mass of the substances formed after the reaction (the products).
• This principle is known as the Law of Conservation of Mass.
• To understand this law, consider the reaction between sodium sulphate and barium chloride solutions.
• When these two solutions are mixed, a white precipitate of barium sulphate is formed along with sodium chloride in solution. The reaction can be represented as
  Sodium sulphate + Barium chloride → Barium sulphate + Sodium chloride
• When the masses of the reactants and products are carefully measured in a closed system, the total mass before and after the reaction is found to be equal. This confirms the Law of Conservation of Mass.
• When a gas is produced in a chemical reaction (such as carbon dioxide from the reaction of vinegar and baking soda), the experiment must be performed in a closed system.
• If the system is open, the escaping gas causes an apparent loss in mass, which may mislead students into thinking mass is not conserved. In reality, the mass is conserved, the gas simply escapes into the surroundings.

Antoine Lavoisier (1743 – 1794) is known as the Father of Modern Chemistry. He proposed the Law of Conservation of Mass, which applies to every chemical reaction. Lavoisier stated that in every operation an equal quantity of matter exists both before and after the operation. His careful experiments using closed containers and weighing made this law an important foundation of modern chemistry.

Example 1.
In a group activity, students placed 4.0 g of calcium carbonate with 2.92 g of hydrochloric acid in a closed container. After the reaction, they measured 1.76 g of carbon dioxide, 0.72 g of water, and 4.44 g of calcium chloride. Verify whether the law of conservation of mass is obeyed.
Solution:
Mass of reactants:
Calcium carbonate = 4.0 g
Hydrochloric acid = 2.92 g
Total mass of reactants = 4.0 + 2.92 = 6.92 g

Mass of products:
Carbon dioxide = 1.76 g
Water = 0.72 g
Calcium chloride = 4.44 g
Total mass of products = 1.76 + 0.72 + 4.44 = 6.92 g
Since, the mass of reactants = mass of products = 6.92 g, the Law of Conservation of Mass is obeyed.

Example 2.
12 g of carbon combines with 32 g of oxygen to form 44 g of carbon dioxide. If 2.4 g of carbon reacts completely with oxygen, how much carbon dioxide will be produced?
Carbon + Oxygen → Carbon dioxide
Solution:
Given, 12 g of carbon → 44 g of carbon dioxide
So, 1 g of carbon produces = \(\frac{44}{12}\) g of carbon dioxide
Thus, 2.4gof carbon produces = \(\frac{44}{12}\) × 2.4 g =8.8 g of carbon dioxide.

→ Law of Constant Proportions:

• This law is given by, the French chemist Joseph Louis Proust. He observed that in any given chemical compound, the elements that make up the compound are always present in a fixed ratio by mass, regardless of the source of the compound or the method by which it was prepared.
• This is known as the Law of Constant Proportions, also called the Law of Definite Proportions or Proust’s Law.
• According to this law “a chemical compound always contains its constituent elements in a fixed mass ratio”.
• For example, water obtained from any source-rivers, borewells, rain, or the ocean when purified, always contains hydrogen and oxygen in the mass ratio of 1 : 8. This means that in 9 g of pure water, there is always 1 g of hydrogen and 8 g of oxygen. This proportion does not change regardless of the source or the method of preparation of water.

Joseph Louis Proust (1754 – 1826) was a prominent French chemist known for his careful experimental work. He contributed to the Law of Definite Proportions by showing that chemical compounds always contain their elements in fixed ratios by mass. His work laid an important foundation that helped shape modern chemistry.

Example 3.
Sodium chloride (NaCl) contains sodium and chlorine in the mass ratio of 23 : 35.5. If 46 g of sodium reacts completely, how much chlorine is needed to form NaCl?
Solution:
Mass of chlorine required = (\(\frac{35.5}{23}\)) × 46 = 71 g

In ancient civilisations, a red pigment 1 obtained from rocks was widely used for painting and colouring objects. In India, it was known as hingula, while in Latin and English it was called cinnabar. Cinnabar Overtime, it was discovered that heating cinnabar yields mercury and sulphur in the approximate proportions of 86.22% and 13.78%, respectively.

It was also observed that grinding mercury and sulphur in this ratio could form cinnabar, but due to the toxic nature of both substances, this method was not widely used.

Dalton’s Atomic Theory

• The two laws, the Law of Conservation of Mass and the Law of Constant Proportions formed the basis for a comprehensive theoretical explanation of matter.
• This explanation was provided by the English scientist John Dalton in 1808, in the form of his Atomic Theory. This theory was the first modern scientific attempt to explain why elements combine in fixed ratios and why mass is conserved during chemical reactions.

John Dalton (1766 – 1844) was born in England. In 1793, he moved to Manchester to teach mathematics, physics and chemistry. He spent most of his life in teaching and researching. In 1808, he presented his atomic theory, which proved to be a turning point in the study of matter. He proposed that atoms are the fundamental building blocks of all matter.

→ Postulates of Dalton’s Atomic Theory:

John Dalton proposed the following postulates

• All matter is made up of very tiny, indivisible particles called atoms, which participate in chemical reactions.
• Atoms are indivisible particles, which cannot be created or destroyed in a chemical reaction.
• Atoms of a given element are identical in mass and chemical properties.
• Atoms of different elements have different masses and chemical properties.
• Atoms combine in the ratio of simple whole numbers to form compounds.
• The relative number and kinds of atoms in a given compound are always constant.

Some of Dalton’s original postulates were later revised in the light of new discoveries. For example, the discovery of subatomic particles (electrons, protons and neutrons) showed that atoms are, in fact, divisible. However, Dalton’s theory remains a crucial milestone because it provided the first systematic, scientific model of atomic behaviour.

→ How Atoms Combine?

• Atoms of elements can combine with each other to form molecules. A molecule may be defined as an electrically neutral entity consisting of more than one atom that is capable of independent existence and shows all the properties of that substance.
• Atoms combine because the total energy of the combined system is lower than the sum of energies of the individual atoms, making the resulting arrangement more stable. The force that holds atoms together in a molecule is called a chemical bond. Atoms generally combine in two ways
  • Sharing of electrons: Share a few or all of their valence electrons with another atom.
  • Transfer of electrons: One or more valence electrons are transferred from one atom to another atom, or accepted from some other atom.
• Some elements, such as helium and neon, exist only as single atoms because their outermost shells are already complete, making them stable without any bonding.

Bonding by Sharing of Electrons (Covalent Bond)

When atoms share one or more pairs of electrons to achieve a stable electronic configuration, the bond formed is called a covalent bond. The compound formed is called a covalent compound.

A. Molecules of Elements
Atoms of the same element can share electrons with each other to form molecules. Some important examples are

→ Formation of Hydrogen Molecule (H₂)

• Hydrogen has atomic number 1. Its only electron is in the K-shell, which can hold a maximum of two electrons. Therefore, each hydrogen atom needs one more electron to become stable.
• To achieve this, two hydrogen atoms each share their single electron with each other, forming a shared pair of electrons. This shared pair holds both atoms together, forming a hydrogen molecule (H₂).
  H + H → H – H (or H₂)
• This type of bond formed by sharing of one electron pair between two atoms is called a single covalent bond, represented by a single line (-) between the atomic symbols.

→ Formation of Chlorine Molecule (Cl₂):

• Chlorine has atomic number 17 and seven electrons in its valence shell. Each. chlorine atom needs one more electron to complete its octet.
• Two chlorine atoms therefore share one electron each, forming a shared pair that holds them together. The chlorine molecule (Cl₂) is thus formed and represented as Cl-Cl.

→ Formation of Oxygen Molecule (O₂):

• Oxygen has atomic number 8 and six electrons in its valence shell. Each oxygen atom requires two more electrons to complete its octet.
• Two oxygen atoms therefore share two electrons each, forming two shared pairs between them.
• This results in a double covalent bond, represented as O = O. The oxygen molecule (O₂) is formed.

Note:
In nitrogen (N₂), nitrogen has five valence electrons and needs three more to complete its octet. Two nitrogen atoms therefore share three electron pairs each, forming a triple covalent bond (N ≡ N). This triple bond makes the nitrogen molecule extremely stable and relatively unreactive under normal conditions.

B. Molecules of Compounds

When atoms of two different elements combine by sharing electrons, they form a covalent compound. Some important examples are

→ Formation of Hydrogen Chloride (HCl):

• Hydrogen needs one electron and chlorine needs one electron to complete their respective stable configurations. Each atom shares its one electron with the other, forming a single covalent bond.
• The resulting molecule, hydrogen chloride (HCl), is a covalent compound represented as H—Cl.

→ Formation of Water (H₂O):

Oxygen needs two electrons to complete its octet, while each hydrogen atom needs only one. Therefore, two hydrogen atoms each share one electron with the same oxygen atom, resulting in two O—H single covalent bonds. The water molecule (H₂O) thus contains one oxygen atom and two hydrogen atoms.

→ Bridging Science and Society:

Atoms release energy through nuclear reactions, which is used in electricity generation, medicine, research and space exploration. In nuclear power plants, this energy produces steam to run turbines, providing a cleaner alternative to fossil fuels. In India, Raja Ramanna played a key role in developing and promoting the peaceful use of nuclear energy for development.

C. Naming Covalent Compounds

• Covalent compounds are named using a prefix system that indicates the number of atoms of each element present. The first element retains its name, while the second element uses the suffix -ide.
• Common prefixes are: mono- (1), di- (2), tri- (3), tetra- (4), penta- (5), hexa- (6), etc. Note that ‘mono-’ is generally omitted for the first element.
• If a prefix ends with ‘o’ or ‘a and the element starts with a vowel the last vowel (for example, monoxide, pentoxide). If the prefix ends with ‘i’, keep it for pronunciation (for example, dioxide trioxide).

Formula	Name
CO	Carbon monoxide
CO2	Carbon dioxide
CS2	Carbon disulphide
PCl3	Phosphorus trichloride
SF6	Sulphur hexafluoride
N2O4	Dinitrogen tetroxide
N2O5	Dinitrogen pentoxide
H2S	Hydrogen sulphide

Note:
When hydrogen is the first element in the formula, no prefix is added before hydrogen, regardless of the number of its atoms. For example, H₂S is named ‘hydrogen sulphide’, not ‘dihydrogen sulphide’. Some compounds are known only by common names: H₂O is ‘water’ and NH₃ is ‘ammonia’.

Bonding by Electron Transfer (Ionic Bond)

When the valence shell of an atom has fewer than four electrons, the atom tends to lose those electrons to achieve a stable configuration. Conversely, atoms with more than four valence electrons tend to gain or share electrons to complete an octet. When one atom donates electrons to another atom, oppositely charged ions are formed and held together by electrostatic attraction. This type of chemical bond is called an ionic bond, and the compounds formed are called ionic compounds.

→ Formation of Sodium Chloride (NaCl):

• Consider the formation of common salt, sodium chloride (NaCl). The atomic number of sodium (Na) is 11, and its electronic configuration is 2, 8, 1. Its valence shell has only one electron.
• When the sodium atom loses its one valence electron, it has 11 protons but only 10 electrons. This makes it a positively charged ion, called a sodium cation, represented as Na⁺.

Chlorine (Cl) has atomic number 17 and electronic configuration 2, 8, 7. It has seven valence electrons and needs one more electron to complete its octet.

When chlorine gains the electron lost by sodium, it has 17 protons but 18 electrons. This makes it a negatively charged ion, called a chloride anion, represented as Cl^–.

The positively charged sodium ion (Na⁺) and the negatively charged chloride ion (Cl^–) are attracted to each other by a strong electrostatic force. This electrostatic force of attraction between oppositely charged ions is what we call an ionic bond. The compound formed, sodium chloride, is electrically neutral because the positive and negative charges balance each other.

Na + Cl → Na⁺ + Cl^– → NaCl

Note:
Ionic compounds, such as sodium chloride, do not exist as simple molecules. Instead, they form large three – dimensional crystal structures. In the NaCl crystal, each Na⁺ ion is surrounded by six Cl^– ions, and each Cl^– ion is surrounded by six Na⁺ ions, arranged in a regular repeating pattern known as a crystal lattice.

→ Naming Ionic Compounds:

• In naming ionic compounds, the name of the cation is written first, followed by the name of the anion. Names of simple anions end with the suffix -ide (e.g. chloride, oxide, sulphide). Generally, metals form cations and non-metals form anions.
• Cations and anions are collectively called ions. Some ions are formed by a combination of two or more atoms. These are called polyatomic ions. Names of polyatomic ions generally do not end with ‘-ide’. Common examples are sulphate (\(\mathrm{SO}_4^{2-}\)), nitrate (\(\mathrm{NO}_3^{-}\)), and hydroxide (OH^–).

Some Common Monoatomic Ions

Some Common Polyatomic Ions

Name of Ion	Formula	Valency
Hydroxide	OH–	1
Nitrate	\(\mathrm{NO}_3^{-}\)	1
Hydrogencarbonate	\(\mathrm{HCO}_3^{-}\)	1
Carbonate	\(\mathrm{CO}_3^{2-}\)	2
Sulphate	\(\mathrm{SO}_4^{2-}\)	2
Ammonium	\(\mathrm{NH}_4^{+}\)	1

Writing Chemical Formulae

There is a convenient shorthand method to write the formulae of chemical compounds using the criss-cross method. This involves exchanging the valencies (or charges) of the combining elements or ions to determine the subscripts in the formula.

→ Writing Chemical Formulae of Covalent Compounds:

Follow these steps to write the chemical formula of a covalent compound
(i) Write the symbols of the constituent elements of the compound.
(ii) Write the valencies of these elements below their symbols.
(iii) Criss-cross the valencies and write them as subscripts after the respective element symbols.
(iv) If the subscripts can be reduced by a common factor, simplify them.

Examples:
→ The formula of hydrogen chloride

The formula of the compound would be HCl.
If the valency is one after criss-crossing, it is not written.

→ The formula of hydrogen sulphide

The formula of the compound would be H₂S.

→ The formula of carbon tetrachloride

The formula of the compound would be CCl₄.

→ Writing Chemical Formulae of Ionic Compounds:

Follow these steps to write the chemical formula of an ionic compound.
(i) Write the symbol of the cation first, followed by the symbol of the anion.
(ii) Write the charges below the respective symbols.
(iii) Criss-cross the charges (only the numbers) as subscripts.
(iv) Simplify by dividing by a common factor if possible.
(v) If a polyatomic ion appears more than once in the formula, enclose it in brackets before adding the subscript.

Note:
The charges on the ions are not indicated in the final formula of the compound. Also, when the valencies of both ions are equal, simplify so that the formula shows the simplest ratio (e.g. Mg₂O₂ simplifies to MgO).

Example:
→ The formula of calcium chloride

Thus, in calcium chloride, there are two chloride ions (Cl^–) for each calcium ion (Ca²⁺). The positive and negative charges must balance each other, and the overall structure must be neutral.

→ The formula of aluminium oxide

→ The formula for magnesium oxide-

Here, the valencies of the two elements are the same. We arrive at the formula Mg₂O₂ but it is simply written as MgO.
This method can also be used to write formulae of compounds of metal with other polyatomic ions, such as calcium carbonate.

Example:
→ The formula for calcium carbonate

Here, the valencies of the two ions are the same.
The formula Ca₂(CO₃)₂ is simply written as CaCO₃, as explained above.

→ The formula of magnesium hydroxide Symbol

Properties of Ionic and Covalent Compounds

Ionic and covalent compounds have distinctly different physical properties. These differences arise directly from the nature of the bonds that hold them together.

A. Solubility:

• Ionic compounds (such as sodium chloride and copper sulphate) are generally soluble in water but insoluble in organic solvents such as kerosene and petrol.
• Covalent compounds (such as camphor and naphthalene) are generally insoluble in water but dissolve readily in organic solvents like kerosene and petrol.
• This difference is explained by the principle ‘like dissolves like’. Water is a polar solvent, so it dissolves polar ionic compounds. Organic solvents are non-polar, so they dissolve non-polar covalent compounds.

B. Electrical Conductivity:

• Ionic compounds do not conduct electricity in the solid state because their ions are fixed in their positions in the crystal lattice and cannot move freely.
• When ionic compounds are dissolved in water, the ions become free to move, allowing them to conduct electricity. Thus, solutions of ionic compounds are good conductors.
• Most covalent compounds, even when dissolved in water (e.g. sugar), do not conduct electricity because they do not produce ions in solution.

C. Melting and Boiling Points:

• Ionic compounds generally have high melting and boiling points due to the strong electrostatic forces of attraction between their oppositely charged ions.
• Covalent compounds generally have low melting and boiling points because the intermolecular forces holding their molecules together are relatively weak.

Comparison of Ionic and Covalent Compounds

Property	Ionic Compounds	Covalent Compounds
Solubility in water	Generally soluble	Generally insoluble
Solubility in organic solvents	Generally insoluble	Generally soluble
Electrical conductivity (solid)	Do not conduct	Do not conduct
Electrical conductivity (solution)	Conduct electricity	Generally do not conduct
Melting and boiling points	High	Low
Bond type	Ionic bond (electron transfer)	Covalent bond (electron sharing)

Molecular Mass of Covalent Compounds

The molecular mass of a covalent compound is the total mass of one molecule of that compound, calculated by adding the atomic masses of all the atoms present in it. Atomic masses are expressed in atomic mass units (u).

Example 4.
Molecular mass of water (H₂O)
Solution:
Atomic masses: H = 1 u; O = 16 u
Molecular mass of H₂O = (1 u × 2) + (16 u × 1)
= 2 + 16 = 18 u.

→ Formula Unit Mass of Ionic Compounds

Since ionic compounds do not form discrete molecules but instead form large crystal lattices, the concept of molecular mass does not apply to them. Instead, we use the term formula unit. A formula unit is the collection of ions in the simplest whole number ratio as represented by the chemical formula. The mass of a formula unit is called the formula unit mass.

The formula unit mass is calculated in the same way as molecular mass — by adding the atomic masses of all the atoms present in the formula unit.

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