Atoms and Molecules - Class 9
Atoms and Molecules
1 Historical Development of Atomic Theory
1.1 Ancient Indian and Greek Philosophers
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Concept of Divisibility
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Ancient Indian and Greek philosophers have always wondered about the unknown and unseen
form of matter. They speculated that matter could be divided into smaller and smaller particles
until the smallest, indivisible unit is reached.
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Contributions of Ancient Indian Philosophers:
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• Around 500 BC, Indian philosopher Maharishi Kanad postulated that if we continue to divide
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matter (padarth), we shall get smaller and smaller particles. Ultimately, a stage will come when
further division will not be possible. He named these smallest particles Parmanu.
• Another Indian philosopher, Pakudha Katyayama, further elaborated this concept by suggesting
that these particles exist in a combined form to give us various forms of matter.
Contributions of Ancient Greek Philosophers:
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• Around the same era, Greek philosophers Democritus and Leucippus suggested that continuous
division of matter would eventually result in the smallest indivisible particles called atoms, meaning
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”indivisible” in Greek.
• Their ideas were based on philosophical considerations, and no experimental validation was possible
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until the eighteenth century.
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1.2 Modern Developments in Atomic Theory
Scientific Recognition
By the end of the eighteenth century, scientists recognized the difference between elements and
compounds and naturally became interested in understanding how and why elements combine.
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Contribution of Antoine L. Lavoisier:
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• Antoine L. Lavoisier, a French chemist, laid the foundation of modern chemical sciences.
• He established two fundamental laws of chemical combination, which provided a systematic ap-
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proach to understanding the combination of elements.
Key Insight
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Lavoisier’s experiments showed that mass is conserved in chemical reactions, which led to the
formulation of the Law of Conservation of Mass.
2 Laws of Chemical Combination
The following two laws of chemical combination were established after much experimentation by Lavoisier
and Joseph L. Proust.
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Atoms and Molecules - Class 9
2.1 Law of Conservation of Mass
Law of Conservation of Mass
Mass can neither be created nor destroyed in a chemical reaction.
Explanation: Is there a change in mass when a chemical reaction takes place? According to the law,
during any physical or chemical change, the total mass of reactants equals the total mass of products.
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Numerical Example
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If 10 g of calcium carbonate is decomposed, it gives 5.6 g of calcium oxide and 4.4 g of carbon
dioxide. Verify the law of conservation of mass.
Solution: Total mass of reactants = 10 g
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Total mass of products = 5.6 + 4.4 = 10 g
Answer: The total mass is conserved.
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2.2 Law of Constant Proportions
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Law of Constant Proportions
In a chemical substance, the elements are always present in definite proportions by mass.
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Explanation: Lavoisier and other scientists noted that compounds always contain elements in the
same proportion by mass, regardless of their source or preparation method.
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Examples:
• In water (H2 O), the ratio of the mass of hydrogen to the mass of oxygen is always 1 : 8.
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• In ammonia (N H3 ), nitrogen and hydrogen are always present in the mass ratio 14 : 3.
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Numerical Example
If 36 g of water is decomposed, determine the masses of hydrogen and oxygen obtained.
Solution: Mass ratio of hydrogen to oxygen in water is 1 : 8.
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Mass of hydrogen = × 36 = 4 g
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Mass of oxygen = × 36 = 32 g
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Answer: The decomposition yields 4 g of hydrogen and 32 g of oxygen.
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3 Dalton’s Atomic Theory
British chemist John Dalton provided the basic theory about the nature of matter, which was based
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on the laws of chemical combination. His atomic theory explained the laws of conservation of mass and
constant proportions.
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3.1 Postulates of Dalton’s Atomic Theory
1. All matter is made of very tiny particles called atoms, which participate in chemical reactions.
2. Atoms are indivisible particles, which cannot be created or destroyed in a chemical reaction.
3. Atoms of a given element are identical in mass and chemical properties.
4. Atoms of different elements have different masses and chemical properties.
5. Atoms combine in the ratio of small whole numbers to form compounds.
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Atoms and Molecules - Class 9
6. The relative number and kinds of atoms are constant in a given compound.
Key Point
Dalton’s atomic theory successfully explained the laws of chemical combination and gave a basic
understanding of atoms, which formed the foundation for modern chemistry.
4 What is an Atom?
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Introduction:
Atoms are the basic building blocks of all matter. Just like a building is made of bricks, all matter is
composed of atoms.
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4.1 Size of Atoms
Atoms are extremely small, much smaller than we can imagine. The radius of an atom is measured in
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nanometers.
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Atomic Radius
1 nm = 10−9 m
Relative Sizes:
Object
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Atom of hydrogen 10−10
Molecule of water 10−9
Molecule of hemoglobin 10−8
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Grain of sand 10−4
Ant 10−3
Apple 10−1
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Atoms are incredibly small but make up everything around us.
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5 Modern-Day Symbols of Atoms of Different Elements
Development of Symbols:
Dalton was the first to use symbols for elements in a specific sense. Modern symbols are based on the
suggestions of Berzilius, who proposed using one or two letters from the element’s name.
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• The first letter is always uppercase, and the second letter (if present) is lowercase.
• Examples:
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– Hydrogen: H
– Aluminium: Al (not AL)
– Cobalt: Co (not CO)
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Elements Derived from Latin Names:
• Iron: Fe (Ferrum)
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• Sodium: Na (Natrium)
• Potassium: K (Kalium)
Element Symbol Element Symbol
Aluminium Al Iron Fe
Boron B Oxygen O
Calcium Ca Zinc Zn
Chlorine Cl Hydrogen H
Silver Ag Potassium K
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Atoms and Molecules - Class 9
6 Atomic Mass
The concept of atomic mass was introduced by Dalton, who proposed that each element has a charac-
teristic atomic mass.
Definition of Atomic Mass Unit (AMU)
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1 atomic mass unit (u) is equal to 12 th the mass of one atom of carbon-12.
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Explanation:
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• The atomic mass of elements is measured relative to carbon-12.
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• Earlier, atomic mass was taken as 1
16 of oxygen, but carbon-12 was later chosen as the standard.
Example
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The atomic masses of some common elements:
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Element Atomic Mass (u)
Hydrogen 1
Carbon 12
Nitrogen 14
Oxygen 16
Sodium
Magnesium n23
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Sulphur 32
Chlorine 35.5
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Calcium 40
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7 How Do Atoms Exist?
Most atoms do not exist independently. They combine to form:
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• Molecules: Atoms combine to form molecules by chemical bonds.
• Ions: Atoms may gain or lose electrons to form charged particles called ions.
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Molecule
A molecule is a group of two or more atoms chemically bonded together.
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Ion
An ion is a charged particle formed when an atom gains or loses electrons.
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8 What is a Molecule?
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Definition:
A molecule is a group of two or more atoms that are chemically bonded together by attractive forces.
A molecule can be defined as the smallest particle of an element or a compound that is capable of
independent existence and shows all the properties of that substance.
Key Points:
• Molecules can consist of atoms of the same element or different elements.
• They are held together by chemical bonds.
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Atoms and Molecules - Class 9
8.1 Molecules of Elements
The molecules of an element are made up of the same type of atoms. Some elements, like argon (Ar) and
helium (He), exist as monoatomic molecules. However, most non-metals exist as polyatomic molecules.
For example:
• Oxygen exists as a diatomic molecule (O2 ).
• Ozone consists of three oxygen atoms (O3 ).
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Definition:
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The number of atoms constituting a molecule is known as its atomicity.
Examples of Atomicity:
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Type of Element Name Atomicity
Non-Metal Argon Monoatomic
Non-Metal Helium Monoatomic
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Non-Metal Oxygen Diatomic
Non-Metal Hydrogen Diatomic
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Non-Metal Nitrogen Diatomic
Non-Metal Chlorine Diatomic
Non-Metal Phosphorus Tetra-atomic
Non-Metal Sulphur Poly-atomic
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Metals and some elements like carbon have complex structures with a large and indefinite number of
atoms bonded together.
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8.2 Molecules of Compounds
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Atoms of different elements combine in definite proportions to form molecules of compounds. Some
common compounds and their atomic compositions are listed below:
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Compound Combining Elements Mass Ratio
Water (H2 O) Hydrogen, Oxygen 1:8
Ammonia (N H3 ) Nitrogen, Hydrogen 14:3
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Carbon Dioxide (CO2 ) Carbon, Oxygen 3:8
Concept
In a molecule of a compound, atoms are present in a fixed ratio by mass, which defines the unique
chemical identity of the compound.
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9 What is an Ion?
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Definition:
Compounds composed of metals and non-metals contain charged species known as ions. Ions may consist
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of:
• A single charged atom (monoatomic ion).
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• A group of atoms with a net charge (polyatomic ion).
Types of Ions:
• A negatively charged ion is called an anion.
• A positively charged ion is called a cation.
For example, in sodium chloride (N aCl), the constituent particles are:
• Positively charged sodium ions (N a+ ).
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Atoms and Molecules - Class 9
• Negatively charged chloride ions (Cl− ).
Polyatomic Ion
A group of atoms carrying a charge is known as a polyatomic ion.
9.1 Examples of Ionic Compounds
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Ionic Compound Constituting Elements Mass Ratio
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Calcium Oxide (CaO) Calcium and Oxygen 5:2
Magnesium Sulphide (M gS) Magnesium and Sulphur 3:4
Sodium Chloride (N aCl) Sodium and Chlorine 23:35.5
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Key Point:
Ions are fundamental to the formation of ionic compounds, where electrostatic forces bind oppositely
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charged ions.
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10 Writing Chemical Formulae
Definition:
The chemical formula of a compound is a symbolic representation of its composition. The chemical
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formulae of different compounds can be written easily by understanding the symbols and combining
capacity of elements.
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10.1 Valency
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Definition:
The combining power (or capacity) of an element is known as its valency. It determines how atoms of
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an element will combine with atoms of another element to form a chemical compound.
Example:
Imagine an octopus with 8 arms and a human with 2 arms. If each arm must hold another, the combi-
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nation can be represented as:
O + 4H → OH4
Key Rule:
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The subscript indicates the number of atoms present in the compound.
10.2 Common Valencies of Elements and Ions
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Valency Name of Ion Symbol Polyatomic Ion Symbol
1 Sodium Na+ Ammonium NH+4
Potassium K+ Hydroxide OH−
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Silver Ag+ Nitrate NO−3
2 Magnesium Mg2+ Carbonate CO2−
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Calcium Ca2+ Sulphate SO2−
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3 Aluminium Al3+ Phosphate PO3−
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Iron (III) Fe3+
10.3 Rules for Writing Chemical Formulae
1. The valencies or charges on the ion must balance.
2. When a compound consists of a metal and a non-metal, the metal is written first, followed by the
non-metal.
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Atoms and Molecules - Class 9
Examples:
• Calcium oxide: CaO
• Sodium chloride: NaCl
• Iron sulphide: FeS
• Copper oxide: CuO
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3. For polyatomic ions, if more than one ion is present, enclose it in brackets and indicate the number
outside. Example:
Mg(OH)2
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If only one ion is present, no brackets are needed. Example:
NaOH
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Concept
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The chemical formula of a compound represents the types and number of atoms present in the
simplest unit of the compound.
11 Crossover Method for Writing Ionic Compounds
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11.1 Steps to Follow
1. Write the symbols of the elements or ions involved.
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2. Determine their valencies (or charges).
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3. Crossover the valencies to the opposite ion.
4. Write the final chemical formula after simplification.
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11.2 Examples
Sodium Chloride (NaCl)
Na+ Cl− ⇒ Na1 Cl1 ⇒ NaCl
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Magnesium Chloride (MgCl2 )
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Mg2+ Cl− ⇒ Mg1 Cl2 ⇒ MgCl2
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Aluminum Oxide (Al2 O3 )
Al3+ O2− ⇒ Al2 O3
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Calcium Hydroxide (Ca(OH)2 )
Ca2+ OH− ⇒ Ca(OH)2
Key Concept
The total positive and negative charges must balance each other in an ionic compound.
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Atoms and Molecules - Class 9
12 Molecular Mass and Formula Unit Mass
12.1 Molecular Mass
Molecular Mass
The molecular mass of a substance is the sum of the atomic masses of all the atoms in a
molecule of the substance. It is expressed in atomic mass units (u).
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Key Points:
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• Molecular mass is calculated by summing the atomic masses of the constituent atoms.
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• It provides the relative mass of a molecule compared to atomic mass units.
• Molecular mass helps in determining the weight of compounds in chemical reactions.
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12.1.1 Example Calculations
Calculation of Molecular Mass of Water (H2 O)
Given:
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Atomic mass of Hydrogen = 1 u
Atomic mass of Oxygen = 16 u
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Solution:
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Molecular mass of water = 2 × 1 + 1 × 16 = 18 u
Answer: The molecular mass of water (H2 O) is 18 atomic mass units.
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Calculation of Molecular Mass of HNO3
Given:
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Atomic mass of Hydrogen = 1 u
Atomic mass of Nitrogen = 14 u
Atomic mass of Oxygen = 16 u
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Solution:
Molecular mass of HNO3 = 1 + 14 + (3 × 16) = 1 + 14 + 48 = 63 u
Answer: The molecular mass of HNO3 is 63 atomic mass units.
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12.2 Formula Unit Mass
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Formula Unit Mass
The formula unit mass of a substance is the sum of the atomic masses of all atoms in a formula
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unit of a compound. This term is used for substances whose constituent particles are ions.
Key Points:
• Formula unit mass applies to ionic compounds.
• It is calculated in the same way as molecular mass.
• It helps in analyzing ionic compounds like NaCl, CaCl2 , etc.
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12.2.1 Example Calculations
Calculation of Formula Unit Mass of NaCl
Given:
Atomic mass of Sodium (Na) = 23 u
Atomic mass of Chlorine (Cl) = 35.5 u
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Solution:
Formula unit mass of NaCl = 1 × 23 + 1 × 35.5 = 58.5 u
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Answer: The formula unit mass of NaCl is 58.5 atomic mass units.
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Calculation of Formula Unit Mass of CaCl2
Given:
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Atomic mass of Calcium (Ca) = 40 u
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Atomic mass of Chlorine (Cl) = 35.5 u
Solution:
Formula unit mass of CaCl2 = 1 × 40 + 2 × 35.5 = 40 + 71 = 111 u
Answer: The formula unit mass of CaCl2 is 111 atomic mass units.
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Key Concept
Molecular mass and formula unit mass help in determining the weight of molecules and formula
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units, crucial for stoichiometric calculations in chemistry.
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