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Guess Papers for 9th Class Chemistry 2024
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Guess Papers for 9th Class Chemistry 2024
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CHAPTER NO. 1
CHEMISTRY’S FUNDAMENTALS.
MULTIPLE OPTIONS
1: There are a large number of naturally occurring elements.
02 08 108 114
2: Which of the following is Benzene’s empirical formula?
C2H2O4 2nd C2H2O C6H6 CH
3: Glucose’s empirical formula is.
CH2O CHO C2HO C2H2O
4: The iron valence in ferrous sulfate is.
1+2 +3 +4 +5
5: Water’s molecular mass is.
18 amu 18 g
Third, 18 mg/kg
6: What is the mass of one mole of water?
2 3 16 18
One amu is the same as.
1.66 x 10-24 mg
1.66 x 10-24 g
1.55 x 10-24 kg
1.66 x 10-23 g
7: Which of the following molecules does not have three atoms?
H2 O3
H2O CO2
8: The following are all examples of triatomic molecules.
H2 O2 H2O CO2
QUESTIONS SHORT
1: What exactly is an element? Use an example to demonstrate
2: Define the term valency. Calculate the value of Na.
3: Define the number Avogadro.
4: Determine the number of gram molecules in 40 g of phosphoric acid.
5: Define the term “atomic mass unit.”
6: What is the difference between a compound and an element?
7: What exactly is meant by “mixture”? Give one specific example.
8: Define relative atomic mass using C-12.
9: Give an example of the molecular formula.
10: Make a note of the chemical formula for water and sugar.
11: Distinguish between molecular mass and formula mass.
12: Give an example of a mole.
LONG QUESTIONS:
1: Describe three or five differences between a compound and a mixture.
2: 3.01 x 1030 (CO2) molecules are contained in a pot. Determine the moles and mass.
3: Explain two sorts of molecules based on atom types.
4: With an example, define atomic number and mass number.
CHAPTER 2
ATOM STRUCTURE
MULTIPLE OPTIONS:
1: So, which of the following shells is made up of three subshells?
Shell O
Shell, N
L stands for shell.
2: M – Shell _ is made up of three subshells.
-Shell M
Shell L Shell N Shell
Shell, O
3: Deuterium is utilized in the manufacturing process.
Hard water, soft water, and heavy water
Clearwater
4: Who invented the proton?
Rutherford
J.J.Thomson
Neil Bohr and Goldstein
5: The P subshell possesses.
a single orbital
2 Orbital
Three Orbital Objects
Orbital Four
6: Who invented the proton?
Rutherford
J.J.Thomson
Neil Bohr and Goldstein
7: Maximum number of electrons in sub-shell “P”
1 4 6 8
SHORT QUESTIONS:
- Record the observations made in Rutherford’s atomic model.
- Contrast the atomic theories proposed by Rutherford and Bohr.
- Specify the electronic configuration of sulfur.
- Provide the electronic configuration of chloride ions (Cl-).
- Define the term “Quantum.”
- Determine the electronic configuration of carbon (12C6) using subshells.
- State the electronic configuration of nitrogen with an atomic number of 7.
- Identify the drawbacks of Rutherford’s atomic model.
- Explain the concept of canal rays.
- List the electronic configuration of aluminum.
- State the atomic number and electronic configuration of phosphorus.
- Specify the electronic configuration of an element with 11 electrons.
LARGE QUESTIONS:
1: How was the neutron discovered? Make a list of the properties.
2: Compare and contrast Rutherford’s and Neil Bohr’s atomic theories.
3: Name four or five Cathode ray characteristics.
CHAPTER NO. 3
PERIODIC TABLE AND PROPERTY PERIODICITY.
MULTIPLE OPTIONS:
1: Who established the atomic number?
Rutherford, Dalton
H. Bohr Mosely
2: How many elements are there in the contemporary periodic table?
3 4 5 6
3: The contemporary periodic table’s foundation is.
a large number
Avogadro’s formula
The atomic number
The quantum number
4: Horizontal lines are referred to as.
Periods
The atomic number
Short time frames
Long periods
5: How many groups are there in the periodic table’s lengthy form?
5 18
6: 10 20 Group 17 is appropriate.
Nobel halogen gases
Metals containing alkali
None
7: The distance between two carbon atoms’ nuclei.
154 p.m., 140 p.m., 110 p.m., and 115 p.m.
8: Nitrogen’s electron negativity is.
2 3 4 5
SHORT QUESTIONS:
- Define the term “periods” and provide the names and elements of the first period.
- Enumerate the elements belonging to the 1st period of the periodic table.
- Explain the trend of ionization energy within both the period and group.
- Define electron affinity and offer an example.
- List the elements found in the 1st group.
- Define Ionization Energy.
- Describe the trend of ionization energy within a period.
- Define electronegativity and state the values for Nitrogen, Oxygen, and Chlorine.
LARGE QUESTIONS:
1: Talk about any three of the contemporary periodic table’s key components.
2: Explain what an atomic radius is. Describe its tendencies within the periodic table’s groupings and eras.
3: The Shielding Effect: Definition. Describe its patterns across periods and groups.
CHAPTER NO. 4
Molecular structure.
MULTIPLE OPTIONS:
1: The total number of electrons involved in a single covalent bond.
2 3 6 8
2: A triple covalent bond has how many electrons?
2 4 6 8
3: A triple bond is an example.
O2 C2H4 N2 NH3
4: Which of the following is a polar molecule?
O2 Cl2 HCl H2
5: The force between molecules is.
Metallic force Covalent force
Intermolecular attraction
The ionic force
6: The consequence of an electron transistor between atoms is.
Metallic Finishing
Ionic confinement
Covalent bonds
Covalent bonding should be coordinated.
7: It is assumed that a bond established between two nonmetals will be.
Ionic Covalent Covalent
Polar covalent bonds
Covalent coordinates
8: Determine which of the following pairs possesses polar covalent bonds.
Cl2 and O2
However, H2O and HCl H2O and N2 H2O and C2H2
SHORT QUESTIONS:
- Define double covalent bonds and offer examples.
- Explain the distinctions between a donor atom and an acceptor atom.
- Define a non-polar covalent bond and provide an example.
- Describe what is meant by bonding electrons.
- Explore the characteristics of triple covalent bonds and present examples.
- Discuss the nature of HF weak solid.
- Define polar covalent bonds and furnish an example.
- Highlight the differences between the ion pair and bond pair of electrons.
- Contrast polar covalent bonds with non-polar covalent bonds.
- Identify the type of covalent bond formed in N2 gas.
- Examine why water possesses a polar covalent bond.
- Explain the concept of a Metallic bond.
LARGE QUESTIONS:
1: Make a list of mental qualities.
2: Give an example of hydrogen bonding.
3: How does a coordinate covalent bond form? Use examples to demonstrate.
4: Name four covalent compound qualities.
CHAPITRE 5
MATTER’S PHYSICAL STATES.
MULTIPLE OPTIONS:
A voltmeter is used to measure atmospheric pressure.
Manometer
Barometer
Lactometer
- How many Pascal is one atmospheric pressure?
a) 101325 106075 10325 10523 - Liquids are __ times denser than gases.
a) 100 1000 10000 100000
SHORT QUESTIONS:
1: What exactly is Charles’ law? Create an equation for it.
2: Explain how temperature affects evaporation.
3: Why does evaporation increase as temperature rises?
4: Give an example of evaporation and define it.
HEAVY QUESTIONS:
1: Boyle’s Law can be confirmed through experimentation.
2: Explain what a boiling point is. Describe how various circumstances affect it.
3: Vapor pressure: what is it? How it alters in response to temperature variations.
4: Name three elements that have an impact on evaporation.
CHAPITRE NO.6
SOLUTIONS
MULTIPLE OPTIONS:
1: The solution’s most important components.
5 3 4 2
2: Which of the following is not a solid in a gas solution?
Smoke in the air
Buttered Brass
Fog
3: A solid solute solution in a solid solvent is an example.
Fog
Brass
4: The Solvent-to-Solute Ratio in Cheese Air Concentration
Solution to solute
solution to solvent
Both a and b are correct.
5: The volume of solute dissolved in 100 grams of solution is denoted as cm3.
% m/m % m/v
% v/m % v/v
6: Which one’s solubility reduces as temperature rises?
Ca(OH)2 KNO3 NaCl AgNO3
7: Which of the following is an example of suspension?
Solution of albumin
Solution for soap
Solution based on starch
Magnesia milk
QUICK QUESTIONS:
1: Define the term aqueous solution. Make a list of its parts.
2: Unsaturated solution should be defined.
3: What exactly is the distinction between a solution and an aqueous solution?
4: What does volume/volume% mean?
5: The distinction between concentrated and dilute solutions.
6: Define the term “saturated solution.”
7: How much KOH is needed to make one molar solution?
8: How to make molar solutions.
HEAVY QUESTIONS:
1: Explain how concentrated solutions are converted into dilute solutions.
2: Make a comparison of suspension and colloid.
3: Write down the four colloidal properties.
CHAPITRE NO.7
Electrical Chemistry
MULTIPLE OPTIONS:
1: The addition of oxygen to a chemical reaction is known as.
Evaporation
Reduced Condensation
Oxidation
2: The process of adding one electron to a material is known as.
Oxidation
Neutralization
Reduction
Ionization
3: Which of the following is not a strong electrolyte?
HCl CH3COOH NaOH H2SO4
4: Which of the following are strong electrolytes?
Sodium Chloride Sugar
Benzene
The acetic acid
5: A powerful electrolyte is an example.
Ca(OH)2 CH3COOH
CN6H6 NaOH
6: That’s not an electrolyte.
Sugar solution; sulfuric acid solution
Lime solution
Sodium chloride solution
7: Corrosion is the most typical example.
As a result, chemical degradation
However, rusting of iron
Aluminum rusting
Tin corrosion
8: The rust formula is.
Fe2O3.NH2O
Fe2O3
H2O Fe(OH)2 Fe(OH)2
SHORT QUESTIONS:
- Provide definitions for oxidation and reduction reactions.
- Define a Redox reaction and offer an example.
- Explain the concept of an electrochemical cell and enumerate its types.
- Define an electrolyte and present an example.
- Identify the electrolyte used in chromium electroplating.
- Define alloy and furnish an example.
- Calculate the oxidation number of sulfur in H2SO4.
- Calculate the oxidation number of chlorine in KClO3.
- Discuss the purpose of galvanizing.
- Explain the concept of electroplating.
- Highlight the distinctions between corrosion and rusting.
HEAVY QUESTIONS:
1: Make a list of four or five rules for allocating an Oxidation number to an element.
2: Explain electroplating. Explain chromium electroplating in detail.
3: What exactly is electroplating? The process of silver electroplating.
4: Explain the redox reactions using two instances.
5: Explain the rusting process of iron.
6: What is corrosion? Write four ways for corrosion prevention.
CHAPITRE NO. 8
REACTIVITY OF CHEMICALS
MULTIPLE OPTIONS:
1: Metals combine to generate ions, each of which has a charge.
Dipositive Unipositive
Try to be optimistic All
2: It is the most reactive metal.
Iron and gold
Cesium
Aluminum
3: Which metal is prone to breaking?
Aluminum Sodium
Selenium
Magnesium
4: Which of the following metals is the lightest?
Mg Ca Li Na Ca
LARGE QUESTIONS:
1: List the following four nonmetal chemical characteristics.
2: get metals. Additionally, list three or four metals’ chemical properties.
SHORT QUESTIONS:
- Enumerate two applications of Sodium.
- Identify the metals recognized for their high malleability and ductility.
- Provide definitions for the malleable and ductile properties of metals.
- List applications of Magnesium.
- Specify two moderately reactive metals.
- Identify the most precious metal.
- Define Metallic Character.
- Outline two practical uses of Gold.
- Name two highly reactive metals.
- Enumerate two physical properties of nonmetals.
- Explain why Sodium is more reactive than magnesium.
- Detail two uses of Silver.
- Elaborate on why platinum is preferred for jewelry making.
- Discuss two chemical properties of non-metals.
Chemistry Guess Papers for 9th Grade 2024
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Guess Papers for 9th Class Chemistry (2024)
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