Atomic Structure and Bonding | Jamb Chemistry

Jamb chemistry key points on The concept of atoms, molecules and ions etc

The Concept of Atoms

1. Atoms are the smallest units of matter that retain the properties of an element.
2. Atoms are made up of subatomic particles: protons, neutrons, and electrons.
3. The concept of indivisible atoms was first proposed by Democritus in ancient Greece.
4. Atoms combine in whole-number ratios to form compounds.
5. Atoms of the same element have identical properties but differ from atoms of other elements.
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The Concept of Molecules

6. Molecules are formed when two or more atoms bond chemically.
7. Diatomic molecules consist of two atoms, e.g., $O_2$ and $H_2$.
8. Molecules can be homonuclear (same element) or heteronuclear (different elements).
9. Molecular structure determines the physical and chemical properties of substances.
10. Water $(H_2O)$ is a common example of a molecular compound.
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The Concept of Ions

11. Ions are charged particles formed when atoms gain or lose electrons.
12. Cations are positively charged ions formed by losing electrons.
13. Anions are negatively charged ions formed by gaining electrons.
14. Ions play critical roles in chemical reactions, such as ionic bonding.
15. Common examples include $Na^+$, $Cl^-$, and $SO_4^{2-}$.
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Contributions to Atomic Theory

Dalton’s Works

16. John Dalton proposed the atomic theory in 1803, suggesting that atoms are indivisible.
17. Dalton’s theory explained the law of definite proportions and the law of multiple proportions.
18. He described atoms as solid spheres, laying the foundation for modern atomic theory.
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Millikan’s Works

19. Robert Millikan measured the charge of the electron using the oil drop experiment in 1909.
20. Millikan’s work established the quantization of electric charge.
21. His findings confirmed the existence and charge of the electron.
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Rutherford’s Works

22. Ernest Rutherford’s gold foil experiment in 1911 demonstrated the presence of a dense nucleus in the atom.
23. Rutherford proposed the nuclear model of the atom, where electrons orbit the nucleus.
24. His work disproved the plum pudding model proposed by Thompson.
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Moseley’s Works

25. Henry Moseley discovered that the atomic number (not atomic mass) defines an element.
26. Moseley’s X-ray experiments helped organize the periodic table by atomic number.
27. He proved the significance of the proton count in an atom’s nucleus.
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Thomson’s Works

28. J.J. Thomson discovered the electron in 1897 through cathode ray experiments.
29. He proposed the plum pudding model, where electrons are embedded in a positively charged matrix.
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Bohr’s Works

30. Niels Bohr developed the Bohr model, suggesting that electrons orbit the nucleus in fixed energy levels.
31. Bohr’s model explained atomic spectra and energy transitions in hydrogen atoms.
32. His work introduced the concept of quantized energy levels.
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Atomic Structure

33. Atoms consist of a nucleus (protons and neutrons) surrounded by electrons in orbitals.
34. The nucleus is positively charged due to protons, while electrons are negatively charged.
35. Neutrons are neutral particles that contribute to an atom’s mass but not its charge.
36. Electrons occupy specific regions around the nucleus called orbitals.
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Electron Configuration

37. Electron configuration describes the arrangement of electrons in an atom’s orbitals.
38. Electrons fill orbitals according to the Aufbau principle, from lower to higher energy levels.
39. The Pauli exclusion principle states that no two electrons in an atom can have identical quantum numbers.
40. Hund’s rule states that electrons occupy degenerate orbitals singly before pairing.
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Atomic Number and Mass Number

41. The atomic number $(Z)$ is the number of protons in an atom’s nucleus and defines the element.
42. The mass number $(A)$ is the total number of protons and neutrons in the nucleus.
43. Electrons equal protons in a neutral atom, maintaining electrical neutrality.
44. Example: Carbon $(Z = 6, A = 12)$ has 6 protons, 6 electrons, and 6 neutrons.
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Isotopes

45. Isotopes are atoms of the same element with different mass numbers due to varying neutron counts.
46. Example: Carbon-12 and Carbon-14 are isotopes of carbon.
47. Isotopes may have different physical properties but identical chemical properties.
48. Radioactive isotopes like Carbon-14 are used in radiometric dating.
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Elements with Atomic Numbers

49. 1 - Hydrogen $(H)$, 2 - Helium $(He)$, 3 - Lithium $(Li)$, 4 - Beryllium $(Be)$.
50. 5 - Boron $(B)$, 6 - Carbon $(C)$, 7 - Nitrogen $(N)$, 8 - Oxygen $(O)$.
51. 9 - Fluorine $(F)$, 10 - Neon $(Ne)$, 11 - Sodium $(Na)$, 12 - Magnesium $(Mg)$.
52. 13 - Aluminum $(Al)$, 14 - Silicon $(Si)$, 15 - Phosphorus $(P)$, 16 - Sulfur $(S)$.
    17 - Chlorine $(Cl)$, 18 - Argon $(Ar)$, 19 - Potassium $(K)$, 20 - Calcium $(Ca)$.
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Shapes of s and p Orbitals

53. The $(s)$-orbital is spherical in shape.
54. The $(p)$-orbital is dumbbell-shaped with lobes along the x, y, or z axes.
55. Electrons in the $s$-orbital are closer to the nucleus compared to $(p)$-orbitals.
56. The $p$-orbital allows for directional bonding in molecules.
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Deducing Protons, Neutrons, and Electrons

57. Protons = Atomic Number $(Z)$.
58. Neutrons = Mass Number $(A)$ - Atomic Number $(Z)$.
59. Electrons = Protons in neutral atoms.
60. Example: Sodium $(A = 23, Z = 11)$ has 11 protons, 12 neutrons, and 11 electrons.
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Rules for Electron Arrangement

61. Electrons occupy orbitals in increasing energy order: $1s, 2s, 2p, 3s, 3p$, etc.
62. Each orbital can hold a maximum of two electrons with opposite spins (Pauli exclusion principle).
63. Degenerate orbitals (e.g., $p$-orbitals) are filled singly before pairing (Hund’s rule).
64. The Aufbau principle guides the order of orbital filling.
65. Example: Oxygen $(Z = 8)$ has the configuration $1s^2 2s^2 2p^4$.
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Additional Insights

66. The atomic structure has evolved through multiple theories and experimental validations.
67. The Bohr model provides a simple yet foundational understanding of atomic energy levels.
68. The quantum mechanical model is the most accurate representation of the atom.
69. J.J. Thomson’s discovery of the electron led to the understanding of subatomic particles.
70. Rutherford’s nuclear model introduced the concept of a dense, positively charged nucleus.
71. Moseley’s work corrected inconsistencies in the periodic table.
72. Millikan’s oil drop experiment provided precise measurements of electron charge.
73. Isotopes have varying applications in medicine, agriculture, and archaeology.
74. The atomic number uniquely identifies an element.
75. Electron configuration determines an atom’s chemical reactivity and bonding.
76. Neutrons provide stability to the nucleus by reducing electrostatic repulsion between protons.
77. Radioactive isotopes decay over time, releasing energy in the process.
78. The periodic table is organized by increasing atomic number.
79. Orbitals represent the probability distribution of finding an electron around the nucleus.
80. Understanding atomic structure is fundamental to advancements in chemistry, physics, and material science.
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Jamb chemistry Key points on The The periodic table and periodicity of elements, etc

The Periodic Table and Families of Elements

1. Definition: The periodic table is a systematic arrangement of elements based on increasing atomic number and recurring chemical properties.
2. Rows: Known as periods, representing energy levels of atoms.
3. Columns: Known as groups, elements in a group have similar chemical properties.
4. Groups: Include alkali metals, alkaline earth metals, halogens, noble gases, and transition metals.
5. Periodic Law: Properties of elements are periodic functions of their atomic numbers.
6. Groups 1–2: Known as s-block elements.
7. Groups 3–12: Known as d-block or transition metals.
8. Groups 13–18: Known as p-block elements.
9. Lanthanides and Actinides: Known as f-block elements, placed separately at the bottom of the table.
10. Significance: The periodic table helps predict chemical behavior and trends.
11. Alkali Metals (Group 1): Highly reactive, soft metals like sodium $(Na)$ and potassium $(K)$.
12. Halogens (Group 17): Nonmetals like chlorine $(Cl)$ and fluorine $(F)$ that are highly reactive.
13. Noble Gases (Group 18): Inert gases like helium $(He)$ and neon $(Ne)$, with full outer electron shells.
14. Transition Metals (Groups 3–12): Elements like iron $(Fe)$ and copper $(Cu)$, known for forming colored compounds and variable oxidation states.
15. Periodic Trends: Include ionization energy, atomic radii, electronegativity, and electron affinity.
16. Atomic Number: Determines the position of elements in the periodic table.
17. Mendeleev's Contribution: Organized elements by atomic mass, leading to the discovery of periodic trends.
18. Modern Periodic Table: Organized by atomic number, as suggested by Moseley.
19. Families: Elements in the same family have similar valence electron configurations.
20. Applications: Used in material science, chemistry, and industrial processes.
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Periodic Properties of Elements

21. Ionization Energy: The energy required to remove an electron from a gaseous atom.
22. Trend Across Periods: Increases across a period due to increased nuclear charge.
23. Trend Down Groups: Decreases down a group as atomic size increases.
24. Example: Ionization energy of fluorine is higher than that of lithium.
25. Electronegativity: A measure of an atom's ability to attract electrons in a bond.
26. Trend Across Periods: Increases across a period due to stronger nuclear pull.
27. Trend Down Groups: Decreases down a group as atomic size increases.
28. Example: Fluorine has the highest electronegativity value (3.98).
29. Electron Affinity: The energy change when an electron is added to a gaseous atom.
30. Trend Across Periods: Becomes more negative (more energy released) across a period.
31. Trend Down Groups: Becomes less negative down a group.
32. Example: Chlorine has a high electron affinity.
33. Atomic Radii: The size of an atom from its nucleus to the outermost electron.
34. Trend Across Periods: Decreases across a period due to increased nuclear charge.
35. Trend Down Groups: Increases down a group as new energy levels are added.
36. Ionic Radii: The size of an ion compared to its parent atom.
37. Cations: Smaller than their parent atoms due to loss of electrons.
38. Anions: Larger than their parent atoms due to electron gain.
39. Example: $Na^+$ is smaller than $Na$, while $Cl^-$ is larger than $Cl$.
40. Reactivity Trends: Metals become more reactive down a group; nonmetals become less reactive.
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Properties of Specific Groups

41. Group 1 (Alkali Metals): Soft, shiny metals with 1 valence electron.
42. Reactivity: Increases down the group due to lower ionization energy.
43. Reaction with Water: Forms hydrogen gas and metal hydroxides.
44. Group 17 (Halogens): Nonmetals with 7 valence electrons, highly reactive.
45. Reactivity: Decreases down the group as electronegativity decreases.
46. Example: Fluorine reacts more vigorously than iodine.
47. Group 18 (Noble Gases): Chemically inert due to full valence shells.
48. Uses: Helium in balloons, neon in lights, argon in welding.
49. Transition Metals: Known for variable oxidation states and forming alloys.
50. Catalytic Properties: Used in industrial processes like the Haber process.
51. Group 2 (Alkaline Earth Metals): Reactive but less so than alkali metals.
52. Example: Magnesium reacts with water to form magnesium hydroxide.
53. Lanthanides: Known as rare earth metals with high magnetic properties.
54. Actinides: Radioactive elements, including uranium and thorium.
55. Metalloids: Elements like boron and silicon with properties of metals and nonmetals.
56. Chemical Reactivity: Varies widely across groups and periods.
57. Periodic Trends in Reactivity: Metals increase reactivity down a group, while nonmetals decrease.
58. Electron Configuration: Explains similarities in chemical behavior within groups.
59. Example: Sodium $(Na)$ and potassium $(K)$ react similarly with water.
60. Applications of Periodic Trends: Predicting element behavior in reactions.
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Isotopy and Mass Numbers

61. Isotopes: Atoms of the same element with different numbers of neutrons.
62. Mass Number: The sum of protons and neutrons in the nucleus.
63. Examples of Isotopes: Carbon-12 and Carbon-14.
64. Radioactive Isotopes: Used in medicine (e.g., Iodine-131 for thyroid treatment).
65. Effect on Atomic Mass: Isotopes cause variation in average atomic mass.
66. Periodic Table Placement: Isotopes occupy the same position since they have the same atomic number.
67. Example: Hydrogen has isotopes: protium, deuterium, and tritium.
68. Relation to Properties: Isotopes have identical chemical properties but may differ physically.
69. Average Atomic Mass: Weighted average based on isotope abundance.
70. Applications: Used in radiocarbon dating and nuclear energy.
71. Mass Spectrometry: Identifies isotopic composition of elements.
72. Isotopy and Stability: Heavier isotopes are often radioactive.
73. Elements Exhibiting Isotopy: Hydrogen, carbon, oxygen, chlorine, and uranium.
74. Chlorine Isotopes: Chlorine-35 and Chlorine-37 exist naturally.
75. Effect on Molecular Mass: Isotopy affects molecular weights in compounds.
76. Atomic Stability: Isotopes with balanced neutron-to-proton ratios are more stable.
77. Periodic Trends in Isotopes: Similarity in properties across isotopes of the same element.
78. Nuclear Properties: Governed by isotopic composition.
79. Atomic Mass Units (AMU): Standardized for isotope comparisons.
80. Significance in Chemistry: Isotopes provide insights into atomic behavior and reactions.
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Reasons for Periodic Variations

81. Across Periods: Properties change due to increasing nuclear charge.
82. Down Groups: Properties vary due to increasing atomic size.
83. Ionization Energy Trends: Influenced by electron shielding and nuclear attraction.
84. Electronegativity Trends: Reflects the ability to attract electrons in bonding.
85. Atomic Radii Trends: Results from nuclear pull and added energy levels.
86. Reactivity Trends: Metals become more reactive down a group, while nonmetals become less.
87. Metallic Character: Increases down a group and decreases across a period.
88. Nonmetallic Character: Decreases down a group and increases across a period.
89. Shielding Effect: Inner electrons reduce nuclear pull on valence electrons.
90. Effective Nuclear Charge: Increases across a period, strengthening bonds.
91. Valence Electrons: Determine chemical properties and bonding patterns.
92. Block Properties: s-block elements are highly reactive, while d-block elements form stable compounds.
93. Periodicity: Recurring trends due to periodicity in electron configuration.
94. Energy Levels: Additional levels down a group reduce attraction for valence electrons.
95. Bonding Trends: Covalent bonding is common in nonmetals, while metals prefer ionic bonding.
96. Group Similarities: Elements in the same group share chemical and physical properties.
97. Transition Metals: Display unique properties like variable oxidation states and colored ions.
98. Inert Pair Effect: Seen in heavy p-block elements due to reluctance of s-electrons to bond.
99. Applications of Trends: Predicting reactions, designing catalysts, and material synthesis.
100. Periodic Table Importance: The periodic table remains a cornerstone of chemistry, guiding predictions and research.
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Jamb chemistry Key Points on Chemical Bonding

Chemical Bonding Overview

1. Definition: Chemical bonding involves the attraction between atoms or ions that enables the formation of chemical compounds.
2. Primary Bond Types: Includes ionic (electrovalent), covalent, and metallic bonds.
3. Purpose: Atoms bond to achieve a stable electron configuration, often resembling noble gases.
4. Octet Rule: Atoms gain, lose, or share electrons to complete an octet in their outer shell.
5. Types of Bonding: Distinguished by electron transfer, sharing, or delocalization.
6. Electronegativity Difference: Determines bond type; large differences favor ionic bonds, while smaller differences favor covalent bonds.
7. Bond Strength: Covalent bonds are directional and typically stronger than ionic bonds.
8. Hybridization: Explains molecular geometry by mixing atomic orbitals.
9. Molecular Properties: The type of bonding affects melting point, boiling point, solubility, and conductivity.
10. Valence Electrons: The number of valence electrons determines bonding capacity.
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Electrovalency (Ionic Bonding)

11. Definition: Electrovalency involves electron transfer from one atom to another, forming ions.
12. Formation: Metals lose electrons to form cations, while nonmetals gain electrons to form anions.
13. Example: Sodium $(Na)$ transfers an electron to chlorine $(Cl)$ to form $Na^+$ and $Cl^-$.
14. Electron Configuration: $Na: 1s^2 2s^2 2p^6 3s^1$ becomes $Na^+$, resembling neon.
15. Properties of Ionic Compounds: High melting and boiling points, solubility in water, and electrical conductivity in molten or aqueous states.
16. Lattice Energy: Determines the strength of ionic bonds; higher lattice energy leads to stronger bonds.
17. Examples: $NaCl$, $MgO$, and $CaF_2$.
18. Electrostatic Forces: Hold oppositely charged ions together in a rigid lattice.
19. Role in Biology: Ionic bonding is essential in nerve impulse transmission and muscle contraction.
20. Limitations: Pure ionic bonds rarely exist; most bonds have some covalent character.
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Covalency (Covalent Bonding)

21. Definition: Covalent bonding involves the sharing of electrons between atoms.
22. Formation: Occurs between nonmetals with similar electronegativities.
23. Electron Sharing: Each atom contributes one or more electrons to the bond.
24. Single, Double, and Triple Bonds: Represent the number of shared electron pairs.
25. Example: Oxygen forms a double bond $(O=O)$ to share two pairs of electrons.
26. Molecular Geometry: Determined by the VSEPR (Valence Shell Electron Pair Repulsion) theory.
27. Polarity: Covalent bonds can be polar or nonpolar, depending on electronegativity differences.
28. Properties of Covalent Compounds: Low melting points, low boiling points, and poor conductivity.
29. Examples: $H_2O$, $CO_2$, $CH_4$.
30. Bond Strength: Determined by bond energy; shorter bonds are stronger.
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Tendency to Attain Noble Gas Configuration

31. Stability Goal: Atoms bond to achieve a full valence shell like noble gases.
32. Examples: Neon $(Ne: 1s^2 2s^2 2p^6)$ and argon $(Ar: 1s^2 2s^2 2p^6 3s^2 3p^6)$.
33. Electron Gain or Loss: Metals lose electrons, and nonmetals gain electrons to achieve noble gas configuration.
34. Electron Sharing: Nonmetals share electrons to complete their octet.
35. Applications: Noble gas stability explains reactivity trends in the periodic table.
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Special Types of Bonding

Hydrogen Bonding

36. Definition: A weak bond between a hydrogen atom and an electronegative atom e.g., $O$, $N$, $(F)$.
37. Formation: Results from the attraction between a hydrogen atom covalently bonded to one electronegative atom and another electronegative atom.
38. Examples: Water $(H_2O)$ and ammonia $(NH_3)$.
39. Biological Importance: Stabilizes DNA and protein structures.
40. Effect on Properties: Causes high boiling points and surface tension in water.
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Metallic Bonding

41. Definition: Involves delocalized electrons shared among a lattice of metal cations.
42. Characteristics: Explains properties like malleability, ductility, and conductivity.
43. Electron Sea Model: Free-moving electrons provide cohesion between positively charged metal ions.
44. Examples: Metals like copper $(Cu)$, aluminum $(Al)$, and iron $(Fe)$.
45. Applications: Essential for electrical wiring, construction, and machinery.
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Coordinate (Dative) Bonding

46. Definition: A covalent bond where both electrons in the bond are donated by one atom.
47. Formation: Occurs in complexes and molecules like $NH_4^+$ and $[Fe(CN)_6]^{3-}$.
48. Examples:
    • $[Fe(CN)_6]^{3-}$: Iron shares electrons with cyanide ligands.
    • $[Cu(NH_3)_4]^{2+}$: Copper forms bonds with ammonia ligands.
    • $[Ag(NH_3)_2]^+$: Silver bonds with ammonia molecules.
49. Hybridization: Explains the geometry of coordination complexes.
50. Applications: Found in biological systems (e.g., hemoglobin) and catalysis.
51. Polarity: Depends on the geometry of the complex.
52. Electron Configuration of Central Metal: Dictates coordination behavior e.g., $Fe^{3+}$.
53. Bond Strength: Dative bonds are typically stronger than regular covalent bonds.
54. Charge Distribution: Determines stability and reactivity of complexes.
55. Uses: Coordination compounds are used in medicines, dyes, and fertilizers.
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Deducing Bond Types from Electron Configurations

56. Electron Configuration: Determines whether an atom will form ionic or covalent bonds.
57. Metals: Tend to lose electrons due to low ionization energy e.g., $Na: 1s^2 2s^2 2p^6 3s^1$.
58. Nonmetals: Tend to gain electrons due to high electronegativity e.g., $Cl: 1s^2 2s^2 2p^6 3s^2 3p^5$.
59. Transition Metals: Form dative bonds due to vacant $d$-orbitals.
60. Bond Polarity: Determined by electronegativity differences between bonded atoms.
61. Examples:
    • $NaCl$: Ionic bond due to large electronegativity difference.
    • $H_2O$: Polar covalent bond due to moderate electronegativity difference.
62. Electron Pairing: Guides covalent bond formation.
63. Hybridization: Explains multiple bonding in molecules like $CO_2$ $(sp)$ and $CH_4$ $(sp^3)$.
64. Coordination Chemistry: Relies on the availability of lone pairs.
65. Reactivity Trends: Atoms with incomplete valence shells are more reactive.
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Relating Bonding to Compound Properties

66. Ionic Compounds: High melting and boiling points due to strong electrostatic forces.
67. Covalent Compounds: Lower melting points and poor conductivity due to localized electrons.
68. Metallic Compounds: Conduct electricity and heat due to delocalized electrons.
69. Polarity: Determines solubility; polar compounds dissolve in polar solvents.
70. Hydrogen Bonding: Enhances boiling points and surface tension.
71. Coordination Compounds: Exhibit unique colors and magnetic properties.
72. Bond Strength: Stronger bonds result in more stable compounds.
73. Reactivity: Covalent compounds with polar bonds are more reactive.
74. Thermal Stability: Ionic compounds are more thermally stable than covalent ones.
75. Flexibility: Metallic bonds allow metals to be malleable and ductile.
76. Conductivity: Depends on free electrons; ionic compounds conduct only in molten or dissolved states.
77. Hardness: Covalent networks (e.g., diamond) are extremely hard.
78. Example Comparisons:
    • $NaCl$: Ionic, soluble in water, conducts electricity.
    • $CH_4$: Covalent, insoluble in water, poor conductor.
79. Biological Importance: Hydrogen bonding stabilizes DNA and proteins.
80. Industrial Applications: Metallic bonds are essential in the construction of tools and machinery.
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Jamb chemistry Key Points on Shapes of simple molecules

Overview of Molecular Shapes

1. Definition: The shape of a molecule is determined by the arrangement of atoms and electron pairs around the central atom.
2. Valence Shell Electron Pair Repulsion (VSEPR) Theory: Predicts molecular shapes based on repulsion between electron pairs.
3. Bond Pairs and Lone Pairs: Bond pairs form covalent bonds, while lone pairs occupy space but do not bond.
4. Electron Geometry vs. Molecular Shape: Electron geometry considers all electron pairs, while molecular shape focuses on bonded atoms.
5. Linear Shape: Atoms align in a straight line with a bond angle of 180°.
6. Non-Linear Shape: Molecules with bent or angular structures have bond angles less than 120°.
7. Tetrahedral Shape: Central atom surrounded by four atoms in a symmetrical arrangement with a bond angle of 109.5°.
8. Pyramidal Shape: A central atom bonded to three atoms and one lone pair, forming a triangular pyramid.
9. Importance of Shape: Determines molecular polarity, reactivity, and interactions.
10. Hybridization and Shape: Hybrid orbitals like $sp, sp^2, sp^3$ influence molecular geometry.
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Linear Molecules (H₂, O₂, Cl₂, HCl, CO₂)

11. Linear Shape: Atoms form a straight line.
12. Bond Angle: 180°.
13. Examples:
    • $H_2$: Two hydrogen atoms share a single bond.
    • $O_2$: Two oxygen atoms share a double bond.
    • $Cl_2$: Two chlorine atoms form a single bond.
14. Carbon Dioxide $(CO_2)$: A central carbon atom forms two double bonds with oxygen atoms.
15. Polarity: Linear molecules are nonpolar if the bonds are symmetrical (e.g., $CO_2$ ).
16. Hybridization:
    • $H_2, O_2, Cl_2$: No hybridization required.
    • ( CO_2 ): Central carbon atom undergoes $sp$ hybridization.
17. Intermolecular Forces: Typically weak van der Waals forces for nonpolar molecules.
18. Physical Properties: Linear molecules like $CO_2$ are gases at room temperature due to weak intermolecular forces.
19. Reactivity: Linear diatomic molecules like $O_2$ are highly reactive.
20. Bonding: Simple bonding in diatomic molecules contrasts with double bonding in $CO_2$.
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Non-Linear Molecule (H₂O)

21. Non-Linear Shape: A bent or angular structure due to lone pairs on the central atom.
22. Bond Angle: Approximately 104.5°.
23. Water $H_2O$: Oxygen is the central atom, with two bond pairs and two lone pairs.
24. Polarity: $H_2O$ is polar because of unequal distribution of charge.
25. Hybridization: Oxygen undergoes $sp^3$ hybridization.
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Tetrahedral Molecule (CH₄)

26. Tetrahedral Shape: Four atoms are symmetrically arranged around the central atom.
27. Bond Angle: 109.5°.
28. Methane $(CH_4)$: Carbon is the central atom, bonded to four hydrogen atoms.
29. Symmetry: High symmetry makes $CH_4$ nonpolar.
30. Hybridization: Carbon undergoes $sp^3$ hybridization.
31. Physical Properties: $CH_4$ is a gas at room temperature due to weak dispersion forces.
32. Chemical Properties: Highly flammable, used as a fuel source.
33. Examples of Tetrahedral Molecules: Methane $(CH_4)$, ammonium ion $(NH_4^+)$.
34. Applications: Tetrahedral geometry is common in organic compounds.
35. Electron Pair Repulsion: Equal repulsion between bond pairs ensures symmetry.
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Pyramidal Molecule (NH₃)

36. Pyramidal Shape: A triangular pyramid structure due to one lone pair on the central atom.
37. Bond Angle: Approximately 107°.
38. Ammonia $(NH_3)$: Nitrogen is the central atom, bonded to three hydrogen atoms with one lone pair.
39. Polarity: $(NH_3)$ is polar due to the lone pair creating an uneven charge distribution.
40. Hybridization: Nitrogen undergoes $sp^3$ hybridization.
41. Intermolecular Forces: Strong hydrogen bonding makes $NH_3$ soluble in water.
42. Chemical Reactivity: Ammonia acts as a weak base and forms ammonium ion $(NH_4^+)$.
43. Examples of Pyramidal Molecules: Ammonia $(NH_3)$, phosphine $(PH_3)$.
44. Biological Importance: Ammonia is essential in nitrogen metabolism.
45. Applications: Used in fertilizers and cleaning agents.
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Differentiating Molecular Shapes

46. Linear vs. Non-Linear: Linear molecules like $CO_2$ have no lone pairs on the central atom, while non-linear molecules like $H_2O$ do.
47. Tetrahedral vs. Pyramidal: Tetrahedral molecules like $CH_4$ have no lone pairs, while pyramidal molecules like $NH_3$ have one lone pair.
48. Bond Angles: Linear (180°), non-linear $(<120°)$, tetrahedral (109.5°), pyramidal (107°).
49. Polarity: Symmetrical shapes (linear and tetrahedral) are often nonpolar, while asymmetrical shapes (non-linear and pyramidal) are polar.
50. Electron Pair Influence: Lone pairs cause repulsion, reducing bond angles and altering molecular shape.
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Jamb chemistry Key Points on Nuclear Chemistry

Radioactivity Overview

1. Definition: Radioactivity is the spontaneous emission of particles or energy from unstable atomic nuclei.
2. Discovery: Henri Becquerel discovered radioactivity in 1896 while studying uranium salts.
3. Types of Radioactivity: Alpha $(\alpha)$, beta $(\beta)$, and gamma $(\gamma)$ radiation.
4. Unstable Nuclei: Emit radiation to attain stability by altering their neutron-to-proton ratio.
5. Nuclear Chemistry: Focuses on reactions involving changes in the nucleus, unlike regular chemical reactions involving electrons.
6. Natural Radioactivity: Occurs in isotopes like uranium-238 and radium-226.
7. Artificial Radioactivity: Induced by bombarding stable nuclei with particles in a laboratory.
8. Applications: Used in medicine, energy production, and archaeology.
9. Units of Measurement: Radioactivity is measured in becquerels (Bq) or curies (Ci).
10. Half-Life: The time required for half the nuclei in a radioactive sample to decay.
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Types and Properties of Radiations

11. Alpha Radiation $(\alpha)$: Consists of 2 protons and 2 neutrons (helium nucleus).
12. Properties of $(\alpha)$: Heavy, low penetration, stopped by paper or skin.
13. Beta Radiation $(\beta)$: High-energy electrons or positrons emitted from the nucleus.
14. Properties of $(\beta)$: Moderate penetration, stopped by aluminum foil.
15. Gamma Radiation $(\gamma)$: High-energy electromagnetic waves emitted with $\alpha$ or $\beta$.
16. Properties of $\gamma$: High penetration, requires thick lead or concrete for shielding.
17. Ionization Power: $\alpha > \beta > \gamma$, with $\alpha$ being the most ionizing.
18. Penetration Power: $\gamma > \beta > \alpha$, with $\gamma$ being the most penetrating.
19. Speed: $\alpha$ particles are slowest, $\gamma$ radiation travels at the speed of light.
20. Example Isotopes:
    • $\alpha$: Uranium-238 $(U)$.
    • $\beta$: Carbon-14 $(C)$.
    • $\gamma$: Cobalt-60 $(Co)$.
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Nuclear Reactions and Equations

21. Definition: Nuclear reactions involve changes in the atomic nucleus, resulting in new elements or isotopes.
22. Types: Includes fission, fusion, and decay processes.
23. Nuclear Decay: Spontaneous emission of $\alpha, \beta,$ or $\gamma$.
24. Fission: Splitting a heavy nucleus (e.g., uranium-235) into smaller nuclei, releasing energy.
25. Fusion: Combining light nuclei (e.g., hydrogen isotopes) to form a heavier nucleus, releasing energy.
26. Balanced Nuclear Equation: Ensures conservation of mass number and atomic number.
27. Example $\alpha$-Decay:
    • $^{238}_{92}U \rightarrow ^{234}_{90}Th + ^4_2He$.
28. Example $\beta$-Decay:
    • $^{14}_6C \rightarrow ^{14}_7N + ^0_{-1}e$.
29. Example $\gamma$-Emission:
    • $^{60}_{27}Co^* \rightarrow ^{60}_{27}Co + \gamma$.
30. Nuclear Stability: Determined by the neutron-to-proton ratio; stable nuclei have optimal ratios.
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Half-Life and Calculations

31. Half-Life Definition: The time required for half the radioactive nuclei in a sample to decay.
32. Formula: $N = N_0 \times (1/2)^{t/T}$, where $N_0$ is the initial amount, $t$ is elapsed time, and $T$ is the half-life.
33. Example Calculation: If 100 g of a substance has a half-life of 5 years, the amount left after 10 years is:
    • $N = 100 \times (1/2)^{10/5} = 25g$.
34. Decay Constant $(\lambda)$: Related to half-life by $T = \ln(2)/\lambda$.
35. Activity: The rate of decay is proportional to the number of radioactive nuclei.
36. Example Problem: If a sample initially contains 50 atoms and its half-life is 2 minutes, how many atoms remain after 6 minutes?
    • $N = 50 \times (1/2)^{6/2} = 6.25$.
37. Real-Life Example: Carbon-14 has a half-life of 5730 years, used in radiocarbon dating.
38. Exponential Decay: Radioactive decay follows an exponential pattern over time.
39. Measurement Techniques: Includes Geiger counters and scintillation detectors.
40. Applications: Used in determining the age of archaeological finds and managing nuclear waste.
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Applications of Radioactivity

41. Medicine:
    • Cancer Treatment: Radiation therapy uses cobalt-60 to target tumors.
    • Diagnostics: Technetium-99m is used in imaging techniques.
42. Energy Production: Nuclear fission powers nuclear reactors.
43. Archaeology: Radiocarbon dating estimates the age of ancient artifacts.
44. Agriculture: Radioisotopes improve crop yields and control pests.
45. Industry: Gamma radiation is used in non-destructive material testing.
46. Sterilization: Radiation sterilizes medical equipment and food.
47. Environmental Monitoring: Detects pollution using radioisotopes.
48. Space Exploration: Radioisotope thermoelectric generators (RTGs) provide energy for spacecraft.
49. Nuclear Weapons: Harness nuclear reactions for destructive purposes.
50. Scientific Research: Radioactive tracers are used to study chemical and biological processes.

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