Chemistry Grade 9 Text Book

Chemistry Grade 9

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Chemistry Fundamentals and Core Concepts

1. Definition and Scope of Chemistry

Chemistry is defined as "the science that deals with the properties, composition, and structure of substances (elements and compounds), the transformations they undergo, and the energy that is released or absorbed during these processes."

• Substances: A particular kind of matter with uniform properties (e.g., gold, water, table salt).
• Matter: A physical substance that occupies space and possesses rest mass (e.g., book, pencil).
• Properties: Attributes, qualities, or characteristics that distinguish one substance from another due to their unique composition and structure. Substances are continuously changing due to external and internal forces, undergoing transformations accompanied by energy changes.

Branches of Modern Chemistry:

The study of modern chemistry is generally broken down into five main disciplines:

• Physical Chemistry: Studies macroscopic and atomic properties, and phenomena in chemical systems, including reaction rates, energy transfers, and molecular-level physical structure.
• Organic Chemistry: (Mentioned in Unit Summary)
• Inorganic Chemistry: (Mentioned in Unit Summary)
• Analytic Chemistry: (Mentioned in Unit Summary)
• Biochemistry: (Mentioned in Unit Summary)

Importance and Scope of Chemistry:

Chemistry affects all aspects of life, both living and non-living, because everything is made of matter. Its scope extends to various sectors including agriculture, medicine, food production, building construction, explaining the natural world, preparing for career opportunities, and fostering informed citizens.

2. Measurements and Scientific Methods in Chemistry

Chemistry is an experimental science, relying heavily on precise measurements and a systematic approach to problem-solving.

2.1. Measurements and Units (SI Units and Prefixes)

The International System (SI) uses a selection of metric units and prefixes to express physical quantities.

SI Base Units:

Seven fundamental units from which all other units are derived.

Base Quantity	Name of Unit	Symbol
Length	Meter	m
Mass	Kilogram	kg
Time	Second	s
Electrical current	Ampere	A
Temperature	Kelvin	K
Amount of substance	Mole	mol
Luminous intensity	Candela	cd

Derived Units:

Units obtained by combining SI base units.

• Volume: "defined as length cubed and has the SI unit of cubic meter (m3)." Commonly used laboratory units include cubic decimeters (dm3) or cubic centimeters (cm3, also written cc). 1 L = 1 dm3 and 1 mL = 1 cm3.
• Density: "The density of an object is its mass per unit volume." Expressed as $d = m/V$ with SI unit kg/m3. Density is a "characteristic property of a material" and can be used to determine purity (e.g., gold).
• Other derived units include Area (m2), Speed (m/s), Acceleration (m/s2), Force (Newton, N), and Energy (Joule, J).
• Temperature: Measures intensity of heat. Heat is energy that flows from hotter to colder bodies. The Celsius and Kelvin scales have the same degree size, with "every Kelvin temperature is 273.15 units above the corresponding Celsius temperature." Fahrenheit to Celsius conversions involve subtracting 32 degrees Fahrenheit before scaling.
• Common Prefixes (SI/Metric Units): Modifiers used in decimal fashion to create larger or smaller units (e.g., Kilo (10^3), Centi (10^-2), Micro (10^-6)).

2.2. Uncertainty, Precision, Accuracy, and Significant Figures

• Uncertainty: It is often "impossible" to obtain the exact value of a quantity under investigation. Uncertainty is an interval within which the result can be found with a given probability.
• Significant Figures: "The meaningful digits in a measured or calculated quantity." When significant figures are used, "the last digit is understood to be uncertain." The amount of uncertainty depends on the measuring device and user's ability. Rules for determining significant figures involve counting non-zero digits, captive zeros, leading zeros (not significant), and trailing zeros (significant if a decimal point is present).
• Accuracy: "how closely a measured value agrees with the correct value."
• Precision: "how closely repeated measurements of the same quantity agree with one another."
• Errors: Can be systematic (consistent bias) or random (unpredictable fluctuations).
• Scientific Notation: Used for writing very large or very small numbers.

2.3. The Scientific Method and Experimental Skills

The scientific method is a systematic approach to solving problems. It involves:

• Observing
• Inferring
• Predicting
• Comparing & contrasting
• Communicating
• Analyzing
• Classifying
• Applying
• Theorizing
• Measuring
• Asking questions
• Developing hypotheses
• Designing experiments
• Interpreting data
• Drawing conclusions
• Making generalizations
• Problem-solving

Experimental skills include proper use of laboratory apparatus (e.g., pipets, graduated cylinders, balances) and adherence to safety rules. Laboratory reports document experimental procedures and findings.

Atomic Structure

3. Historical Development of Atomic Theory

• Early Philosophers (400 B.C.): Proposed that atoms were the "fundamental building blocks of matter" and "uncutable" or "non-divisible" (from Greek 'a-tomos'). These were speculations, not experimentally testable.
• Dalton’s Atomic Theory (1808): John Dalton, a British school teacher, proposed the first scientifically based atomic theory, viewing the atom as the "ultimate particle of matter." Key postulates included:
  • Matter consists of indivisible atoms.
  • All atoms of a given element are identical in all properties (later disproven by isotopes).
  • Atoms of different elements have different properties.
  • Compounds are formed by the combination of atoms of different elements in simple whole-number ratios.
  • Atoms are completely solid, homogeneous, and have no internal structure. (Later disproven by discovery of subatomic particles).
  • Atoms are different in their sizes, shapes, and weight.
  • Atoms cannot be created or destroyed in chemical reactions.
• J.J. Thomson (1897): Discovered the electron using the Crooke's Discharge Tube (Cathode Ray Tube). Cathode rays are streams of negatively charged particles. He also determined the electron's charge-to-mass ratio. Thomson proposed the Plum Pudding Model (1904), where electrons were embedded in a positively charged mass, forming an electrically neutral atom.
• Robert Millikan (1909): Through his oil drop experiment, Millikan determined the "charge and mass of the electron." He found the magnitude of the charge by measuring electric force and field on tiny oil droplets.
• Ernest Rutherford (1909-1911): Disproved the plum pudding model with his gold foil experiment (not detailed in this excerpt but implied by conclusions).

  Rutherford's Atomic Model (Planetary Model):

  • "The atom is mostly composed of empty space."
  • "The whole positive charge and mass of the atom are concentrated in a small central part known as the nucleus."
  • The nucleus is extremely small (diameter 10^-13 cm) compared to the atom (10^-8 cm).
  • "The electrons existing outside the nucleus rotate around the nucleus with high velocities to counterbalance the electrostatic forces of attraction between protons and electrons," similar to planetary motion.
  • Rutherford also demonstrated that strong nuclear forces "hold the protons and the neutrons together in the nucleus."

4. Subatomic Particles

• Protons: "one of the three sub-atomic particles found in the nucleus of an atom." They have a "positive electrical charge of one (+1) and a mass of 1.0073 atomic mass unit (amu)."
• Neutrons: "subatomic particles found inside the nucleus of an atom." They have "no charge." Their mass (1.0087 amu) is "a little greater than the mass of a proton." Together with protons, they "make up virtually all of the mass of an atom."
• Electrons: Negatively charged particles that orbit the nucleus.

5. Atomic Number, Mass Number, and Isotopes

• Atomic Number (Z): "The number of protons in an atom is called its atomic number (Z)." This number is "unique for atoms of a given element." "All atoms of an element have identical number of protons, and every element has a diverse number of protons in its atoms." For a neutral atom, the number of electrons equals the number of protons.
• Mass Number (A): "The total number of protons and neutrons in its nucleus." Also called the total number of nucleons. Calculated as: Mass number = (number of protons) + (number of neutrons).
• Ions: "Any atom or molecule with a net charge, either positive or negative." Formed by ionization, the process of a neutral atom losing or gaining valence electrons.
  • Net electric charge = number of protons – number of electrons
  • Anions: Atoms or groups of atoms with a net negative charge (more electrons than protons, e.g., Cl-, SO4^2-).
  • Cations: Atoms or groups of atoms with a net positive charge (more protons than electrons, e.g., Na+, Ca2+).
• Isotopes: "Atoms of the same element having the same atomic number but different mass numbers." This means they have the same number of protons but different numbers of neutrons.
  • Example: Hydrogen isotopes (Protium (1 proton, 0 neutrons), Deuterium (1 proton, 1 neutron), Tritium (1 proton, 2 neutrons)).
  • The existence of isotopes contradicts Dalton's postulate that all atoms of a given element are identical in mass.
  • Elements in nature exist as "constant uniform mixtures of their naturally occurring isotopes" in specific relative abundances.

6. Atomic Mass and Average Atomic Mass

• Atomic Mass Unit (amu) or Dalton (Da): Masses of single atoms are very small. Scientists use carbon-12 as a reference standard; "one atom of carbon-12 is assigned a mass of 12 atomic mass units (amu)." "1 amu corresponds to 1.660539040×10−24 g."
• Average Atomic Mass: "a weighted average of the masses of all the isotopes." Calculated as: Average Atomic Mass = (%isotope 1)(mass of isotope1) + (%isotope2)(mass of isotope 2) + ....

7. Electronic Configuration and Valence Electrons

• Electronic Configuration: "The assignment of all the electrons in an atom into specific shells or orbitals (s, p,d, f)."
• Electron Shells/Energy Levels: Electrons occupy specific energy levels or shells around the nucleus. The maximum number of electrons in a shell (n) is given by 2n^2.
  • K shell (n=1): 2 electrons
  • L shell (n=2): 8 electrons
  • M shell (n=3): 18 electrons
  • N shell (n=4): 32 electrons

  General filling pattern emphasizes that the outermost shell (valence shell) usually holds a maximum of 8 electrons, the penultimate shell 18, and the anti-penultimate shell 32.

• Valence Electrons: "The electrons that occupy the outermost shell of an atom." They are "the most easily lost, and the ones that determine the element’s chemical properties, and how an atom will react." The purpose of identifying valence electrons is to understand chemical reactivity.
• Octet Rule: Atoms achieve stability by having 8 electrons in their valence shell (except for hydrogen and helium, which aim for 2). Atoms can satisfy the octet rule by:
  • Losing their valence electrons.
  • Gaining valence electrons from other elements.
  • Sharing their valence electrons with other atoms.

Periodic Classification of Elements

8. Organization of the Periodic Table

• Elements are arranged in the periodic table by increasing atomic number.
• Periods (Horizontal Rows): Indicate the number of electron shells (main energy levels) an atom possesses. Elements in the same period have the same number of main shells.
• Groups (Vertical Columns): Elements in a given group "have the same number of outermost shell electrons" (valence electrons) and "similar chemical properties." For main group elements (A groups), "the group number equals the number of valence electrons."
• Blocks (s, p, d, f): Elements are classified into blocks based on the orbital where the last electron enters.
  • s-block elements: Last electron enters s-orbital. General configuration: ns1-2. (Groups IA and IIA).
  • p-block elements: Last electron enters p-orbital. General configuration: ns2np1-6. (Groups IIIA to VIIIA). Helium (He) is an exception, classified as s-block by configuration but p-block (Group VIII) by chemical behavior.
  • d-block elements: Last electron enters d-orbital. General configuration: ns0-2(n-1)d1-10. (Transition metals).
  • f-block elements: Last electron enters f-orbital. General configuration: ns2(n-1)d(0-1)(n-2)f(1-14). (Lanthanides and Actinides).

9. Major Trends in the Periodic Table

Four major trends are observed in the periodic table: atomic size, ionization energy, electron affinity, and electronegativity.

• Atomic Size (Atomic Radius):
  • Across a Period (Left to Right): Atomic radius generally decreases. This is due to increasing nuclear charge pulling outer electrons closer to the nucleus, despite increasing electron count.
  • Down a Group (Top to Bottom): Atomic radius generally increases. This is because additional electron shells are added, placing outer electrons further from the nucleus.
• Ionization Energy: "The amount of energy required to remove an electron from a neutral atom in its gaseous state."
  • Across a Period (Left to Right): Ionization energy generally increases. This is because increasing nuclear charge holds electrons more tightly.
  • Down a Group (Top to Bottom): Ionization energy generally decreases. This is due to increasing atomic size and shielding, making it easier to remove outer electrons.
• Electron Affinity: "the change in energy (in kJ/mole) of a neutral atom (in the gaseous phase) when an electron is added to the atom to form a negative ion." It reflects "the neutral atom’s likelihood of gaining an electron."
  • Nonmetals tend to have "negative" electron affinities (exothermic process), meaning they release energy when gaining an electron. They "have a greater electron affinity than metals" because they have more valence electrons and their valence shell is closer to the nucleus.
  • Across a Period (Left to Right): Electron affinity generally increases (becomes more negative).
  • Down a Group (Top to Bottom): Electron affinity generally decreases (becomes less negative).
  • Halogens (Group VIIA) have high electron affinities due to their tendency to gain one electron to achieve a stable octet. Noble gases have "extremely low (almost zero) electron affinities" because their stable electron configurations resist adding electrons.
• Electronegativity: "The tendency of an atom in a molecule to attract the shared pair of electrons towards itself." It is a dimensionless property.
  • Across a Period (Left to Right): Electronegativity generally increases.
  • Down a Group (Top to Bottom): Electronegativity generally decreases.

Chemical Bonding

10. Types of Chemical Bonds

Chemical bonds are formed when atoms associate by sharing, losing, or gaining valence electrons to achieve stability (often following the octet rule). The type of bond depends on the electronegativity difference between the bonded atoms.

10.1. Ionic Bonding

• Definition: "A bond formed by two oppositely charged ions due to electrostatic force." Occurs between "metallic and non-metallic elements."
• Formation: Involves the complete transfer of valence electrons from one atom (metal) to another (nonmetal) to form ions. Metals tend to lose electrons to form cations, while nonmetals gain electrons to form anions.
• Properties of Ionic Compounds:
  • Form Crystals: Ionic compounds typically exist as "crystalline solid form." Ions are held together by strong electrostatic forces in a "three-dimensional structure" (e.g., NaCl crystal).
  • High Melting and Boiling Points: Considerable heat energy is needed to overcome the strong electrostatic attractions between ions, requiring "high temperature to melt and boil ionic compounds."
  • Hard and Brittle: Due to the strong forces holding ions in a rigid structure, but brittle because a shift can align like charges, leading to repulsion and shattering.
  • Soluble in Polar Solvents: Follows the "like dissolves like" rule. "Polar solvents dissolve polar compounds." Water, methanol, and ethanol are common examples of polar solvents.
  • Conduct Electricity (Molten or Dissolved): Ions are fixed in a solid crystal lattice but "free to move" when molten or dissolved in water, allowing them to carry an electric current.

10.2. Covalent Bonding

• Definition: "formed when two atoms share one or more electron pairs." Occurs between identical atoms or atoms with insufficient electronegativity difference for electron transfer.
• Types of Covalent Bonds (based on shared electron pairs):
  • Single bond: "one pair of electrons is shared." (e.g., H-H in H2)
  • Double bond: "two pairs of electrons" shared. (e.g., O=O in O2)
  • Triple bond: "three pairs of electrons" shared. (e.g., N≡N in N2)

  "Bonds sharing more than one pair of electrons are called multiple covalent bonds."

• Lewis Dot Formula: Represents the chemical symbol of the element with valence electrons as dots surrounding it. Shared pairs can be shown as dots or solid lines.
• Polarity of Covalent Bonds:
  • Nonpolar Covalent Bond: Occurs when "molecules share electrons equally in a covalent bond, there is no net electrical charge across the molecule." This typically happens between "atoms of the same element, which have the same electronegativity."
  • Polar Covalent Bond: Forms when the "shared electron pair between the two atoms gets displaced more towards" the more electronegative atom, creating partial positive (δ+) and partial negative (δ-) charges. The molecule is then called a "polar molecule." Examples include HF, H2O, NH3.
• Electronegativity Difference as a Guide for Bond Type:
  • Ionic Bond: Electronegativity difference ≥ 2.0.
  • Polar Covalent Bond: Electronegativity difference in the range of 0.5-2.0.
  • Nonpolar Covalent Bond (pure covalent bond): Electronegativity difference < 0.5.

10.3. Coordinate Covalent Bond

Definition: "a covalent bond where the electron pair is provided by only one of the bonded atoms but shared by both atoms after bond formation."

10.4. Metallic Bonding

• Definition: Occurs "among the same metal atoms."
• Electron Sea Model: "an array of positive ions in a sea of electrons."
• Valence electrons become "delocalized," meaning they are "detached from its parent atom" and "freely move within these molecular outermost shells."
• The metal is held together by "strong forces of attraction between the positive nuclei and the delocalized electrons."
• Properties explained by Metallic Bonding: The free-flowing electrons explain high electrical conductivity (as electrons can "conduct electrical change when an electric field is applied") and thermal conductivity, as well as malleability and ductility of metals.