Comprehensive Grade 9 Chemistry Study for Grade 9 Chemistry Student

Textbook Metadata, Contributors, and Care Guidelines

This Grade 9 Chemistry Student Textbook is published by the Ministry of Education of the Federal Democratic Republic of Ethiopia. The development of this educational resource was led by writers Sisay Tadesse (Ph.D.) and Tegene Tesfaye (Ph.D.), and edited by a team comprising Curriculum Editor Tesfaye Semela (Ph.D.), Language Editor Kenenisa Beresa (M.A.), and Content Editor Ahmed Awel (M.Sc.). The book features illustrations by Abinet Tilahun (M.Sc.) and design by Konno B. Hirbaye (M.Sc.). Evaluation was conducted by Legesse Adane (Ph.D.), Nega Gichile (B.Sc., M.A.), Sefiw Melesse (M.Sc.), and Tolessa Mergo (B.Sc., M.Sc.). The publication was supported by the General Education Quality Improvement Program for Equity (GEQIP-E), with funding from international partners including the World Bank, the UK’s FCDO, Finland’s Ministry for Foreign Affairs, the Norwegian Embassy, UNICEF, the Global Partnership for Education, and the Danish Ministry of Foreign Affairs. It was printed in August 2023 by Gravity Group Ind LLC in Sharjah, United Arab Emirates, under contract number MOE/GEQIP-E/LICB/G-01/23. The ISBN for this volume is 978-9999-06-001-1.

Students are instructed to take significant care of this textbook as it is school property. Specific care guidelines include covering the book with protective materials to prevent surface damage and storing it exclusively in clean, dry environments. Users must ensure their hands are clean before handling the pages and should avoid writing on the covers or interior. It is recommended to use paper or cardboard bookmarks rather than folding page corners or tearing out illustrations. For structural longevity, new books should be laid on their backs and opened a few pages at a time while pressing lightly along the bound edge to maintain the integrity of the adhesive. Any existing tears should be repaired immediately with tape or paste, and the book must be packed carefully in school bags to avoid crushing.

Definition, Scope, and Branches of Chemistry

Chemistry is defined as the scientific study of the properties, composition, and structure of substances, including both elements and compounds. This discipline examines the transformations that matter undergoes and the energy captured or released throughout these processes. Matter is defined as any physical substance that occupies space and possesses rest mass, exemplified by common items such as books, pencils, and televisions. A substance is a specific kind of matter with uniform properties, such as gold, silver, water, or soap. The properties of a substance refer to its unique attributes or characteristics, which allow it to be distinguished from others. Composition relates to the specific nature of a substance’s ingredients, such as the chemical combination of sodium and chlorine in table salt. Structure describes the arrangement and relationships between the constituent parts of a complex system, such as the ordered layout of a building's roof, walls, and floors.

The scope of modern chemistry is vast, given that both living and non-living things are composed of matter. It is generally categorized into five primary disciplines. Physical chemistry focuses on macroscopic and atomic properties and phenomena within chemical systems, involving the study of reaction rates and energy transfers. Organic chemistry is the study of substances containing carbon, which is highly abundant and forms over twenty million known chemicals. Inorganic chemistry deals with substances not primarily based on carbon, such as minerals and materials used in energy technology. Analytical chemistry is concerned with the composition of matter, focusing on separating and quantifying chemicals. Biochemistry is the study of chemical processes within living organisms, ranging from basic cellular functions to the investigation of disease treatments. Beyond providing useful materials, the study of chemistry is also critical for understanding environmental challenges caused by toxic substances like fluorochlorohydrocarbons and nitrogen oxides.

The Relationship Between Chemistry and Other Natural Sciences

Chemistry is often referred to as the central science because it links various fields through the study of matter. Science itself is a process of learning about the natural universe through observation, testing, and model generation. While science is divided into fields for focus, there is significant overlap. Biochemistry sits at the intersection of chemistry and biology, focusing on metabolic processes and the isolation of compounds from medicinal plants. Geochemistry bridges geology and chemistry by studying the distribution and abundance of chemical compounds and isotopes within geologic environments. Physical chemistry and chemical physics use techniques from atomic and molecular physics to investigate chemical systems, focusing on the interactions of matter and energy. Furthermore, medicinal chemistry links chemistry and medicine through the design, synthesis, and development of pharmaceutical drugs. These overlaps demonstrate that high-level understanding in one field often requires foundational knowledge in chemistry.

The Role of Chemistry in Production, Society, and Industry

In agriculture, chemistry has revolutionized crop yields through the creation of chemical fertilizers such as calcium super phosphate, urea, ammonium sulphate, and sodium nitrate. It is also responsible for the manufacture of pesticides, including fungicides, herbicides, and insecticides, which mitigate crop damage. In the food sector, chemistry provides preservatives to prolong the lifespan of products and methods to detect adulterants, ensuring food quality and nutritional content. In medicine, it provides life-saving drugs like cisplatin and Taxol for cancer therapy, and AZT for managing HIV-AIDS. It also provides essential medical categories such as analgesics for pain relief, anesthetics for surgery, antibiotics for infections, and antiseptics for wound care. Disinfectants and sanitizers are also critical for public health, particularly in managing pandemics like COVID-19. In building construction, chemistry provides resources such as high-grade steel, glass, and durable cement, which are vital for projects like the Grand Ethiopian Renaissance Dam (GERD), a $6,450\,MW$ hydropower project.

Ethiopia hosts several medium and large-scale chemical industries that convert natural resources into finished products. The Chorra Gas and Chemical Products factory in Addis Ababa produces plastics and petroleum products, as well as aluminum sulphate and sulphuric acid. The Ziway Caustic Soda Factory produces sodium hydroxide, while the Abijata Soda Ash Factory in Bulbula produces trona ($Na_3H(CO_3)_2 \times 2H_2O$). Repi Soap and Detergent P.L.C. and Etab Laundry Soap Factory are major producers of cleaning agents. The Adami Tulu Pesticide Processing Plant formulates essential agrochemicals like malathion and diazinon. Other notable enterprises include the Nefas Silk Paints Factory, Tadesse Filatea PLC which produces soap and infant milk formula, and various pharmaceutical labs in Addis Ababa and Adigrat. The Ethiopian government is currently expanding these capabilities through the development of industrial parks to boost national revenue and production.

Measurements, Units, and Temperature Scales

Chemistry is an experimental science that relies heavily on the International System of Units (SI). There are seven base quantities: Length (meter, m), Mass (kilogram, kg), Time (second, s), Electrical current (ampere, A), Temperature (kelvin, K), Amount of substance (mole, mol), and Luminous intensity (candela, cd). Metric prefixes are used to modify these units by powers of ten, ranging from Tera ($10^{12}$) and Giga ($10^9$) down to Nano ($10^{-9}$) and Pico ($10^{-12}$). Common laboratory devices facilitate measurements: meter sticks for length, balances for mass, and graduated cylinders, pipettes, or burettes for volume. Macroscopic properties can be measured directly, while microscopic properties require indirect methods. In calculations, volume is often measured in liters ($L$), where $1\,L = 1\,dm^3 = 1000\,cm^3$, and $1\,mL = 1\,cm^3$.

Temperature measures the intensity of heat or the kinetic state of atoms. While Celsius ($^{\circ}C$) and Kelvin ($K$) are common in science, Fahrenheit ($^{\circ}F$) is also used. The conversion between Celsius and Kelvin follows the formula $K = ^{\circ}C + 273.15$. To convert between Celsius and Fahrenheit, the following formulas apply: $^{\circ}F = (1.8 \times ^{\circ}C) + 32$ and $^{\circ}C = \frac{^{\circ}F - 32}{1.8}$. Note that the Celsius and Kelvin scales have the same degree size, but different starting points, meaning a temperature change of $34^{\circ}C$ is identical to a change of $34\,K$. Heat is defined as energy that flows spontaneously from a hotter body to a colder body. Density is a derived unit representing mass per unit volume, expressed as $d = \frac{m}{V}$. Derived units also include area ($m^2$), speed ($m/s$), force ($kg \times m/s^2$ or Newtons), and energy ($kg \times m^2/s^2$ or Joules).

Uncertainty, Precision, and Accuracy in Scientific Data

Every measurement involves inherent uncertainty, categorized as systematic or random. Systematic uncertainties consistently shift values in one direction, often due to miscalibrated instruments or reaction times, and can theoretically be eliminated. Random uncertainties are unpredictable variations that arise from the estimated portion of a measurement and can be reduced by repeated trials but never fully removed. Precision refers to how closely multiple measurements of the same quantity agree with one another (reproducibility), while accuracy is the closeness of a single measurement to the true or accepted value. Scientists aim for both. Uncertainty is often reported as a range, such as $64\,mm \times 2.5\,mm$, or as a percentage calculated by multiplying the absolute uncertainty divided by the measurement by $100$. For instance, weighing $23.25\,g$ on a balance with an uncertainty of $0.05\,g$ results in a percent uncertainty of $0.2\%$.

Significant figures indicate the exactness of a measurement. The rules include: all non-zero digits are significant; zeros between non-zero digits are significant; leading zeros are not significant; and trailing zeros are significant if a decimal point is present. In addition and subtraction, the final answer must have the same number of decimal places as the measurement with the fewest decimal places. In multiplication and division, the result must have the same number of significant figures as the original number with the smallest number of significant figures. Scientific notation or exponential notation is used for very large or small numbers, such as writing $602,000,000,000,000,000,000,000$ as $6.02 \times 10^{23}$. This avoids ambiguity in the number of significant figures, particularly with trailing zeros in large whole numbers.

The Scientific Method and Laboratory Safety

The scientific method is a rigorous process used to investigate natural phenomena. It begins with Observation and the formulation of a Question. This leads to Data Collection and the creation of a Hypothesis, which is an educated guess about the cause or relation of the observation. Testing the Hypothesis is done through experimentation to see if the hypothesis is supported or contradicted. Finally, Analysis and Conclusion involve interpreting experimental results. If a hypothesis is repeatedly verified without discrepancy, it may become a scientific Theory. A scientific Law differs in that it is a generalized description of an observation (like the Law of Gravity) but does not necessarily explain the underlying "why."

Laboratory safety is paramount and based on common sense and protocol. Students must wear approved eye protection and avoid eating or smoking in the lab. Unauthorized experiments are prohibited, and one must never work alone. Chemicals should be handled carefully; for example, one must always pour acid into water to avoid violent spatters, never the reverse. Long hair and loose clothing must be confined to prevent fire hazards. In the event of skin contact with chemicals, affected areas should be flushed with water immediately. For volume transfers, mouth suction on pipettes is strictly forbidden; suction devices must be used. Laboratory reports should be neat, written in ink, and include pre-lab preparation, documented procedures, detailed observations using correct significant figures, and graphs. Graphs must have labeled axes (abscissa for x-axis, ordinate for y-axis) and a descriptive title.

Fundamental Laws of Chemical Reactions and Early Atomic Theory

Chemical reactions are governed by three fundamental laws of combination. The Law of Conservation of Mass, discovered by Antoine Lavoisier in 1789, states that mass is neither created nor destroyed in a chemical reaction; the total mass of reactants equals the total mass of products in a closed system. The Law of Definite Proportions (Proust’s Law) states that a chemical compound always contains the same fixed proportions of its constituent elements by mass. For instance, water ($H_2O$) always contains hydrogen and oxygen in a mass ratio of $1:8$. The Law of Multiple Proportions, proposed by John Dalton, states that when two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in ratios of small whole numbers. For example, in $CO$ and $CO_2$, the masses of oxygen combining with $1\,g$ of carbon are in a $1:2$ ratio.

The history of atomic theory began with Democritus and Leucippus around $440\,BC$, who proposed that matter is made of indivisible bits called "atomos." They believed atoms were solid, homogeneous, and separated by a void. However, Aristotle rejected this, and the idea lay dormant for two millennia. John Dalton revived the concept in 1808 based on scientific evidence. His theory stated that all elements are made of atoms; atoms of the same element are identical in mass and size; atoms of different elements have different properties; and atoms combine in simple whole numbers to form compounds. While some of Dalton's points were later modified (such as the existence of isotopes and subatomic particles), his work provided the first modern scientific framework for chemistry.

Discovery of Subatomic Particles and Atomic Models

The internal structure of the atom was revealed through various high-profile experiments. J.J. Thomson discovered the electron in 1897 using a cathode ray tube. He found that cathode rays travel in straight lines, possess particle mass (they can rotate a paddle wheel), and are negatively charged (deflected toward positive plates). Thomson proposed the "Plum Pudding" model, where electrons are embedded in a positive sphere like seeds in a watermelon. Robert Millikan refined this in 1909 via the Oil Drop Experiment, determining the electron's charge to be $-1.59 \times 10^{-19}$ Coulombs and its mass to be approximately $9 \times 10^{-31}\,kg$. The proton was predicted by Eugene Goldstein in 1886 using anode rays (canal rays), which are positively charged gaseous ions moving toward a perforated cathode.

Ernest Rutherford’s Gold Foil Experiment in 1911 revolutionized atomic understanding. By bombarding thin gold foil with alpha particles ($2$ protons, $2$ neutrons), he observed that while most passed through, a tiny fraction deflected at large angles or bounced back. He concluded that the atom's mass and positive charge are concentrated in a tiny central region called the nucleus. This led to the Planetary Model of the atom. In 1913, Niels Bohr improved this by proposing that electrons orbit the nucleus in specific, stationary energy levels ($K, L, M, N$). James Chadwick completed the basic model in 1932 by discovering the neutron, a neutral particle in the nucleus with mass nearly equal to a proton. Modern Atomic Theory now acknowledges that atoms can be subdivided into these particles and that isotopes have different masses but the same chemical identity.

Atomic Number, Mass Number, and Isotopes

An element is defined by its atomic number ($Z$), which represents the number of protons in its nucleus. In a neutral atom, the number of protons equals the number of electrons. The mass number ($A$) is the sum of protons and neutrons in the nucleus ($A = Z + N$). These are written in nuclear symbols as $^A_Z X$. For example, Carbon-12 is represented as $^{12}_6 C$. Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons. Hydrogen has three primary isotopes: Protium ($1$ proton, $0$ neutrons), Deuterium ($1$ proton, $1$ neutron), and Tritium ($1$ proton, $2$ neutrons). Because isotopes exist as natural mixtures, chemists use the average atomic mass, which is a weighted average of the masses of all isotopes based on their natural percent abundance.

The calculation for average atomic mass follows the formula: $\text{Average Mass} = \frac{(\% \text{isotope 1} \times \text{mass 1}) + (\% \text{isotope 2} \times \text{mass 2})}{100}$. For example, Lithium has two isotopes: Li-7 ($93\%$) and Li-6 ($7\%$). The average mass is calculated as $[(7 \times 0.93) + (6 \times 0.07)] = 6.93\,amu$. An atomic mass unit ($amu$) is defined as exactly one-twelfth the mass of a Carbon-12 atom. Electrons are arranged in principal energy levels or shells, where the maximum capacity follows the rule $2n^2$. Thus, the first shell holds $2$, the second $8$, the third $18$, and the fourth $32$ electrons. The outermost shell, known as the valence shell, can hold a maximum of $8$ electrons, determining the element's chemical reactivity.

Periodic Classification of Elements and Trends

The periodic table evolved from early attempts to group elements. Johann Dobereiner identified "Triads" where the middle element's mass was the average of the other two (e.g., Li, Na, K). John Newlands proposed the "Law of Octaves," noting similarities every eighth element, similar to musical notes. Dmitri Mendeleev created the first widely accepted periodic table in 1871, arranging elements by atomic weight and leaving gaps for undiscovered elements like Gallium and Germanium. The Modern Periodic Table, based on Henry Moseley’s work, arranges elements by atomic number. It consists of 18 vertical groups and 7 horizontal periods. Elements in a group share the same number of valence electrons and similar chemical properties, while elements in a period share the same number of electron shells.

Major periodic trends allow for the prediction of element behavior. Atomic radius increases down a group (due to increasing shells) and decreases from left to right across a period (due to increasing nuclear pull). Ionization energy, the energy required to remove an electron, increases across a period and decreases down a group. Electron affinity measures the energy change when an atom gains an electron; it generally increases across a period. Electronegativity is the tendency of an atom in a molecule to attract shared electrons toward itself. Defined by the Pauling Scale, Fluorine is the most electronegative ($4.0$) and Cesium is the least ($0.7$). Metals generally have low electronegativity and low ionization energy, while non-metals have high values in both categories.

Chemical Bonding: Ionic, Covalent, and Metallic

Chemical bonding occurs as atoms seek to obey the Octet Rule, reaching a stable state with eight valence electrons (resembling a noble gas). Ionic bonding involves the total transfer of electrons from a metal (which becomes a positive cation) to a non-metal (which becomes a negative anion). The resulting electrostatic attraction creates ionic compounds. These substances typically form crystal lattices, have high melting and boiling points, are hard and brittle, and conduct electricity only when molten or dissolved in water. Lewis dot structures represent these bonds by showing the symbol and valence electrons as dots. For instance, Sodium ($2, 8, 1$) transfers one electron to Chlorine ($2, 8, 7$) to form $NaCl$.

Covalent bonding involves the sharing of electron pairs between non-metal atoms. Single bonds share one pair, double bonds share two, and triple bonds share three. Non-polar covalent bonds occur between identical atoms or those with minimal electronegativity differences ($< 0.5$). Polar covalent bonds ($0.5$ to $2.0$) involve unequal sharing, creating partial positive ($\delta+$) and negative ($\delta-$) charges, as seen in $HCl$. Coordinate covalent (dative) bonds occur when both shared electrons come from a single atom, such as the lone pair on nitrogen in ammonia ($NH_3$) sharing with boron in $BF_3$. Metallic bonding is described by the "Sea of Electrons" model, where positive metal ions are surrounded by delocalized valence electrons. This model explains why metals are excellent conductors of heat and electricity, as the electrons are free to move throughout the structure.

Questions & Discussion

The textbook includes various startup activities and group discussion prompts designed for classroom interaction. Early in the curriculum, students are asked to discuss the composition of air, water, and clothing to understand that everything is made of matter. In measurements, students participate in exercises identifying traditional Ethiopian indigenous measurement methods for length, mass, and time compared to standardized SI units. During atomic theory lessons, students are prompted to imagine splitting a seed to its smallest parts to conceptualize levels of division. Discussion regarding the Grand Ethiopian Renaissance Dam (GERD) focuses on the practical application of building materials like cement and steel. Periodic table activities include using a calendar of Nehassie 2013 E.C. as a metaphor for repeating periodic patterns. Interaction also involves identifying chemical industries in one's local town and researching solution methods for environmental pollution caused by local chemical waste. These prompts serve to connect laboratory chemistry with the social and economic development of Ethiopia.

Lab Experiment: Measurements and Density

The primary lab experiment focuses on learning to use common measuring instruments like the balance, graduated cylinder, and pipette. Students weigh a 50 mL beaker across three different balances to determine the average mass and evaluate the sensitivity of each instrument. Volume measurements are conducted by comparing the accuracy of a pipette (marked "TD" for To Deliver) with a graduated cylinder and a beaker. The experiment involves measuring specific aliquots of water and using the density of water to verify the volume. Students must determine the density of a cylindrical metal bar (Aluminum, Copper, or Brass) using two methods: Method I involves water displacement in a graduated cylinder, and Method II involves measuring dimensions with a ruler and using the formula $V = \pi r^2 h$. Results are evaluated for relative error and relative average deviation to distinguish between precision and accuracy. The experiment concludes with determining the density of a salt solution using both a volumetric pipette and a graduated cylinder, discussing why the metal bar or salt solution occupies a larger volume based on the context of matter density.