Introduction to Chemistry PDF
Document Details
Uploaded by SelfSatisfactionAllusion
Tags
Summary
This document provides a foundational introduction to chemistry, explaining its core concepts, branches (like organic, inorganic, and physical chemistry), and diverse applications in various fields. It also outlines the scientific method, crucial for conducting chemical research.
Full Transcript
Introduction to Chemistry I. CHEMISTRY AS A SCIENCE What is Chemistry? - Chemistry is a branch of natural science that deals principally with the properties of matter, the changes they undergo, and the natural laws that describe these changes. It is often referred to as the "central science" as it...
Introduction to Chemistry I. CHEMISTRY AS A SCIENCE What is Chemistry? - Chemistry is a branch of natural science that deals principally with the properties of matter, the changes they undergo, and the natural laws that describe these changes. It is often referred to as the "central science" as it connects other sciences like physics, biology, and geology. Key Points: Chemistry is the study of the composition, structure, properties, and behavior of matter. It is a fundamental science that underpins many other scientific disciplines. Chemistry plays a critical role in understanding and shaping the world around us. II. BRANCHES OF CHEMISTRY Organic Chemistry - The study of carbon-based compounds, including those found in living organisms. This includes the study of organic reactions, structures, properties, polymers, drugs, and fuels. Inorganic Chemistry - The study of compounds not covered by organic chemistry, such as ionic compounds, organometallic compounds, minerals, cluster compounds, and solid-state compounds. Analytical Chemistry - The study of the composition and properties of matter, and the development of tools to measure those properties. This includes quantitative and qualitative analysis, separations, extractions, distillation, and spectroscopy. Physical Chemistry - The application of thermodynamics and quantum mechanics to chemistry. This branch focuses on the physical properties of matter and the relationships between them. Biochemistry - The study of the chemical processes that occur within living organisms, including the study of proteins, nucleic acids, lipids, and carbohydrates. Other Branches - Astrochemistry, Chemical kinetics, Green chemistry, Nuclear chemistry, Environmental chemistry, and more. III. THE SCOPE OF CHEMISTRY Chemistry has a wide range of applications and impacts many aspects of our lives: Medicine: Developing new drugs, understanding biological processes, and improving medical treatments. Environment: Studying environmental pollution, developing sustainable technologies, and understanding natural processes. Industry: Improving manufacturing processes, creating new materials, and developing innovative products. Food and Agriculture: Enhancing food production, improving crop yields, and developing new food processing techniques. Materials Science: Designing and creating new materials with desired properties. Energy: Developing alternative energy sources, improving energy efficiency, and understanding energy- related processes. Forensics: Analyzing evidence and solving crimes through the application of chemical principles. Space Exploration : Understanding the chemical composition of celestial bodies and developing materials for space travel. IV. CHEMISTRY IN HUMAN LIFE Chemistry plays a crucial role in our daily lives, both in obvious and subtle ways: Daily Life Applications: Cleaning agents, cooking processes, and preservatives. Innovative Technologies: Development of smart materials, biodegradable plastics, and advanced medical devices. V. THE SCIENTIFIC METHOD - The scientific method is an empirical method for acquiring knowledge. systematic relies on observation, experimentation, and evidence-based data collection Obtain verifiable and replicable results Ensures objectivity and consistency in research. Allows for reproducibility Fundamental for advancing scientific knowledge and technology. WHY USE THE SCIENTIFIC METHOD? Eliminate Bias: By following a structured approach, personal biases and preconceived notions are minimized. Structured Problem: Provides a clear path from observation to conclusion, ensuring that every step is logically connected. Validation of Hypothesis: Allows scientists to test hypotheses and either support or refute them based on empirical evidence. VI. STEPS OF THE SCIENTIFIC METHOD 1. OBSERVATION The process begins with observing phenomena or gathering data through the senses or instruments. E.g. You noticed that a plant grows faster in sunlight than in the shade 2. QUESTION Based on the observation, a question is formulated that seeks to explain what was observed. E.g. Why does the plant grow faster in sunlight? 3. HYPOTHESIS A hypothesis is an educated guess or a tentative explanation that can be tested through experimentation. E.g. The plant grows faster in sunlight because sunlight provides the energy needed for photosynthesis. 4. EXPERIMENT Design and conduct an experiment to test the hypothesis. CONTROL VS. EXPERIMENTAL GROUPS Control group: Does not receive the experimental treatment Experimental group: Receives the treatment VARIABLES Independent Variable: The factor that is changed or manipulated Dependent Variable: The factor being measured ✓ E.g. Grow two sets of plants, one in sunlight and one in the shade, and measure their growth over time. 5. ANALYSIS Analyze the data collected during the experiment to determine if it supports or refutes the hypothesis. Use statistical tools to interpret data. Look for patterns, correlations, or significant differences. E.g. Comparing the growth rates of plants in sunlight versus shade and determining if the difference is statistically significant. 6. CONCLUSION Based on the analysis, the hypothesis is either accepted, rejected, or modified. - Accept Hypothesis - Reject Hypothesis - Modify Hypothesis E.g. Conclude that sunlight does indeed increase plant growth, or identify other factors that might be at play. 7. COMMUNICATION Sharing the results with the scientific community through publications, presentations, or reports. Allows others to verify the results, apply the findings, or build upon the research. VII. IMPORTANCE OF SCIENTIFIC METHOD Scientific Discovery: By following a structured method, scientists can discover new chemical reactions, elements, or compounds. Reproducibility: Other scientists must be able to replicate experiments and obtain similar results to validate findings. Peer Review: Before research is accepted, it undergoes peer review, where other experts evaluate the methods and conclusions. Practical Applications: The scientific method has led to breakthroughs in medicine, technology, and environmental science, all of which have practical benefits in everyday life. Exercise Assign each step of the scientific method to the corresponding situations: Increasing the temperature increases the reaction rate. Hypothesis Noticing that a reaction occurs faster when the solution is heated. Observation Determine whether the hypothesis is supported or not. Conclusion Plot the reaction time against temperature and analyze the trend. Analysis How does temperature affect the rate of the reaction? Question Share findings in a lab report or presentation. Communication Conduct the reaction at different temperatures (e.g., 20°C, 40°C, 60°C) and measure the time taken for the reaction to complete. Experiment MEASUREMENT OBJECTIVES: T o learn how to identify and use significant figures in measurements and calculations to accurately represent the precision of data in scientific work. To apply scientific notation as a method for expressing very large or very small numbers efficiently and accurately in scientific calculations. To master the rules for rounding off numbers in various contexts, ensuring accuracy and consistency in numerical data presentation and calculations. To explore different systems of measurement, understand their units, and convert between them to enhance accuracy and clarity in scientific and everyday applications. Measurements: M easurement is the quantification of attributes of an object or event, which can be used to compare with other objects or events. An important part of most experiments involves the determination (often, the estimation) of quantity, volume, dimensions, capacity, or extent of something—these determinations are measurements. In many cases, some sort of scale is used to determine a value such as this. In these cases, estimations rather than exact determinations need to be made. SI Units: (Système International d’Unités) Physical Quantity Name of Unit Abbreviation Mass kilogram kg Length meter m Time second s Temperature kelvin K Amount of Substance mole mol Electric Current ampere A Luminous Intensity candela cd Pressure pascal Pa Derived Units: A derived unit is a unit that results from a mathematical combination of SI base units. Quantity Symbol Unit Abbreviation Derivation Area A square meter m² length × width Volume V cubic meter m³ length × width × height Density D kilograms per cubic meter kg/m³ mass / volume Concentration c moles per liter mol/L amount / volume Speed v meters per second m/s distance / time Acceleration a eters per second m m/s² velocity / time squared Force F newton N or kg·m/s² mass × acceleration Energy E joule J or kg·m²/s² force × length Conversion of Measurements: Length: km = 0.62137 mi 1 1 mi = 5280 ft = 1.6093 km 1 m = 1.0936 yd 1 in = 2.54 cm 1 Å = 10⁻¹⁰ m Volume: L = 10⁻³ m³ = 1 dm³ = 1000 cm³ = 1.0567 qt 1 1 gal = 4 qt = 3.78541 L 1 cm³ = 1 mL 1 in³ = 16.4 cm³ Mass: 1 kg = 2.2046 lb 1 lb = 453.49 g = 16 oz 1 amu = 1.66 × 10⁻²⁴ g Pressure: 1 Pa = 1 N/m² = 1 kg/m·s² 1 atm = 101.325 Pa = 760 torr = 14.70 lb/in² 1 bar = 10⁵ Pa Temperature: 0 K = -273.14°C = -459.67°F K = °C + 273.15 °C = 5/9(°F - 32) Energy: 1 J = 0.2390 cal 1 cal = 4.184 J 1 eV = 1.602 × 10⁻¹⁹ J Measurement Tools: Physical Quantity Measurement Tools Mass mechanical scale, digital scale Length r uler, measuring tape, caliper, odometer Time clock, timer, stopwatch Temperature thermometer, pyrometer mount of A mass spectrometer Substance Electric Current ammeter, multimeter Luminous Intensity photometer Pressure barometer, pressure gauge Dimensional Analysis (Conversion Factors): T he term "dimensional analysis" refers to a procedure that yields the conversion of units and follows the general formula: 𝐺𝑖𝑣𝑒𝑛𝑈𝑛𝑖𝑡𝑠(𝐷𝑒𝑠𝑖𝑟𝑒𝑑𝑈 𝑛𝑖𝑡𝑠/𝐺𝑖𝑣𝑒𝑛𝑈𝑛𝑖𝑡𝑠{𝑐𝑜𝑛𝑣𝑒𝑟𝑠𝑖𝑜𝑛𝑓 𝑎𝑐𝑡𝑜𝑟}) = 𝐷 𝑒𝑠𝑖𝑟𝑒𝑑𝑈𝑛𝑖𝑡𝑠 .A dvantages of Using Dimensional Analysis: Reinforces the use of units of measurement. Eliminates the need for memorizing specific formulas for most problems. For example:How many moles of H₂O are present in 27.03 g H₂O? . Identify the conversion factor between grams and moles. 1 2. Multiply the given quantity by the appropriate factor. o solve the example "How many moles of H₂O are present in 27.03 g H₂O?", we need to use the molar mass of water (H₂O) T to convert from grams to moles. Step 1: Find the molar mass of H₂O Hydrogen (H) has an atomic mass of approximately 1.008 g/mol. Oxygen (O) has an atomic mass of approximately 16.00 g/mol. The molar mass of H₂O is: 𝑀𝑜𝑙𝑎𝑟𝑚𝑎𝑠𝑠𝑜𝑓𝐻₂𝑂 = (2×1. 008𝑔/ 𝑚𝑜𝑙) + 16. 00𝑔/ 𝑚𝑜𝑙 = 18. 016𝑔/𝑚𝑜𝑙 Step 2: Set up the dimensional analysis o convert 27.03 g of H₂O to moles, divide the mass by the molar mass: T 𝑀𝑜𝑙𝑒𝑠𝑜𝑓𝐻 ₂𝑂 = 18. 016𝑔/ 𝑚 𝑜𝑙27. 03𝑔𝐻₂𝑂 = 1. 501𝑚𝑜𝑙𝐻₂ 𝑂 Final Answer: There are 1.501 moles of H₂O in 27.03 g of H₂O. Accuracy vs. Precision: A ccuracyrefers to the proximity of a measurementto the true value of a quantity. Precisionrefers to the proximity of several measurementsto each other. For a measured quantity, we can generally improve its accuracy by making more measurements. Measured Quantities and Uncertainty: T he measured quantity, for example, 3.7, is an estimation. We are more confident in the "3" than in the "7" (we are sure of the 3, but not so sure of the 7). Whenever possible, you should estimate a measured quantity to one decimal place smaller than the smallest graduation on a scale. Uncertainty in Measured Quantities: W hen measuring, for example, how much an apple weighs, the mass can be measured on a balance. The balance might be able to report quantities in grams, milligrams, etc. Let’s say the apple has a true mass of 55.51 g. The balance we are using reports mass to the nearest gram and has an uncertainty of +/- 0.5 g. The balance indicates a mass of 56 g The measured quantity (56 g) is true to some extent and misleading to some extent. The quantity indicated (56 g) means that the apple has a true mass which should lie within the range 56 +/- 0.5 g (or between 55.5 g and 56.5 g). Significant Figures: figures with some degree of reliability 1. All non-zero digits are significant (e.g., 1.644 has 4 significant figures). 2. Zeros between two non-zero figures are themselves significant (e.g., 1.6044 has 5 significant figures). 3. Zeros at the beginning of a number are never significant (e.g., 0.005 has 1 significant figure). 4. Zeros at the end of a number are significant if a decimal point is written in the number (e.g., 1500. has 4 significant figures; 1500.0 has 5 significant figures). For the number 1500, assume there are two significant figures unless further context or a decimal point clarifies the number of significant digits. Exact Quantities: In certain cases, some situations will utilize relationships that are exact, defined quantities. ○ For example, a dozen is defined as exactly 12 objects (eggs, cars, donuts, whatever...) ○ 1 km is defined as exactly 1000 m. ○ 1 minute is defined as exactly 60 seconds. Each of these relationships involves aninfinite numberof significant figuresfollowing the decimal placewhen being used in a calculation. Relationships between metric units are exact (e.g. 1 m = 1000 mm, exactly) Relationships between imperial units are exact (e.g. 1 yd = 3 ft, exactly) Relationships between metric and imperial units are not exact (e.g. 1.00 in = 2.54 cm) Rules for Mathematical Operations: Addition and Subtraction: Limit the reported answer to the rightmost column that all numbers have significant figures in common. he result should be reported to the least precise decimal place (rightmost column) that all numbers share. T Example:12. 11 + 18. 0 + 1. 013 = ? 12.11 (2 decimal places) 18.0 (1 decimal place) 1.013 (3 decimal places) Since 18.0 has only 1 decimal place, the final answer should be rounded to 1 decimal place: 1. 1 Answer: 31.1 12. 11 + 18. 0 + 1. 013 = 31. 123→3 Multiplication and Division: C ount the number of significant figures in each number being multiplied or divided and then limit the significant figures in the answer to the lowest count. he result should be reported with the same number of significant figures as the number in the calculation with the fewest T significant figures. Example(Multiplication):4. 56×1. 4 = ? 4.56 has 3 significant figures. 1.4 has 2 significant figures. The result will be rounded to 2 significant figures: 6. 4 Answer: 6.4 4. 56×1. 4 = 6. 384→ Example(Division):25. 32÷1. 1 = ? 25.32 has 4 significant figures. 1.1 has 2 significant figures. The result should be rounded to 2 significant figures: 3 Answer: 23 25. 32÷1. 1 = 23. 018→2 Logarithm: W hen you take the logarithm of a number, keep as many significant figures to the right of the decimal point as there are significant figures in the original number. When taking the logarithm of a number, keep as many significant figures to the right of the decimal point as there are in the original number. Example:𝑙𝑜𝑔(3. 00) = ? 3.00 has 3 significant figures. The answer should have 3 decimal places. 𝑙𝑜𝑔(3. 00) = 0. 477121→0. 477 Answer: 0.477 Antilogarithm: W hen you take the antilogarithm of a number, the result should have the same number of significant figures as the number of significant decimal places in the principal value. When taking the antilogarithm of a number, the result should have the same number of significant figures as the number of decimal places in the original number. 1.754 xample:𝐴𝑛𝑡𝑖𝑙𝑜𝑔(1. 754) = 10 E = ? The number 1.754 has 3 decimal places, so the result should have 3 significant figures. 1.754 10 56. 8 Answer: 56.8 = 56. 75446054→ Rules for Rounding Off Numbers: If the digit to be dropped is greater than 5, the last retained digit is increased by one. ○ For example, 12.6 is rounded to 13. If the digit to be dropped is less than 5, the last remaining digit is left as it is. ○ For example, 12.4 is rounded to 12. If the digit to be dropped is 5, and if any digit following it is not zero, the last remaining digit is increased by one. ○ For example, 12.51 is rounded to 13. If the digit to be dropped is 5 and is followed only by zeroes, the last remaining digit is increased by one if it is odd, but left as it is if even. ○ For example, 11.5 is rounded to 12, 12.5 is rounded to 12. This rule means that if the digit to be dropped is 5 followed only by zeroes, the result is always rounded to the even digit. The rationale is to avoid bias in rounding: half of the time we round up, half the time we round down. Scientific Notations: S cientific notation ensures accuracy and reduces error in small or large numbers, facilitating easier interpretation for those unfamiliar with large or small numbers. Formula: 𝑏 𝑎×1 Where: a is a number or decimal number where 1 ≤ ∣a∣ < 10 b is the power of 10 required so that the scientific notation is mathematically equivalent to the original number. Calculations with Scientific Notation: Addition/Subtraction: Align the exponents before adding or subtracting. Multiplication/Division: Multiply/divide the numbers and add/subtract the exponents. Matter Objectives Identify and classify different types of matter and explain their physical and chemical properties. Differentiate between physical and chemical changes and describe the law of conservation of mass. Explain the historical development of atomic models and describe the structure of atoms, including subatomic particles and quantum theory. Apply the rules for writing electron configurations and orbital diagrams for elements. Illustrate Lewis structures for molecules and ions, including resonance structures and exceptions to the octet rule. What is Matter? M atteris anything that has mass and occupies space. Examples of matter include: ○ Tree ○ Cloud ○ Smoke ○ Air Momentum is not matter. Matter in Chemistry M atter is fundamental to chemistry. It serves as the: ○ Building blocks of the universe ○ Chemical reactions ○ Properties and classifications ○ Understanding phases and states ○ Real-world applications Classification of Matter Matter is broadly classified into two categories: ○ Pure Substances: Composed of a single type of atomor molecule. ○ Elements: Cannot be broken down into simpler substances(e.g., gold, oxygen). ○ Compounds: Formed by the chemical combination of twoor more elements (e.g., water, salt). ○ Mixtures: Composed of two or more substances thatare not chemically bonded. ○ Homogeneous Mixtures: Have a uniform composition throughout(e.g., saltwater). ○ Solutions: A special type of homogeneous mixture whereone substance is dissolved in another (e.g., sugar dissolved in water). ○ Heterogeneous Mixtures: Have a non-uniform composition(e.g., sand and water). Properties of Matter A. Physical Properties P hysical properties describe a substance's characteristics that can be observed or measured without changing its chemical composition. Examples of physical properties include: Property Example Color Blue, red, green Hardness Diamond is hard, butter is soft Malleability Gold can be hammered into thin sheets Solubility Sugar dissolves in water Electrical Conductivity Copper conducts electricity Density Water is less dense than oil Melting Point The temperature at which a solid turns into a liquid Boiling Point The temperature at which a liquid turns into a gas B. Chemical Properties C hemical properties describe a substance's potential to undergo a chemical change or reaction based on its composition. Examples of chemical properties include: Property Example Flammability Wood burns easily Reactivity with Acids Metals react with acids to produce hydrogen gas Corrosiveness Acids corrode metals pH A measure of acidity or alkalinity Heat of Combustion The amount of heat released when a substance burns Toxicity The ability of a substance to harm living organisms Chemical Stability The tendency of a substance to resist decomposition Oxidation States The number of electrons an atom gains or loses in a chemical reaction Extensive vs. Intensive Properties Extensive Properties: Depend on theamount of matterpresent (e.g., mass, volume, density). ○ ○ Intensive Properties: Depend only on thetype of matterpresent (e.g., color, temperature, solubility). States of Matter Matter exists in four main physical states: ○ Solid: Particles are tightly packed and arranged ina regular pattern, with limited movement. ○ Liquid: Particles are closer together than in a gas,but still have some movement. ○ Gas: Particles are far apart and move freely. ○ Plasma: A high-energy state of matter where atomsare ionized, forming a mixture of electrons and nuclei. Changes Between States of Matter T ransitions between states of matter involve energy changes. Enthalpy of System: The total energy content of the system. Increasing Energy: Transitions from solid to liquid, liquid to gas, and gas to plasma require an increase in energy (enthalpy). Transitions: ○ Solid to Liquid:Melting ○ Liquid to Solid:Freezing ○ Solid to Gas:Sublimation ○ Gas to Solid:Deposition ○ Liquid to Gas:BoilingorEvaporationorVaporization ○ Gas to Liquid:Condensation ○ Gas to Plasma:Recombination ○ Plasma to Gas:Ionization Changes in Matter A. Physical Changes P hysical changes alter the form or appearance of a substance but do not change its chemical composition. Characteristics of physical changes include: ○ No new substance is formed. ○ The change is reversible. ○ Examples include changes in color, temperature, or texture. B. Chemical Changes C hemical changes result in the formation of new substances with different chemical compositions. Characteristics of chemical changes include: ○ Changes in color, formation of gas or precipitate, changes in odor, temperature, or the production of light. C. Law of Conservation of Mass L aw of Conservation of Mass: Mass is neither created nor destroyed in a closed system during a chemical or physical change. Matter can change form through physical and chemical changes, but through any of these changes, matter is conserved. Example: The reaction between hydrogen and oxygen to form water: ○ 2H2 + O2 → 2H2O ○ Reactants, reaction, product Atomic Structure and Quantum Theory Historical Development of Atomic Theory Sphere (1803 | John Dalton) ○ Atoms are hard spheres ○ Atoms of the same element are the same ○ Combinations of different atoms make compounds Plum Pudding (1904 | J J Thompson) ○ Thomson discovered the electron in 1897 ○ Atoms are electrons scattered through a sea of positive charge Nuclear (1911 | Ernest Rutherford) ○ The atom is mostly empty space with a positive middle ○ He fired alpha particles at gold foil, some deflected but most passed through Planetary (1913 | Neils Bohr) ○ Electrons move around the nucleus in orbits of specific energy and size ○ Electrons fit into specific energy levels Quantum (1926 | Erwin Schrodinger) ○ There are clouds called orbits where electrons exist ○ Electrons do not orbit in a predetermined path around the nucleus Subatomic Particles ○ roton - positive electrical charge P ○ Electron - negative electrical charge ○ Neutron - neutral ○ Nucleus - protons and neutrons Particle Actual mass (g) Relative mass (amu) Relative charge proton 1. 6726𝑥10 −24 1. 007 +1 neutron 1. 6749𝑥10 −24 1. 008 0 electron 9. 108𝑥10 −28 −4 5. 54𝑥10 -1 Additional Notes: P rotonsare found in the nucleus of an atom and contributeto the atomic number, determining the element's identity. Neutronsare also located in the nucleus and contributeto the mass of the atom. Their number can vary in isotopes of the same element. Electronsorbit the nucleus in various energy levels(shells) and are responsible for chemical bonding and reactions. Quantum Mechanical Model of the Atom W ○ ave-Particle Duality: Electrons exhibit both wave-likeand particle-like properties. ○ Heisenberg Uncertainty Principle: It is impossibleto simultaneously know both the position and momentum of an electron with perfect accuracy. ○ Schrödinger's Equation and Electron Orbitals: A mathematicalequation that describes the behavior of electrons in atoms. Additional Knowledge: Electron Orbitals: Orbitals are categorized into types(s, p, d, f) based on their shapes and energy levels: ○ s-orbital:Spherical in shape. ○ p-orbital:Dumbbell-shaped. ○ d and f-orbitals:More complex shapes. Quantum Numbers P ○ rincipal Quantum Number (𝑛) : Describes the electron'senergy level. ○ Azimuthal Quantum Number (𝑙, 𝑛 − 1): Describes theshape of the electron's orbital. ○ Magnetic Quantum Number (𝑚𝑙, − 𝑙, ..., − 1, 0, 1, ..., 𝑙 ) : Describes the orientation of the electron's orbitalin pace. s ○ Spin Quantum Number (𝑚𝑠, + 1 /2, − 1/2 ): Describesthe intrinsic angular momentum of the electron. . 1 rincipal Quantum Number (n): P ★ Describes the energy level or shell an electron is in. ★ Can be any positive integer (n = 1, 2, 3,...). ★ Higher n values mean higher energy levels and larger orbitals. 2. Azimuthal Quantum Number (l): ★ Describes the shape of the orbital (s, p, d, f). ★ Values range from 0 to n-1: ○ l=0 → s-orbital (spherical) ○ l=1 → p-orbital (dumbbell-shaped) ○ l=2 → d-orbital ○ l=3 → f-orbital . 3 Magnetic Quantum Number (ml): ★ Describes the orientation of the orbital in space. ★ Values range from −l to +l (e.g., for l = 1, ml can be -1, 0, or +1). 4. Electron Spin Quantum Number (ms): ★ Describes the electron’s spin direction, which can be +1/2 or -1/2. Electron Configuration A ufbau Principle: Electrons occupy orbitals in orderof increasing energy. The Diagonal Rule: A mnemonic device used to rememberthe order of filling orbitals in electron configurations. Pauli's Exclusion Principle: No two electrons in thesame atom can have identical values for all four of their quantum numbers. ○ no more than two electrons can occupy the same orbital and ○ two electrons in the same orbital must have occupies spins Hund's Rule: When assigning electrons to orbitals,an electron first seeks to fill all the orbitals with similar energy. Writing Electron Configurations E lectron configurations are shorthand representations of the distribution of electrons in an atom. Examples: ○ Potassium (K): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ ○ Hydrogen (H): 1s¹ ○ Oxygen (O): 1s² 2s² 2p⁴ Orbital Diagrams A n orbital diagram, or orbital filling diagram, is a type of notation which illustrates an atom’s electron distribution and electron spin with orbitals. Orbital diagrams are visual representations of the electron distribution and electron spin within orbitals. How to Draw Orbital Diagrams . D 1 etermine the Electron Configuration: You'll need to know the electron configuration of the element. 2. Represent Orbitals: Draw boxes to represent each orbital. ○ s orbitals: Draw one box for the 's' subshell (e.g., 1s, 2s). ○ p orbitals: Draw three boxes for the 'p' subshell (e.g., 2p, 3p). ○ d orbitals: Draw five boxes for the 'd' subshell (e.g., 3d, 4d). ○ f orbitals: Draw seven boxes for the 'f' subshell (e.g., 4f, 5f). 3. Fill in Electrons: Add arrows to represent the electrons. ○ Hund's Rule: Fill each orbital with one electron (up arrow) before pairing them (down arrow). ○ Pauli Exclusion Principle: Each orbital can hold a maximum of two electrons with opposite spins (up and down arrows). Example: Oxygen (O) E lectron Configuration: 1s² 2s² 2p⁴ Orbital Diagram: ○ 1s: ↑↓ ○ 2s: ↑↓ ○ 2p: ↑ ↑ ↑ Example: Fluorine (F) E lectron Configuration: 1s² 2s² 2p⁵ Orbital Diagram: ○ 1s: ↑↓ ○ 2s: ↑↓ ○ 2p: ↑↓ ↑ ↑ Example: Neon (Ne) E lectron Configuration: 1s² 2s² 2p⁶ Orbital Diagram: ○ 1s: ↑↓ ○ 2s: ↑↓ ○ 2p: ↑↓ ↑↓ ↑↓ Noble Gas Notation A way of writing an abbreviated electron configuration, with the noble gas substituting the beginning energy levels and orbital-filled shells. Noble gas notation is a shorthand way of writing electron configurations using the noble gas symbol to represent the filled inner shells.
