General Chemistry Chapter 1 Notes Review

Janssen COVID-19 Vaccine Overview

General Information

  • Type: The Janssen COVID-19 Vaccine is a recombinant adenosine type 26 (Ad26) vector vaccine, meaning it uses a modified common cold virus to deliver genetic instructions for the SARS-CoV-2 spike protein into human cells. It is presented as a sterile suspension for intramuscular injection.

  • Authorization: This vaccine is designated for use under Emergency Use Authorization (EUA) by the U.S. Food and Drug Administration (FDA) for active immunization to prevent Coronavirus Disease 2019 (COVID-19) caused by the SARS-CoV-2 virus in individuals 18 years of age and older.

  • Dosage: Each multi-dose vial contains a target of 10 doses, with each dose being 0.5 mL. It is typically administered as a single-dose regimen.

  • Storage: Unopened vials should be stored frozen at -25^ ext{o} ext{C} to -15^ ext{o} ext{C} (-13^ ext{o} ext{F} to 5^ ext{o} ext{F}) until their expiration date. After the first puncture, the vial must be held at a refrigerated temperature range of 2^ ext{o} ext{C} to 8^ ext{o} ext{C} (36^ ext{o} ext{F} to 46^ ext{o} ext{F}) and discarded after 6 hours or 6 withdrawals, whichever comes first.

  • No Preservative: The vaccine formulation does not contain preservatives, which necessitates careful aseptic technique during preparation and administration to prevent contamination.

  • Expiration Date: For precise expiration information, consult the official website at www.vaxcheck.jnj, as expiration dates may vary by lot and can be extended based on stability data.

Chapter Outline

  1. Classification of Matter

  2. Properties of Matter

  3. Matter and Energy

  4. The Scientific Method, Hypotheses, Theories, and Laws

  5. The International System of Units

  6. Significant Digits

  7. Dimensional Analysis

  8. Density

  9. Temperature Scales

Section 1.1: Classification of Matter

Definition of Matter

  • Chemistry is fundamentally defined as the comprehensive study of matter and energy, including their composition, structure, properties, and the changes they undergo.

  • Matter: Anything that possesses mass and occupies space. This definition encompasses all physical substances, from microscopic atoms to macroscopic galaxies.

  • Elements: The most basic, pure forms of matter that consist of only one type of atom. Each element is distinguished by a unique number of protons in its nucleus (its atomic number) and cannot be broken down into simpler substances by ordinary chemical means. Examples include hydrogen (H), oxygen (O), and gold (Au).

  • Building Blocks: All the diverse forms of matter observed on Earth are constructed from approximately one hundred naturally occurring elements, which are systematically organized on the periodic table.

Pure Substances

  • Atom: The absolute smallest unit of an element that still retains the unique chemical identity and characteristics of that element. Atoms are composed of subatomic particles: protons, neutrons, and electrons.

  • Chemical Bonds: Atoms of the same or different elements can interact through attractive forces called chemical bonds, leading to the formation of molecules or more complex structures. These bonds result from the sharing or transfer of electrons between atoms.

  • Compounds: Chemical combinations of two or more different elements bonded together in a fixed, definite ratio. Compounds possess distinct physical and chemical properties that are entirely different from those of their constituent elements. For instance, water (H2O) is a liquid at room temperature, while its constituent elements, hydrogen (H2) and oxygen (O_2), are gases.

  • Types: Elements and compounds are both categorized as pure substances because they have homogenous compositions and uniform properties throughout.

Examples and Mixtures

Example 1.1

  • Identify mixtures:

    a. A container with separate samples of iron (Fe), copper (Cu), and zinc (Zn). This represents three distinct elements, not a mixture in the context of being combined.

    b. A flask containing hydrogen gas (H2) and water vapor (H2O). This contains one unreactive element (hydrogen) and one compound (water) physically mixed.

    c. A solution of table salt (NaCl) dissolved in water (H2O) and sugar (C{12}H{22}O{11}). This contains two different compounds (salt and sugar) entirely dissolved within another compound (water), forming a homogeneous mixture.

Mixtures

  • Mixtures consist of two or more pure substances that are physically combined, meaning their components are not chemically bonded to each other. Each component retains its individual chemical identity.

  • Separation: The individual components of a mixture can often be separated from one another by physical processes, such as filtration, distillation, decantation, magnetism, evaporation, or chromatography, without undergoing chemical change.

  • Definite Compositions: Unlike compounds, mixtures do not have fixed or definite compositions; the proportions of their components can vary widely. Mixtures can be broadly classified as either heterogeneous or homogeneous.

Heterogeneous Mixtures

  • In heterogeneous mixtures, the components are not uniformly distributed throughout the sample. Different samples taken from the same mixture may exhibit different compositions and properties, and often, the individual components are visibly distinguishable. Examples include sand and water, oil and vinegar, granite, or a bowl of mixed nuts.

Homogeneous Mixtures

  • Also commonly referred to as solutions. In homogeneous mixtures, the components are uniformly mixed at a molecular level, resulting in a consistent composition and appearance throughout the entire sample. These can exist in any phase (solid, liquid, or gas).

  • Examples: Air (a gas-phase solution primarily of nitrogen and oxygen), metal alloys like brass (a solid-phase solution of copper and zinc), and aqueous solutions (where water acts as the major component, such as saltwater or sugar water).

Table 1.1: Classification of Matter

  • Classification: This table provides practical examples to illustrate the classification of matter:

    • Pure Substances: Hydrogen (an Element), Water (a Compound composed of hydrogen and oxygen in a fixed 2:1 ratio by atoms).

    • Mixtures: Oil and water (a clearly visible two-phase Heterogeneous mixture), Brass (a uniform Homogeneous mixture or alloy of copper and zinc).

Figure 1.2: Classification of Matter

  • Structure of Matter: This illustrative diagram visually represents the hierarchical classification of matter, detailing the relationships between fundamental components such as pure substances, mixtures, elements, and compounds, showcasing how they branch from broader categories.

Section 1.2: Properties of Matter

Definition of Properties

  • Every substance possesses a distinct set of characteristics, known as properties, that enable it to be identified and differentiated from other substances.

  • Physical Properties: These are characteristics that can be observed, measured, or described without altering the chemical composition or identity of the substance. Examples include color, odor, density, melting point, boiling point, hardness, texture, and state of matter (solid, liquid, gas).

  • Chemical Properties: These properties describe a substance's potential to undergo chemical changes or reactions, thereby transforming into different substances. They explain how a substance reacts with other substances or breaks down. Examples include flammability (ability to burn), reactivity with acids or bases, oxidation (rusting), and decomposition.

Extensive and Intensive Properties

  • Extensive Properties: These properties are dependent on the amount of matter present in a sample. They change with the quantity of the substance. Examples include mass, volume, length, and total energy.

  • Intensive Properties: These properties are independent of the amount of matter present; their value does not change regardless of the sample size. Intensive properties are particularly useful for identifying substances because they are intrinsic to the material itself. Examples include density, melting point, boiling point, temperature, color, hardness, and concentration.

Example 1.5

Example Questions:

  • a. If sample A weighs twice as much as sample B, and both are unknown substances, can you identify them?

    • Answer: No, you cannot identify them based on weight alone. Weight is an extensive property, meaning it depends on the amount of material. Knowing the weight ratio doesn't provide information about the intrinsic nature of the substances.

  • b. If sample A is magnetic and sample B is a white powder that glows under ultraviolet light, can you identify them?

    • Answer: Yes, you can potentially identify them based on these observations. Magnetic attraction (a specific response to a magnetic field) and fluorescence (glowing under UV light) are intensive properties. These characteristics are inherent to the material regardless of its quantity and can serve as strong clues for identification (e.g., sample A could be iron, and sample B could be a specific mineral like calcite or zinc sulfide).

Properties of Compounds and Mixtures

  • The properties of compounds are fixed and uniquely different from those of their constituent elements. For instance, water (H2O) extinguishes fires, while its component elements, hydrogen (H2) and oxygen (O_2), are highly flammable and support combustion, respectively. This change in properties is a result of the formation of new chemical bonds.

  • The properties of mixtures, conversely, are not fixed but rather a blend or combination of the properties of their individual components. The overall properties of a mixture often depend directly on the proportion of each component present. For example, saltwater will have a different boiling point and density than pure water, and these properties will vary with the amount of salt dissolved.

Changing Matter

Physical Changes

  • A physical change occurs when a substance alters its physical appearance or state (e.g., shape, size, phase) but does not change its chemical composition or create new substances. The chemical identity of the substance remains the same.

  • Examples: Melting of ice into liquid water (still H_2O), boiling water into steam, dissolving sugar in water, cutting paper, bending a metal wire, or crushing a rock. These changes are often reversible.

Chemical Changes

  • A chemical change, also known as a chemical reaction, involves the transformation of reactants into entirely new products with different chemical compositions and properties. This process typically involves the breaking and forming of chemical bonds.

  • Examples: Iron rusting (iron reacts with oxygen to form iron oxide), burning wood (cellulose reacts with oxygen to form carbon dioxide, water, and ash), cooking an egg, digestion of food, or the reaction of baking soda and vinegar.

Table 1.2: Properties of Iron, Sulfur, and Iron–Sulfur Compound

  • This table serves to list and compare the distinct physical properties (e.g., phase at room temperature, luster, magnetism, color, density, solubility) of individual elements like iron and sulfur with those of the compound formed when they chemically react, iron sulfide. It visually demonstrates how a compound’s properties are different from its constituent elements.

Chemical Bond Changes

  • Chemical reactions always involve changes in chemical bonds. The process of breaking existing chemical bonds in reactants absorbs energy from the surroundings (endothermic process). Conversely, the process of forming new chemical bonds in products releases energy into the surroundings (exothermic process). The overall energy change of a reaction is the net difference between the energy absorbed and released.

Section 1.3: Matter and Energy

Mass

  • Mass is a fundamental measure of the amount of matter in an object. It represents the object's resistance to acceleration (inertia) and is a constant intrinsic property regardless of its location. Unlike weight, which is a measure of the gravitational force exerted on an object and varies with location (e.g., on Earth vs. the Moon), mass remains unchanged.

Energy

  • Energy is defined as the capacity to do work (apply a force over a distance) or to produce heat. It is a fundamental property of matter and radiation.

  • Governed by the Law of Conservation of Energy (also known as the First Law of Thermodynamics), which states that in an isolated system, energy cannot be created or destroyed. Instead, it can only be transformed from one form to another or transferred from one system to another. Various forms of energy include thermal (heat), chemical (stored in bonds), nuclear (stored in atomic nuclei), mechanical (kinetic and potential), electrical, light (electromagnetic radiation), and sound energy.

Example 1.7

  • This example effectively describes various energy conversions in common scenarios. For instance, in a flashlight, chemical energy stored in batteries is converted into electrical energy, which then transforms into light energy (to illuminate) and thermal energy (some heat generation). Similarly, in a phone, chemical energy in the battery is converted to electrical energy, which powers computations, communication, and produces light (screen), sound, and heat.

Section 1.4: The Scientific Method, Hypotheses, Theories, and Laws

The Scientific Method

  • The scientific method is a rigorous, systematic, and empirical approach to acquiring knowledge and understanding the natural world. It combines careful observation, the formulation of testable hypotheses, and systematic experimentation to validate or refute these hypotheses.

Steps in the Scientific Method:

  1. Identify observed data/Ask a Question: Begin with an observation that leads to a question about a natural phenomenon. This often involves collecting initial qualitative or quantitative data.

  2. Formulate hypotheses (possible explanations): Based on the observations, propose a tentative, testable explanation or prediction for the phenomenon. A good hypothesis is specific and falsifiable.

  3. Test hypotheses through experimentation: Design and conduct controlled experiments to gather empirical data that will either support or contradict the hypothesis. This step involves careful data collection, analysis, and interpretation.

Hypotheses, Theories, and Laws

  • Hypothesis: An initial explanation or proposed explanation for an observable phenomenon. It is an educated guess based on prior observations or knowledge, which must be testable through experimentation. A hypothesis is a starting point for further investigation.

  • Theory: A well-substantiated, comprehensive, and unifying explanation for a broad range of related phenomena. A theory is not merely a guess; it is extensively supported by a vast body of evidence from numerous experiments and observations. Theories can be modified or refined as new evidence emerges but are rarely disproven entirely (e.g., the Theory of Evolution, Kinetic Molecular Theory).

  • Law: A descriptive statement that reliably summarizes an observed natural phenomenon or a set of phenomena, often expressed mathematically. Scientific laws describe what happens under certain conditions but generally do not explain why it happens (e.g., Law of Gravity, Law of Conservation of Mass).

Figure 1.8: The Scientific Method

  • This diagram visually illustrates the iterative and cyclical nature of the scientific method, showing the logical flow from initial observations to asking questions, formulating hypotheses, designing experiments, analyzing results, and ultimately leading to the formation or revision of theories. It emphasizes that scientific understanding is always subject to refinement.

Law of Conservation of Mass

  • The Law of Conservation of Mass, proposed by Antoine Lavoisier in 1785, states that in any closed system undergoing a chemical reaction, the total mass of the reactants before the reaction must equal the total mass of the products after the reaction. Mass is neither created nor destroyed during a chemical change.

  • Recent theories indicate its exceptions in nuclear reactions, where a measurable amount of mass can be converted into energy, as described by Einstein's mass-energy equivalence equation (E=mc^2). However, for typical chemical reactions, the law holds true to a very high degree of accuracy.

Section 1.5: The International System of Units

SI Base Units

  • The International System of Units (SI) is a modern form of the metric system, establishing a universally accepted set of seven fundamental base units from which all other units can be derived. These fundamental units encapsulate core physical quantities:

    • Length (meter, m)

    • Mass (kilogram, kg)

    • Time (second, s)

    • Electric Current (ampere, A)

    • Thermodynamic Temperature (kelvin, K)

    • Amount of Substance (mole, mol)

    • Luminous Intensity (candela, cd)

Common Conversions:

  • These are crucial for comparing and exchanging scientific data:

    • 1 meter (m) = 100 centimeters (cm) = 1000 millimeters (mm)

    • 1 kilogram (kg) \approx 2.2046 pounds (lb) (for mass conversion between metric and imperial systems)

    • 1 liter (L) = 1000 milliliters (mL) = 1000 cubic centimeters (cm^3)

Table 1.3: SI Base Units

  • This table comprehensively details the seven fundamental physical quantities, their respective SI base units, and their standardized symbols. It serves as a concise reference for the foundational measurements in science.

SI Prefixes

  • SI prefixes are standardized multipliers that can be attached to SI base units (and derived units) to indicate decimal multiples or submultiples of a unit. They simplify the expression of very large or very small measurements.

  • For example, 'kilo-' (k) represents 10^3 (one thousand), so 1 kilogram is 1000 grams. 'Milli-' (m) represents 10^{-3} (one-thousandth), so 1 millimeter is 0.001 meter.

Table 1.4: SI Prefixes

  • This table provides a list of common SI prefixes, their corresponding symbols, and their magnitudes (powers of 10). It is an essential tool for understanding the scale of measurements in scientific contexts.

SI Derived Units for Area and Volume

  • Derived units are combinations of two or more SI base units through multiplication or division, used to measure complex quantities.

Definitions:

  • Area is a measure of a two-dimensional surface and is expressed in square units (e.g., m^2, cm^2). It is calculated by multiplying two length measurements, such as length imes width.

  • Volume is a measure of the three-dimensional space occupied by a substance and is expressed in cubic units (e.g., m^3, cm^3). It is calculated by multiplying three length measurements, such as length imes width imes height. Other derived units include density (mass/volume) and speed (distance/time).

Table 1.5: SI Derived Units

  • This table concisely outlines several important derived units (e.g., for force, pressure, energy, power) and their associations with the fundamental SI base units from which they are constructed. It demonstrates how complex measurements are built from simpler ones.

Section 1.6: Significant Digits

Measurement Principles

  • Precise and reliable scientific measurements fundamentally depend on careful execution and often involve repetitions to determine accurate values. All measurements have a degree of uncertainty, meaning the last digit is estimated.

Definitions:

  • Accuracy: Refers to how close a measured value is to the true or accepted actual value. High accuracy implies minimal systematic error.

  • Precision: Refers to the reproducibility of a series of measurements, or how close multiple measurements of the same quantity are to one another. High precision implies minimal random error.

Rules for Significant Digits

  • Significant digits (or significant figures) are all the digits in a measurement that are known with certainty plus one estimated digit. They communicate the precision of a measurement.

  • Counted as significant:

    • All nonzero digits (1-9) are always significant (e.g., 23.45 has 4 sig figs).

    • Zeros located between nonzero digits (sandwich zeros) are significant (e.g., 1005 has 4 sig figs, 10.02 has 4 sig figs).

    • Trailing zeros (at the end of the number) are significant only if the number contains a decimal point (e.g., 12.00 has 4 sig figs, 120. has 3 sig figs). If there's no decimal, trailing zeros are often not significant unless explicitly marked (e.g., 120 has 2 sig figs).

    • Exact numbers (from counting or definitions) have an infinite number of significant figures.

  • Not counted as significant:

    • Leading zeros (zeros to the left of all nonzero digits) are never significant; they only indicate the position of the decimal point (e.g., 0.0025 has 2 sig figs).

Section Review of Significant Digits

  • This section reviews the critical practices for making and reporting scientific measurements, including the proper identification and calculation of significant digits, which are essential for reflecting the reliability and precision of experimental data in arithmetic operations (addition/subtraction and multiplication/division).

Section 1.7: Dimensional Analysis

Overview

  • Dimensional analysis, also known as the factor-label method, is a powerful problem-solving technique in chemistry and physics that treats units as algebraic quantities. It uses conversion factors (ratios of equivalent units) to convert quantities from one unit to another. This systematic approach ensures that calculations are set up correctly, as units must cancel out to yield the desired units for the final answer.

  • For example, calculating wages based on time involves converting hours to minutes or vice versa using conversion factors like rac{60 \text{ minutes}}{1 \text{ hour}}. This method makes unit conversions clear and minimizes errors.

Section Review of Dimensional Analysis

  • This review emphasizes the core concept of canceling units, much like canceling variables in algebra. It highlights how this technique provides a robust framework for correctly setting up and solving problems involving unit conversions and various quantitative relationships, ensuring dimensional consistency.

Section 1.8: Density

Density Definition

  • Density is an intrinsic physical property of matter that quantifies the mass of a substance contained per unit of its volume. It is expressed by the formula:

    d = \frac{m}{V}

    where d is density, m is mass, and V is volume. Common units for density include grams per milliliter (g/mL) or grams per cubic centimeter (g/cm^3) for liquids and solids, and grams per liter (g/L) for gases.

  • Characteristics: Density is an intensive property, meaning it's independent of the amount of substance. This makes it useful for identifying unknown substances. Objects with lower density will float in fluids that are denser; conversely, objects with higher density will sink in less dense fluids, illustrating the principle of buoyancy.

Densities of Common Substances

  • This section provides a comparative understanding of densities, often including values for various common materials such as metals (e.g., gold: \approx 19.3 \text{ g/cm}^3, copper: \approx 8.96 \text{ g/cm}^3), water (approximately 1.00 \text{ g/cm}^3 at 4^ ext{o} ext{C}), and other chemicals. Such data is crucial for identification purposes, understanding material behavior, and practical applications like ship buoyancy or hot air balloons.

Section Review of Density

  • This summary recaps the fundamental principles of density calculation, its role as a key intensive property, and its widespread utility in identifying substances and predicting how materials will behave in various fluids.

Section 1.9: Temperature Scales

Temperature Representation

  • Temperature is a measure of the average kinetic energy of the particles within a substance. Several scales are used to represent temperature:

    • Fahrenheit (^ ext{o} ext{F}): Water freezes at 32^ ext{o} ext{F} and boils at 212^ ext{o} ext{F}.

    • Celsius (^ ext{o} ext{C}): Water freezes at 0^ ext{o} ext{C} and boils at 100^ ext{o} ext{C}, making it a more convenient scale for scientific work due to its 100-degree interval between phase changes.

    • Kelvin (K): The absolute temperature scale where 0 K (absolute zero) represents the theoretical point at which all atomic motion ceases. Water freezes at 273.15 K and boils at 373.15 K. There are no negative temperatures on the Kelvin scale, and its units are the same size as Celsius degrees.

Conversion between Scales

  • Precise conversion equations are essential to facilitate accurate temperature changes among different scales:

    • Celsius to Fahrenheit: ^ ext{o} ext{F} = (^ ext{o} ext{C} \times \frac{9}{5}) + 32

    • Fahrenheit to Celsius: ^ ext{o} ext{C} = (^ ext{o} ext{F} - 32) \times \frac{5}{9}

    • Celsius to Kelvin: K = ^ ext{o} ext{C} + 273.15

    • Kelvin to Celsius: ^ ext{o} ext{C} = K - 273.15

Example 1.31

  • This example demonstrates the practical application of temperature conversion calculations. To convert normal human body temperature from Fahrenheit to Celsius:

    Given: 98.6^ ext{o} ext{F}

    Applying the formula: ^ ext{o} ext{C} = (98.6 - 32) \times \frac{5}{9}

    ^ ext{o} ext{C} = (66.6) \times \frac{5}{9}

    ^ ext{o} ext{C} = 37.0^ ext{o} ext{C}

    This shows that 98.6^ ext{o} ext{F} is equivalent to 37.0^ ext{o} ext{C}.

Section Review of Temperature Scales

  • This section provides a concise summary of the different temperature scales (Fahrenheit, Celsius, Kelvin), their respective reference points, and the necessary conversion processes. Understanding these scales and conversions is fundamental for accurate scientific measurement and communication.