Physics Chapter on J.J. Thomson's Discoveries

Part 1: Development of Atomic Theory

• An atom is the smallest entity that retains an element's identity (the simplest form of substance)
• The atom is a solid sphere that cannot be divided into smaller particles or pieces (Democritus, 400 B.C.)
• The idea of atoms emerged during the time of Democritus.
• Greek word "atomos" means uncuttable

John Dalton's Atomic Theory (Laws)

• Law of Conservation of Mass: The total mass of reactants before a chemical reaction is exactly equal to the products after the reaction is completed.
• Law of Definite (or Constant) Proportions: The proportion by mass of the elements in a given compound is always the same.

John Dalton's Atomic Theory (Proven True)

• Elements are made of indivisible and indestructible particles called atoms.
• An element (like silver) consists of silver atoms.
• Atoms cannot be split or destroyed to form other atoms.
• This idea is supported by the Law of Conservation of Mass.
• Compounds are formed by joining atoms of two or more elements.
• The ratio of atoms combining to form a compound is always in a whole number ratio.
• Example: A water molecule has one oxygen atom and two hydrogen atoms (ratio 1:2)

John Dalton's Atomic Theory (Proven False)

• The atoms of a certain element are identical (same properties and appearance).
• Different elements have different atoms (different masses).
• Each atom of an element has a different mass from other elements.

Joseph John (J.J.) Thomson

• English physicist who conducted experiments on cathode ray tubes.
• Determined the deflection of cathode rays when exposed to an external magnetic field.
• Discovered electrons through research on cathode rays.

Thomson's Cathode Ray Experiment (1897)

• Cathode ray tube is made of glass. Most of the air has been sucked out of it.
• The negatively charged part (cathode) releases an invisible ray as electricity is transmitted through the tube.
• The cathode ray is attracted to the positively charged part (anode) of the tube.
• A coating on the glass or the tube causes it to produce a bright light.
• Cathode rays move toward the positive plate and are deflected away from the negative plate.
• Experiments with different cathodes gave the same set of particles and behavior.

Thomson's Plum Pudding Model (1904)

• From a spherical model, ideas on the atomic model shifted to a plum pudding model.
• The larger sphere is the atom.
• Smaller spheres embedded on the larger sphere represent the corpuscles (electrons) in the 'fluid'.

Additional Inputs from Thomson

• Calculated the charge-to-mass ratio of the electron.
• The ratio is 1.76 x 10⁸ coulombs (C)/gram.
• 1 coulomb (C) is the SI unit for measuring charge.
• 1 C = 6.24 x 10¹⁸ electrons

Robert Millikan

• American physicist and professor at the University of Chicago.
• Performed the famous oil drop experiment.
• Calculated the approximate mass of an electron using the mass-charge ratio.

Millikan's Oil Drop Experiment (1909)

• Charge of an electron = 1.602 x 10⁻¹⁹ C
• Charge-to-mass ratio = 1.76 x 10⁸ C/g
• Mass of an electron = 9.10 x 10⁻²⁸ g

Radioactivity

• Pioneers: Henri Becquerel, Pierre Curie, and Marie Curie
• Types of rays: Alpha (positive), Beta (negative), and Gamma (neutral)

Ernest Rutherford

• Physicist from New Zealand who conducted experiments on alpha rays.
• Considered the "Father of Nuclear Physics" and a Nobel Prize recipient.
• Famous for his Gold foil experiment.

Rutherford's Gold Foil Experiment (1910)

• Most alpha particles went straight through the gold foil.
• A few alpha rays showed scattering. These rays had a positive charge.
• Some particles bounced back, implying a small, dense, positively-charged nucleus in the atom.

Rutherford's Nuclear Model (1911)

• The atom is similar to a small solar system where electrons orbit a positively charged nucleus.

The Quantum Theory

• By the beginning of the 20th century, the wave model of light was universally accepted by scientists.

Max Planck

• Predicted accurately how the spectrum of radiation emitted by an object changes with its temperature.
• Any energy emitted or absorbed by an object is restricted to pieces of particular sizes (quantum).
• Energies absorbed or emitted by atoms are quantized.
• Discovered the Planck Constant, which is a proportionality between frequency (v) and energy.

The Planck Constant (1900) and Flame Test

• Energy = h x v, where:
  • h = Planck's constant = 6.626 x 10⁻³⁴ J•s
  • v = frequency

Albert Einstein

• Saw Planck's idea in a new way of thinking about light.
• Known for discovering the photoelectric effect.
• Recipient of the Nobel Prize for Physics.

Photoelectric Effect (1905)

• Electrons are ejected from the surface of a metal when light shines on the metal.
• For each metal, a minimum frequency of light is needed to release electrons.

Line Spectra

• A spectrum that contains only certain colors or wavelengths.
• Also known as atomic emission spectrum.
• Atomic fingerprint (unique to each element).

Neils Bohr

• Agreed with Rutherford's model of the atom but recognized its flaws.
• Attended Rutherford's lectures in 1911.
• Connected Planck's theory with Rutherford's planetary model.

Bohr's Solar System Model (1913)

• Emission from hydrogen atoms is restricted to certain frequencies.
• The lone electron can possess only specific energies corresponding to particular orbits around the nucleus.
• Allowed orbits and quantum leaps.

Louis de Broglie

• A French graduate student (Ph.D.).
• Wrote the shortest dissertation in history (3 pages).
• Used Einstein's famous mass-energy relationship (E = mc²) to explain matter waves.

de Broglie's Equation (1924)

• The relationship between the wavelength (λ) of an object, its mass (m), its velocity (v), and Planck's constant (h).
• λ= h/mv
• The larger the mass or velocity, the smaller the wavelength.
• Both light and electrons behave as waves and particles (wave-particle duality)
• Energy is quantized

Erwin Schrodinger

• Austrian physicist.
• Derived an equation that describes electrons as waves in three-dimensional space.
• Described what orbitals are.

Schrodinger's Electron Cloud Model (1926)

• An orbital is a well-defined region of three-dimensional space that can be inhabited by an electron.
• Orbitals come in different shapes.
• Dense regions of orbitals form the electron cloud model, the model used today.

Evolution of Atomic Models

• Summarizes the history of atomic models from solid sphere to quantum model.

Part 2: Atomic Structure

• Details the characteristics of protons, neutrons, and electrons.

Three Subatomic Particles

• Proton: Symbol (p⁺), Charge (1+), Mass (1.007 AMU, 1.673 x 10⁻²⁴ g), Location (nucleus)
• Neutron: Symbol (n⁰), Charge (0), Mass (1.009 AMU, 1.675 x 10⁻²⁴ g), Location (nucleus)
• Electron: Symbol (e⁻), Charge (1-), Mass (5.486 x 10⁻⁴ AMU, 9.109 x 10⁻²⁸ g), Location (outside the nucleus)

Atomic Number vs. Mass Number

• Atomic Number (Z): Denotes the number of protons within an element's nucleus. Z is unique for each element.
• Mass Number (A): Amount of neutrons and protons found in an atom's nucleus. It is roughly the same as the total mass of protons and neutrons (the mass of the electrons is negligible)

Symbols

• A = Mass number (proton + neutron)
• Z = Atomic number (proton)
• Number of neutrons = mass number - atomic number (A-Z)

Example 1

• Explanation of atomic number (protons) and mass number (protons + neutrons) for the element Sodium (23Na¹¹). This example shows how to find the number of protons and neutrons from the atomic symbol and explains that the number of protons is equal to the number of electrons.

Example 2

• Explanation of how to determine the number of protons and neutrons for Chlorine (³⁵Cl¹⁷). This example shows how to find the number of protons and neutrons from the atomic symbol and explains that the number of protons is equal to the number of electrons.

What are Ions?

• A charged atom is either positively or negatively charged.
• When an electron is lost, the atom has more protons than electrons (cation). Example: Na⁺
• When an atom gains electrons, it has more electrons than protons (anion). Example: Cl⁻

Part 3: Isotopes, Isotones, and Isobars

• Isotopes: Atoms of the same element with different atomic masses (same number of protons, different numbers of neutrons). Example: Protium, deuterium, and tritium.
• Isotones: Atoms with different atomic masses and atomic numbers but have the same number of neutrons.
• Isobars: Atoms with the same mass number but different atomic numbers.

Element's Atomic Weight (Formula)

• Atomic weight = ∑[(isotope mass) x (fractional isotope abundance over all isotopes of the element)]

Example 1 (Solution)

• Detailed steps calculating the average atomic mass of carbon, given the masses and abundances of its isotopes. Demonstrates how to compute the average mass when fractional abundances are provided.

Example 2 (Solution)

• Detailed steps calculating the average atomic mass of chlorine, given the masses and abundances of its isotopes. This example provides a solution demonstrating how to compute the average atomic mass given the masses and fractional abundances of isotopes.

Example 3

• Calculation of boron's average atomic mass, utilizing the masses and fractional abundances of its isotopes to demonstrate how to compute the average mass when fractional abundances are provided.

Part 4: Properties and Changes in Matter

• Properties: Observable characteristics of a substance that give it a unique identity (physical or chemical). Physical properties include extensive (dependent on the amount of substance—e.g., mass, volume) and intensive (not dependent on the amount of substance—e.g., density, color). Chemical properties describe how a substance changes its identity during a reaction.

Physical Properties (Examples)

• Includes various examples of physical properties such as temperature, structure, color, taste, odor, boiling point, freezing point, hardness, conductivity, solubility, and density.

Chemical Properties (Examples)

• Examples of chemical properties, including the reaction of Copper with Nitric Acid, illustrating changes in substance identity with chemical reactions.

Physical and Chemical Properties (Copper)

• Physical properties: Malleable and ductile, reddish-brown metallic luster, melting point of 1083 °C, boiling point of 2570 °C.
• Chemical properties: Slow formation of a blue-green carbonate in moist air, reaction with acids (nitric or sulfuric), and slow formation of a deep blue solution in aqueous ammonia

Extensive vs. Intensive Physical Properties

• Extensive properties depend on the amount of matter (mass, volume, length).
• Intensive properties do not depend on the amount of matter (density, melting point, boiling point, color).

Extensive vs. Intensive Properties (Example)

• Comparison of extensive (mass, volume) and intensive (color, melting point) properties for sulfur crystals and sulfur powder.

Physical vs. Chemical Changes

• Physical change: Changes in physical appearance or phase without altering the substance's composition (e.g. melting, freezing, condensation, decantation).
• Chemical change: Changes that result in the formation of new substances with different properties and composition (e.g., burning of paper, reaction of acids and bases).

Part 5: Classification of Matter

• Matter is classified as mixtures or pure substances.
• Mixtures are combined physically and can be separated (e.g., heterogeneous, homogenous (solutions)).
• Pure substances are chemically combined with fixed proportions and are elements or compounds (identified using chemical formulas).

Pure Substances

• A substance with a definite or fixed composition. Examples include elements and compounds.
• Elements: Simplest form of a substance (e.g., oxygen, iron).
• Compounds: Combinations of two or more elements in definite proportions (e.g., water, salt).

Pure Substances (Elements)

• A pure substance composed of only one kind of atom.
• Defined by its atomic number (number of protons).
• The simplest form of a substance.

Elements (Metals)

• Typically solids at room temperature.
• Have shiny or metallic luster.
• Can conduct heat and electricity.
• Malleable and ductile.
• Usually sonorous.
• Have high density.
• Have high melting points.

Elements (Non-Metals)

• Usually solids or gases at room temperature.
• Have dull luster.
• Poor conductors of heat and electricity.
• Brittle materials.
• Non-sonorous.
• Have low density.
• Have low melting points.

Elements (Metalloids)

• Share characteristics of both metals and nonmetals.
• Found along the zig-zag line in the periodic table.
• Commonly used as semiconductors.

Pure Substances (Compounds)

• A pure substance composed of two or more elements chemically combined.
• The elements in a compound are in fixed proportions.
• Compound components cannot be separated by ordinary physical means (e.g. water, table salt, baking soda, household bleach, milk of magnesia).

Compounds (Examples)

• Includes examples of compounds (water, table sugar, baking soda, table salt, household bleach, and milk of magnesia).

Mixtures

• Matter with variable composition.
• A blend of two or more pure substances.
• Physical properties of constituents are retained in a mixture.

Homogenous vs. Heterogeneous Mixtures

• Homogeneous mixtures (solutions): Do not have visibly different parts (e.g., lemonade, gasoline, steel), composed of a solvent and a solute, and demonstrate solubility or miscibility.
• Heterogeneous mixtures: Have visibly different parts (e.g., dirt, blood, milk).

Types of Solution Systems / Kinds of Mixtures

• Gas-gas, Gas-liquid, Liquid-liquid, Solid-liquid, Solid-solid
• Examples: Air, air in scuba tanks, seawater.
• Suspension, colloid, and solution, categorized by particle size (nanometers), and separation characteristics (e.g., settling, transparency).

Tyndall Effect

• Light scattering in colloidal dispersion.
• Used to determine whether a mixture is a true solution or a colloid (light scattering).

Mixture Separation Technique (Filtration)

• Used to separate an insoluble solid component from a mixture using a filter. Filter openings are smaller than the size of solid particles.

Mixture Separation Technique (Decantation)

• Separating liquid components from settled solids.

Mixture Separation Technique (Evaporation)

• Separating a solid from a liquid mixture by evaporating the liquid component through heating.

Mixture Separation Technique (Distillation)

• Separating liquid components of a mixture based on differing boiling points. Liquid with a lower boiling point evaporates first.

Mixture Separation Technique (Flotation)

• Separating components of a mixture based on density differences using air supply.

Mixture Separation Technique (Magnetism)

• Separating magnetic metal components from a non-magnetic mixture.

Mixture Separation Technique (Centrifugation)

• Separating components of a mixture based on density differences using centrifugation.

Mixture Separation Technique (Chromatography)

• Separating components of a mixture based on their relative movement through a mobile and a stationary phase.

Part 6: Acids, Bases, and Salts

• Acids: Substances that release hydrogen ions (H⁺) when dissolved in water.
• Bases: Substances that release hydroxide ions (OH⁻) when dissolved in water.
• Salts: Products of the neutralization reaction between an acid and a base.

Properties of Acids

• Form H⁺ ions when dissolved in water.
• Most acid formulas start with "H" (e.g., H₂CO₃, HCl, CH₃COOH).
• React with active metals to form hydrogen gas.
• Taste sour (but not all acids are safe to taste).
• Turn blue litmus paper red, and bromothymol blue to yellow.
• Corrosive materials.
• pH value less than 7.

Properties of Bases

• Form OH⁻ ions when dissolved in water.
• Most base formulas end with "OH" (e.g., Mg(OH)₂, Al(OH)₃, NaOH).
• Not reactive to metals.
• Taste bitter (but not all bases are safe to taste).
• Turn red litmus paper blue and bromothymol blue remains blue.
• Feel slippery when in contact with skin.
• pH value greater than 7.

pH Scale

• Developed by Soren Sorensen.
• A 14-point scale that measures the concentration of hydrogen ions (H⁺) in a solution.
• pH = -log[H⁺]
• Midpoint (pH 7) is neutral (pure water).
• Lower than 7 is acidic; higher than 7 is basic.

pOH Scale

• Counterpart to the pH scale for hydroxide ions (OH⁻)
• pOH = -log[OH⁻]
• A 14 point scale, Lower than 7 is basic; higher than 7 is acidic.

Neutralization

• Chemical reaction between an acid and a base.
• A salt and water are the products of neutralization.
• Salt: A compound formed from a cation and an anion (examples include metal and non-metal ions).

Indicators

• Substances that change color when added to an acid or a base.
• Colors of indicators can indicate if a substance is an acid or base (e.g., red litmus, blue litmus, bromothymol blue, phenolphthalein, methyl red, red cabbage).
• Indicators have a pH range (examples provided).

Part 7: Forces

• Force: Push or a pull (SI unit = Newton).
• Vector quantity (magnitude and direction).
• Contact force: Interaction of two objects in direct physical contact.
• Examples include normal force (perpendicular to the surface), applied force, tension (forces directed toward ropes, etc), spring force (Hooke's Law: F = -kx), air resistance, and friction.
• Action-at-a-distance force: Interaction of two objects without physical contact. Examples include gravitational force, magnetic force, and electric force.

Part 8: Newton's Laws of Motion

• Newton's First Law (Law of Inertia): An object at rest stays at rest, and an object in motion stays in motion with the same speed and in the same direction unless acted upon by an unbalanced force.
• Newton's Second Law (Law of Acceleration): The acceleration of an object depends on the net force acting on the object and the mass of the object. Fnet = ma.
• Newton's Third Law (Law of Interaction): For every action, there is an equal and opposite reaction.

Part 9: Fluids

• Fluids: Substances that flow and continuously deform under shear force (liquids and gases).
• Density: Mass per unit volume (ρ=m/V).
• Specific gravity: Ratio of the density of the substance to the density of water at 4°C.

Density (Exercises)

• Exercises demonstrating calculation of density related to mass and volume, using the formula (ρ = m/V) for various objects (ball, fluid, rectangular box, block of aluminum).

Triple Beam Balance (TBB)

• Instrument used precisely measure the mass.
• Tips for using it: Place object on pan; adjust riders on beams to balance the pointer, read the values of the riders to get the total mass of the object.

Density of Solids

• Calculating density of regular and irregular solids.

Density of Liquids

• Calculating density of liquids by measuring the mass of the graduated cylinder + liquid; with the initial mass, and final masses, then deriving the volume to calculate the density.

Pressure

• Pressure is a scalar quantity that measures the amount of force per unit area (P= F/A).
• Standard unit for pressure is Pascals (Pa).
• Fluids exert pressure.

Pressure (Exercises)

• Calculating pressure using the formula P = F/A, including examples related to different situations (person standing, a box, and a performer balancing on a cane).

Science Instruments Involving Pressure

• Mercury barometer: Measures atmospheric pressure (height of the mercury column).
• Sphygmomanometer: Measures blood pressure (force exerted by blood on arterial walls).

Systolic vs. Diastolic Blood Pressure

• Systolic pressure: Maximum pressure when heart beats.
• Diastolic pressure: Pressure when heart is resting between beats.

Fluid Pressure

• Pressure = (density) (gravitational constant) (height/depth).

Pressure (Exercise)

• Calculations involving pressure in a water tower and the comparison of dam strength based on the height and width of the dam.

End of Integrated Science 2 Quarter 2 Long Test Reviewer

Test your knowledge on the fundamental discoveries made by J.J. Thomson, focusing on the charge of the particles he identified. This quiz will challenge your understanding of atomic structure and early physics concepts. Dive in to see how much you know about this pivotal moment in scientific history!