Chemistry - Atomic Theory and Structure + Light and Electrons

Evidence for atoms

• Law Conservation of Mass (Antoine Lavoisier) - Total mass remains constant during chemical reaction

• Law of Constant Composition (Joseph Proust) - A pure compound always contains the same ratio of elements by mass

• Law of Multiple Proportions (John Dalton) - For 2 compounds with the same 2 elements, the ratio of mass ratios is a small whole number. Mass ratios leading to whole numbers leads to the idea that elements are made of whole bits (atoms)

Early Models of Atomic Structure

• Dalton’s Indivisible Atom

  • Early 1800’s, based on his meteorological hobby

• Thomsen’s Plum Pudding model

  • 1897, based on work with Cathode Ray Tubes (CRT’s)

• Rutherford’s Nuclear model

  • 1911, based on results of Gold Foil experiment

Dalton’s Indivisible Atom

• “Atomos” is Greek for indivisible

• Different properties of atoms due to different sizes and shapes

• Based on decades of meteorological study

Plum pudding model

• Based off Thomson’s experiment with Cathode Rays

• Idea that protons and electrons were scattered in an atom

Rutherford’s Nuclear model

• Atom has a positively charged nucleus that is surrounded by electrons

• Alpha particles passing far from the nucleus are slightly deflected while alpha particles directly surrounding the nucleus are deflected at large angles

Atoms

The smallest particle of an element that still retains its properties

Which particles account for most of an atom’s mass?

Neutrons and protons

Cathode-ray tube experiment - Thompsen

• Discovery of electrons which changed Dalton’s Individual Atom

• Sending a stream of unknown particles, or cathode rays, from a negative electrode (cathode) to a positive electrode (anode) through an almost airless glass tube

• Result #1: Cathode Ray beam always flowed from neg. cathode toward pos. anode, and deflected toward the positive plate.

• Conclusion #1: The Cathode Ray beam is composed of negatively charged particles.

• Result #2: Got the same beam no matter what metal was used for electrodes

• Conclusion #2: All atoms contained these negative particles. Also, since atoms are neutral, they must also contain a positive piece.

• Result #3: He could never create an “Anode Ray beam” of positive particles.

• Conclusion #3: The positive part of the atom was very massive and immobile.

• Mass to change ratio then used it to figure out that charged particles are much less than hydrogen atoms, the lightest known atom

How was an electric field used to determine the charge of a cathode ray?

Because the cathode ray was deflected toward the positively charged plate by an electric field, the particles in the ray must have a negative charge.

Gold foil experiment - Rutherford

• Discovery of nucleus

• He struck gold foil with alpha particles and noted where flashes occurred

• Result #1: Most alphas shot straight through or were slightly deflected

• Conclusion #1: The atom is mostly empty space.

• Result #2: A very small percentage of alphas was completely reflected, which was where the nucleus was.

• Conclusion #2: The nucleus is very small and very dense and has protons

• Conclusion #3: The electron must be orbiting the nucleus and are held within the atom because of their attraction to the protons in the nucleus

Discovery of neutrons

Coworker of Rutherford, James Chadwick

Atoms

• Made up of electrons, protons, and neutrons

• Spherically shaped with a nucleus of positive surrounded by electrons

• Electrons travel through empty space surrounding nucleus

• Electrons held within atom because of attraction to proton

• Nucleus is most of atom’s mass

• Number of protons is equal number of electrons because atoms are electrically neutral

• Increasing mass - neutron, proton, electron

Chemical name 

• Name of the element

• Ex: Hydrogen

Chemical symbol

• Letter of element

• Ex: H

Atomic number

Number of protons and number of electrons

Atomic mass

• Entire mass of the atom

• Neutron, protons, electrons 

Mass number

Number of protons and neutrons

Atomic mass unit (amu)

• 1/12 of the mass of a carbon-12 atom

• It’s nearly equal to the mass of a single proton or a single neutron

• Makes it easier for chemists to measure the mass of the atom

• To find take atomic mass times percent abundance

Isotopes

Atoms with the same number of protons but not neutrons

What do the superscript and subscript in the notation 90/40 K represent?

90 represents the mass number and 40 represents the atomic number

Ions

• The number of electrons in any given element is also not constant.

• Has an electrical charge because it has a different number of protons and electrons

What is an electrons chemical behavior related to?

Arrangement of electrons in its atom

Electromagnetic radiation

A form of every that exhibits wavelike behavior as it travels through space

Properties of light waves or electromagnetic radiation

• Wavelength

  • The shortest distance between two equivalent points (crest to crest) on a continuous wave

  • Usually measured in meters, nanometers, or centimeters.

  • Expressed as λ

• Frequency

  • The number of waves that pass per second

  • Hz, unit of frequency, equals one wave per second

  • Expressed as v

• Amplitude

  • The wave’s height from the origin to a trough

  • Wavelength and frequency do not affect the amplitude of the wave

• Speed of light

  • All electromagnetic waves, including visible light, travel at a speed of 2.998 x 10 to the power of 8 m/s

  • Expressed as c

  • Product of wavelength and its frequency

Visible light

White light passing through a prism separates into a continuous spectrum of colors of red, orange, yellow, green, blue, indigo, and violet

Electromagnetic spectrum

All forms of electromagnetic radiation with only differences being in the types of radiation being their frequencies and wavelengths

Quantum Concept

• Matter can gain or lose energy only in discrete, small amounts called quanta

• Explains how light emitted by heated objects change color (thus frequency and wavelength) as their temperature and kinetic energy increase

Quantum

The minimum amount of energy that can be gained or lost by an atom

Ground state

The lowest allowable energy state of an atom

The energy of light

• Elight = hv

• h = Planck’s constant = 6.626 x 10-34 J*s

• v = frequency

• Energy increases as frequency increases

The energy of a light wave is proportional to its frequency

• High frequency = short wavelength = high energy

• Low frequency = long wavelength = low energy

• High energy = short wave length = high frequency 

• Radio to microwave to infrared to UV to x-rays to gamma rays

How did Einstein explain the photoelectric effect?

• Photoelectric - the emission of electrons from a material's surface when light of a high enough frequency shines on it

• Light has a dual nature. A beam of light has wavelike and particlelike properties can

• Said that it can be thought of as a beam of as a beam of bundles of energy called photons (massless particle that carries a quantum of energy)

• Proposed that the energy of a photon must have a certain threshold value to cause the ejection of photoelectron from the surface of the metal

We can add energy to atoms

• Heat, light, electrical, …

• Electrons in the atom absorb that energy

• They move farther from nucleus, “orbit” faster

Electrons do NOT like to be in this excited state (addition of energy), so …

• The excited electron drops back to its lower energy ground state

• When it does, it emits energy, ALWAYS IN THE FORM OF LIGHT

• We can calculate the energy of that light (E = hλ)

• This E corresponds to the difference in energy between the ground state and the excited state.

Electrons can only have certain energies/speeds

Their energies are discontinuous

Heisenberg Uncertainty Principle

The position and momentum of a particle cannot be determined simultaneously.

Dalton’s Billiard Ball model

Describes atoms as solid, indivisible, and indestructible spheres, much like billiard balls

Thomson’s Plum Pudding model

• Based off Thomson’s experiment with Cathode Rays

• Idea that protons and electrons were scattered in an atom

Rutherford’s Nuclear model

• Atom has a positively charged nucleus that is surrounded by electrons

• Alpha particles passing far from the nucleus are slightly deflected while alpha particles directly surrounding the nucleus are deflected at large angles

Bohr’s Planetary model

Electrons orbiting a positively charged nucleus in fixed, circular paths called orbits, similar to planets orbiting the sun

What happens when an atom absorbs light, heat, or electrical energy?

Its e-’s jump to a higher energy level.

Electrons are limited to energy levels

Each level is divided into subshells (s < p < d < f

How is an atomic spectrum made?

• That light is passed through a prism (or diffraction grating) and is broken down into individual colors

• That creates a spectrum, a fingerprint for that atom

• Atomic spectrum - Each wavelength/frequency/energy of light corresponds to an e- transition

What happens when e- gets the chance to drop down energy?

• It emits light energy 

Energy light

• Energylight = h ∙ ν

• (h = Plank’s constant, ν = frequency)

How many orbitals does each subshell have?

s has 1 orbital, p has 3 orbitals, d has 5 orbitals, f has 7 orbitals

How many electrons can be in an orbital?

2

The Aufbau Principles

• Rule #1 – Lowest energy orbitals fill first (Aufbau Principle)

• Rule #2 – Electrons spread out when possible (Hund’s Rule)

• Rule #3 – Only 2 e-’s per orbital. Up and Down arrows are used to represent opposite spin of electrons in an orbital (Related to Pauli Exclusion Principle)