unit 4: atomic theory structure of the atom (& radioactivity)

Unit 4: Atomic Theory

Structure of the Atom (& Radioactivity)

Early Atomic TheoriesModels of the Atom

Date scientist discovery________________ 100 BC Democritus/Greeks concept of the atom 1770 Antoine Lavosier Law of conservation of mass 1800 Joseph Proust Law of definite proportions 1803 John Dalton Law of multiple proportions Atomic Model I 1880 William Crookes Cathode Rays (electrons) 1885 Goldstein Canal Rays (protons) 1900 J.J. Thomson Plum Pudding Model

Electron Atomic Model II

1909 Ernest Rutherford nucleus of atom Atomic Model III 1913 Niels Bohr Planetary Model Atomic Model IV 1920- Schroedinger/Planck/ Modern or Wave Model Present DeBroglie/Einstein/etc. Atomic Model V

Early Atomic Theories Atomists and Democritus

Greeks approx. 2,500 years ago Matter was made up of atoms

“atomos” or “Indivisible” particles Seashell experiment—broken into

smaller & smaller pieces

Early Atomic Theories John Dalton

1766-1844; returned to theory of atoms Atoms are like billiard balls (solid spheres)

which can’t be broken down further 4 major postulates

1. All elements are composed of atoms2. Atoms of the same element are identical3. Atoms can physically mix or chemically

combine in simple whole number ratios4. Reactions occur when atoms separate, join,

or rearrange

Dalton’s Model of the Atom

No subatomic particles!

Early Atomic Theories

William Crookes Developed Crookes tube (a.k.a cathode

ray tube) in 1870’s First evidence for existence of electrons

because you could “see” electrons flow and confirm their existence.

Tube is precursor to today’s TV picture tubes

Building the Atom – The Electron J.J. Thomson

Discovered electron in 1897 Discovered positively charged particles

surrounded by electrons Found the ratio of the charge of an

electron to its mass to be 1/1837

Thomson’s Cathode Ray Tube Experiment J.J. Thomson

Video: Cathode Ray Tube Demo

Building the Atom – The Electron J.J. Thomson

Cathode ray tube experiments – advancement of Crookes tube

“plum-pudding model”

Thomson’s Model of the Atom (Plum Pudding Model)

Millikan’s Oil Drop Experiment Robert Millikan

Oil drop experiment Determined the charge and mass of an

electron

Video: Millikan's Oil Drop Experiment

Building the Atom – The Nucleus Ernest Rutherford

Discovered nucleus (dense core of atom) in 1911

Famous Gold foil experiment

Quote from E.R.’s Lab Notebook “It is about as incredible as if you had fired

a 15-inch shell at a piece of tissue paper and it came back and hit you.” -ER

Rutherford’s Gold Foil Experiment

Video Clip: Rutherford Gold Foil Experiment

Rutherford’s Model of the Atom

Building the Atom – The Neutron James Chadwick

Discovered the neutron (no charge, but same mass as proton)

Neutrons help disperse the strong repulsion of positive charges

Nucleus diameter = 10-5 nm Atom diameter = 10-1 nm If Nucleus = basketball -->

then Atom = 6 miles wide!

Building the Atom Niels Bohr

Improved on Rutherford’s work “Planetary model”- positive center is

surrounded by electrons in defined orbits circling the center

Bohr Model of the Atom (Planetary Model)

Bohr Model of the Atom Vocab. Energy level – the location where an

electron is found at a set distance from the nucleus dependent on the amount of energy it has

Ground state – the typical energy level where an electron is found; lowest energy

Excited state – an energy level higher than the ground state for an electron; temporary condition

Schrödinger Model (Quantum Mechanical Model) Quantum Mechanical Model

Erwin Schrödinger; Mathematical model Electron locations are based on probability Electrons are not particles, but waves!

Interactive Simulation: try it!

Defined: Orbital – region where an electron is likely to

be found 90% of the time

Schrödinger Model of the Atom(Quantum Mechanical Model)

Atomic TheoryAtom – the smallest particle of

matter that retains its properties. Smallest individual unit of an element

One atom of hydrogen is different from one atom of carbon.

Subatomic particles – the component parts of an atom: proton, neutron, and electron

Atomic Theory Ion - atom with the same number of

protons but a different number of electrons i.e. an atom with a charge!

If the atom has a (+) charge it has more protons than electrons.

If the atom has a (-) charge it has more electrons than protons.

Subatomic ParticlesSubatomic

ParticleMass and

Abbreviation Charge Location Discoverer

Protonp+

Mass =1 amu+1 Nucleus ----

Neutron nMass =1 amu 0 Nucleus Chadwick

in 1932

Electrone-

Mass ≈ zero amu

-1

Electron cloud

(outside nucleus)

Thomson in 1897

Atomic SymbolsAtomic symbol – the letter or

letters that represent an element.

13

Al

26.981

Atomic number

Atom symbol

Atomic mass or weight

Atomic NumberAtomic number = the number of protons in the nucleus.

(same for every atom of that element)

13Al

26.981

Atomic number

Atom symbol

Atomic mass or weight

Mass Number

Mass Number = # protons + # neutrons

A Boron atom can have:5 p + 5 n = 10 amu

Named as boron-10

 Mass number

Atomic number

Calculations w/ Subatomic ParticlesAtomic number = # of protonsMass number = # of protons + # of

neutrons

(For a neutral atom): # of protons = # of electrons

(For a charged ion): Charge = #p+ - #e-

Isotope Notation Isotope (Isotopic Notation)

Mass #

Atomic #

Atomic Symbol

Z

AX

Example: Uranium-238

Example  

Example  

Sample ProblemWrite the atomic symbols for the following:

The isotope of carbon with a mass of 13

The nuclear symbol when A = 92 and the number of neutrons = 146.

31

Isotopes Isotope – atoms of the same element with

different numbers of neutrons (different mass numbers) Example: Carbon-12 Carbon-14

Atomic mass – weighted average of the masses of all the isotopes of an element

Atomic Mass

The weighted average is the addition of the contributions from each isotope.

Isotopic Abundance is the percent or fraction of each isotope found in nature.

33

Most Abundant Isotope

34

13

Al

26.981

Atomic number

Atom symbol

Atomic mass or weight

Usually can round atomic mass on the periodic table to nearest whole number

(but not always!!)

Example: Determine the average atomic mass of magnesium which has three isotopes with the following masses: 23.98 amu (78.6%), 24.98 amu (10.1%), 25.98 amu (11.3%).

1) Multiply the mass number of the isotope by the

decimal value of the percent for that isotope.

2) Add the relative masses of all of the isotopes to

get the atomic mass of the element.

35

Example: Determine the average atomic mass of magnesium which has three isotopes with the following masses: 23.98 amu (78.6%), 24.98 amu (10.1%), 25.98 amu (11.3%).

36

Now You Try!If 90% of the beryllium in the world has a mass number of 9 and only 10% has a mass number of 10, what is the atomic mass of beryllium?

Radioactivity - Vocabulary

Radioactivity - the spontaneous emission of radiation from a substance

Radiation - rays and/or particles emitted from radioactive material

Nuclear reactions - reactions involving changes in an atom’s nucleus

Radioactivity Radioactive isotopes are unstable

These isotopes decay over time by emitting particles and are transformed into other elements

Particles emitted: Alpha (α) particles: helium nuclei Beta (β) particles: High speed electrons Gamma (γ) rays: high energy light

39

Types of Radiation – α particles

Alpha radiation - stream of high energy alpha particles Alpha particles consist of 2 protons and 2

neutrons and are identical to helium-4 nucleus.

Symbol: 4He 2+

2

Not much penetrating power, travel a few centimeters, stopped by paper, no health hazard

Types of Radiation – β particles

Beta radiation - High speed electrons To form beta radiation a neutron splits into

a proton and an electron The proton stays in nucleus and the

electron propels out at high speed.

Symbol: 0e- 0β

-1 -1

100 times more penetrating then alpha radiation, pass through clothing to damage skin

Types of Radiation – Radiation

Gamma radiation Similar to X rays Doesn’t consist of particles (instead, high

energy light) Symbol:

0 0

Penetrates deeply into solid material, body tissue, stopped by Pb or concrete, dangerous

End of Unit 4!