Chapters 4 and 5 The Structure of the Atom And Electrons in Atoms

Chapters 4 and 5

The Structure of the AtomAnd

Electrons in Atoms

Early Theories of Matter

Democritus (460-370 B.C.) Named atom (atomos)

Early Theories of Matter

Aristotle (384-322 B.C.)

Early Theories of Matter

John Dalton (1766-1844) First Atomic Theory

Defining an Atom The smallest

particle of an element that retains the properties of the element.

About 1 X 10-10 m in diameter.

Can be seen with a scanning tunneling microscope.

Discovering the Electron

William Crookes (1800’s)

Discovering the Electron J.J. Thomson (late 1890’s)

Determined the charge-to-mass ratio

Mass must be less than a hydrogen atom

Plum Pudding Model of atom

Discovering the Electron

Robert Millikan (1909) Determined charge

of electron 1/1840 mass of a hydrogen atom

The Nuclear Atom

Ernest Rutherford (1911)

The Nuclear Atom Atom contains:

Mostly empty space

Tiny, dense nucleus which is positively charged

Creates nuclear model of atom

Other Subatomic Particles

Rutherford (1920) Concluded nucleus contains proton Proton as equal but opposite charge

of electron James Chadwick (1932)

Discovered neutron Neutron has no charge

Subatomic Particles

How Atoms Differ

Moseley (shortly after Gold Foil) Atoms of each element contain a

unique number of protons Atomic Number= #protons

Identifies the atom

Isotopes Isotopes – atoms that contain the

same number of protons but different number of neutrons.

Most elements contain a mixture of isotopes.

The relative abundance of each isotope is constant.

Isotopes

Mass Number = #protons + #neutrons

Simple Practice

AtomicNumber

MassNumber

# of Protons

# of neutrons

# of electrons

Mg 25 12

Zn 30 35

Be 4 9

Hg 120 80

12 13 12

65 30 30

4 5 4

80 200 80

Mass of Atoms Atomic mass unit – 1/12 of a carbon-

12 atom. Atomic Mass – weighted average

mass of the isotopes of that element.

Calculating Atomic Masses

6X has mass of 6.015 amu and abundance of 7.50%. 7X has mass of 7.016 amu and abundance of 92.5%.

(6.015)(.0750) + (7.016)(.925) = 6.94 amu

More Challenging Problems!

Cu-63 has a mass of 62.940 amu and an abundance of 69.17%. Find the mass and abundance of the other isotope.

Boron has two isotopes with the masses of 10.013 amu and 11.009 amu. Find the abundance of each isotope.

Radioactivity Nuclear Reactions – changes an atom’s

nucleus. Atom changes into a new element Due to unstable nuclei

Radiation contains rays and particles emitted from a radioactive material.

Radioactive decay is the spontaneous emission of radiation.

Types of Radiation

Types of Radiation

Radiation

Type

Symbol Mass (amu)

Charge

Alpha or 4 2+

Beta e- or 1/1840 1-

Gamma 00 0 0

He2

4

1

0

Nuclear Reactions

Mass numbers and Atomic numbers on both sides of the reaction must be equal

Practice Problem:

_______1

0

6

14

C

Chapter 5

Electrons in Atoms

Electromagnetic Radiation

Electromagnetic Radiation is a form of energy that has wave-like behavior.

4 properties of waves: wavelength, amplitude, speed and frequency.

Properties of Waves Frequency()- number of waves that pass a

given point per second. (hertz or 1/s or s-1) Speed (c)- is constant for all waves. 3 x 108

m/s

Calculating Properties of Waves c= What is the frequency of light with a

wavelength of 5.80 x 10-7 m?

A radio station broadcasts with a frequency of 104.3 MHz. What is the wavelength of the broadcast?

Particle Nature of Light Max Planck (1900) discovered that

matter can gain or lose energy in small, specific amounts called quanta.

Equantum= h Planck’s Constant (h)=6.626 x 10-34J·s

Practice Problems

What is the energy of a wave with a frequency of 6.25 x 1019Hz?

What is the frequency of a wave that contains 8.64 x 10-18J of energy?

A wave contains 4.62 x 10-15J of energy. Determine its wavelength.

Photoelectric Effect Photoelectric effect – electrons are emitted

from a metal’s surface when light of a certain frequency shines on it.

Frequency (color) of light, not brightness of light determines if electrons are emitted.

Einstein (1905)- light has wave-like properties but is also a stream of tiny particles or bundles of energy called photons.

Photon – a piece of EM with no mass and carries a quantum of energy.

Atomic Emission Spectrum When atoms absorb energy they

become excited. Atomic Emission Spectrum- unique set of

frequencies emitted by excited atoms.

Bohr Model of the Atom Bohr (1913) proposed why the emission

spectrum of hydrogen is not continuous. Electrons can have only certain “energy

states” Ground State - the lowest allowable

energy state. Excited State – energy state of an

electron when it gains energy

Bohr Model of the Atom

Electrons as Waves Louis de Broglie (1924) thought

Bohr’s model had electrons having similar properties to waves.

de Broglie equation:

m

h

Predicts that all moving particles have wave properties.

Heisenberg Uncertainty Principle When viewing an

electron, a photon of light hits it and changes the velocity and position of the electron.

It is impossible to know precisely both the velocity and position of a particle at the same time.

Quantum Mechanical Model of the Atom Schrödinger (1926) derived

an equation that treated hydrogen’s electron as a wave.

Allows electron to have only certain energy but does not give path of electron.

Atomic orbital – a 3-D region around the nucleus in which the electron can be found 90% of the time.