Defining Atoms & Electrons in Atoms Democritus (460-370 BC) Originated idea of the atom.

Defining Atoms

&Electron

s in Atoms

Democritus (460-370 BC)

Originated idea of the atom

John Dalton (1766 - 1844)

1803 Dalton’s Atomic Theory

Several changes have been made to Dalton’s theory.

Modern Atomic Theory

• Dalton said:Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties

• Modern theory states:Atoms of an element have a characteristic average mass which is unique to that element.

Modern Atomic Theory

Dalton said:Atoms cannot be subdivided, created, or destroyed

Modern theory states:Atoms cannot be subdivided, created, or destroyed in ordinary chemical reactions. However, these changes CAN occur in nuclear reactions!

J.J. Thomson (1856-1940)

1897 Discovered the electron(“plum pudding” model)

Mass of the Electron

1909 – Robert Millikan determines the mass of the electron.

The oil drop apparatus

Mass of the electron is 9.109 x 10-31 kg

Conclusions from the Study of Electrons

• Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons.

• Atoms are neutral, so there must be positive particles in the atom to balance the negative charge of the electrons.

• • Electrons have so little mass that atoms

must contain other particles that account for most of the mass

Rutherford (1871-1937)

1911 Discovered the nucleus(gold foil experiment)

Rutherford’s Findings• Most of the particles passed right

through • A few particles were deflected • VERY FEW were greatly deflected

The nucleus is small The nucleus is dense The nucleus is positively charged

“Like howitzer shells bouncing off of tissue paper!”

Conclusions:

Since opposite charges attract each other, why don’t the

electrons fall into the nucleus?

Niels Bohr (1885-1962)

1913 proposed Planetary model

Neils Bohr

I pictured electrons orbiting the nucleus much like planets orbiting the sun.

But I was wrong! They’re more like bees around a hive.

WRONG!!!

The Bohr Model of the Atom

Based upon the work of several men, a new mathematical model was developed to describe the

structure of the atom.

1926 The Quantum-Mechanical

Model

Louis de Broglie (1892-1987)Werner Heisenberg (1901-1976)Erwin Schrodinger (1887-1961)

Atomic Number• Atomic number of an element is the number

of protons in the nucleus of each atom of that element.

Element Atomic # # of protons

# of electrons

Carbon 6

Phosphorus 15

Gold 79

• In a neutral atom:# electrons = # protons

6

15

79

6

15

79

The Atomic Scale• Most of the mass of the atom is in the

nucleus (protons and neutrons)

• Electrons are found outside of the nucleus (the electron cloud)

• Most of the volume of the atom is the electron cloud.

Particle Charge Mass (g) Location

Electron e- -1

9.109 x 10-28

(1/1840 amu)

Electron cloud

Proton p+ +1

1.673 x 10-24

(1 amu)Nucleus

Neutronn 0

1.675 x 10-24

(1 amu)Nucleus

Atomic Particles

Reading the Periodic Table

Atomic Number 3

Li6.941Lithium

Element Symbol

Element Name

Atomic Mass

# p+ # e- (in a neutral atom)

# p+ + # n0

# n0 = Atomic Mass - #p+

The Quark…

Oops…wrong Quark!

The Atomic Scale• Most of the mass of the

atom is in the nucleus (protons and neutrons)

• Electrons are found outside of the nucleus (the electron cloud)

• Most of the volume of the atom is empty space

“q” is a particle called a “quark”

About Quarks…• Protons and neutrons are

NOT fundamental particles.

• Protons are made of two “up” quarks and one “down” quark.

• Neutrons are made of one “up” quark and two “down” quarks.

• Quarks are held together by “gluons”

Isotopes• Elements occur in

nature as mixtures of isotopes.

• Isotopes are atoms of the same element that differ in the number of neutrons

Mass Number

Mass # = p+ + n0

• Mass number is the number of protons and neutrons in the nucleus of an isotope.

Symbols of Isotopes

Carbon-12

Atomic Mass C126Atomic Number

146

Carbon-14

Atomic Mass CAtomic Number

3517

Atomic Number Chlorine-35Atomic Mass Cl

Mass Number

Mass # = protons + + neutons0

Element p+ n0 e- Mass #

Oxygen-16 8

33 42

-31 15

29 64

Neon-20 10

8 8 16

Arsenic -75 33 75

Phosphorus 15 3116

35 29Copper-64

10 10 20

Mass Number

Mass # = p+ + n0

Isotope p+ n0 e- Mass #

Oxygen - 10

- 33 42

-31 15

8 8 1818

Arsenic 75 33 75

Phosphorus 15 3116

O188

As7533

P3115

Isotope Protons

Electrons

Neutrons

Nucleus

Hydrogen–1

(protium)

1 1 0

Hydrogen-2

(deuterium)

1 1 1

Hydrogen-3

(tritium)

1 1 2

Isotopes…Again (must be on the test)

Isotopes are atoms of the same element having different masses due to varying numbers of neutrons.

Isotope SymbolComposition of

the nucleus % in nature

Carbon-12

12C6 protons 6 neutrons 98.89%

Carbon-13

13C6 protons 7 neutrons 1.11%

Carbon-14

14C6 protons 8 neutrons <0.01%

• Atomic mass is the weighted average of all the naturally occuring isotopes of that element.

• Multiply the mass of each isotope by its natural abundance, expressed as a decimal, and then add the products.

Carbon = 12.011

Atomic Masses

Isotope name Isotope mass

(amu) percentage

Silver-107 106.90509 51.86

Silver-109 108.90470 remainder

Silver has two naturally occurring isotopes:

Find the missing percentage.

Find the average atomic mass of an atom of silver.

Silicon has three naturally occurring isotopes:

Look over the data before you begin the problem. Estimate the value of the answer before you begin the calculation. Will the weighted average be closer to 28,  29, or 30?

Find the average atomic mass of silicon.

Isotope nameIsotope mass

(amu)Relative

AbundanceSilicon-28 27.98 92.21Silicon-29 28.98 4.70Silicon-30 29.97 3.09

Iron has four naturally occurring isotopes:

Estimate the average mass.

Find the average atomic mass of iron.

Isotope nameIsotope

abundanceIsotope mass

(amu)Iron-54 5.90% 53.94Iron-56 91.72% 55.93Iron-57 2.10% 56.94Iron-58 0.280% 57.93

The Periodic Table

• Period: horizontal rows of the periodic table.

• Group or Family: vertical columns of the periodic table. Elements within a group have similar chemical and physical properties.

Period(rows)

Group or Family(columns)

The discovery of the STM's ability to image variations in the density distribution of surface state electrons created in the artists a compulsion to have complete control of not only the atomic landscape, but the electronic landscape also. Here they have positioned 48 iron atoms into a circular ring in order to "corral" some surface state electrons and force them into "quantum" states of the circular structure. The ripples in the ring of atoms are the density distribution of a particular set of quantum states of the corral. The artists were delighted to discover that they could predict what goes on in the corral by solving the classic eigenvalue problem in quantum mechanics -- a particle in a hard-wall box.

Atomic Structure

• Where are the electrons?

• What are electron shells?

• How many electrons per shell?

Electrons are in shells that circles the nucleus at light speed.

Electron Shells The letter n, represents the electron shell.

Number of electrons that can fit in a shell:

2n2

Electron Shell 1 can hold:

2e-

Electron Shell 2 can hold:

8e-

Electron Shell 3 can hold:

18e-

Electron Shell 4 can hold: 32e-

Energy Level 1 can hold:

2e-

Energy Level 2 can hold:

8e-

Energy Level 3 can hold:

18e-

Energy Level 4 can hold:

32e-

Things to remember:• The element’s period # = the # of electron shells

• There can only be 2 e- in the first energy level

All atoms want to have 8 electrons on their outer shells

Valence Electrons

• Valence electrons – Electrons on highest energy level / highest electron shell.

COOL FACT about atoms!

• Most of the volume (space) of an atom is made up of electrons

• Electrons have very little mass and take up very little space

• SO, atoms are mainly empty space

• We are made of atoms SO we are empty mainly made up of Empty Space

Sulfur

Nucleus

# Neutrons:

# Protons:

# Electrons:

Orbital shapes are defined as the surface that contains 90% of the total electron probability.

An orbital is a region within an atom where thereis a probability of finding an electron. This is a probability diagram for the s orbital in the first energy level…

The s orbital has a spherical shape centered around the origin of the three axes in space.

s orbital shape

s sublevelshave 1 orbital

There are three dumbbell-shaped p orbitals in each energy level above n = 1, each assigned to its own axis (x, y and z) in space.

P orbital shape

Things get a bit more complicated with the five d orbitals that are found in the d sublevels beginning with n = 3. To remember the shapes, think of “double dumbells”

…and a “dumbell with a donut”!

d orbital shapes

Shape of f orbitals

http://micro.magnet.fsu.edu/electromag/java/atomicorbitals/

Orbital Notation

• Tell where electrons are arranged in s, p, d, and f sublevel orbitals in each level around the nuclei of atoms.

• Use boxes to represent orbitals• Use or to represent e-

Aufbau Principle

• Electrons occupy lowest energy orbitals first

Like This

Pauli Exclusion Principle

Only two electrons can occupy one orbital… and they must have opposite spin.

Wolfgang Pauli

Like This p

orbitals

Hund’s Rule

• One electron enters each orbital until all the orbitals contain one electron with the same spin direction…

…then they pair up.

3p4

Principal Energy Level

Sublevel

# of e-

Writing Electron Configurations

Describes e- location.

Orbital filling table

Electron configuration of the elements of the first three series

Irregular configurations of Cr and Cu

Chromium steals a 4s electron to half fill its 3d sublevel

Copper steals a 4s electron to FILL its 3d sublevel

Wave-Particle Duality

His son, George Thomson won the Nobel prize for describing the wave-like nature of the electron.

The electron is an energy

wave!

The electron is a particle!

JJ Thomson won the Nobel prize for describing the electron as a particle.

The Wave-like Electron

Louis deBroglie

The electron propagates through space as an

energy wave. To understand the atom, one

must understand the behavior of

electromagnetic waves.

The Electromagnetic Spectrum and Light

• Wavelength - The distance between two consecutive peaks of a wave.

• Frequency - The number of cycles in a certain period of time… measured in cycles per second, or Hertz (Hz).

• 1Hz = 1/sec = 1 sec -1

c = C = speed of light, a constant (3.00 x 108 m/s)

= frequency, in units of hertz (hz, sec-1)

= wavelength, in meters

Electromagnetic radiation propagates through space as a wave moving at the speed of light.

Types of electromagnetic radiation:

Ultraviolet Rays

E = hE = Energy, in units of Joules (kg·m2/s2)

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

= frequency, in units of hertz (hz, sec-

1)

The energy (E ) of electromagnetic radiation is directly proportional to the frequency () of the radiation.

Long Wavelength

= Low Frequency

= Low ENERGY

Short Wavelength

= High Frequency

= High ENERGY

Wavelength Table

You broke your big toe!  The x ray they take of your toe uses waves that have a length 2.19 x 10-

10m.     1. What is the speed of the wave in m/s?

2. What is the frequency of the x ray?

3. What is the Energy of the photons?

…produces all of the colors in a continuous spectrum

Spectroscopic analysis of the visible spectrum…

Refraction of White Light

…produces a “bright line” spectrum

Spectroscopic analysis of the hydrogen spectrum…

Atomic Spectrum of Hydrogen

Atomic SpectraWhen atoms absorb energy electrons move into higher energy levels…

…these electrons lose energy by emitting light when they return to lower energy

levels.

No two elements have the same emission spectrum

Iron

Hydrogen

• Ground State: the lowest possible energy of the electron

• Excitation of the electron by absorbing energy raises it from the ground state to an excited state

• A quantum of energy (in the form of light) is emitted when the electron drops back to a lower energy level.

• The light emitted by an electron moving from a higher to a lower energy level has a frequency directly proportional to the energy change of the electron.

Heisenberg Uncertainty Principle

You can find out where the electron is, but not where it is going.

OR…

You can find out where the electron is going, but not where it is!

“One cannot simultaneously determine both the position and momentum of an electron.”

Werner Heisenberg

Quantum of energy hits electron

electron elevates to another level.

As electron falls back to ground state, light is emitted

Electron transitions involve jumps of definite amounts of energy.

This produces bands of light with definite

wavelengths.

Electron excitation & emission

• http://micro.magnet.fsu.edu/primer/java/scienceopticsu/exciteemit/index.html

Flame Tests

Pickle Light