Historyand Progression

of Atomic Theory

Structure of the Atom

• An atom is the smallest particle of an element that retains the chemical properties of that element.

• The nucleus is a very small region located at the center of an atom.

• The nucleus is made up of at least one positively charged particle called a proton and usually one or more neutral particles called neutrons.

• Surrounding the nucleus is a region occupied by negatively charged particles called electrons.

• Protons, neutrons, and electrons are often referred to as subatomic particles.

Democritus

This is the Greek philosopher Democritus who began the search for a description of matter more than 2400 years ago.

He asked: Could matter be divided into smaller and smaller pieces forever, or was there a limit to the number of times a piece of matter could be divided?

400 BC

DEMOCRITUS’ ATOMIC THEORY 400BC

Democritus asserted that space contained an infinite number of particles

• Named atomos, “not cutting” or "indivisible” •Atoms are eternal and invisible; absolutely small, so small that their size cannot be diminished; totally full and incompressible.•Atoms are homogeneous, differing only in shape, arrangement, position, and number

Atomos

To Democritus, atoms were small, hard particles that were all made of the same material but were different shapes and sizes.

Atoms were infinite in number, always moving and capable of joining together.

Dalton’s Theory (early 1800’s)

He deduced that all elements are composed of atoms. Atoms are indivisible and indestructible particles.

Atoms of the same element are exactly alike.

Atoms of different elements are different.

Compounds are formed by the joining of atoms of two or more elements in specific ratios.

In chemical reactions, atoms are combined, separated, and rearranged

Explained the Law of the Conservation of Mass

Matter is neither nor destroyed during ordinary chemical reactions or physical changes

Modern Atomic Theory

Atomic Theory has been modified We now know that atoms are divisible

into smaller particles And that elements can have atoms with

different masses (isotopes)

J.J. Thomson (English physicist1897)

Used Cathode Ray tube to determine the presence of

– (electrons) and + (protons) particles.

http://highered.mcgraw-hill.com/olcweb/cgi/pluginpop.cgi?it=swf::100%::100%::/sites/dl/free/0072512644/117354/01_Cathode_Ray_Tube.swf::Cathode%20Ray%20Tube

Current passed through a glass tube filled with various gases at low pressure from the cathode to the anode (terminals of the voltage source)

Stream of particles glowed (cathode ray) Cathode rays deflected away from negatively charged objects Led to the hypothesis that cathode ray particles are negatively

charged These negatively charged particles were named electrons

J.J. Thomson

Plum Pudding Model -- the structure of an atom is something like pudding. He assumed that the basic body of an atom is a spherical object containing electrons & protons randomly confined in homogeneous jellylike material. Positive charges cancel the negative charges.

This model was soon disproved

http://www.learnerstv.com/animation/animationcategory.php?cat=Chemistry

Ernest Rutherford (Eng.1911)

Atoms have a central positive nucleus surrounded by negative orbiting electrons.

This idea was the result of his famous Gold Foil Experiment(see next slide). This experiment involved the firing of radioactive alpha particles through gold foil.

Alpha particles = + charged particles about 4x mass of H This model suggested that most of the mass of the atom

was contained in the small nucleus, and that the rest of the atom was mostly empty space.

Most of particles passed straight through the foil but approximately 1 in 8000 were deflected.

Conclusion- volume of the nucleus very small compared to the volume of the atom

To put the size or the nucleus in perspective, if the nucleus of an atom was the size of a marble, the atom would be the size of a football field!

Rutherford’s Gold Foil Alpha Scattering Experiment

Alpha scattering due to repulsion from the nucleus

Nucleus-Positive Charged Center The nucleus is far too large in this drawing.

Niels Bohr (Danish 1913)

The Bohr Model is probably familiar to us as the "planetary model" of the atom is used to symbolize atomic energy.

Electrons orbit the nucleus much like planets orbiting the Sun. Electrons travel in certain orbits at specific distances from the

nucleus with definite energy levels (energy shells or energy levels) Lower energy state near the nucleus and higher energy state as

you move away from the nucleus Electrons gain or lose energy by jumping between orbits,

absorbing or emitting electromagnetic radiation    Think of rungs on a ladder, except the distance between the rungs is not constant!

Bohr Model of the Atom

Bohr Atom used as logo for Atomic Energy Commission

Chadwick (English 1932)

James Chadwick discovered a third type of particle, which he named the Neutron. Neutrons help to reduce the repulsion

between protons and stabilize the atom's nucleus.

Neutrons always reside in the nucleus of atoms and they are about the same mass and size as protons.

Neutrons do not have any electrical charge; they are electrically neutral

Chadwick’s atom w/ P+, e- & N+/-

Table 1 (p.76)

Particle Relative Electric Charge

Mass Number

Relative Mass (amu)

Actual Mass (kg)

Electron (e-)

-1 0 0.0005486

9.109E-31

Proton (p+)

+1 1 1.007276 1.673E-27

Neutron (n0)

0 1 1.008665 1.675E-27

Cloud Theory (1920’-1930’s)

Quantum Mechanics

Quantum Model of the Atom

Louis de Broglie (French scientist 1924) Hypothesized that electrons have wavelike

properties Electrons exist at certain frequencies

corresponding to specific energies Confirmed experimentally by investigating

diffraction (bending of wave as it passes through small opening) and interference (waves overlap)

Cloud Theory

Based on the work of many scientists Based on the mathematical approach of

Quantum Mechanics Electrons are assigned regions of space

(Orbitals) not pathways (Orbits) Electrons are moving around the nucleus

rapidly in no predictable path producing a cloud of e-’s over time. Think of a rapidly moving fan blade.

Cloud Theory of Today

Electron Cloud

Quantum Theory

Describes mathematically the wave properties of electrons and other very small particles

Schrodinger Wave Equation (Austrian 1926)- showed mathematically how electrons can exist as waves of certain energies and therefore certain frequencies

Heisenberg Uncertainty Principle(Heisenberg- German 1927)

Electrons act as particles and waves How can you locate electrons in an atom?

Observe interference with photons This knocks the electron off course

“It is impossible to determine simultaneously both the position and velocity of an electron or any other particle”

Atomic Number

• Atoms of different elements have different numbers of protons.

• Atoms of the same element all have the same number of protons.

• The atomic number (Z) of an element is the number of protons of each atom of that element.

Atomic Number

Mass Number

The mass number is the total number of protons and neutrons that make up the nucleus of an isotope.

How to find # of Protons

Number of protons in an atom is ALWAYS equal to the Atomic Number

Number of Electrons (e-)

Atoms – Number of protons and electrons are equal (overall neutral charge)

Ions Loss of electron(s) makes positive ions

Gain of electron(s) makes negative ions

Loss or Gain of Electrons

Atom Ion Na Na+1 + 1e- (LOSS) +11 +11 -11 -10 0 net +1 net charge

Cl2 + 2e- 2Cl- (GAIN) +17 +17 -17 -18 0 net -1 net

How to calculate # of Neutrons

Atomic Mass (rounded to integer) - Atomic Number ---------------------------------------------------- Number of Neutrons in the nucleus

Atomic Mass – Atomic Number = # Neutrons

Atom Contents

Protons (p+) – always equal to Atomic # Electrons (e-)

In Atoms – Same as the # of Protons In Ions – Net charge after e-’s have been lost

or gained in an attempt to become stable Loss of e-’s = Positive charge Gain of e-’s = Negative charge

Neutrons (n0) = Atomic Mass – Atomic #

Isotopes

Isotopes

• Isotopes are atoms of the same element that have different masses.

• The isotopes of a particular element all have the same number of protons and electrons but different numbers of neutrons.

• Most of the elements consist of mixtures of isotopes.

Designating Isotopes

• Hyphen notation: The mass number is written with a hyphen after the name of the element.

• uranium-235

• Nuclear symbol: The superscript indicates the mass number and the subscript indicates the atomic number.

235 92 U

• The number of neutrons is found by subtracting the atomic number from the mass number.

mass number − atomic number = number of neutrons

235 (protons + neutrons) − 92 protons = 143 neutrons

• Nuclide is a general term for a specific isotope of an element.

Sample Problem AHow many protons, electrons, and neutrons are there in an atom of chlorine-37?

Solution: atomic number = number of protons = number of electronsmass number = number of neutrons + number of protons# P+= 17# e-= 17# N0= mass number - #P+= 37 – 17 = 20

• Law of definite proportions: a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound

Law of multiple proportions: if two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers

Law of Conservation of Mass

Relative Atomic Masses

• The standard used by scientists to compare units of atomic mass is the carbon-12 atom, which has been arbitrarily assigned a mass of exactly 12 atomic mass units, or 12 amu.

• One atomic mass unit, or 1 amu, is exactly 1/12 the mass of a carbon-12 atom.

• The atomic mass of any atom is determined by comparing it with the mass of the carbon-12 atom.

Average Atomic Masses of Elements

• Average atomic mass is the weighted average of the atomic masses of the naturally occurring isotopes of an element.

Calculating Average Atomic Mass• The average atomic mass of an element

depends on both the mass and the relative abundance of each of the element’s isotopes.

Average Atomic Masses of Elements, continued

Copper consists of 69.15% copper-63, which has an atomic mass of 62.929 amu, and 30.85% copper-65, which has an atomic mass of 64.927 amu.

The average atomic mass of copper can be calculated by multiplying the atomic mass of each isotope by its relative abundance (expressed in decimal form) and adding the results.

Average Atomic Masses of Elements, continued

(0.6915 × 62.929 601 amu) + (0.3085 × 64.927 794 amu) = 63.55 amu

The calculated average atomic mass of naturally occurring copper is 63.55 amu.

Hotel Head to Toe Activity

How many floors were there in the hotel? How many double beds were in the s

rooms? The p rooms? The d rooms? The f rooms?

How many guests would each room hold? What was special about how guests had to

sleep in the beds?

We will now compare our hotel and its rooms and guests with electrons filling orbitals in atoms

Guests = electrons

Floors = Energy levels (1, 2, 3, 4, 5, 6, or 7)

Rooms = Energy sub-levels (s, p, d, and f) Double beds = Orbitals The way guests sleep in the bed =

Electron spin

Quantum Numbers

Quantum numbers specify the properties of atomic orbitals and electrons in orbitals Energy level- principal quantum number (n) Orbital shape- angular momentum quantum

number (l) Orientation around the nucleus- magnetic

quantum number (m) Electron spin- spin quantum number

Principal Quantum Number (n)

Indicated main energy level for the electron

Values of 1, 2, 3, … n=1 at the orbital closest to the nucleus,

lowest energy level As n gets larger, electron’s energy and

distance from the nucleus increase

Angular Momentum Quantum Number (l)

Orbital shape; sublevels

Letter (sublevel)

s (sphere)

p (dumb-bell)

d

f

Magnetic Quantum Number (m)

Orientation around the nucleus Orbitals with same shape but different location

in 3-D coordinate system

Only one s orbital Three p orbitals; px, py, and pz

Five d orbitals Seven f orbitals

Orbitals shapes

S or spherical orbital

Orbital Shapes

P orbitals or perpendicular orbitals y x

z

Orbital Shapes D orbitals or Diagonal orbitals

Spin Quantum Number (+1/2 or -1/2)

Each orbital can contain up to two electrons

The two electrons must have different spin states

Aufbau Rule

An electron occupies the lowest-energy orbital that can receive it.

Example: Take the least number of steps to get to your hotel room. Start filling rooms on the first floor and then the second floor. Sometimes it takes fewer steps to go up a floor to an s room, than to walk to the end of the hall on the current floor (d or f room)

Hund’s Rule

"Orbitals of equal energy are each occupied by one electron before any is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin"

Example: One person fills each bed in the room before doubling up!

Example: each p orbital at a particular energy level must receive one electron before they can have 2 electrons

Pauli Exclusion Principle

No two electrons in the same atom can have the same four quantum numbers Think of your home and address Street number, city, zip code, state

Two electrons is the same orbital must have different spin quantum numbers (opposite spin)

Example: sleep head to toe

Homework

Electron Configuration packet