Atomic TheoryChemistry 11

Crash Course

Back in the day…

• Aristotle proposed that all matter is made up of 4 elements with 4 different properties:

The fifth element is Aether, the material that fills the region of the universe above the terrestrial sphere

Democritus

Democritus suggested that all matter was made of tiny indivisible particles called atoms. (Greek “atoma”)

These ideas were based on philosophical reasoning rather than experimentation and empirical observations

This concept has been accepted by scientists since it elegantly explains new discoveries in the field of chemistry.

Ancient Greeks were the first to come up with the idea of atoms.

In the dark ages, the idea of atoms was frowned upon. Not much progress was made.

ATOMS? What’s that? OFF with

your HEAD!

In the early 1800’s, John Dalton came up with the ATOMIC THEORY.

His main points were:

Dalton’s Atomic Theory1) All matter is made of atoms. Atoms

are indivisible and indestructible.2) All atoms of a given element are

identical in mass and properties 3) Compounds are formed by a

combination of two or more different kinds of atoms.

4) A chemical reaction is a rearrangement of atoms.

Compounds are composed of little “balls” called atoms, joined together by “bonds” to form molecules.

An indestructible “Dalton” atom

Ha Ha! You can’t break me!!!!

Thomson’s Plum Pudding Model

Ernest Rutherford

Rutherford’s Scattering Experiment

Rutherford could not explain why the electron didn't’t fall into the nucleus and destroy the atom.

?

Neils Bohr

I think I can help!

Atomic Number and Atomic Mass

• The elements are differentiated from one another by the numbers of protons in the nucleus.

• Atomic Number:

• The number of protons in the nucleus.

• A neutral atom has no charge, therefore:

In a neutral atom:Number of Protons = Number of Electrons

• For Example:

How many electrons are possessed by the following?

N3+

Ca2+

Br-

If electrons are added or subtracted, the resulting particles is an ION.

• Since both neutrons and protons have a mass of 1.0u, the atomic mass of an atom will be found by their combined totals.

• Isotope:

• Species having the same atomic number, but different atomic masses (same # of protons, different number of neutrons).

For Example:

• For Example:

• Find the number of protons, neutrons, and electrons possessed by the following:

• For Example:

Experiments show that chlorine is a mixture which is 75.77% Cl-35, and 24.23% Cl-37. If the precise molar mass of Cl-35 is 34.968853 g/mol and of Cl-37 is 36.965903 g/mol, what is the average molar mass of the chlorine atoms in such a mixture?

The molar masses given on the periodic table are found by calculating the

average mass of a sample containing a mixture of isotopes.

• You may also use the atomic mass to calculate the average. The average mass will be less exact, but still satisfactory.

Homework:

• Do:

• Introduction to Atomic Theory W.S.

The Electronic Structure of the Atom

• When a hydrogen atom is irradiated by energy, some of the energy is absorbed then reemitted as light

• If the light is passed through a prism, a “line spectrum” is observed.

• In 1913, Niels Bohr proposed a model that explained why the observed line spectrum for Hydrogen looks that way it does.

• He proposed that:

• The electron in hydrogen can only exist in specific energy states. These energy states are associated with specific circular orbits which the electron can occupy around the atom.

• When an electron absorbs energy, it instantaneously moves from one orbit to another.

• The greater the energy, the farther the orbit is from the nucleus.

• ENERGY LEVEL:

A specific amount of energy which an electron in an atom can possess.

• The energy levels of hydrogen have the pattern below (“n” is the number of the energy level).

• The observed spectrum represents energy level differences occurring when an electron gives off energy and drops from a higher energy level.

• The energy difference between two different energy levels is called the QUANTUM of energy associated with the transition between the two levels.

• A few years after Bohr published his theories, several changes were made to his ideas.

• The idea of electrons orbiting along specific path in a well defined orbit had to be abandoned.

• Instead, different electrons, depending on their energies, occupy particular regions of space called “orbitals”.

The Energy Level Diagram for Hydrogen

• The lowest sets of energy levels for hydrogen are as follows:

• Each dash represents the energy possessed by a particular orbital in the atom.

• The letter s, p, d, and f refer to the four “types” of orbitals.

• Shell

• The set of all orbitals having the same n value

• For example:

The 3rd shell of the 3s, 3p, and 3d orbitals.

• Subshell:

• A set of orbitals of the same type.

• For Example:

The set of five 3d orbitals in the 3rd shell is a subshell.

•Some notes…• All the orbitals for a hydrogen atom with a given value of

n have the same energy (not true for atoms with more than one electron).

• Rules governing which type of orbitals can occur:

• For a given value of “n”, certain types of orbitals are possible• For n = 1: only the s type is possible

• For n = 2: the s and p types are possible

• For n = 3: the s, p, and d types are possible

• For n = 4: the s, p, d, and f types are possible

• An s type subshell consists of ONE s orbital

• A p type subshell consists of THREE p orbitals

• A d type subshell consists of FIVE d orbitals

• An f type subshell consists of SEVEN f orbitals

The Energy Level Diagram for Polyelectronic Atoms

• The energy level diagram must be modified to describe any other atom.

• The following diagram applies to ALL polyelectronic atoms (atoms having more than one electron)

Electron Configurations• The addition of electrons to the orbitals follows three

simple rules:

• As atomic number increases, electrons are added to the available orbitals. To ensure LOWEST POSSIBLE ENERGY for the atom, electrons are added to the orbitals having the lowest energy FIRST.

• A maximum of TWO electrons can be placed in each subshell.

• When electrons occupy subshells of equal energy, they must be singly occupied with electrons having parallel spins. 2nd electrons are then added to each subshell so each electron has opposite spin

n = 7 7s

n = 6 6s 6p 6d

n = 5 5s 5p 5d 5f

n = 4 4s 4p 4d 4f

n = 3 3s 3p 3d

n = 2 2s 2p

n = 1 1s

Writing Electron Configurations for Neutral Atoms

• ELECTRON CONFIGURATION:

Describes which orbitals in an atom contain electrons and how many electrons are in each orbit

Where are the Electrons?

Tryski…• Predict the electron configuration of the following:

• Si

• Tc

• Ca

• Zr

• Ga

Core Notation

• The electrons belonging to an atom can be broken into two subsets:

• The CORE electrons.

• The OUTER electrons.

The CORE of an atom is the set of electrons with the configuration of the

nearest noble gas having an atomic number LESS than that of the atom

being considered.

The OUTER electrons are all those outside the core. Since the core electrons are not involved in chemical reactions, the are excluded from the electron configuration.

For Example:

Al

becomes

• Write the following using core notation:

• Zr

• Ga

• Co

Homework:

• Do:

• The Periodic Table and Stuff W.S. #1-4

Writing Electron Configurations for Ions• Anions:

• Add elections to the last unfilled subshell, starting where the neutral atom left off.

• For Example:

Oxygen: [He] 2s22p4 + 2e-

Sulphur: [Ne] 3s23p4 + 2e-

• Cations:

• 2 Rules

1. Electrons in the outermost shells (largest n value) are removed first.

2. If there are electrons in both the s and p orbitals of the outermost shell, the electrons in the p orbitals are removed first.

p electrons BEFORE s electrons BEFORE d electrons

• For Example:

Tin: [Kr] 5s24d105p2 2e-

Tin: [Kr] 5s24d105p2 4e-

• Ru3+

• Sb3+

• S2-

• N3-

Predicting Number of Valence Electrons

• Valence Electrons:

• Electrons that can take place in chemical reactions

• Are all the electron in atom EXCEPT:

• Core electrons.

• In filled d or f subshells

• For Example:

• Al ([Ne] 3s23p1)

• Ga ([Ar] 4s23d104p1)

• Pb ([Xe] 6s24f145d106p2)

• Xe ([5s24d105p6)