The Atom Atomic Theory - SD33sss.sd33.bc.ca/sites/default/files/Chemistry 11 Honours - 5-1... · 20140520 1 Atomic Theory Chemistry 11 The Atom !The concept of a discrete unit that

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Atomic Theory Chemistry 11

The Atom �  The concept of a discrete unit that makes up all

matter has been around for centuries.

�  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.

In addition… �  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.

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- 10 electrons

Ca2+ 18 electrons

Br- 36 electrons

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

What about the electrons?

�  For Example: �  Find the number of protons, neutrons, and electrons

possessed by the following:

�

1327Al

�

3375As

13 protons, 14 neutrons, 13 electrons

33 protons, 42 neutrons, 33 electrons

�  Isotope: �  Species having the same atomic number, but different

atomic masses (same # of protons, different number of neutrons).

For Example:

= ORDINARY HYDROGEN (called “protium”).

= DEUTERIUM (sometimes call “heavy” hydrogen).

= TRITIUM (called “radioactive” hydrogen). �

11H

�

12H

�

13H

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

average mass of a sample containing a mixture of isotopes.

�  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?

�  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.

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The Periodic Table!

�  As real elements became discovered, the Greek ideas of Air, Earth, Fire, and Water had to be abandoned.

�  Scientists needed an elegant, easy to use method of accessing all the information about the elements.

Major Divisions Within the Periodic Table

�  Period:

The set of elements in a given row going across the table.

�  Group or Family:

The set of elements in a given column going up and down the table.

�  Two important trends appear in the periodic table:

�  There are several groups, rows, and “blocks” of elements: In summary:

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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 the 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.

Listen  here,  I  say!  

�  For Hydrogen: �  Lyman Series:

�  Wavelengths in the UV spectrum of the hydrogen atom.

�  Results from electrons dropping into the n = 1 orbit.

�  Balmer series: �  Wavelengths in the visible light spectrum

of the hydrogen atom.

�  Results from electrons dropping into the n = 2 orbit.

�  Paschen Series: �  Wavelengths in the infrared spectrum of

the hydrogen atom �  Results from electrons dropping into the n

= 3 orbit. �  Brackett Series:

�  Wavelengths in the infrared spectrum of the hydrogen atom

�  Resuls from electrons dropping into the n = 4 orbit.

�  Pfund Series: �  Wavelengths in the infrared spectrum of

the hydrogen atom

�  Results from electrons dropping into the n = 5 orbit.

�  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).

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�  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 a 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 (more to come later).

�  Shell: �  The set of all orbitals having the same n value. �  For Example:

The 3rd shell consists 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 types 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 orbital consists of ONE s subshell. �  A p type orbital consists of THREE p subshells. �  A d type orbital consists of FIVE d subshells. �  An f type orbital consists of SEVEN f subshell.

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).

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ELECTRON CONFIGURATIONS �  The addition of electrons to the orbitals follows three simple

rules:

�  Aufbau Principle: �  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. 

�  Pauli Exclusion Principle: �  A maximum of TWO electrons can be placed in each subshell.

�  Hunds Rule: �  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.

Writing Electron Configurations for Neutral Atoms�

�  ELECTRON CONFIGERATION: Describes which orbitals in an atom contain electrons and how many electrons are in each orbit.

How do we do this?

Tryski… �  Predict the electron configuration of the following:

�  Si 1s22s22p63s23p2

�  Tc 1s22s22p63s23p64s23d104p65s24d5

�  Ca 1s22s22p63s23p64s2

�  Zr 1s22s22p63s23p64s23d104p65s24d2

�  Ga 1s22s22p63s23p64s23d104p1

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.

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

the electron configuration.

For Example:

Al 1s22s22p63s23p1

becomes:

[Ne]3s23p1

�  Write the following using core notation:

�  Zr ([Kr]5s24d2)

�  Ga ([Ar]4s23d104p1)

�  Co ([Ar]4s23d7)

Homework: �  Do:

�  The Periodic Table and Stuff W.S. #1-4 �  Anions:

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

For Example:

Oxygen: [He] 2s2 2p4 → [He] 2s22p6

Sulphur: [Ne] 3s2 3p4 → [Ne] 3s23p6

Writing Electron Configurations for Ions �

�  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

Outermost electrons are removed preferentially. Also, e- in the highest energy outermost orbital require the

least amount of energy to be completely removed from the atom.

�  For Example:

Tin: [Kr] 5s2 4d10 5p2 → [Kr] 5s24d10

Tin: [Kr] 5s2 4d10 5p2 → [Kr] 4d10

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�  Ru3+ [Kr]4d5

�  Sb3+ [Kr]5s24d10

�  S2- [Ne]3s23p6

�  N3- [He]2s22p6

I pity the fool who doesn’t do these examples!

�  2 exceptions to the configurations of elements up to Kr: �  Cr ([Ar] 4s2 3d4) → “3d4” is one e- short of a half

filled subshell. �  Cu ([Ar] 4s2 3d9) → “3d9” is one e- short of a filled

subshell.

�  The actual configurations for Cr and Cu are found to be: �  Cr ([Ar] 4s1 3d5) → “4s1” and “3d5” are two half filled

subshells. �  Cu ([Ar] 4s1 3d10) → “4s1” is a half filled subshell,

and “3d10” is a filled subshell.

�  Therefore:

A filled or exactly half filled d- subshell is especially stable.

�  Because of this extra stability, an atom or ion that is one e- short of a “d5” or “d10” configuration will shift an e- from the s- subshell having the highest energy into the unfilled d- subshell.

Predicting Number of Valence Electrons

�  Valence Electrons: �  Electrons that can take place in chemical reactions.

�  Are all the electrons in the atom EXCEPT: �  Core electrons.

�  In filled d or f subshells.

�  For Example:

�  Al([Ne] 3s2 3p1) has 3 valence electrons:

→ “3s2 3p1”

�  Ga([Ar] 4s2 3d10 4p1) has 3 valence electrons:

→ Omit “3d10” b/c filled

�  Pb([Xe] 6s2 4f14 5d10 6p2) has 4 valence electrons: → Omit “4f14” and “5d10” b/c filled 

�  Xe([Kr] 5s2 4d10 5p6) has ZERO valence electrons: → Noble gas configuration

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Quantum Numbers �  A way of giving each electron in an atom a specific address.

�  The four quantum numbers n, ℓ, mℓ, and ms specify the complete and unique quantum state of a single electron in an atom.

�  Gives the primary energy level.

�  The further from the nucleus, the larger the value of n.

�  n = 1, 2, 3, 4, 5, 6, 7.

�  Specifies the shape of an orbital with a particular principal quantum number.

�  The secondary quantum number divides the shells into smaller groups of orbitals called SUBSHELLS.

�  Usually, a letter code is used to identify ℓ to avoid confusion with n. �  ℓ = 0 refers to an s-subshell (sharp). �  ℓ = 1 refers to a p-subshell (principal). �  ℓ = 2 refers to a d-subshell (diffuse). �  ℓ = 3 refers to an f-subshell (fundamental).

Where: = (n − 1) and 0 ≤ ≤ (n − 1)

�  Describes the orientation of the orbital in space, where:

�  This gives each orbital in a subshell a unique name.

�  For Example: �  There is only one orbital in an s-subshell (ℓ = 0) because mℓ

can only have one number, 0.

�  In a p-subshell (ℓ = 1), you have three orbitals that can be uniquely named by mℓ as +1, 0, and -1.

m = −,...0...,+

�  Specifies the orientation of the spin axis of an electron.

�  An electron can spin in only one of two directions (sometimes called up and down).

�  Therefore:

�  The Pauli Exclusion Principle states that no two electrons in the same atom can have identical values for all four of their quantum numbers.

�  This means that no more than two electrons can occupy the same orbital, and that two electrons in the same orbital must have opposite spins.

ms = + 12 or − 1

2

�  When an electron spins, it creates a magnetic field which can be oriented in one of two directions.

�  For two electrons in the same orbital, the spins must be opposite to each other; the spins are said to be paired.

�  These substances are not attracted to magnets and are said to be DIAMAGNETIC.

�  Atoms with more electrons that spin in one direction than another contain unpaired electrons.

�  These substances are weakly attracted to magnets and are said to be PARAMAGNETIC.

n ℓ Orbital mℓ # of orbitals # of electrons 1 0 1s 0 1 2

2 0 2s 0 1 2

1 2p -1, 0, 1 3 6

3

0 3s 0 1 2

1 3p -1, 0, 1 3 6

2 3d -2, -1, 0, 1, 2 5 10

4

0 4s 0 1 2

1 4p -1, 0, 1 3 6

2 4d -2, -1, 0, 1, 2 5 10

3 4f -3, -2, -1, 0, 1, 2, 3 7 14

Possible Quantum Numbers:

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Peanut butter jelly time!

�  Predict the number of orbitals in the fourth shell, that is, for n = 4.

�  Give the label for each of these orbitals.

�  How many subshells are in each of these orbitals?

Homework: �  Do:

�  The Periodic Table and Stuff W.S. #5-21

�  Study for your quiz!!! �  Electron Configurations and Quantum Numbers

What is on the Exam? �  History of the Atom

�  The Atom �  Atomic Number and Mass

�  # of Protons, Neutrons, and Electrons

�  Isotopes

�  The Periodic Table

�  The Electronic Structure of the Atoms �  Theory

�  Configurations (Neutral, Ions, Core)

�  Exceptions

�  Valence Electrons

�  Quantum Numbers