Atom-the smallest particle of an element that can exist.

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JF Basic Chemistry Tutorial Atomic Theory
Shane Plunkett
plunkes_at_tcd.ie
Matter -anything that had mass and takes up
space -classified into 3 types elements,
compounds and mixtures
Element -consists of only one kind of atom. So,
an element is a pure substance.

• Atom -the smallest particle of an element that
  can exist.
• -made up of three subatomic particles protons,
  neutrons and electrons.

How do we know this?
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Theories through time.
1890s J.J. Thompson - The cathode ray
experiment negative particles discharged were
named electrons and denoted e-
Early 20th Century The Plum Pudding model! atoms
made up of blobs of a positively charged jelly
with electrons suspended in it
1908 Ernest Rutherford The Gold Leaf
Experiment 2 students Geiger and Marsden shot
a-particles (positively charged particles) at a
thin piece of gold foil. Most passed through,
some bounced back.

• Conclusions Cant be a plum pudding structure!
• - central positively charged point and a large
  volume of empty space

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Central point called the atomic nucleus and this
is where all positively charged protons are
located.
Atoms do not have an overall charge. There must
be the same number of electrons as there are
protons.
Definition The number of protons in an atomic
nucleus is called the atomic number (Z) of the
element. Note the atomic number of an element is
the smaller number

• If we take a look at the periodic table, we
  see that it is made up of vertical blocks (these
  are called groups) and horizontal rows (these are
  called periods).

e.g. look at Oxygen in the Periodic Table
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• Questions
•  
• Give the name, atomic number and group number of
  the element
• with the following Z value 

(a) Z16 Sulfur Group VI
(b) Z8 Oxygen Group VI
(c) Z19 Potassium Group I
(d) Z13 Aluminium Group III
(e) Z32 Germanium Group IV
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What about neutrons?
A mass spectrometer is an instrument used by
chemists to determine the mass of a given
element. Not all atoms of a particular element
have the same mass.
Example Neon (Ne) gives 3 types of atoms, each
heavier than the last.
Explanation There must be a third type of
subatomic particle that contributes to this
change in mass
 Neutrons (n) have no electric charge, but have
the same mass as protons.
Definition Atoms of a single element that
differ from each other in mass are called
isotopes

• Isotopes have the same atomic number but
  different atomic masses
• They must have the same number of electrons and
  protons but a
• different number of neutrons

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Definition The mass number or atomic mass (A) is
the total number of protons and neutrons in a
nucleus of an atom. Note the mass number or
atomic mass of an element is the larger number
As we go across the rows of the periodic table,
the mass of the nucleus increases.
 Example
Hydrogen Atomic number 1
This means a hydrogen atom possesses one electron
and one proton
Mass number 1.0079
No. of neutrons in the hydrogen atom Mass no.
Atomic no. 1.0079 1 ? 0
This means that the hydrogen atom has one
electron, one proton and no neutrons.

• If we move onto the next element, Helium, we
  find
• Atomic number 2, i.e. 2 electrons and 2 protons
  in each He atom
• Mass number 4.00
• No. of neutrons in the helium atom Mass no.
  Atomic no. 2
• Each helium atom has 2 electrons, 2 protons and 2
  neutrons

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Sample Question Given Z (the atomic number), how
many protons, neutrons, and electrons are in the
following elements?
(a) Z6 (Carbon)
? have 6 protons and 6 electrons
number of neutrons mass number atomic
number 12 6 6 neutrons
(b) Z12 (Magnesium)
? have 12 protons and 12 electrons number of
neutrons 24 12 12 neutrons
(c) Z17 (Chlorine)
? have 17 protons and 17 electrons number of
neutrons 35 17 18 neutrons

• Now you try! Z 30 Z16 Z35

Z 30, Ans. 30 electrons, 30 protons, 35
neutrons Z 16, Ans. 16 electrons, 16 protons,
16 neutrons Z 35, Ans. 35 electrons, 35
protons, 45 neutrons
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Sample Question
The three naturally occuring isotopes of argon
are 36Ar, 38Ar and 40Ar. How many protons,
neutrons and electrons are present in each?
Note that in isotopes, the mass numbers can
change, but the atomic number always remains the
same. It is only the number of neutrons that can
change.
Mass numbers are 36, 38 and 40 respectively.
From the Periodic Table, we know that the atomic
number of argon is 18. This means that each
isotope has 18 protons and 18 electrons.

•   Number of neutrons mass number atomic
  number
•  
• 36Ar no. of neutrons 36 18 18 neutrons
• 38Ar no. of neutrons 38 18 20 neutrons
• 40Ar no. of neutrons 40 18 22 neutrons

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We have so far built up a picture of the atom
that looks like this
   
Atomic Nucleus -contains the protons and
neutrons
Empty space
Electrons are found somewhere in here
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We know that electrons orbit the atomic nucleus.
Can we be more definite about where electrons can
be found?
YES!
There is a probability associated with finding an
electron in a particular location around the
nucleus.
Definition An orbital is a region of space
where the probability of finding an electron is
large. 

• Four types of orbitals s, p, d and f. We will
  look at the first 3.

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Quantum Numbers of Atomic Orbitals Atomic
orbitals have 4 quantum numbers associated with
them
1. The Prinicipal Quantum Number (n)
-positive, whole number (n1, 2, 3)
-tells you about the size of the orbital, i.e.,
the distance from the nucleus
-tells you about the energy of the orbital the
bigger the number, the higher the energy level
-the orbitals form a series of shells (like the
layers of an onion). Shells of higher n
surround shells of lower n.
2. The Angular Momentum Quantum Number (l)
-can be a whole number with values from 0 to n-1
-tells you about the shape of the orbital, i.e.
if it is an s, p or d orbital
-you can get its values when you know n,
e.g. if n1, then l can only equal 0, so you
have an s-orbital
if n2, then l can equal 0 or 1, so you have s-
and p-orbitals
if n3, then l can equal 0, 1 or 2, so can have
s-, p- and d-orbitals
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3. The Magnetic Quantum Number (ml)
-whole number from l through 0 to l
- labels the orbitals and tells you how many of
each orbital type you have
4. The Spin Magnetic Quantum Number (ms)
-electrons behave like spinning spheres and this
behaviour is called spin

• spin can be clockwise (usually represented as
  an arrow pointing upwards (?) and is given the
  value ½

• or anticlockwise represented as an arrow
  pointing downwards (?) and is given the value
  -½

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Sample Question Determine the total number of
orbitals in a shell with a principle quantum
number n 3.
Step 1 Determine the angular momentum quantum
numbers associated with n 3
Angular momentum quantum number, l 0, 1, 2, ,
n -1
For n 3, there are l 0, 1, 2 subshells
For l 0, we have an s-orbital
l 1, we have a p-orbital
l 2, we have a d-orbital

• To find out how many of each orbital we have, use
  the magnetic
• quantum number, ml

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Step 2 Use the Magnetic Quantum Number to
determine the number of each type of orbital we
have
Magnetic quantum number, ml l, l -1, l -2, , -l
For l0 ml0
This means we have one s-orbital
For l1 ml1, 1-1, -1
1, 0, -1
This means we have three p-orbitals
For l2 ml 2, 2-1, 2-2, 2-3, -2
2, 1, 0, -1, -2
This means we have five d-orbitals
So, in total, we have 135 9 orbitals in the
shell with n3.

• Now you try!
• How many orbitals are there in the shell with
  n2?
• Ans. 4 orbitals (1s-orbital and 3 p-orbitals)

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What do these orbitals look like?
s-orbitals -spherical in shape

p-orbitals -a cloud with two lobes on opposite
sides of the nucleus.

-the two lobes are separated by a planar region
called a nodal plane
-there are 3 p-orbitals which point along the x,
y, and z axes respectively

px
pz
py
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• d-orbitals
• -there are five d-orbitals named after the
  directions in which they point

dx2 y2
dxz
dxy
dz2
dyz
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There are rules for how electrons will
accommodate these orbitals.

• Rule 1 Electrons will occupy the orbital of
  lowest available energy.
• The energy of orbitals increases in the order
  sltpltd,f

Rule 2 Each orbital can accommodate a maximum of
two electrons.
The two electrons in an orbital must have
opposite spins (i.e. one must be clockwise and
the other must be anticlockwise).
This is called the Pauli Exclusion Principle.
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Example Helium, He Atomic number 2Lowest
energy orbital available is first s-orbital,
denoted the 1s orbital
 

• Because we have 2 electrons and each orbital can
  have 2 electrons, we
• can place both electrons of He into the 1s
  orbital

Nucleus containing 2 protons and 2 neutrons
?
Electron
?
Electron
1s orbital
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The name we give to the build-up of these
electrons in orbitals is the electronic
configuration

• So, the electronic configuration of the He atom
  is 1s2, where the
• superscript 2 indicates the number of electrons
  present in the orbital.

Question How many electrons are present in the
hydrogen atom and how would you write its
electronic configuration?
Answer Hydrogen has one electron and an
electronic configuration of 1s1
After the 1s orbital, next comes the 2s orbital.
Again this can hold 2 electrons
Question How many electrons are present in the
lithium atom and how would you write its
electronic configuration?
Answer Lithium has 3 electrons and an
electronic configuration of 1s2 2s1
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Once the 2s orbital has been filled, the 2p
orbital is now filled

• Example Boron, B Atomic number 5

First, the 1s and 2s orbitals are filled. This
accounts for 4 out of the 5 electrons.
Because the 3s orbital is too high in energy and
we can only place two electrons in each orbital,
the final electron goes into the 2p orbital
Hence, the electronic configuration for the
ground state boron atom is 1s2 2s2 2p1
The third rule we need to know is called Hunds
Rule and states that electrons will occupy
degenerate orbitals (those of the same energy)
singly before pairing (happens in buses).
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ExampleThe atomic number of carbon is 6 and the
electronic configuration is 1s2 2s2 2p2

• The 2p orbital fills up as follows

??
??
?
?
?
1s2
2s2
2p2
Note! The 2p orbitals have been filled up
singly!
??
??
??
?
1s2
2s2
2p2
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QuestionWhat are the electronic configurations
of the following elements (a) Nitrogen, (b)
Magnesium, (c) Chlorine, (d) Silicon, (e)
Potassium?

• The elements we have dealt with so far have only
  contained s and p
• orbitals

These elements make up the s and p blocks of the
periodic table
The first two groups are called the s-block and
the last six groups are called the p-block
In between these, we find the d-block or the
transition metal elements
The d-block elements are so called because this
is where we begin to fill the d-orbitals
5 d-orbitals ? can accommodate 10 electrons
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QuestionWhat are the electronic configurations
of the following elements (a) Nitrogen, (b)
Magnesium, (c) Chlorine, (d) Silicon, (e)
Potassium?

• Answer
• (a) Nitrogen has 7 electrons and an electronic
  configuration of 1s2 2s2 2p3
• (b) Magnesium has 12 electrons and an electronic
  configuration of 1s2 2s2 2p6 3s2
• (c) Chlorine has 17 electrons and an electronic
  configuration of 1s2 2s2 2p6 3s2 3p5
• (d) Silicon has 14 electrons and an electronic
  configuration of 1s2 2s2 2p6 3s2 3p2
• (e) Potassium has 19 electrons and an electronic
  configuration of 1s2 2s2 2p6 3s2 3p6 4s1

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ExampleThe atomic number of Iron is 26 ?we have
26 electrons

• We begin to fill up the orbitals as follows

??
??
??
??
??
??
??
??
??
2s2
2p6
3s2
3p6
1s2
We have filled 18 out of 26 electrons
Now, we face a problem!
As n increases, the sublevel energies get closer
together. This results in the overlap of some
sublevels and here we have such a case.
The 4s sublevel is slightly lower in energy that
the 3d orbital, so it is filled first.
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??
??
?
?
?
?
4s2
3d6
Question
What is the electronic configuration of (a)
Vanadium, (b) Cobalt, (c) Zinc?
Answer (a) Vanadium has 23 electrons and an
electronic configuration of 1s2 2s2 2p6 3s2 3p6
4s2 3d3 (b) Cobalt has 27 electrons and an
electronic configuration of 1s2 2s2 2p6 3s2 3p6
4s2 3d7 (c) Zinc has 30 electrons and an
electronic configuration of 1s2 2s2 2p6 3s2 3p6
4s2 3d10
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Light is emitted in discrete or definite packets
called quanta or phonons
The nature of light and Atomic Spectra

• Frequency of light, ?, increases proportionally
  with increase in energy, E

?E h?
where h is Plancks constant, 6.63 10-34Js
Electrons possess both kinetic and potential
energy
Electrons can behave like rocks on a cliff when
the rock falls it gives up potential energy
?
If electron falls towards nucleus, they give up
potential energy
When excited electrons in atoms fall from a high
energy level to a low energy state, light is
emitted with a specific frequency or colour
Electrons must occupy certain energy levels
like steps on a ladder
If an electron absorbs a photon or quantum of
light, it is elevated to a higher energy level
an excited state
When the electron falls to lower energy state, it
gives up this energy in specific quanta
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Atoms with all electrons in their lowest energy
levels are said to be in their ground state

• The energy difference between any two levels may
  be found by using
• Rydbergs constant

?E Efinal Einitial 2.18 10-18 J
x
where 2.18 10-18J is the Rydberg constant and n
is a positive integer representing the shell
number
We can convert the units of the Rydberg constant
from Joules to metres-1
? c/?
?E h?
?E 2.18 10-18 J
x
Combine these to give
h Plancks constant
?E hc 2.18 10-18 J ?
x
c Speed of light
? wavelength
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? 1 2.18 10-18J ? hc
x
(2.18 10-18 J) .
(6.626 10-34Js)(3.00 108 ms-1)
x
1.0967 107 m-1
x
Example
What is the wavelength of a photon emitted during
a transition from the ninitial 5 state to the
nfinal 2 state in the hydrogen atom?
?E R
2.18 10-18 J (1/4 1/25)
4.58 10-19 J
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To calculate the wavelength of the photon

• ? c/? ch/?E

(3.00 108 ms-1)(6.63 10-34Js) 4.58
10-19 J
4.34 10-7m
For wavelengths, must convert to nanometres
4.34 10-7 m (1 109 nm)/(1 m)
434 nm
Question
What is the wavelength of a photon emitted during
a transition from ninitial 4 to nfinal 2
state in the H atom?
Answer 487 nm