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All matter is composed of atoms.
Understanding the structure of
atoms is critical to understanding
the properties of matter
HISTORY OF THE ATOM
1808
John Dalton
suggested that all matter was made up of
tiny spheres that were able to bounce around
with perfect elasticity and called them
ATOMS
DALTONS ATOMIC THEORY
Particle
Mass
(g)
Charge
(Coulombs)
Electron (e-) 9.1 x 10-28 -1.6 x 10-19
Proton (p)
-24
1.67 x 10
Neutron (n) 1.67 x 10-24
+1.6 x 10
0
-19
Charge
(units)
-1
+1
0
mass p = mass n = 1840 x mass e-
HISTORY OF THE ATOM
1898
Joseph John Thompson
found that atoms could sometimes eject a far
smaller negative particle which he called an
ELECTRON
HISTORY OF THE ATOM
1910
Ernest Rutherford
oversaw Geiger and Marsden carrying out his
famous experiment.
they fired Helium nuclei at a piece of gold foil
which was only a few atoms thick.
they found that although most of them
passed through. About 1 in 10,000 hit
Rutherford’s Model of the Atom
atomic radius ~ 100 pm = 1 x 10-10 m
nuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m
Atoms are composed of
-protons – positively charged particles
-neutrons – neutral particles
-electrons – negatively charged particles
Protons and neutrons are located in the
nucleus. Electrons are found in orbitals
surrounding the nucleus.
Atomic Structure
HELIUM ATOM
Shell
proton
+
electron
N
N
+
-
neutron
Atomic Structure
Every different atom has a characteristic
number of protons in the nucleus.
atomic number = number of protons
Atoms with the same atomic number
have the same chemical properties and
belong to the same element.
Atomic Structure
Each proton and neutron has a mass of approximately 1 dalton.
The sum of protons and neutrons is the atom’s atomic mass.
Isotopes – atoms of the same element that have different atomic
mass numbers due to different numbers of neutrons.
ATOMIC STRUCTURE
Atomic mass
the number of protons and
neutrons in an atom
4
Atomic number
the number of protons in an atom
2
He
number of electrons = number of protons
ATOMIC NUMBER (Z) = number of protons in nucleus
MASS NUMBER (A) = number of protons + number of neutrons
= atomic number (Z) + number of neutrons
ISOTOPS are atoms of the same element (X) with different numbers
of neutrons in the nucleus
Mass Number
A
ZX
Atomic Number
1
1H
235
92
2
1H
U
Element Symbol
(D)
238
92
3
1H
U
(T)
1
1
H
2
1
H (D)
3
1
H (T)
* An element can have several naturally occurring isotopes.
* These isotopes of a element behave in the same way.
* In calculating the relative atomic mass of an element with
isotopes, the relative mass and proportion or percentage of
each is taken into account.
Isotope
Relative isotopic
mass
Relative abundance
(%)
Cl
34.969
75.80
Cl
36.966
24.20
Ar = (relative isotopic mass X1 % abundance) + relative isotopic mass X2 % abundance)
100
Ar (Cl) = (34.969 X 75.8) + ( 36.966 X 24.2)
100
Ar (Cl) = 2650.65 + 894.58
100
Ar (Cl) = 35.45 amu (atomic mass unit)
HISTORY OF THE ATOM
1913
Niels Bohr
studied under Rutherford at the Victoria
University in Manchester.
Bohr refined Rutherford's idea by adding
that the electrons were in orbits. Rather
like planets orbiting the sun. With each
orbit only able to contain a set number of
electrons.
MULTIELECTRON ATOMS
Quantum Number
* The principal quantum number, n, describes the energy
level on which the orbital resides.
* Largest E difference is between E levels
* The values of n are integers > 0.
* 1, 2, 3,...n.
Azimuthal Quantum
Number, l
* defines shape of the subshell
* Allowed values of l are integers ranging from 0 to n − 1.
Value of l
Type of subshell
0
s
1
p
2
d
3
f
So each of these letters corresponds to a shape of orbital.
*Describes the three-dimensional orientation of
the orbital.
*Values are integers ranging from -l to l:
−l ≤ ml ≤ l.
*Therefore, on any given energy level, there can
be up to:
* 1 s (l=0) orbital (ml=0),
* 3 p (l=1) orbitals, (ml=-1,0,1)
* 5 d (l=2) orbitals, (ml=-2,-1,0,1,2)
* 7 f (l=3) orbitals, (ml=-3,-2,-1,0,1,2,3)
The energy of electron only depends on the value of n
shell = all orbitals with the same value of n
subshell = all orbitals with the same value of n and l
an orbital is fully defined by three quantum numbers, n, l, and ml
Each shell of QN = n
contains n subshells
n = 1, one subshell
n= 2, two subshells, etc
Each subshell of QN = l,
contains 2l + 1 orbitals
l = 0, 2(0) + 1 = 1
l = 1, 2(1) + 1 = 3
*Value of l = 0.
*Spherical in shape.
*Radius of sphere
increases with
increasing value of n.
*Value of l = 1.
*Have two lobes with a nodal plane between
them.
Note: always 3 p orbitals for a given n
*Value of l is 2.
*2 nodal planes
*Four of the five
orbitals have 4
lobes; the other
resembles a p
orbital with a
doughnut around
the center.
Note: always 5 d orbitals for a given n.
*This leads to a fourth
quantum number, the
spin quantum number
ms.
*The spin quantum
number has only 2 values
+1/2 and -1/2
*Describes magnetic
field vector of electron
*
*No two electrons in the
same atom can have
exactly the same energy.
*For example, no two
electrons in the same
atom can have identical
sets of quantum
numbers.
*Name of each electron unique
*Name consists of four numbers:
*n,l,ml,ms
*Example:
*Mr. George Herbert Walker Bush
*We must learn to name our
electrons
*Unlike people, there is a lot in the
“name” of an electron.
*Distribution of all
electrons in an atom.
*Consist of
* Number denoting the
energy level.
* Letter denoting the type
of orbital.
* Superscript denoting the
number of electrons in
those orbitals.
*Each box represents
one orbital.
*Half-arrows represent
the electrons.
*The direction of the
arrow represents the
spin of the electron.
1s22s1
(of maximum multiplicity)
NOT:
“For degenerate orbitals,
the lowest energy is
attained when the
number of electrons with
the same spin is
maximized.”
*We fill orbitals in increasing order of energy.
*Different blocks on the periodic table, then
correspond to different types of orbitals.
* Remember:
The periodic
table was arranged the way
it was based on chemical
properties.
* Totally empirical, until now.
Based only on observation.
Short cut for writing electron configurations
Some irregularities
occur when there
are enough
electrons to half-fill
s and d orbitals on a
given row.
For instance, the
electron
configuration for
Chromium, is
[Ar] 4s1 3d5
rather than the
expected
[Ar] 4s2 3d4.
Atomic Structure
Neutral atoms have the same number of
protons and electrons.
Ions are charged atoms.
-cations – have more protons than
electrons and are positively charged
-anions – have more electrons than
protons and are negatively charged
An ion is formed when an atom, or group of atoms, has a
net positive or negative charge (why?).
If a neutral atom looses one or more electrons
it becomes a cation.
Na
11 protons
11 electrons
Na+
11 protons
10 electrons
If a neutral atom gains one or more electrons
it becomes an anion.
Cl
17 protons
17 electrons
Cl-
17 protons
18 electrons