Transcript 1s 2 2s 2 2p 6 - Lawndale High School

Chapter 5
Electrons in Atoms
Ms. Wang
Lawndale High School
Section 5.1 – Models of the Atom
In 1897 J. J. Thomson discovered the electron
Observed that a magnet
deflected the straight paths of
the cathode rays
Atoms were known to be electrically neutral
which meant that there had to be some
positively charged matter
to balance the negative charges
Ernest Rutherford’s
experiment disproved the
plum pudding model of the
atom and suggested that
there was a positively
charged nucleus because
many of the alpha particles
hit the thin gold foil and
bounced back
BUT, Rutherford’s
atomic model could
not explain the
chemical properties
of elements
The Bohr Model
In 1913, Niels Bohr came up with a new
model (Bohr was a student of Rutherford)
He noticed that light given
out when atoms were
heated always had
specific amounts of
energy, so he proposed
that electrons in an
atom must be orbiting
the nucleus and can
reside only in fixed
energy levels.
Energy Levels

Energy levels – fixed energy that
an electron can have
This
is similar to steps of a ladder
Quantum
– amount of energy
required to move an electron from
one energy level to another energy
level (to be quantized)
The Quantum Mechanics View
of the Atom
The quantum
mechanical model that
scientist use today does
not describe the exact
path an electron takes
around the nucleus but
more concerned with
the probability of finding
an electron in a certain
place.
Atomic Orbitals

Atomic Orbitals – a region of space in
which there is a high probability of
finding an electron
Each
energy sublevel corresponds to
an orbital of different shape describing
where the electron is likely to be found
Labeling Electrons in Atoms

Each electron in an atom is assigned a set of
four quantum numbers. These help to determine
the highest probability of finding the electrons.
Three
of these numbers (n, l, m) give the location
of the electron
The
fourth (s) describes the orientation of an electron in
an orbital.
Quantum letters can be thought of like the
numbers and letters on a concert ticket
Labeling Electrons in Atoms
Probable Location
of e-
Probability
Probable location
of Finding Beyonce
Energy level (n)
High Probability
Hotel Floor
Sublevel (l)
Higher
Probability
Highest
probability
Wing
Orbitals (m)
Room
n= principal quantum number
Used to describe the energy of the
electron. The farther away from nucleus,
the higher the energy
The n quantum number can
have values = 1, 2, 3, …. n

n = 1 can hold 2 electrons
n = 2 can hold 8 electrons
n = 3 can hold 18 electrons
n = 4 can hold 32 electrons
Draw the electron shell diagram for Beryllium.
Be has 4 electrons
Nucleus
Be
Electrons
Draw the electron shell diagram for
Nitrogen. N has 7 electrons
N
Draw the electron shell diagrams
for these elements





Nickel
Aluminum
Argon
Carbon
Calcium

What does n represent?

How many electrons can each n hold?
l = sublevel


Provides a code for the shape of orbitals
They are designated by letters
• l =0, 1, 2, (n-1)
l
0
1
2
3
letter
s
p
d
f
Answer these questions
If n = 1 what does l =? Which letter does
that correspond to?
 If n = 2 what does l = Which letter does
that correspond to?

If n = 3 what does l =? Which letter does
that correspond to?

• If n= 4 what does l =? Which letter does
the correspond to?
Sublevels Available
Principal Energy
Level
1
1s
2
2s2p
3
3s3p3d
4
4s4p4d4f
5
5s5p5d5f5g
6
6s6p6d6f6g6h
For principal energy level 3, there are 3 sublevels
s < p< d <f in energy
m=magnetic quantum number

Used to describe each orbital within a
sublevel
Sublevel
Orbitals Available
Number or
electrons in the
sublevel
s
1=s
2
P
3 = px, py, pz
6
d
5 = dxy, dxz, dyz,
dx2 – y2, dz2
10
Section 5.2 – Electron
Configurations

Each orbital holds 2 electrons
Filling order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d,
5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p


Example He = 2 electrons
1s2
 Example Li = 3 electrons
1s22s1

Example B = 5 electrons
1s22s22p1
Practice Problems
Write electron configurations for the following atoms
1.
2.
3.
4.
Li
N
Be
C
5. P
6. Si
7. Mg
8. Al
Electron Configurations can be
written in terms of noble gases
To save space, configurations can be written
in terms of noble gases


Example 1: Ne = 1s22s22p6
S = 1s22s22p63s23p4
Or
S=
[Ne] 3s23p4
Example 2: Ar = 1s22s22p63s23p6
Mn = 1s22s22p63s23p64s23d5
Mn =
[Ar]
4s23d5
Reading the Periodic Table
Locating Electrons in Atoms
So far we have discussed 3 quantum numbers

n= principal quantum level (principal energy level)

l= Sublevel

m = magnetic quantum number (shape of orbitals)
1s2
n
Number of electrons in sublevel
l
s = spin
When an electron moves, it generates a
magnetic field.
 s describes the direction of electron spin
around its axis.


They must spin in opposite directions

Spin= up

There are two values of s: +1/2 and -1/2
down
Orbital Diagrams

The electron configuration gives the
number of electrons in each sublevel but
does not show how the orbitals of a
sublevel are occupied by the electrons.
Orbital Diagrams

They are used to show how electrons are
distributed within sublevels.
Each orbital is represented by a box and
each electron is represented by an arrow.
 The direction of the spin is represented by
the direction of the arrow
Example: Boron 1s22s22p1

1s
2s
2p
Orbital Diagrams
Steps to writing orbital diagrams:ex F (Z=9)
1.
Write the electron configuration
1s22s22p5
2. Construct an orbital filling diagram using boxes
for each orbital
1s 2s
2p
3. Use arrows to represent the electrons in each
orbital.
1s
2s
2p
Aufbau Principle


Electrons must occupy the orbital with the
lowest energy first
Example: Oxygen 1s22s22p4
1s
2s
2p
1s
2s
2p
Pauli Exclusion Principle



An atomic orbital may describe at most
two electrons
The 2 electrons must have opposite spins
Example: Oxygen 1s22s22p4
1s
2s
2p
1s
2s
2p
Hund’s Rule


Orbitals of equal energy are each occupied
by one electron before any pairing occurs
Example: Oxygen 1s22s22p4
1s
2s
2p
1s
2s
2p
Draw orbital diagrams for these
elements
1.
2.
3.
4.
Li
N
Be
C
5. P
6. Si
7. Mg
8. Al
Section 5.3 - Atomic Spectra


When atoms absorb energy, electrons
move into higher energy levels
These electrons lose energy by emitting
light when they return to lower energy
levels

Atomic Emission Spectrum – the discrete
lines representing the frequencies of light
emitted by an element
Atomic Spectra



Each discrete line in an emission spectrum
corresponds to one exact frequency of
light emitted by the atom
Ground State – lowest possible energy of
the electron in the Bohr model
The light emitted by an electron moving
from higher to a lower energy level has a
frequency directly proportional to the
energy change of the electron
Homework
Chapter 5 Assessment Page 148
#’s 22-24, 27, 29, 30-39,
50-53, 57, 60, 68, 70-72