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Chapter 5 Electrons in Atoms

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Title: Chapter 5 Electrons in Atoms


1
Chapter 5Electrons in Atoms
  • Ms. Wang
  • Lawndale High School

2
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
3
  • Atoms were known to be electrically neutral which
    meant that there had to be some positively
    charged matter
  • to balance the negative charges

4
  • Ernest Rutherfords 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, Rutherfords atomic model could not explain
the chemical properties of elements
5
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.

6
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)

7
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.

8
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

9
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.

10
Quantum letters can be thought of like the
numbers and letters on a concert ticket
11
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 Wing
Orbitals (m) Highest probability Room
12
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
13
  • 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
14
Draw the electron shell diagrams for these
elements
  • Nickel
  • Aluminum
  • Argon
  • Carbon
  • Calcium
  • What does n represent?
  • How many electrons can each n hold?

15
l sublevel
  • Provides a code for the shape of orbitals
  • They are designated by letters
  • l 0, 1, 2, (n-1)

l letter
0 s
1 p
2 d
3 f
16
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?

17
Principal Energy Level Sublevels Available
1 1s
2 2s2p
3 3s3p3d
4 4s4p4d4f
5 5s5p5d5f5g
6 6s6p6d6f6g6h
For principal energy level 3, there are 3
sublevels
s lt plt d ltf in energy
18
mmagnetic quantum number
  • Used to describe each orbital within a sublevel

Sublevel Orbitals Available
s 1 s
P 3 px, py, pz
d 5 dxy, dxz, dyz, dx2 y2, dz2
Number or electrons in the sublevel
2
6
10
19
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20
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

21
Practice Problems
Write electron configurations for the following
atoms
  • Li 5. P
  • N 6. Si
  • Be 7. Mg
  • C 8. Al

22
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

23
Reading the Periodic Table
24
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
25
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 down
  • There are two values of s 1/2 and -1/2

26
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.

27
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
2p
2s
1s
28
Orbital Diagrams
  • Steps to writing orbital diagramsex F (Z9)
  • Write the electron configuration
  • 1s22s22p5
  • 2. Construct an orbital filling diagram using
    boxes for each orbital
  • 3. Use arrows to represent the electrons in each
    orbital.

2p
2s
1s
2p
2s
1s
29
Aufbau Principle
  • Electrons must occupy the orbital with the lowest
    energy first
  • Example Oxygen 1s22s22p4

2p
2p
2s
2s
1s
1s
30
Pauli Exclusion Principle
  • An atomic orbital may describe at most two
    electrons
  • The 2 electrons must have opposite spins
  • Example Oxygen 1s22s22p4

2p
2p
2s
2s
1s
1s
31
Hunds Rule
  • Orbitals of equal energy are each occupied by one
    electron before any pairing occurs
  • Example Oxygen 1s22s22p4

2p
2p
2s
2s
1s
1s
32
Draw orbital diagrams for these elements
  • Li 5. P
  • N 6. Si
  • Be 7. Mg
  • C 8. Al

33
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

34
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

35
Homework
  • Chapter 5 Assessment Page 148
  • s 22-24, 27, 29, 30-39,
  • 50-53, 57, 60, 68, 70-72
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