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Chapter 5Electrons 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 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

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 Wing

Orbitals (m) Highest probability 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 letter

0 s

1 p

2 d

3 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?

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

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

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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

- Li 5. P
- N 6. Si
- Be 7. Mg
- C 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 down

- There are two values of s 1/2 and -1/2

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

2p

2s

1s

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

Aufbau Principle

- Electrons must occupy the orbital with the lowest

energy first - Example Oxygen 1s22s22p4

2p

2p

2s

2s

1s

1s

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

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

Draw orbital diagrams for these elements

- Li 5. P
- N 6. Si
- Be 7. Mg
- C 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