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Chapter 5 Electrons in Atoms. Different colors of light are associated with the movement of electrons. 5.1 Models of the Atom. Key Concepts What was inadequate about Rutherford’s atomic model? What was the new assumption in the Bohr model of the atom?

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chapter 5 electrons in atoms

Chapter 5Electrons in Atoms

Different colors of light are

associated with the movement of

electrons

5 1 models of the atom
5.1 Models of the Atom
  • Key Concepts
    • What was inadequate about Rutherford’s atomic model?
    • What was the new assumption in the Bohr model of the atom?
    • What does the quantum mechanical model determine about the electrons in an atom?
    • How do sublevels of principal energy levels differ?
slide3
Rutherford (1897)
    • Electrons move around a

nucleus made of protons and neutrons

    • Could not explain the _______________ properties of elements
      • Ex: why do metals change colors when

heated?

chemical

slide4
Bohr (1913) proposed that

electrons are found in specific

paths, or _________, around a nucleus

    • Each orbit has a __________ energy level
      • Ex: rungs of a ladder

The higher an electron is on the

ladder, the farther it is away from

the nucleus

orbitals

fixed

slide5
Quantum  the amount of energy needed to move an electron from one energy level to another
    • Moving up energy levels = ____________ energy
    • Moving down energy levels = _______ energy (light)
    • Energy levels are not equally spaced
      • It takes less energy to

move between higher rungs

or energy levels than lower

ones

absorb

emit

slide6
Bohr’s model only worked for hydrogen
  • De Broglie (1923) proposed that small particles, like electrons, have _____________ properties
  • Schrodinger (1926) devised a mathematical calculation to describe the wave-like movement of electrons
  • _________________________
    • Determines the allowed energies an electron can have and how _________ it is to find the electron at various locations around the nucleus

wave-like

Quantum Mechanical Model

likely

slide7

Electron may be found

anywhere within the

electron cloud

Movement of electron is similar

to movement of a propeller blade

Blade may be anywhere in the

blurry region

slide8

principle quantum numbers

  • The energy levels in the quantum mechanical model are labeled by _____________________________ (n)
    • n = 1, 2, 3, 4, 5, 6…
  • Each principle energy level has sublevels that correspond to different cloud shapes, called _______________________

Name: Principle Quantum Number

Symbol: n

Represents: energy levels

Values: any integer ≥ 1

atomic orbitals

slide9

spherical

4 Different Atomic Orbitals

  • S orbital = _____________ (1)
  • P orbitals = __________________ (3)

3. D orbitals = ____________________ (5)

4. F orbitals = ___________________ (7)

dumbbells

four leaf clovers

Complex shapes

slide10

Orbitals

get bigger

because they

can hold more

electrons

slide11

Notice that n = # of sublevels and # of different shapes

      • The order of the shapes moves from lowest energy to highest energy (s, p, d, f…)
    • Example: n = 2; sublevel shapes = s & p
    • Name the shape(s) be for the following:
      • n = 1
      • n = 3
      • n = 4

s

s, p, d

s, p , d, f

slide12

Notice the number of orbitals per shape

    • s = 1 - d = 5
    • p = 3 - f = 7
  • If n = 3, then we will have 3 different shapes (s, p & d)
  • The s shape has 1 orbital available
  • The p shape has 3 orbitals available
  • And, the d shape has 5 orbitals available
  • How many orbitals does the 3rd energy level have?

It has 9 orbitals

slide13

What is the relationship between n = 3 and 9 available orbitals?

    • 9 = 32
    • So, we can say that n2 = # of orbitals available
  • We represent each orbital by listing the energy level followed by the orbital shape

1s

n = 1, orbital shape = sphere

3d

n = 3, orbital shape = clover

slide14

How many orbitals are in the 3s sublevel?

    • Well, we know that n = 3 and that the orbital shape = s
      • The 3rd energy level can have 9 total orbitals between s, p, & d…but we are only looking the s orbital

 There is only 1 orbital in the 3s sublevel

  • Each orbital available can hold 2 electrons
    • s = 1 orbital = 2 electrons
    • p = 3 orbitals = 6 electrons
    • d = 5 orbitals = 10 electrons
    • f = 7 orbitals = 14 electrons
slide15

So, when n = 4, we can have…

    • s = 1 orbital = 2 electrons
    • p = 3 orbitals = 6 electrons
    • d = 5 orbitals = 10 electrons
    • f = 7 orbitals = 14 electrons

32 total electrons = 2n2

slide16
# of electrons at each main level = 2n2

n = 1 _____ electrons

n = 2 _____ electrons

n = 3 _____ electrons

n = 4 _____ electrons

2

8

18

32

key answer 1
Key Answer #1
  • What was inadequate about Rutherford’s atomic model?

It did not explain the chemical properties of elements

key answer 2
Key Answer #2
  • What was the new assumption in the Bohr model of the atom?

An electron is found only in specific orbits

key answer 3
Key Answer #3
  • What does the quantum mechanical model determine about the electrons in an atom?

The likelihood or probability of finding electrons in

various locations around the nucleus

key answer 4
Key Answer #4
  • How do sublevels of principal energy levels differ?

The sublevels are the different shapes of the electron clouds

5 2 electron arrangement in atoms
5.2 Electron Arrangement in Atoms
  • Key Concepts
    • What are the three rules for writing the electron configurations of the elements?
    • Why do actual electron configurations for some elements differ from those assigned using the aufbau principle?
slide22
Unstable arrangements become more stable by ____________ energy
  • Does the rock formation in the picture look stable? _______
  • What will it eventually do to become more stable?

losing

No

It will fall or rearrange

slide23
In an atom, _________ and the ___________ interact to make the most stable arrangement possible
  • The ways in which electrons are arranged around the nuclei of atoms are called __________ ____________________

electrons

nucleus

electron

configurations

slide24
Three rules tell you how to find the electron

configurations of atoms

1. ______________________

2. ______________________

3. ______________________

Aufbau Principle

Pauli Exclusion Principle

Hund’s Rule

slide25

lowest

  • Aufbau Principle – electrons occupy the orbitals of __________ energy first

Each box represents

an orbital

Will electrons occupy the 2s orbital or the 2p orbital

first? __________

2s

slide26
What do the three boxes by 2p represent?
  • Which is higher, 3d or 4s?

The three dumbbell orbitals = px, py, pz

3d

slide27

hold

  • Pauli Exclusion Principle – an atomic orbital may ________ at most two electrons that have ____________ spins

opposite

Counterclockwise

Clockwise

= ________________________

= ___________________

slide28

same

  • Hund’s Rule – electrons occupy orbitals of the same energy in a way that makes the number of electrons with the ______ spin direction as large as possible
    • Orbitals in the same sublevel have ________ energy levels
      • They all must have electrons with the same spin first
      • Practice – Put three electrons into the 2p sublevel

equal

2p

x

y

z

slide29
Practice – put four electrons into the 2p sublevel
  • Put five electrons into the 2p sublevel

2p

x

y

z

2p

x

y

z

slide30

unpaired

  • If an orbital only has one electron in its box, it is called ______________

Ex: How many unpaired electrons are in the following sublevel?

4

3d

xy

xz

yz

x2-y2

z2

slide31
We use all three rules to write a shorthand method for showing the electron configurations of an atom
    • It consists of…
      • The _____________ of the main energy levels
      • The ____________ of the atomic orbital
      • _____________ indicating the # of electrons

Number (n)

letter

Superscript

slide32

1

  • Practice: Hydrogen
    • How many electrons does hydrogen have? ___

1s1

_______________

slide33

2

  • Practice: Helium
    • How many electrons does helium have? ____

1s2

_______________

slide34

3

  • Practice: Lithium
    • How many electrons does lithium have? ____

1s22s1

slide35

6

  • Practice: Carbon
    • How many electrons does carbon have? ____

1s22s22p2

slide36

16

  • Practice: Sulfur
    • How many electrons does sulfur have? ____

1s22s22p63s23p4

slide37

22

  • Practice: Titanium
    • How many electrons does titanium have? ____

1s22s22p63s23p64s23d2

slide40

--- s block

--- p block

slide41

--- s block

--- p block

--- d block

slide42
Exceptional Electron Configurations
    • Some atoms do not follow the aufbau principle

Ex: Copper – 29 electrons _________________

Actually is __________________

Ex: Chromium – 24 electrons_______________

Actually is ________________

1s22s22p63s23p64s23d9

1s22s22p63s23p64s13d10

1s22s22p63s23p64s23d4

1s22s22p63s23p64s13d5

slide43
These atoms are trying to become more ______________ by having sublevels that are least ______ full.
  • Filled and half-filled energy sublevels are more stable than partially-filled energy sublevels.

stable

half

key answer 11
Key Answer #1
  • What are the three rules for writing the electron configuration of elements?

Aufbau Principle

Pauli Exclusion Principle

Hund’s Rule

key answer 21
Key Answer #2
  • Why do the actual electron configurations for some elements differ from those assigned using the aufbau principle?

They are trying to become more stable

slide46
What is the basis for exceptions to the aufbau diagram?

Filled and half-filled orbitals are more stable than partially filled orbitals

5 3 physics and the quantum mechanical model
5.3 Physics and the Quantum Mechanical Model
  • Key Concepts
    • How does quantum mechanics differ from classical mechanics?
    • How are the wavelengths and frequency of light related?
    • What causes atomic emission spectra?
slide48

light

waves

  • The quantum mechanical model grew out of the study of ___________, which also moves in ___________
  • If all objects have a wavelike motion, why can’t we see this for ordinary objects like baseballs or trains?
  • Classical Mechanics – describes the motion of __________ objects
  • Quantum Mechanics – describes the motion of ________________ particles and __________

Mass has to be very small in order to detect wavelength

large

subatomic

atoms

slide49
Complete Wave Cycle
    • Starts at _________
    • Moves to its ___________ point and __________ point
    • Returns to ___________

zero

highest

lowest

zero

slide50

zero

height

crest

  • Amplitude
    • a wave’s _________ from ______ to the __________
  • Wavelength (λ)
    • Distances between the _____________

crests

slide51

time

  • Frequency (v)
    • Number of wave cycles to pass a certain point per unit of _________
    • Unit = ______________________
      • (SI) = ____________ or ______

Cycles per second

Hertz (Hz)

S-1 (1/s)

slide52

Speed of light

  • The product of wavelength (λ) and frequency (v) always equals a constant (c), the ____________________

c = 2.998 x 108 m/s

λ x v = c

slide53

inversely

  • Wavelength (λ) and frequency are ____________ proportional
  • What happens to frequency when wavelength decreases?

Frequency increases

slide54
Sample Problem 5.1
    • Calculate the wavelength (λ) of the yellow light emitted by the sodium lamp shown above if the frequency of the radiation is 5.10 x 1014 Hz.

c = λv

c = λ

v

2.998 x 108 m/s

5.10 x 1014/s

= 5.88 x 10-7 m

slide55

electromagnetic

  • According to the wave model, light consists of ___________________ waves
  • Electromagnetic radiation includes…
    • _____________________
    • _____________________
    • _____________________
    • _____________________
    • _____________________
    • _____________________
    • _____________________
    • ______________________

Radio waves

Radar

Microwaves

Infared

Visible light

Ultraviolet light

X - Rays

G amma rays

slide56

Radio waves

  • Which type of electromagnetic radiation has the lowest frequency?
slide57
The different frequencies of visible light can be seen when sunlight is passed through a ________

prism

ROYGBIV

slide58

Red

  • What color in the visible spectrum has the highest wavelength?
slide59

higher

  • Atoms emit light when their electrons move from a __________ energy level to a _________ energy level
      • Higher energy levels = ________ state
        • n = 2, 3, 4, 5, …
      • Lowest energy level = _________ state
        • n = 1

lower

excited

ground

slide60
The frequencies of light emitted by each element that separate into distinct lines is called its ____________________________
    • Unique to each element – like a ______________

Atomic emission spectra

fingerprint

Each gas glows a

different color

key answer 12
Key Answer #1
  • How are wavelength and frequency of light related?

They are inversely proportional

key answer 22
Key Answer #2
  • What causes atomic emission spectra?

Electrons of atoms moving from an

excited state to a lower state

key answer 31
Key Answer #3
  • How does quantum mechanics differ from classical mechanics?

Quantum mechanics – small objects

Classical mechanics – large objects

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