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Chapter 5 Electrons in AtomsPowerPoint Presentation

Chapter 5 Electrons in Atoms

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Chapter 5 Electrons in Atoms

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Chapter 5Electrons in Atoms

Different colors of light are

associated with the movement of

electrons

- 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?

- 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?

- Ex: why do metals change colors when

- Electrons move around a

chemical

- 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

- Ex: rungs of a ladder

- Each orbit has a __________ energy level

orbitals

fixed

- 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

- It takes less energy to

absorb

emit

- 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

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

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

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

Orbitals

get bigger

because they

can hold more

electrons

- 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

- 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

- 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

- 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

- The 3rd energy level can have 9 total orbitals between s, p, & d…but we are only looking the s orbital

- Well, we know that n = 3 and that the orbital shape = s
- 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

- 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

- # of electrons at each main level = 2n2
n = 1_____ electrons

n = 2_____ electrons

n = 3_____ electrons

n = 4_____ electrons

2

8

18

32

- What was inadequate about Rutherford’s atomic model?

It did not explain the chemical properties of elements

- What was the new assumption in the Bohr model of the atom?

An electron is found only in specific orbits

- 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

- How do sublevels of principal energy levels differ?

The sublevels are the different shapes of the electron clouds

- 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?

- 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

- 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

Three rules tell you how to find the electron

configurations of atoms

1. ______________________

2. ______________________

3. ______________________

Aufbau Principle

Pauli Exclusion Principle

Hund’s Rule

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

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

The three dumbbell orbitals = px, py, pz

3d

hold

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

opposite

Counterclockwise

Clockwise

= ________________________

= ___________________

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

- Orbitals in the same sublevel have ________ energy levels

equal

2p

x

y

z

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

2p

x

y

z

2p

x

y

z

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

- 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

- It consists of…

Number (n)

letter

Superscript

1

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

1s1

_______________

2

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

1s2

_______________

3

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

1s22s1

6

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

1s22s22p2

16

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

1s22s22p63s23p4

22

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

1s22s22p63s23p64s23d2

LETS COLOR!!!

--- s block

--- s block

--- p block

--- s block

--- p block

--- d block

- Exceptional Electron Configurations
- Some atoms do not follow the aufbau principle
Ex: Copper – 29 electrons _________________

Actually is __________________

Ex: Chromium – 24 electrons_______________

Actually is ________________

- Some atoms do not follow the aufbau principle

1s22s22p63s23p64s23d9

1s22s22p63s23p64s13d10

1s22s22p63s23p64s23d4

1s22s22p63s23p64s13d5

- 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

- What are the three rules for writing the electron configuration of elements?

Aufbau Principle

Pauli Exclusion Principle

Hund’s Rule

- Why do the actual electron configurations for some elements differ from those assigned using the aufbau principle?

They are trying to become more stable

- What is the basis for exceptions to the aufbau diagram?
Filled and half-filled orbitals are more stable than partially filled orbitals

- Key Concepts
- How does quantum mechanics differ from classical mechanics?
- How are the wavelengths and frequency of light related?
- What causes atomic emission spectra?

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

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

zero

highest

lowest

zero

zero

height

crest

- Amplitude
- a wave’s _________ from ______ to the __________

- Wavelength (λ)
- Distances between the _____________

crests

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)

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

inversely

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

Frequency increases

- 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

- 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.

= 5.88 x 10-7 m

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

Radio waves

- Which type of electromagnetic radiation has the lowest frequency?

- The different frequencies of visible light can be seen when sunlight is passed through a ________

prism

ROYGBIV

Red

- What color in the visible spectrum has the highest wavelength?

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

- Higher energy levels = ________ state

lower

excited

ground

- 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

- How are wavelength and frequency of light related?

They are inversely proportional

- What causes atomic emission spectra?

Electrons of atoms moving from an

excited state to a lower state

- How does quantum mechanics differ from classical mechanics?

Quantum mechanics – small objects

Classical mechanics – large objects