Chapter 5 electrons in atoms
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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

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Chapter 5Electrons 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?

    • 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?

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

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

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

  • 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


Key Answer #1

  • What was inadequate about Rutherford’s atomic model?

It did not explain the chemical properties of elements


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

  • 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

  • How do sublevels of principal energy levels differ?

The sublevels are the different shapes of the electron clouds


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?


  • 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

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

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 ________________

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


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 #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


  • 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

  • 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

= 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

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


Key Answer #1

  • How are wavelength and frequency of light related?

They are inversely proportional


Key Answer #2

  • What causes atomic emission spectra?

Electrons of atoms moving from an

excited state to a lower state


Key Answer #3

  • How does quantum mechanics differ from classical mechanics?

Quantum mechanics – small objects

Classical mechanics – large objects


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