chapter 9 electrons in atoms and the periodic table · the bohr model of the atom: electron orbits...

Chapter 9

Electrons in Atoms

and the

Periodic Table

2

Waves

•Amplitude-- the height of the wave.

•Wavelength (λ) ---distance from one crest to

another

3

Frequency of a Wave

• Frequency (n) is the number of complete cycles

per given time.

Units are hertz (Hz), or cycles/s = s-1.

1 Hz = 1 s-1

• Higher Frequency– many cycles per second

• Lower Frequency--- less cycles per second

4

Low Frequency Wave

High Frequency Wave

l

l

l

5

The Nature of Light

• Light --- can be defined as a form of energy

and a form of electromagnetic waves.

• Electromagnetic waves

Has an Electrical and magnetic component

interacting perpendicularly

6

Electromagnetic Waves

• The properties of a wave are determined by:

wave speed

height (amplitude)

wavelength

frequency.

• All electromagnetic waves move through

space at the same, constant speed.

3.00 x 108 m/s = The speed of light, c.

7 Tro's "Introductory Chemistry",

Chapter 9

Electromagnetic Spectrum


Chapter 9

Types of Electromagnetic Radiation

• Classified by the Wavelength Radiowaves = l > 0.01 m.

Low frequency and energy.

Microwaves = 10-4m < l < 10-2 m.

Infrared (IR) = 8 x 10-7 < l < 10-5 m.

Visible = 4 x 10-7 < l < 8 x 10-7 m.

ROYGBIV.

Ultraviolet (UV) = 10-8 < l < 4 x 10-7 m.

X-rays = 10-10 < l < 10-8 m.

Gamma rays = l < 10-10.

High frequency and energy.

9

The Electromagnetic Spectrum • Light passed through a prism is separated into all its

colors. This is called a continuous spectrum.

• The color of the light is determined by its wavelength.


Chapter 9

Color

• The color of light is determined by its wavelength.

• White light is a mixture of all the colors of visible light.

RedOrangeYellowGreenBlueViolet.

• When an object absorbs some of the wavelengths of

white light while reflecting others, it appears colored.

The observed color is predominantly the colors reflected.


Chapter 9

Light’s Relationship to Matter

• Atoms can acquire extra energy, but

they must eventually release it.

• When atoms emit energy, it usually is

released in the form of light.

• However, atoms don’t emit all colors,

only very specific wavelengths.

Because energy is quantized


Chapter 9

Emission Spectrum


Chapter 9

Spectra

14

The Bohr Model of the Atom:

Electron Orbits • In the Bohr model, electrons travel in orbits

around the nucleus.

• Bohr’s major idea was that the energy of the atom was it is quantized.

The integer, n, is called a

quantum number.

n is directly proportional to energy


Chapter 9

The Bohr Model of the Atom:

Ground and Excited States

Ground state: The least energy state an atom can exist in.

Excited State: When an atom absorbs energy.

When the atom gains energy, the electron leaps to a higher energy orbit. The atom is less stable in an excited state and so it will release the extra energy to return to the ground state. Energy is released in the form of light.


Chapter 9

The Bohr Model of the Atom


Chapter 9

The Quantum-Mechanical Model

of the Atom

• Bohr model accurately predicts the spectrum of

hydrogen.

• However, it fails when applied to multi-electron

atoms.

• The quantum-mechanical model was proposed.

• According to this model it is possible to locate

regions in an atom where there is a higher

probability of finding an electron

• The region is called an orbital


Chapter 9

The Quantum-Mechanical Model:

Quantum Numbers

• In Schrödinger’s wave equation,

there are 3 integers, called

quantum numbers, that

quantize the energy.

• The principal quantum

number, n, specifies the main

energy level for the orbital.

19

Quantum Numbers, Continued

• The quantum-mechnical model uses parameters known

as quantum numbers to describe the structure of the

atom.

The principal quantum number designates the energy level.

• Each energy level is made of orbitals

• The quantum number that designates the orbital is often

given a letter.

s, p, d, f.

The number of orbital types = the principal quantum number.

20

Shells and Subshells

21

Orbital Types and Shapes:

s-orbital

s-orbital has a spherical shape

-has only one subshell

22

p-orbital has a dumbell shape

-has only three subshells


p-orbital

23


d-orbital

-has only five subshells

Electron Configurations

• Electron configuration-- distribution of electrons

in an atom

• Each energy shell and subshell has a maximum

number of electrons it can hold.

24

Orbital(or shell) Maximum number of

electrons

s 2

p 6

d 10

f 14

25

Filling an Orbital with Electrons

• Each orbital may have a maximum of 2 electrons.

Pauli Exclusion principle.

• When filling orbitals that have the same energy,

place one electron in each before completing pairs.

Hund’s rule.

• When filling, put electrons in lower energy orbitals first. e.g. a 1s orbital will be filled before the 2s

Aufbau principle.

• When two electrons are in the same orbital, they must have opposite spins.

26

Filling Orbitals According to Energy

Levels

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d

7s

1s orbital is the

lowest, so filling

should start there,

then 2s, 2p, 3s, 3p ….

Example:

Kr = 36 electrons = 1s22s22p63s23p64s23d104p6

27

Orbital Diagrams

• Orbital is represented as a square and the

electrons in that orbital as arrows.

Orbital with

1 electron

Unoccupied

orbital

Orbital with

2 electrons

28

Electron Configurations


Chapter 9

1. Determine the atomic number of the element

from the periodic table.

This gives the number of protons and electrons in

the atom.

Mg Z = 12, so Mg has 12 protons and 12 electrons.

Example—Write the Ground State

Orbital Diagram and Electron

Configuration of Magnesium.


Chapter 9

2. Draw 9 boxes to represent the first 3 energy

levels s and p orbitals.

1s 2s 2p 3s 3p



Configuration of Magnesium,

Continued.

31

3. Add one electron to each box in a set, then

pair the electrons before going to the next set

until you use all the electrons.

• When paired, put in opposite arrows.

1s 2s 2p 3s 3p




Continued.


Chapter 9




Continued. 4. Use the diagram to write the electron

configuration.

Write the number of electrons in each set as a

superscript next to the name of the orbital set.

1s22s22p63s2

1s 2s 2p 3s 3p

33

Example—Write the Full Ground State Electron

Configuration of Sc3+.

Sc Z = 21, therefore 21 e−

therefore Sc3+ has 18 e−

s subshell holds 2 e−

p subshell holds 6 e−

d subshell holds 10 e−

f subshell holds 14 e−

Therefore the electron configuration is 1s22s22p63s23p6

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d

7s

34

Practice—Write the Full Ground State Orbital

Diagram and Electron Configuration of Potassium.

35

Practice—Write the Full Ground State Orbital

Diagram and Electron Configuration of F−.

36

Valence Electrons

• The electrons in all the subshells with the highest principal energy shells are called the valence electrons.

• Electrons in lower energy shells are called core electrons.

• Valence electrons determine the chemical and physical properties of elements


Chapter 9

Valence Electrons, Continued

Rb = 37 electrons = 1s22s22p63s23p64s23d104p65s1

• The highest principal energy shell of Rb that contains

electrons is the 5th, therefore, Rb has 1 valence

electron and 36 core electrons.

Kr = 36 electrons = 1s22s22p63s23p64s23d104p6

• The highest principal energy shell of Kr that contains

electrons is the 4th, therefore, Kr has 8 valence

electrons and 28 core electrons.

38

Practice—Determine the Number and Types

of Electrons in an Arsenic, As, Atom.

39

Electron Configuration and the

Periodic Table • The periodic table is divided into blocks—

s-, p-, d- and f-blocks.

• For the s- and p-blocks

the period number corresponds to the principal

energy level of the valence electrons.

• For the d-block

the period number corresponds to the principal

energy level minus 1 of the valence electrons.

40

s1

s2

d1 d2 d

3 d4 d5 d6 d7 d8 d9 d10

s2 p1 p2 p3 p4 p5

p6

f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14 f14d1

1

2

3

4

5

6

7

41

Electron Configuration from

the Periodic Table

P = [Ne]3s23p3

P has five valence electrons.

3p3

P

Ne

1

2

3

4

5

6

7

1A

2A 3A 4A 5A 6A 7A

8A

3s2

Use the periodic table to write

the electron configuration of P.

42

Electron Configuration from

the Periodic Table

As = [Ar]4s23d104p3

As has five valence electrons.

4s2

Ar 3d10

4p3

As

1

2

3

4

5

6

7

1A

2A 3A 4A 5A 6A 7A

8A

43

Transition Elements • For the d-block metals, the principal energy level is one

less than the period number

Zn

Z = 30, period 4, group 2B

[Ar] 4s23d10

Sn

Z = 50, period 5, group 4B

[Kr] 5s24d105p2

44

Practice—Use the Periodic Table to Write the Short

Electron Configuration of Na and Te

• Na (at. no. 11).

• Te (at. no. 52).

• Ca

• Rb


Chapter 9

Practice Cont’d—Use the Periodic Table to Write

the Short Electron Configuration of Ca and Rb

• Cl

• Ba


Chapter 9

Practice Cont’d—Use the Periodic Table to Write

the Short Electron Configuration of Cl and Ba

47

Periodic Trends in Properties of

Elements

• Effective nuclear charge

• Atomic size (atomic radius)

• Ionization energy

• Metallic character

Two important questions to be answered are:

• What happens when you go across a period (left to right)?

• What happens when you go down the group (top to bottom)?

48

Trends in Effective Nuclear Charge

• Effective nuclear charge--- The net positive charge experienced by an electron in a mutielectron atom

• Important Questions!!

What happens to the effective nuclear charge as:

- we go down a group?

-we go across a period?

49

4 p+

2e-

2e-

12 p+

2e-

8e-

2e-

Be (4p+ and 4e-)

going down a group

e.g., Group IIA

Mg (12p+ and 12e-)

Ca (20p+ and 20e-)

16 p+

2e-

8e-

2e-

8e-

Effective nuclear

charge decreases

down the group!!

50

Li (3p+ and 3e-)

What happens to effective nuclear charge as we move across

the period?

Be (4p+ and 4e-) B (5p+ and 5e-)

6 p+

2e-

4e-

C (6p+ and 6e-)

8 p+

2e-

6e-

O (8p+ and 8e-)

10 p+

2e-

8e-

Ne (10p+ and 10e-)

2e-

1e-

3 p+

2e-

2e-

4 p+

2e-

3e-

5 p+

Effective nuclear charge increases across the period!!

51

Trends in Atomic Size (Atomic Radius)

• atomic radius increases down group

valence shell farther from nucleus

effective nuclear charge drops

• atomic radius decreases across period (left to right)

electrons added to same valence shell

effective nuclear charge increases

52

4 p+

2e-

2e-

12 p+

2e-

8e-

2e-

Be (4p+ and 4e-)

Group IIA

Mg (12p+ and 12e-)

Ca (20p+ and 20e-)

16 p+

2e-

8e-

2e-

8e-


Chapter 9

Li (3p+ and 3e-)

Period 2

Be (4p+ and 4e-) B (5p+ and 5e-)

6 p+

2e-

4e-

C (6p+ and 6e-)

8 p+

2e-

6e-

O (8p+ and 8e-)

10 p+

2e-

8e-

Ne (10p+ and 10e-)

2e-

1e-

3 p+

2e-

2e-

4 p+

2e-

3e-

5 p+

54

Trends in Atomic Size, Continued


Chapter 9

Example—Choose the

Larger Atom in Each Pair.

1. N or F

2. C or Ge

3. N or Al


Chapter 9

1. N or F

2. C or Ge

3. N or Al

1. N or F

2. C or Ge, Ge is further down

1. N or F

2. C or Ge

3. N or Al, Al is further down & left

1. N or F, N is further left


Larger Atom in Each Pair, Continued.

57

Ionization Energy

• Minimum energy needed to remove an electron

from an atom.

Gas state.

Endothermic process.

Valence electron easiest to remove.

M(g) + 1st IE M1+(g) + 1 e-

M+1(g) + 2nd IE M2+(g) + 1 e-

First ionization energy = energy to remove electron from

neutral atom; 2nd IE = energy to remove from +1 ion; etc.

58

General Trends in

Ionization Energy

• The larger the effective nuclear charge on the electron, the more energy it takes to remove it.

• The farther the electron is from the nucleus, the less energy it takes to remove it.

• First IE generally increases across the period.

effective nuclear charge increases

• First IE decreases down the group.

valence electron farther from nucleus


Chapter 9

Trends in Ionization Energy, Continued

60

1. Al or S

2. As or Sb, As is further up

1. Al or S

2. As or Sb

3. N or Si, N is further up

1. Al or S

2. As or Sb

3. N or Si

Example—Choose the Atom in Each Pair

with the Higher First Ionization Energy

1. Al or S, S is further right

61

Metallic Character • How well an element’s properties match the

general properties of a metal.

• Metals: Malleable and ductile as solids. shiny, lustrous, and reflect light. conduct heat and electricity. Form cations in solution. Lose electrons in reactions—oxidized.


Chapter 9

Metallic Character, Continued

• In general, metals are found on the left of

the periodic table and nonmetals on the

right.

• As you traverse left to right across the

period, the elements become less metallic.

• As you traverse down a column, the

elements become more metallic.


Chapter 9

Trends in Metallic Character


Chapter 9

1. Sn or Te

2. P or Sb, Sb is further down

1. Sn or Te

2. P or Sb

3. Ge or In, In is further down & left

1. Sn or Te

2. P or Sb

3. Ge or In

1. Sn or Te, Sn is further left


More Metallic Element in Each Pair

65

Recommended Study Problems Chapter 9

NB: Study problems are used to check the student’s understanding

of the lecture material. Students are EXPECTED TO BE ABLE

TO SOLVE ALL THE SUGGESTED STUDY PROBLEMS.

If you encounter any problems, please talk to your professor or seek

help at the HACC-Gettysburg learning center.

Questions from text book Chapter 9

5, 7, 13, 25, 27, 29, 33,37, 39, 43, 45, 49, 51, 55, 57, 61, 63, 65, 71,

77, 79, 85, 91, 95, 99, 101, 105, 107

ANSWERS

-The answers to the odd-numbered study problems are found at

the back of your textbook

chapter 9 electrons in atoms and the periodic table · the bohr model of the atom: electron orbits...

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