Quantum Theory, Electrons,

1
Unit 2

• Quantum Theory, Electrons, The Periodic Table
• Chapters 4 5

2
Chapter 4

• Arrangement of Electrons in Atoms

3
Properties of Light
Chapter 4 Section 1 The Development of a New
Atomic Model

• Sometimes light behaves like waves, and other
  times like particles.
• Visible light is a kind of electromagnetic
  radiation, which is a form of energy that
  exhibits wavelike behavior as it travels
  through space.

4
The Electromagnetic Spectrum
Chapter 4 Section 1 The Development of a New
Atomic Model

• Together, all the forms of electromagnetic
  radiation form the electromagnetic spectrum.

5
Visible Light
Chapter 4 Section 1 The Development of a New
Atomic Model

• Visible Light is the narrow band of
  electromagnetic radiation that we can see.
• It consists of a range of waves with various
  wavelengths.

  Visible Spectrum Color Wavelength Red  
               700 - 650 nm Orange            649
- 580 nm Yellow             579 - 575
nm Green              574 - 490
nm Blue                 489 - 455 nm Indigo  
          454 - 425 nm Violet              424 -
400 nm
6
The Speed of Light
Chapter 4 Section 1 The Development of a New
Atomic Model

• The constant, c, equals the speed of light, and
  it is a fundamental constant of the universe.
• All waves in the electromagnetic spectrum
  travel atthe speed of light, c 3 x 108 m/s.

7
Properties of Light (continued)
Chapter 4 Section 1 The Development of a New
Atomic Model

• Wavelength (?) is the distance between
  corresponding points on adjacent waves.
• Frequency (?) is defined as the number of waves
  that pass a given point in a specific time,
  usually one second.

8
Wavelength vs. Frequency
Chapter 4 Section 1 The Development of a New
Atomic Model

• Wavelength (?) is inversely proportional to
  frequency (?). In other words, when ? increases,
  ? decreases,and vice versa.

9
Wavelength vs. Frequency (continued)
Chapter 4 Section 1 The Development of a New
Atomic Model

• The relationship between wavelength and frequency
  is described by the equation
• c ??
• Where c is a constant (always the same number)
  equal to 3x108 m/s.
• ? is the wavelength (in m). Problem-solving
  hint 1 nm 10-9 m.
• ? is the frequency (in s-1 or Hz).

10
The Speed of LightSample Problem
Chapter 4 Section 1 The Development of a New
Atomic Model

• A photon of light has a frequency of 4.4x1014 Hz.
  Calculate its wavelength. Does it fall within
  the visible spectrum? If so, what color is it?
• Solution
• Use the equation c ??
• 3x108m/s ? (4.4x1014Hz)

  Visible Spectrum Color Wavelength Red  
               700 - 650 nm Orange            649
- 580 nm Yellow             579 - 575
nm Green              574 - 490
nm Blue                 489 - 455 nm Indigo  
          454 - 425 nm Violet              424 -
400 nm
3 x 108m/s
?
6.8 x10-7m
4.4 x 1014Hz
? 680 x 10-9 m, or
Yes, it is red light
? 680 nm
11
The Photoelectric Effect
Chapter 4 Section 1 The Development of a New
Atomic Model

• By the early 1900s, scientists observed
  interactions of light and matter that couldnt be
  explained by wave theory.
• The photoelectric effect refers to the emission
  of electrons from a metal when light shines on
  the metal.

12
The Particle Description of Light
Chapter 4 Section 1 The Development of a New
Atomic Model

• A quantum of energy is the minimum quantity of
  energy that can be lost or gained by an atom.
• A photon is a particle of electromagnetic
  radiation having zero mass and carrying a
  quantum of energy.
• The energy of a particular photon is directly
  proportional to the frequency of the radiation.

13
Energy States
Chapter 4 Section 1 The Development of a New
Atomic Model

• Ground state The lowest energy state of an
  atom.
• Excited state an atom has a higher potential
  energy than it has in its ground state.
• When an excited atom returns to its ground
  state, it gives off energy in the form of
  electromagnetic radiation.

14
Hydrogens Line Emission Spectrum
Chapter 4 Section 1 The Development of a New
Atomic Model
15
The Bohr Model
Chapter 4 Section 1 The Development of a New
Atomic Model

• In 1913, Danish physicist Niels Bohr proposed a
  hydrogen-atom model that linked the atoms
  electron to its line-emission spectrum.

16
The Bohr Model (continued)
Chapter 4 Section 1 The Development of a New
Atomic Model

• According to the Bohr model, the electron can
  circle the nucleus only in allowed paths, or
  orbits.
• The energy of the electron is higher when it is
  in orbits that are farther from the nucleus.

17
Electrons as Waves
Chapter 4 Section 2 The Quantum Model of the
Atom

• In 1924, French scientist Louis de Broglie
  suggested that electrons act like waves confined
  to the space around an atomic nucleus.
• It followed that the electron waves could exist
  only at specific frequencies corresponding to
  the quantized energies of Bohrs orbits.

18
The Heisenberg Uncertainty Principle
Chapter 4 Section 2 The Quantum Model of the
Atom

• In 1927, German physicist Werner Heisenberg
  realized that an attemptto locate an electron
  with a photon knocks the electron off its
  course.
• The Heisenberg uncertainty principle states that
  it is impossible to determine simultaneously
  both theposition and velocity of an electronor
  any other very small particle.

19
The Schrödinger Wave Equation
Chapter 4 Section 2 The Quantum Model of the
Atom

• In 1926, Austrian physicist Erwin Schrödinger
  developedan equation that treated electrons in
  atoms as waves.
• Together with Heisenberg and others, Schrödinger
  laid the foundation for modern quantum theory.

20
Quantum Theory
Chapter 4 Section 2 The Quantum Model of the
Atom

• Quantum theory describes mathematically the wave
  properties of electrons and other very small
  particles.
• There are four different types of quantum
  numbers used
• Principal quantum (n) energy level.
• Angular momentum quantum (l) - sublevel.
• Magnetic quantum (m) - orbital.
• Spin quantum (s).

21
Principal Quantum Number
Chapter 4 Section 2 The Quantum Model of the
Atom

• The Principal Quantum Number (n) indicates the
  main energy level occupied by an electron.
• As n increases, the electrons energy and its
  distance from the nucleus increases.

22
Angular Momentum Quantum Number
Chapter 4 Section 2 The Quantum Model of the
Atom

• The Angular Momentum Quantum Number (l) (also
  called the sublevel) indicates the shape of the
  orbital.
• The number of sublevels allowed for each energy
  level is equal to n.

s orbital sphere
p orbital dumbbell
d orbital cloverleaf
f orbital complex
23
Magnetic Quantum Number
Chapter 4 Section 2 The Quantum Model of the
Atom

• The Magnetic Quantum Number (m) indicates the
  orientation of an orbital around the nucleus.

Sublevel Orbitals
s 1
p 3
d 5
f 7
24
Spin Quantum Number
Chapter 4 Section 2 The Quantum Model of the
Atom

• The Spin Quantum Number (s) indicates the
  fundamental spin state of an electron in an
  orbital.
• There are only two possible values for s, ½
  and ½.
• A single orbital can hold a maximum of two
  electrons, but the electrons must have
  opposite spin states.

25
Quantum Numbers Overview
Chapter 4 Section 2 The Quantum Model of the
Atom
26
Electron Configuration Rules
Chapter 4 Section 3 Electron Configuration

• According to the Aufbau Principle, an electron
  occupies the lowest-energy orbital that can
  receive it.
• The order of increasing energyis shown on
  thevertical axis. Eachbox represents
  anorbital. (diagram on pg. 111)

27
Electron Configuration Rules (continued)
Chapter 4 Section 3 Electron Configuration

• According to the Pauli exclusion principle, no
  two electrons in the same atom can have the
  same setof four quantum numbers.

28
Electron Configuration Rules (continued)
Chapter 4 Section 3 Electron Configuration

• According to Hunds Rule,orbitals of equal
  energy areeach occupied by one electron before
  any orbital is occupied by a second electron,
  and all electrons in singly occupied orbitals
  must have the same spin state.

Correct
Wrong
Wrong
29
Orbital Notation
Chapter 4 Section 3 Electron Configuration

• An orbital containing one electron is represented
  as
• An orbital containing two electrons is
  represented as
• The lines are labeled with the principal quantum
  number and sublevel letter. For example, the
  orbital notation for helium is written as
  follows

He
1s
30
Orbital NotationSample Problem 1
Chapter 4 Section 3 Electron Configuration

• Write the orbital notation for Carbon.
• Solution
• Carbon is atomic number 6, so it has 6 electrons.
• The first two electrons go in the 1s orbital.
• The next two electrons go in the 2s orbital.
• The final two electrons go in the 2p orbitals.

Carbon
1s
2s
2p
31
Electron Configuration Notation
Chapter 4 Section 3 Electron Configuration

• Electron-configuration notation eliminates the
  lines and arrows of orbital notation.
• Instead, the number of electrons in a sublevel is
  shown by a superscript.
• Example Carbon

32
Blocks of the Periodic Table
Chapter 4 Section 3 Electron Configuration
s
p
d
f
33
Electron Configuration NotationSample Problem 1
Chapter 4 Section 3 Electron Configuration

• Write electron configuration for Selenium (Se).
• How many unpaired electrons are in an atom of
  Selenium?
• Solution
• a.
• b. Only consider the 4p4 electrons, since all
  electrons will be paired in filled orbitals.

1s2
2s2
2p6
3s2
3p6
4s2
3d10
4p4
2 electrons are unpaired
4p
34
Noble Gas Notation
Chapter 4 Section 3 Electron Configuration

• The Group 18 elements (He, Ne, Ar, Kr, Xe, and
  Rn) are called the noble gases.
• Noble gas notation is an abbreviated electron
  configuration.
• Use square brackets around the noble gas at the
  end of the prior period to replace part of the
  configuration.
• Example Calcium

Electron Configuration
Noble Gas Notation
1s22s22p63s23p64s2
Ar4s2
35
Noble Gas NotationSample Problem 1
Chapter 4 Section 3 Electron Configuration

• Write the noble gas notation for Gold (Au).
• How many inner-shell electrons does this atom
  have?
• Solution
• a.
• The outer shell is the one with the highest .
• There are 2 e- in energy level 6 (6s2). All
  the rest are inner-shell electrons.

Xe
6s2
4f14
5d9
77 inner-shell e-
79 total e-
- 2 outer-shell e-
36
Chapter 5

• The Periodic Law

37
Mendeleev and Periodicity
Chapter 5 Section 1 History of the Periodic
Table

• The first periodic table of the elements was
  published in 1869 by Russian chemist Dmitri
  Mendeleev.
• Mendeleev left empty spaces in his table and
  predicted elements that would fill3 of the
  spaces.
• By 1886, all 3 of these elements had been
  discovered.

38
Mosley and the Periodic Law
Chapter 5 Section 1 History of the Periodic
Table

• In 1911, the English scientist Henry Moseley
  discovered that the elements fit into
  patternsbetter when they were arranged
  according to atomic number, rather than atomic
  weight.
• The Periodic Law states that the physical and
  chemical properties of the elements are periodic
  functions of their atomic numbers.

39
The Periodic Table
Chapter 5 Section 2 Electron Configuration and
the Periodic Table

• Elements in the periodic table are arranged into
  vertical columns, called groups or families, that
  share similar chemical properties.
• Elements arealso organizedhorizontally in
  rows, or periods.

40
Group 1 Alkali Metals
Chapter 5 Section 2 Electron Configuration and
the Periodic Table

• Group 1 elements are called alkali metals.
• Alkali metals have a silvery appearance and are
  soft enough to cut with a knife.
• They are extremely reactive and are not found in
  nature as free elements.
• They must be stored under oil or kerosene.

41
Group 2 Alkaline Earth Metals
Chapter 5 Section 2 Electron Configuration and
the Periodic Table

• Elements in group 2 are known as the alkaline
  earth metals.
• Group 2 metals are harder, denserand stronger
  than alkali metals, and have higher melting
  points.
• Less reactive than group 1, but still too
  reactive to be found in nature as free
  elements.

42
Group 17 Halogens
Chapter 5 Section 2 Electron Configuration and
the Periodic Table

• Elements in group 17 are known as the halogens.
• Halogens are the most reactive nonmetals,
  reacting vigorously with metals to form salts

• Most halogens exist in nature as diatomic
  molecules (i.e. F2, Cl2, Br2 and I2.)

43
Group 18 Noble Gases
Chapter 5 Section 2 Electron Configuration and
the Periodic Table

• Elements in group 18 are known as noble gases.
• They are completely non-reactive and dont form
  compounds under normalconditions.
• A new group was added to the periodic table in
  1898 for the noble gases.

44
d-block Transition Metals
Chapter 5 Section 2 Electron Configuration and
the Periodic Table

• Elements in the d-block arecalled transition
  metals.
• They have typical metallic properties such as
  conduction of electricity and high luster.
• Less reactive than group 1 and 2 elements.
• Some (i.e. platinum gold) are so unreactive
  they usually dont form compounds.

45
f-block Lanthanides Actinides
Chapter 5 Section 2 Electron Configuration and
the Periodic Table

• Elements in the period 6 of the f-block are
  called lanthanides (or rare-earth).
• Lanthanides are shiny metals similar in
  reactivity to alkaline earth metals.
• Elements in period 7 of the f-block are called
  actinides.
• Actinides are all radioactive, and many of them
  are known only as man-made elements.

46
Atomic Radii
Chapter 5 Section 3 Electron Configuration and
Periodic Properties

• Atomic radius one-half the distance between the
  nuclei of identical atoms that are bonded
  together.

Group 1

• Atomic radii tend to increase as you go down a
  group because electrons occupy successively
  higher energy levels farther away from the
  nucleus.

47
Atomic Radii (continued)
Chapter 5 Section 3 Electron Configuration and
Periodic Properties

• Atomic radii tend to decrease as you go across a
  period because as more electrons are added they
  are pulled closer to the more highly charged
  nucleus.

Period 2
48
Atomic RadiiSample Problem
Chapter 5 Section 3 Electron Configuration and
Periodic Properties

• Of the elements Mg, Cl, Na, and P, which has the
  largest atomic radius? Explain.
• Solution
• Na has the largest radius.
• All of the elements are in the 3rd period, and
  atomic radii decrease across a period.

49
Ionization Energy
Chapter 5 Section 3 Electron Configuration and
Periodic Properties

• An ion is an atom of group of bonded atoms that
  has a positive or negative charge.
• The energy required to remove an electron froma
  neutral atom of an element is called the
  ionization energy (IE).
• Ionization energy tends to increase across each
  period because a higher nuclear charge more
  strongly attracts electrons in the same energy
  level.

50
Ionization Energy (continued)
Chapter 5 Section 3 Electron Configuration and
Periodic Properties

• Ionization energy tends to decrease down each
  group because electrons farther from the nucleus
  are removed more easily.

51
Ionization EnergySample Problem
Chapter 5 Section 3 Electron Configuration and
Periodic Properties

• Consider two elements, A and B. A has an IE of
  419 kJ/mol. B has an IE of 1000 kJ/mol. Which
  element is more likely to be in the s block?
  Which will be in the p block? Which is more
  likely to form a positive ion?
• Solution
• Element A is most likely to be in the s-block
  since IE increases across the periods.
• Element B would most likely lie at the end of a
  period in the p block.
• Element A is more likely to form a positive ion
  since it has a much lower IE than B.

52
Electron Affinity and Electronegativity
Chapter 5 Section 3 Electron Configuration and
Periodic Properties

• Electron affinity is the energy change that
  occurs when an electron is acquired by a neutral
  atom.
• Electronegativity is a measure of the ability of
  an atom in a chemical compound to attract
  electrons from another atom in the compound.
• Electronegativity applies to atoms in a compound,
  while electron affinity is a property of isolated
  atoms.

53
Electron Affinity and Electronegativity
(continued)
Chapter 5 Section 3 Electron Configuration and
Periodic Properties

• Electron affinity and electronegativity both tend
  to increase across periods, and decrease (or stay
  the same) down a group.

54
ElectronegativitySample Problem
Chapter 5 Section 3 Electron Configuration and
Periodic Properties

• Of the elements Ga, Br, and Ca, which has the
  highest electronegativity? Explain .
• Solution
• All of these elements are in the fourth period.
• Br has the highest atomic number and is farthest
  to the right in the period.
• Br would have the highest electronegativity since
  electronegativity increases across a period.