Quantum Theory, Electrons, - PowerPoint PPT Presentation

Loading...

PPT – Quantum Theory, Electrons, PowerPoint presentation | free to download - id: 6fc2bd-YWQ0Y



Loading


The Adobe Flash plugin is needed to view this content

Get the plugin now

View by Category
About This Presentation
Title:

Quantum Theory, Electrons,

Description:

Unit 2 Quantum Theory, Electrons, & The Periodic Table Chapters 4 & 5 – PowerPoint PPT presentation

Number of Views:21
Avg rating:3.0/5.0
Slides: 55
Provided by: Stan4156
Learn more at: http://www.wtps.org
Category:

less

Write a Comment
User Comments (0)
Transcript and Presenter's Notes

Title: 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.
About PowerShow.com