Transcript Unit 1 Section 4 - Atomic Structure PPT

Unit 1 Section 4
Atomic
Structure
Chemical Foundations:
Elements, Atoms, and
Ions
Textbook - Chapter 4 & 11 Table of Contents
4.1
4.2
4.4
4.5
4.6
4.7
11.1
11.2
11.3
11.4
11.5
11.6
11.7
11.8
11.9
11.10
The Elements
Symbols for the Elements
Formulas of Compounds
The Structure of the Atom
Introduction to the Modern Concept of Atomic Structure
Isotopes (Calculating Average Atomic Mass)
Rutherford’s Atom
Electromagnetic Radiation and Energy
Emission
The Energy Levels of Hydrogen
The Bohr Model of the Atom
The Wave Mechanical Model
The Orbitals
The Wave Mechanical Model: Further Development
Electron Arrangements
Electron Configurations
2
Chapter 4
Table of Contents
Objectives
• To learn how a formula describes a compound’s
composition.
• To study the atom’s structure
3
Section 4.1
The Elements
•
•
Substances that cannot be broken down by
simple chemical means
118 known: 88 found in nature, others are man
made.
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4
Section 4.1
The Elements
Element Abundance
Most abundant elements in the universe:
hydrogen H
60%
helium
He 37%
Most abundant elements in the entire earth:
iron
Fe
35%
silicon
oxygen
O
30%
Si
Most abundant element in earth’s crust:
oxygen
O
49.2%
aluminum
silicon
Si
25.7%
Al
15%
7.50%
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5
Section 4.1
The Elements
Element Abundance
Naturally radioactive elements:
uranium
U
radium
Ra
radon
Rn
Polonium Po
Elements most abundant in human body:
oxygen
O
nitrogen
N
carbon
C
phosphorus
P
hydrogen H
calcium
Ca
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6
Section 4.1
The Elements
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7
Section 4.1 & Section 4.9
The Elements
Diatomic Elements
Diatomic elements – elements that occur naturally
paired as two atoms per molecule
- 7 diatomic elements
H2
Cl2
F2
N2
Br2
O2
I2
GEN-U-INE
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8
Section 4.2
Symbols for the Elements
•
•
•
Each element has a unique one- or two-letter symbol.
First letter is always capitalized and the second is not.
The symbol usually consists of the first one or two
letters of the element’s name.

•
Examples:
Oxygen
Krypton
O
Kr
Sometimes the symbol is taken from the element’s
original Latin or Greek name.

Examples:
Gold Au
Lead Pb
aurum
plumbum
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9
Section 4.2
Symbols for the Elements
Names and Symbols of the Most Common Elements
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10
Section 4.4
Formulas of Compounds
Chemical Formulas Describe Compounds
•
Compound – distinct substance that is
composed of the atoms of two or more
elements and always contains exactly the
same relative masses of those elements.
•
Chemical Formulas – expresses the types of
atoms and the number of each type in each
unit (molecule) of a given compound.
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11
Section 4.4
Formulas of Compounds
Rules for Writing Formulas
1. Symbol  Tells which atoms are present in
compounds
2. Subscript  The number of each type of atom
3. Coefficient  (number in front) tells the number of
molecules
Coefficient
4
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12
Section 4.3
Atomic
Dalton’sTheory
Atomic Theory
Scientists of the 18th century learned that:
1. Most natural materials are mixtures of pure
substances.
2. Pure substances are either elements or
combinations of elements called compounds.
3. A given compound always contains the same
proportions (by mass) of the elements.
 Law of Constant Composition
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13
Section 4.3
Atomic
Dalton’sTheory
Atomic Theory
Law of Constant Composition
•
A given compound always has the same
composition, regardless of where it comes
from.

Water always contains 8 g of oxygen for every 1 g of
hydrogen.
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14
Section 4.5
4.3
The
Structure
of Theory
the Atom
Dalton’s
Atomic
Objectives
• Recognize that science is a progressive endeavor that
reevaluates and extends what is already known.
• Compare and contrast historical models of the atom.
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Section 4.5
4.6
The
Structure of
Introduction
tothe
theAtom
Modern Concept of Atomic Structure
The atom contains:
•
Electrons – found outside
the nucleus; negatively
charged (-1)
•
Protons – found in the
nucleus; positive charge
(+1)
•
Neutrons – found in the
nucleus; no charge
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16
Section 4.5
Section
4.1
The
of the Atom
TheStructure
Elements
William Thomson (Plum Pudding Model)
• Plum pudding model

A physicist, Thompson, believed an
atom was filled with positively charged
material and the electrons were evenly
distributed throughout.
- This model of the atom turned out to be
short-lived, due to the work of Ernest
Rutherford (1871–1937).
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17
4.2
Section 4.5
4.1
The Elements
Ernest Rutherford’s Portrait
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Section 4.5
4.1
Ernest
Rutherford
The Elements
• Discovered the nucleus (1911)
• Stated protons were inside of the nucleus
• Gold-Foil Experiment
 Radioactive alpha particles “shot” through gold foil
- A majority of particles passed straight through foil
- A small fraction of particles bounced off gold foil at
large angles or bounced straight back
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Section 4.5
Section
4.5
4.1
The Elements
Rutherford’s
alpha particle scattering experiment.
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5.5
20
Section 4.5
4.1
Rutherford’s
Atomic Model
The Elements
Rutherford discovered that:
• The atom contains a tiny dense center called the
nucleus
– the volume is about 1/10 trillionth the volume of the
atom
• The nucleus is positively charged
– the amount of positive charge of the nucleus balances
the negative charge of the electrons
• The electrons move around in the empty space of
the atom surrounding the nucleus
21
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Copyright©2004 by Houghton
Section 4.5
4.1
Rutherford’s
possible model
The Elements
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22
Section 4.6
Introduction to the Modern Concept of Atomic Structure
Why do different atoms have different chemical
properties?
•
•
•
The chemistry of an atom arises from its
electrons.
Electrons are the parts of atoms that
“intermingle” when atoms combine to form
molecules.
It is the number of electrons that really
determines chemical behavior.
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23
Section 4.6
Introduction to the Modern Concept of Atomic Structure
Rutherford Atomic Model
Rutherford’s model turned out to be incomplete.
•
The Rutherford atomic model had to be revised in order
to explain the chemical properties of elements.
Return to TOC
Section 4.6
Introduction to the Modern Concept of Atomic Structure
The Modern Atom
• We know atoms are composed of three main
pieces - protons, neutrons and electrons
• The nucleus contains protons and neutrons
• The nucleus is only about 10-13 cm in diameter
• The electrons move outside the nucleus with an
average distance of about 10-8 cm
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25
Section 4.6
Introduction to the Modern Concept of Atomic Structure
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26
Section 4.3
Creating
Dalton’sModels
Atomic Theory
1) Create a drawing that illustrates Thomson’s view
of an atom.
2) Based on what you learned about atoms in other
science classes, create a diagram that represents
the structure of an atom.
3) Explain how your two drawings are different.
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Section 4.7
Isotopes
Objective
• To learn about atomic number, mass number, and
isotopes.
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28
Section 4.7
Isotopes
•Why are atoms with different numbers of neutrons still
considered to be the same element?
•Despite differences in the number of neutrons,
isotopes are chemically alike. They have
identical numbers of electrons, which determine
chemical behavior.
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Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.
Section 4.7
Isotopes
Atomic number and Mass
Atomic number – the number of protons in the
nucleus
Mass number = protons + neutrons
•The mass listed in the periodic table is the
average atomic mass
•It is a weighted average of the atomic masses of
naturally occurring isotopes
Mass Number
Atomic Number
A
ZX
Element Symbol
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30
Section 4.7
Isotopes
Atomic number and Mass
• Remember that atoms are electrically neutral.
 In an atom, protons = electrons
• Protons, neutrons, and electrons can be
calculated from atomic number and mass
number.
How many protons, electrons, and neutrons
are in each atom?
boron and sodium
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31
Section 4.7
Isotopes
Isotopes
•
•
•
1
1H
Isotopes - atoms with the same number of
protons but different numbers of neutrons.
Show almost identical chemical properties;
chemistry of atom is due to its electrons.
In nature most elements contain mixtures of
isotopes.
2
1H
(D)
3
1H
(T)
235
92
U
238
92
U
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32
Section 4.7
Isotopes
A
Z
•
•
•
X
X = the symbol of the element
Z = the atomic number (# of protons)
A = the mass number or atomic mass unit (amu)
(# of protons and neutrons)
A – Z = n (number of neutrons)
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33
Section 4.7
Isotopes
Two Isotopes of Sodium
A – Z = n (number of neutrons)
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34
Section 4.7
Isotopes
Isotopes – An Example
14
6
•
•
•
C
C = the symbol for
carbon
6 = the atomic number
(6 protons)
14 = the mass number
(6 protons and 8
neutrons)
12
6
C
• C = the symbol for
carbon
• 6 = the atomic number
(6 protons)
• 12 = the mass number
(6 protons and 6
neutrons)
A – Z = n (number of neutrons)
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35
Section 4.7
Isotopes
Exercise
A certain isotope X contains 23 protons and 28
neutrons.
• What is the mass number of this isotope?
• Identify the element.
Mass Number = 51
Vanadium
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36
Section 4.7
Isotopes
Atomic Mass Unit – How is it calculated?
• The value shown in the periodic table is the average
atomic mass
 It is a weighted average of the atomic masses of
naturally occurring isotopes
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37
Section 4.7
Isotopes
Atomic Mass Unit – How is it calculated?
• For example: Chlorine has two isotopes
 Chlorine-35 and Chlorine-37
 The abundance is:
 Cl-35 has an amu of 34.9689 with an abundance
of 75.771%
 Cl-37 has an amu of 36.9659 with an abundance
of 24.229%
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38
Section 4.7
Isotopes
Calculating Atomic Mass Unit
• Cl-35 has an amu of 34.9689 with an abundance
of 75.771%
• Cl-37 has an amu of 36.9659 with an abundance
of 24.229%
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39
Section 4.7
Isotopes
Mass Spectroscopy
• an analytical tool used for measuring the molecular
mass of a sample
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40
Section 4.7
Isotopes
Mass Spectrum for the element Boron
• The number of isotopes: The two peaks in the mass spectrum
shows that there are 2 isotopes of boron – with relative isotopic
masses of 10 and 11
• The abundance of the isotopes: Can be determined by the height of
the peak.
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41
Section 4.7
Isotopes
Chapter
11
Modern
Atomic
Theory
Copyright© by Houghton Mifflin Company. All rights reserved.
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Section 4.7
Electrons
in
Atoms
Isotopes
Rutherford’s model has some limitations
It did not explain the chemical properties of the elements
It did not address the electrons
• What are electrons doing? – How are they arranged
& how do they move?
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43
Section 4.7
Electrons
in
Atoms
Isotopes
Figure 11.1: The Rutherford atom.
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44
Section 4.7
Electrons
in
Atoms
Isotopes
• To understand the next development, we must
understand some properties of light..
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45
Section 4.7
11.2 Electromagnetic Radiation and Energy
Isotopes
By the year 1900, there was enough experimental evidence
to convince scientists that light consisted of waves.
• The wavelength, represented by (the Greek letter
lambda), is the distance between two wave peaks.
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46
Section 4.7
11.2 Electromagnetic Radiation and Energy
Isotopes
• The frequency, represented by (the Greek letter nu),
is the number of wave peaks that pass a certain point
per unit of time.
• The SI unit of waves per second is called the hertz (Hz).
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47
Section 4.7
Isotopes
• The frequency ( ) and wavelength ( ) of light are
inversely proportional to each other.
• As the wavelength increases, the frequency decreases.
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48
Section 4.7
Isotopes
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Section 4.7
Isotopes
Rutherfo
https://www.youtube.com/watch?v=cfXzwh3KadE
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50
Section 4.7
11.2 Electromagnetic Radiation and Energy
Isotopes
According to the wave model, light consists of
electromagnetic waves.
Electromagnetic radiation - a form of energy that
exhibits wavelike behavior as it travels through
space.
– All electromagnetic radiation travels at the
speed of light:
c = 3.0 X108 m/s
Electromagnetic radiation includes radio waves,
microwaves, infrared waves, visible light,
ultraviolet waves, X-rays, and gamma rays.
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51
Section 4.7
Isotopes
Visible light of different wavelengths can be
separated into a spectrum of colors.
In the visible spectrum, red light has the longest
wavelength and the lowest frequency.
Violet light has the shortest wavelength and the
highest frequency.
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52
Section 4.7
Wave
Description
of
Light
Isotopes
• Equation relating frequency and
wavelength:
c = 
c = speed of light
 = wavelength
 = frequency
• c is constant, so is , so as frequency
increases, wavelength decreases
(inversely proportional).
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53
Light
as
a
Wave:
Problems
Isotopes
Section 4.7
1) What is the frequency of light if its
wavelength is 4.34 X 10-7 m?
2) What is the wavelength of a wave with
a frequency of 1019 Hz (s-1)?
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54
Section 4.7
Energy
is
Quantized
Isotopes
• In 1900 a physcist named Max Planck proposed
that matter does not emit electromagnetic
energy continuously.
• Max Planck suggested that the object emits
energy in small specific amounts called quanta.
• Quantum - the minimum quantity of energy that
can be lost or gained by an atom.
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55
56
Section 4.7
Light
as
a
Particle:
Isotopes
Planck’s Equation
Energy (Joules) of
a quantum of
radiation
Frequency (Hz)
E = h
Plank’s Constant = 6.626 x 10-34 J•s
Video - http://study.com/academy/lesson/what-is-a-photon-definition-energy-wavelength.html
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Section 4.7
Einstein
Isotopes
• Next Albert Einstein suggested that the
electromagnetic (em) spectrum is itself quantized.
• Einstein proposed that em radiation can be
viewed as photons
• Photon - a particle of electromagnetic radiation
carrying a quantum of energy
– A photon is like an energy packet
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58
Section 4.7
Electrons
in
Atoms
Isotopes
Ephoton = h
1. What is the energy of a photon with a
frequency of 9x1014 Hz?
2. What is the frequency of a photon with an
energy of 5x10-22 J?
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59
Section 4.7
Wave-Particle
Isotopes
Duality
•In 1905 Albert Einstein introduced the
idea of that electromagnetic radiation has
“wave-particle duality”
Light has wavelike particles
and
Light can be thought of as a stream of
particles where each particle carries a
quantum of energy.
Video - http://study.com/academy/lesson/what-is-a-photon-definition-energy-wavelength.html
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Section 4.7
11.4 Energy Levels
Isotopes
•
•
•
Ground State: the lowest energy state of an
atom.
Excited State: when an atom contains excess
energy (has higher potential energy).
When an excited atom returns to ground
state it gives off the energy it has gained as
electromagnetic radiation. Example: Neon
signs
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Section 4.7
Isotopes
Absorption
• An electron absorbs
energy (photon) and
moves from the
ground state to an
excited state.
E
4
E
3
E
2
E
1
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Section 4.7
What goes up…must come down!
Isotopes
Emission
• When an electron in the excited state returns to
the ground state it emits a photon.
E
4
E
3
E
2
E
= h =E -E
photon
3 1
E
1
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Section 4.7
Absorption
Isotopes
and Emission
• Emitting photons creates light or
electromagnetic radiation
• Electromagnetic radiation in the
visible light spectrum has color!
• These photons have wavelengths
that correspond to their color.
VIDEO - http://study.com/academy/lesson/the-bohr-model-and-atomic-spectra.html
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Section 4.7
Isotopes
In an actual atom…
• Emission and
absorption happen
simultaneously and
on different energy
levels.
• Emission produces
a line-emission
spectrum
E
4
E
3
E
2
E
1
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Section 4.7
Line
Isotopes
Emission Spectrum
• Line-Emission Spectrum- a beam of
light separated into a series of
specific frequencies (and therefore
specific wavelengths) of visible
light.
– produced when electrons fall back to
ground state
– the photons emitted in the fall give off
specific patterns (colors) of light
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Section 4.7
Isotopes
The Hydrogen-Atom Line Emission Spectrum
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Section 4.7
Isotopes
The Hydrogen-Atom Line Emission Spectrum
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Section 4.7
11.5
The
Bohr
Model
Isotopes
In 1913, Niels Bohr develops a new atomic
model
Bohr stated that the electrons orbit the
nucleus like the planets orbit the sun.
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69
Section 4.7
Isotopes
Each possible electron orbit in Bohr’s model
has a fixed energy.
The fixed energies an electron can have are
called energy levels.
Each energy level further from the nucleus is of
greater energy
Energy levels are regions where an electron is
likely to be moving
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70
Section 4.7
Isotopes
The rungs on this ladder are like
the energy levels in Bohr’s model.
A person on a ladder cannot
stand between the rungs.
Similarly, the electrons in an
atom cannot exist between energy
levels.
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71
Section 4.7
Isotopes
The Rutherford model could not explain why
elements that have been heated to higher
temperatures give off different colors of light.
The Bohr model explains how the energy levels
of electrons in an atom change when the atom
emits light.
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72
Section 4.7
Bohr’s
model
Isotopes
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73
Section 4.7
Bohr’s Model of the Atom
Isotopes
• Bohr’s explained emission spectra using the idea of
fixed orbits.
• This idea is close but not true…
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Section 4.7
11.6 The Wave (Quantum) Mechanical
Isotopes
Model
Unfortunately, Bohr’s model did not apply
to other atoms
That led scientists to question his model
They wondered why the electron had to be
located in a precise orbit
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75
Section
4.7
The
Isotopes
Wave (Quantum) Mechanical
Model
• That led to the Heisenburg uncertainty
principle
states that it’s impossible to determine the
position and velocity of an electron at the
same time
Video clip - https://www.youtube.com/watch?v=H-AlfuvjPYM
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76
The Wave (Quantum) Mechanical
Isotopes
Model
Section 4.7
• Further developments led to the wave
(quantum) mechanical model
The wave (quantum) mechanical
model describes mathematically the
position of electrons in an atom
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77
The Wave (Quantum) Mechanical
Isotopes
Model
Section 4.7
Like the Bohr model, the quantum
mechanical model of the atom restricts the
energy of electrons to certain values.
However, the quantum mechanical model
does not specify an exact path the electron
takes around the nucleus.
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78
The Wave (Quantum) Mechanical
Isotopes
Model
Section 4.7
• Erwin Schrodinger (1887-1961) developed
the quantum mechanical model further
It determined the allowed energies an
electron can have
Also, Schrodinger developed an equation
to determine how likely it is to find an
electron in a particular location around the
nucleus of an atom.
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79
The Wave (Quantum) Mechanical
Isotopes
Model
Section 4.7
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80
The Wave (Quantum) Mechanical
Isotopes
Model
Section 4.7
Rutherfo
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81
Section 4.7
Isotopes
Concept Check
TRUE or FALSE
In the modern atomic model, electrons are
moving around the nucleus in a circular path.
a) True
b) False
VIDEOCLIP - https://www.youtube.com/watch?v=H-AlfuvjPYM
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Copyright © Cengage Learning. All rights reserved
82
Section 4.7
11.7
The
Orbitals
Isotopes
• The model explained that an electron
exists in certain regions called orbitals.
A 3D region around the nucleus that
indicates the probable location of an
electron
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83
Section 4.7
Quantum
Numbers
Isotopes
The s orbitals are
spherical.
The p orbitals are
dumbbell shaped.
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84
Section 4.7
11.7 The Orbitals - Quantum Numbers
Isotopes
• In order to specify the properties of atomic
orbitals and electrons in orbitals, chemists use
quantum numbers
Principle quantum number (n)
Angular momentum quantum number (l)
Magnetic quantum number (m)
Spin quantum number
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85
Section 4.7
Hydrogen
Isotopes
Energy Levels
• Hydrogen has discrete
energy levels.
 Called principal energy
levels
Principal energy levels
 Labeled as n = 1, n=2,
3, 4, and so forth
 The energy level corresponds
to the periods of the periodic
table
 1st energy level = period 1
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Copyright © Cengage Learning. All rights reserved
86
Section 4.7
Sublevels
Isotopes
• Energy levels can be divided into
sublevels
s, p, d, f are sublevels
• Orbitals exist in sublevels
 An orbital can be empty or it can contain
one or two electrons
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87
Section 4.7
Isotopes
• Each principal energy level is divided into
sublevels.
 Labeled with numbers and letters
 Indicate the shape of the orbital
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88
Section 4.7
Figure 11.18: Principal levels can be divided into
sublevels
Isotopes
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89
Section 4.7
Quantum
Numbers
Isotopes
The spin quantum number indicates the
spin of the electron
• It may be thought of as clockwise or
counterclockwise.
• A vertical arrow indicates an electron
and its direction of spin (  or  )
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90
Section 4.7
Quantum
Numbers
Isotopes
Rutherfo
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91
Section 4.7
11.9
and
11.10
Electron
Arrangement
Isotopes
To describe the arrangement of the electrons
in an atom we use electron configuration
To describe spin we use orbital notation
•Remember – an atom tends to assume the
lowest energy configuration possible
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92
Section 4.7
Electron
Arrangement
Isotopes
To determine electron configuration, follow
three simple rules
1.The Aufbau principle states that an electron
occupies the lowest energy orbital that can
receive it
oWe fill lowest to highest energy
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93
Section 4.7
Orbital
Diagram
Isotopes
Rutherfo
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94
Section 4.7
Electron
Arrangment
Isotopes
2. The Pauli Exclusion principle states that
an atomic orbital can hold a maximum of
two electrons, and those two electrons
must have opposite spins.
3. Hunds Rule states that orbitals of the
same energy must be occupied by one
electron before it can be occupied by a
second electron.
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95
Section 4.7
Isotopes
Hund’s Rule
Orbital
Diagrams
Three electrons would occupy three orbitals of
equal energy as follows.
Electrons then occupy each orbital so that their
spins are paired with the first electron in the
orbital.
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Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.
Orbital Diagram and
Electron Configuration
Isotopes
Look
at the orbital filling diagram of the oxygen atom.
Section 4.7
• An oxygen
atom contains
eight electrons.
Electron Configurations of Selected Elements
Element
1s
2s
2px 2py 2pz
3s
Electron
configuration
H
1s1
He
1s2
Li
1s22s1
C
1s22s22p2
N
1s22s22p3
O
1s22s22p4
F
1s22s22p5
Ne
1s22s22p6
Na
1s22s22p63s1
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Copyright © Pearson Education, Inc., or its affiliates. All Rights
Section 4.7
Orbital Diagram and Electron Configuration
Isotopes
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98
Section 4.7
DO NOW
Isotopes
Draw the orbital diagrams for the following:
• Nitrogen
• Cobalt
1s 2s
2p
3s
3p
4s
3d
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99
Section 4.7
11.10
Electron
Configuration
Isotopes
Orbitals, Sublevels & Electrons
• for a many electron atom, build-up the energy
levels, filling each orbital in succession by energy
2
2
2
6
2
6
10
6
2
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d
6
2
14
10
6
2
14
10
< 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d < 7p
10
6
- each s sublevel has 1 s orbital (can hold 2 e-)
- each p sublevel has 3 p orbitals (can hold 6 e-)
- each d sublevel has 5 d orbitals (can hold 10 e-)
- each f sublevel has 7 f orbitals (can hold 14 e-)
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100
Section 4.7
H Atom
Isotopes
• Electron configuration – electron arrangement
1s2
• Orbital diagram – orbital is a box grouped by
sublevel containing arrow(s) to represent
electrons
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101
Section 4.7
Isotopes
Li Atom
Orbital diagram
Electron configuration
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102
Section 4.7
O
Atom
Isotopes
• The lowest energy configuration for an atom is the one
having the maximum number of unpaired electrons in a
particular set of degenerate (same energy) orbitals.
Electron Configuration
Orbital Diagram
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103
Section 4.7
Orbitals
Isotopes
• Parts of the periodic table corresponds to
each orbital shape
Groups 1 & 2 – s block
Groups 13-18 – p block
Groups 3-12 – d block
Bottom two rows – f block
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104
Section 4.7
Orbitals
Isotopes
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105
Section 4.7
Isotopes
• The electron configurations in the sublevel last
occupied for the first eighteen elements.
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106
Section 4.7
11. 10 Abbreviating Electron Configuration
Isotopes
To avoid writing the inner-level electrons, we
often abbreviate the configurations.
For example:
for sodium
1s22s22p63s1

[Ne]3s1
for titanium
1s22s22p63s23p64s23d2

[Ar]4s23d2
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107
Section 4.7
Abbreviating
Electron
Configuration
Isotopes
Try abbreviating for:
a) Silicon
1s22s22p63s23p2
b) Vanadium
1s22s22p63s23p64s23d3
c) Strontium
1s22s22p63s23p64s23d104p65s2
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108
Section 4.7
Isotopes
Do Now
What is the electron configuration for Germanium?
How many electrons can the p sublevel hold?
How many electrons can principal energy level 3
hold?
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109
Section 4.7
Isotopes
Do Now
Write the noble gas configuration for the following
elements.
1. Calcium
[Ar] 3s2
2. Iodine
[Kr] 5s24d105p5
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110
Section 4.7
Valence
electrons
Isotopes
• Highest energy level for any atom is called
the valence shell
– electrons in the valence shell are called
valence electrons
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111
Section 4.7
Valence
electrons
Isotopes
Example:
• Carbon’s (with an electron configuration of
1s22s22p2) highest energy level is principal
energy level 2.
• There are 2 electrons from the 2s sublevel
and 2 electrons from the 2p sublevel.
• 2 + 2 = 4  4 valence electrons for carbon
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112
Section 4.7
Valence
electrons
Isotopes
Example:
• Lithium’s (with an electron configuration of
1s22s1) highest energy level is principal
energy level 2.
• There is 1 electron from the 2s sublevel.
• 1 valence electron for lithium
How many valence electrons are there in Ca?
I?
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113