Mec chapter 2

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Chapter 2
The Structure of the
Atom and the Periodic
Table

Denniston
Topping
Caret
7th Edition

2.1 Composition of the Atom
• Atom - the basic structural unit of an
element
• The smallest unit of an element that
retains the chemical properties of
that element

2.1 Composition of the Atom Electrons, Protons, and Neutrons
• Atoms consist of three primary particles
• electrons
• protons
• neutrons
• Nucleus - small, dense, positively
charged region in the center of the atom
- protons - positively charged particles
- neutrons - uncharged particles

2.1 Composition of the Atom Characteristics of Atomic
Particles
• Electrons are negatively charged particles
located outside of the nucleus of an atom
• Protons and electrons have charges that
are equal in magnitude but opposite in sign
• A neutral atom that has no electrical
charge has the same number of protons
and electrons
• Electrons move very rapidly in a relatively
large volume of space while the nucleus is
small and dense

2.1 Composition of the Atom Symbolic Representation of
an Element
Charge of
particle
Mass

A C
Z X
Atomic Symbol of
number the atom

• Atomic number (Z) - the number of
protons in the atom
• Mass number (A) - sum of the number of
protons and neutrons

Atomic Calculations
number of protons + number of neutrons = mass number

number of neutrons = mass number - number of protons

number of protons = number of electrons IF positive and
negative charges cancel, the atom charge = 0

Atomic Composition Calculations
Calculate the number of protons, neutrons,
and electrons in each of the following:

11
5 B
55
26 Fe

2.1 Composition of the Atom Isotopes
• Isotopes - atoms of the same element
having different masses
– contain same number of protons 4
– contain different numbers of neutrons
Isotopes of Hydrogen

Hydrogen Deuterium Tritium
(Hydrogen - 1) (Hydrogen - 2) (Hydrogen - 3)

Isotopic Calculations
• Isotopes of the same element have identical
chemical properties
• Some isotopes are radioactive
• Find chlorine on the periodic table
• What is the atomic number of chlorine?
17
• What is the mass given?
35.45
• This is not the mass number of an isotope

2.1 Composition of the Atom Atomic Mass
• What is this number: 35.34?
• The atomic mass - the weighted average of
the masses of all the isotopes that make up
chlorine
• Chlorine consists of chlorine-35 and
chlorine-37 in a 3:1 ratio
• Weighted average is an average corrected
by the relative amounts of each isotope
present in nature

Atomic Mass Calculation
Calculate the atomic mass of naturally
occurring chlorine if 75.77% of chlorine
atoms are chlorine-35 and 24.23% of
chlorine atoms are chlorine-37
Step 1: convert the percentage to a decimal
fraction:
0.7577 chlorine-35
0.2423 chlorine-37

Step 2: multiply the decimal fraction by the
mass of that isotope to obtain the isotope
contribution to the atomic mass:
For chlorine-35:
0.7577 x 35.00 amu = 26.52 amu
For chlorine-37
0.2423 x 37.00 amu = 8.965 amu
Step 3: sum these partial weights to get the
weighted average atomic mass of chlorine:
26.52 amu + 8.965 amu = 35.49 amu

Atomic Mass Determination
• Nitrogen consists of two naturally occurring
isotopes
– 99.63% nitrogen-14 with a mass of 14.003 amu
– 0.37% nitrogen-15 with a mass of 15.000 amu
• What is the atomic mass of nitrogen?

Ions and Charges
• Ions - electrically charged particles that
result from a gain or loss of one or more
electrons by the parent atom
• Cation - positively charged
– results from the loss of electrons
– 23Na  23Na+ + 1e-
• Anion - negatively charged
– results from the gain of electrons
– 19F + 1e-  19F-

2.1 Composition of the Atom Calculating Subatomic Particles
in Ions
• How many protons, neutrons, and electrons
are in the following ions?

39 +
19 K

32 2-
16 S
24 2+
12 Mg

2.2 Development of Atomic
Theory
• Dalton’s Atomic Theory - the first
experimentally based theory of atomic
structure of the atom

Postulates of Dalton’s Atomic Theory
2.2 Development of

1. All matter consists of tiny particles
Atomic Theory

called atoms
2. An atom cannot be created, divided,
destroyed, or converted to any other
type of atom
3. Atoms of a particular element have
identical properties

2.2 Development of
4. Atoms of different elements have
different properties
Atomic Theory

5. Atoms of different elements
combine in simple whole-number
ratios to produce compounds (stable
aggregates of atoms)
6. Chemical change involves joining,
separating, or rearranging atoms
Postulates 1, 4, 5, and 6 are still regarded
as true.

Subatomic Particles:
Electrons, Protons, and Neutrons
2.2 Development of

• Electrons were the first subatomic
Atomic Theory

particles to be discovered using the
cathode ray tube.

Indicated that the
particles were
negatively charged.

Evidence for Protons and
2.2 Development of

Neutrons
Atomic Theory

• Protons were the next particle to be discovered,
by Goldstein
– Protons have the same size charge but opposite in sign
– A proton is 1,837 times as heavy as an electron
• Neutrons
– Postulated to exist in 1920’s but not demonstrated to
exist until 1932
– Almost the same mass as the proton

2.4 The Periodic Law and the
Periodic Table
• Dmitri Mendeleev and Lothar Meyer - two
scientists working independently developed
the precursor to our modern periodic table
• They noticed that as you list elements in
order of atomic mass, there is a distinct
regular variation of their properties
• Periodic law - the physical and chemical
properties of the elements are periodic
functions of their atomic numbers

2.4 The Periodic Law
and the Periodic Table
Classification of the Elements

Important Biological Elements

Parts of the Periodic Table
• Period - a horizontal row of elements in
the periodic table. They contain 2, 8, 8,
18, 18, and 32 elements
• Group - also called families, and are
columns of elements in the periodic table.
• Elements in a particular group or family
share many similarities, as in a human
family.

Families of the Periodic Table
• Representative elements - Group A
elements
• Transition elements - Group B
elements
• Alkali metals - Group IA
• Alkaline earth metals - group IIA
• Halogens - group VIIA
• Noble gases - group VIIIA

Category Classification of
and the Periodic Table Elements

• Metals - elements that tend to lose
electrons during chemical change,
forming positive ions
• Nonmetals - a substance whose atoms
tend to gain electrons during chemical
change, forming negative ions
• Metalloids - have properties intermediate
between metals and nonmetals

Classification of Elements
and the Periodic Table Metals
• Metals:
– A substance whose atoms tend to lose
electrons during chemical change
– Elements found primarily in the left 2/3 of
the periodic table
• Properties:
– High thermal and electrical conductivities
– High malleability and ductility
– Metallic luster
– Solid at room temperature

and the Periodic Table Nonmetals

• Nonmetals:
– A substance whose atoms may gain
electrons, forming negative ions
– Elements found in the right 1/3 of the
periodic table
• Properties:
– Brittle
– Powdery solids or gases
– Opposite of metal properties

and the Periodic Table Metalloids

• Metalloids:
– Elements that form a narrow diagonal band
in the periodic table between metals and
nonmetals
• Properties are somewhat between those
of metals and nonmetals
• Also called semimetals

Atomic Number and Atomic Mass

• Atomic Number:
– The number of protons in the nucleus of
an atom of an element
– Nuclear charge or positive charge from
the nucleus
• Most periodic tables give the element
symbol, atomic number, and atomic
mass

Element Information in the
and the Periodic Table Periodic Table

20 atomic number
Ca symbol
Calcium name
40.08 atomic mass

Using the Periodic Table
• Identify the group and period to
which each of the following belongs:
a. P
b. Cr
c. Element 30
• How many elements are found in
period 6?
• How many elements are in group
VA?

2.5 Electron Arrangement and
the Periodic Table
• The electron arrangement is the primary
factor in understanding how atoms join
together to form compounds
• Electron configuration - describes the
arrangement of electrons in atoms
• Valence electrons - outermost electrons
– The electrons involved in chemical bonding

2.5 Electron Arrangement
and the Periodic Table Valence Electrons
• The number of valence electrons is the
group number for the representative
elements
• The period number gives the energy
level (n) of the valence shell for all
elements

2.5 Electron Arrangement Valence Electrons and Energy
and the Periodic Table Level
• How many valence electrons does Fluorine
have?
– 7 valence electrons

• What is the energy level of these electrons?
– Energy level is n = 2

Energy Level
Electron Arrangement by

2.5 Electron Arrangement Valence Electrons - Detail
and the Periodic Table • What is the total number of electrons in
fluorine?
– Atomic number = 9
– 9 protons and 9 electrons
• 7 electrons in the valence shell, (n = 2 energy level),
so where are the other two electrons?
– In n = 1 energy level
– Level n = 1 holds only two electrons

Determining Electron Arrangement
and the Periodic Table List the total number of electrons, total number of
valence electrons, and energy level of the valence
electrons for silicon.
1. Find silicon in the periodic table
• Group IVA
• Period 3
• Atomic number = 14
1. Atomic number = number of electrons
in an atom
• Silicon has 14 electrons

Determining Electron Arrangement #2
and the Periodic Table List the total number of electrons, total number of
valence electrons, and energy level of the valence
electrons for silicon.
3. As silicon is in Group IV, only 4 of its 14
electrons are valence electrons
• Group IVA = number of valence electrons
3. Energy levels:
• n = 1 holds 2 electrons
• n = 2 holds 8 electrons (total of 10)
• n = 3 holds remaining 4 electrons (total = 14)

Determining Electron Arrangement
and the Periodic Table Practice
List the total number of electrons, total
number of valence electrons, and energy
level of the valence electrons for:
• Na
• Mg
• S
• Cl
• Ar

Energy Levels and Subshells
PRINCIPAL ENERGY LEVELS
• n = 1, 2, 3, …
• The larger the value of n, the higher the energy
level and the farther away from the nucleus the
electrons are
• The number of sublevels in a principal energy
level is equal to n
– in n = 1, there is one sublevel
– in n = 2, there are two sublevels

2.5 Electron Arrangement Principal Energy Levels
and the Periodic Table • The electron capacity of a principal
energy level (or total electrons it can hold) is
2(n)2
– n = 1 can hold 2(1)2 = 2 electrons
– n = 2 can hold 2(2)2 = 8 electrons

• How many electrons can be in the n = 3
level?
– 2(3)2 = 18

• Compare the formula with periodic table…..

n = 1, 2(1)2 = 2
n = 2, 2(2)2 = 8
n = 3, 2(3)2 = 18
n = 4, 2(4)2 = 32

2.5 Electron Arrangement Sublevels
and the Periodic Table • Sublevel: a set of energy-equal orbitals
within a principal energy level
• Subshells increase in energy:
s<p<d<f
• Electrons in 3d subshell have more energy
than electrons in the 3p subshell
• Specify both the principal energy level and a
subshell when describing the location of an
electron

Sublevels in Each Energy Level
Principle energy Possible
level (n) subshells
1 1s

2 2s, 2p

3 3s, 3p, 3d

4 4s, 4p, 4d, 4f

2.5 Electron Arrangement Orbitals
• Orbital - a specific region of a sublevel
containing a maximum of two electrons
• Orbitals are named by their sublevel and
principal energy level
– 1s, 2s, 3s, 2p, etc.
• Each type of orbital has a characteristic
shape
– s is spherically symmetrical
– p has a shape much like a dumbbell

Orbital Shapes
• s is spherically
symmetrical

• Each p has a shape much like a dumbbell,
differing in the direction extending into space

Number of
2.5 Electron Arrangement Subshell
orbitals
s 1

p 3

d 5

f 7

•How many electrons can be in the 4d

•10

Quantum Mechanical Model
and the Periodic Table Shell 4
• Each orbital within a
sublevel contains a 4f •• •• •• •• •• •• ••
maximum of 2

Increasing Energy
electrons
4d •• •• •• •• ••
• Energy increases as n,
shell number Sublevel
increases, but ALSO 4p •• •• ••
increases as you move
from s to p to d to f Orbital
sublevels 4s ••

Electron

2.5 Electron Arrangement Electron Spin
and the Periodic Table • Electron configuration - the
arrangement of electrons in atomic
orbitals
• Aufbau principle - or building up
principle helps determine the electron
configuration
– Electrons fill the lowest-energy orbital that
is available first
– Remember s<p<d<f in energy
– When the orbital contains two electrons,
the electrons are said to be paired

Electron Filling Order

2.5 Electron Arrangement Rules for Writing Electron
and the Periodic Table Configurations
• Obtain the total number of electrons in the atom
from the atomic number
• Electrons in atoms occupy the lowest energy
orbitals that are available – 1s first
• Each principal energy level, n contains only n
sublevels
• Each sublevel is composed of orbitals
• No more than 2 electrons in any orbital
• Maximum number of electrons in any principal
energy level is 2(n)2

2.5 Electron Arrangement Electron Distribution
and the Periodic Table • This table lists the number of electrons in each
shell for the first 20 elements
• Note that 3rd shell stops filling at 8 electrons even though
it could hold more

Orbital Energy-level Diagram

Writing Electron Configurations
• H • Li
– Hydrogen has – Lithium has 3
only 1 electron electrons
– It is in the – First two have
lowest energy configuration
level & lowest of Helium – 1s2
orbital – 3rd is in the
– Indicate orbital of
number of lowest energy
electrons with a in n=2
superscript – 1s2 2s1
– 1s1

Electron Configuration Examples
• Give the complete electron
configuration of each element
– Be

–N

– Na

– Cl

– Ag

and the Periodic Table The Shell Model and Chemical
Properties
• As we explore the model placing electrons
in shells, we will see that the pattern which
emerges from this placement correlates well
with a pattern for various chemical
properties
• We will see that all elements in a group
have the same number of electrons in their
outermost (or valence) shell

2.5 Electron Arrangement Groups Have Similar Chemical
and the Periodic Table Properties and Appearances
• Examples of different elements that
have similar properties and are all in
group VA
– Nitrogen
– Phosphorus
– Arsenic
– Antimony
– Bismuth

2.5 Electron Arrangement Shorthand Electron
and the Periodic Table Configurations
• Uses noble gas symbols to represent the
inner shell and the outer shell or valance
shell is written after
• Aluminum- full electron configuration is:
1s22s22p63s23p1

What noble gas configuration is this?
•Neon
•Configuration is written: [Ne]3s23p1

and the Periodic Table • Remember:
– How many subshells are in each
principle energy level?
– There are n subshells in the n principle
energy level.
– How many orbitals are in each
subshell?
– s has 1, p has 3, d has 5, and f has 7
– How many electrons fit in each orbital?
– 2

2.5 Electron Arrangement Shorthand Electron
and the Periodic Table Configuration Examples

• N

• S

• Ti

• Sn

2.5 Electron Arrangement Classification of Elements
and the Periodic Table According to the Type of
Subshells Being Filled

Use this breakdown of the Periodic Table and you can
write the configuration of any element.

and the Periodic Table Classification of Elements –
by Group
• Representative element: An element in which the
distinguishing electron is found in an s or p
subshell
• Distinguishing electron: The last or highest-
energy electron found in an element
• Transition element: An element in which the
distinguishing electron is found in a d subshell
• Inner-transition element: An element in which
the distinguishing electron is found in a f
subshell

2.6 The Octet Rule
• The noble gases are extremely stable
– Called inert as they don’t readily bond to other
elements
• The stability is due to a full complement of
valence electrons in the outermost s and p
sublevels:
– 2 electrons in the 1s of Helium
– the s and p subshells are full in the outermost
shell of the other noble gases (eight electrons)

Octet of Electrons
2.6 The Octet Rule

• Elements in families other than the noble
gases are more reactive
– Strive to achieve a more stable electron
configuration
– Change the number of electrons in the atom to
result in full s and p sublevels
• Stable electron configuration is called the
“noble gas” configuration

2.6 The Octet Rule The Octet Rule
• Octet rule - elements usually react in such a way
as to attain the electron configuration of the noble
gas closest to them in the periodic table
– Elements on the right side of the table move right to the
next noble gas
– Elements on the left side move “backwards” to the
noble gas of the previous row
• Atoms will gain, lose or share electrons in
chemical reactions to attain this more stable
energy state

2.6 The Octet Rule Ion Formation and the Octet Rule
• Metallic elements tend to form positively
charged ions called cations
• Metals tend to lose all their valence
electrons to obtain a configuration of the
noble gas
Na Na+ + e-
Sodium atom Sodium ion
11e-, 1 valence e- 10e-
[Ne]3s1 [Ne]

2.6 The Octet Rule Ion Formation and the Octet Rule
• All atoms of a group lose the same number of
electrons
• Resulting ion has the same number of electrons as
the nearest (previous) noble gas atom

Al Al3+ + 3e-
Aluminum atom Aluminum ion
[Ne]3s23p1 [Ne]

Isoelectronic
• Isoelectronic - atoms of different elements having
2.6 The Octet Rule
the same electron configuration (same number of
electrons)
• Nonmetallic elements, located on the right side of
the periodic table, tend to form negatively charged
ions called anions
• Nonmetals tend to gain electrons so they become
isoelectronic with its nearest noble gas neighbor
located in the same period to the right
O + 2e- O2-
Oxygen atom Oxide ion
[He]2s22p4 [He]2s22p6 or [Ne]

2.6 The Octet Rule Using the Octet Rule
• The octet rule is very helpful in predicting
the charges of ions in the representative
elements
• Transition metals still tend to lose electrons
to become cations but predicting the charge
is not as easy
• Transition metals often form more than one
stable ion
– Iron forming Fe2+ and Fe3+ is a common example

Examples Using the Octet Rule
2.6 The Octet Rule

• Give the charge of the • Which of the
most probable ion following pairs of
resulting from these atoms and ions are
elements isoelectronic?
– Ca – Cl-, Ar
– Sr – Na+, Ne
– S – Mg2+, Na+
– P – O2-, F-

2.7 Trends in the Periodic Table
• Many atomic properties correlate with
electronic structure and so also with their
position in the periodic table
– atomic size
– ion size
– ionization energy
– electron affinity

Atomic Size
2.7 Trends in the Periodic

• The size of an element increases, moving
down from top to bottom of a group
• The valence shell is higher in energy and
Table

farther from the nucleus traveling down the
group
• The size of an element decreases from left
to right across a period
• The increase in magnitude of positive charge
in nucleus pulls the electrons closer to the
nucleus

Table Variation in Size of Atoms

Cation Size
Cations are smaller than their parent atom
• More protons than electrons creates an increased
nuclear charge
• Extra protons pull the remaining electrons closer
to the nucleus
Table

• Ions with multiple positive charges are even
smaller than the corresponding monopositive
ions
– Which would be smaller, Fe2+ or Fe3+? Fe3+
• When a cation is formed isoelectronic with a
noble gas the valence shell is lost, decreasing the
diameter of the ion relative to the parent atom

Anion Size

Anions are larger than their parent
atom.
• Anions have more electrons than protons
Table

• Excess negative charge reduces the pull
of the nucleus on each individual electron
• Ions with multiple negative charges are
even larger than the corresponding
monopositive ions

Relative Size of Select Ions and
Their Parent Atoms
Table

2.7 Trends in the Periodic Ionization Energy
• Ionization energy - The energy required to
remove an electron from an isolated atom
• The magnitude of ionization energy
Table

correlates with the strength of the attractive
force between the nucleus and the
outermost electron
• The lower the ionization energy, the easier
it is to form a cation
ionization energy + Na  Na+ + e-

Ionization Energy of Select Elements
Table

• Ionization decreases down a family as the
outermost electrons are farther from the nucleus
• Ionization increases across a period because the
outermost electrons are more tightly held
• Why would the noble gases be so unreactive?

2.7 Trends in the Periodic Electron Affinity
• Electron affinity - The energy released
when a single electron is added to an
isolated atom
Table

• Electron affinity gives information about
the ease of anion formation
– Large electron affinity indicates an atom
becomes more stable as it forms an anion

Br + e–  Br– + energy

2.7 Trends in the Periodic Periodic Trends in Electron
Affinity
• Electron affinity
generally
Table

decreases down a
group
• Electron affinity
generally increases
across a period

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