Transcript 2013 The Periodic Table

The Periodic Table
The Modern Periodic Table
 The
modern periodic table is based
on the atomic numbers of the
elements.
The Modern Periodic Table
 The
elements are arranged in order of
increasing atomic number.
 The physical and chemical properties
of the elements repeat in a regular
pattern when they are arranged in
order of increasing atomic number.
The Periodic Table
 Elements
in vertical columns showed
similar properties.
For example, alkaline
earth metals have
high melting points
and low density and
are silver in color,
ductile, and malleable.
Periodicity
 This
repeated pattern is an example
of periodicity in the properties of
elements.
 Periodicity is the tendency to recur at
regular intervals.
The Periodic Table
 On
the periodic table a period,
sometimes also called a series,
consists of the elements in a
horizontal row.
The Periodic Table
 There
are 7 periods in the table.
The Periodic Table
A
group, sometimes also called a
family, consists of the elements in a
vertical column.
Metals, Nonmetals, and Metalloids
 There
are 3 main regions of the table:
metals, nonmetals and metalloids.
Metals
 The
metals are in blue.
Metals
 Metals
are elements that have luster
(are shiny), conduct heat and
electricity, and usually bend without
breaking (malleable).
 Metals are also ductile (can be drawn
out into a wire).
Metals
 Most
metals have one, two, or three
valence electrons.
 Metals tend to lose electrons in order
to achieve the stability of a filled
octet.
Metals
 All
metals except mercury are solids
at room temperature; in fact, most
have extremely high melting points.
Metal Reactivity
A
metal’s reactivity is its ability to
react with another substance.
 Reactivity for metals increases as
you go down a
group and left
across a period.
Metal Reactivity
Metal Reactivity
1. Consult the “Activity Series of Metals”
in the Chemistry Reference Tables to
determine the more active metal.
a) cobalt (Co) or manganese (Mn)
(manganese)
b) barium (Ba) or sodium (Na)
(barium)
Nonmetals
 The
nonmetals are in yellow.
Nonmetals
 Although
the majority of the
elements in the periodic table are
metals, many nonmetals are
abundant in nature.
Nonmetals
nonmetals don’t conduct
electricity and are much poorer
conductors of heat than metals.
 Most
Nonmetals
 Many
are gases at room temperature;
those that are solids lack the luster
of metals and are brittle.
 Their melting points tend to be lower
than those of metals.
Nonmetals
 With
the exception of carbon,
nonmetals have five, six, seven, or
eight valence electrons.
 Nonmetals tend to gain electrons in
order to achieve the stability of a
filled octet.
Nonmetal Reactivity
A
nonmetal’s reactivity is its ability
to react with another substance.
 Reactivity for nonmetals increases
as you go left to right and up the
periodic table.
Nonmetal Reactivity
2. Consult the “Activity Series of
Halogens” in the Chemistry Reference
Tables to determine the less active
nonmetal.
a) fluorine (F2) or chlorine (Cl2)
(chlorine)
b) chlorine (Cl2) or iodine (I2)
(iodine)
Metalloids
 The
metalloids are in pink.
Metalloids
 Metalloids
have some chemical and
physical properties of metals and
other properties of nonmetals.
 In the periodic table, the metalloids
lie along the border between metals
and nonmetals.
Atomic Size and Metal Reactivity
 With
metals the greater the tendency
to lose electrons, the more reactive
the metal is.
Atomic Size and Metal Reactivity
 Larger
atoms have electrons orbiting
farther away from the nucleus.
 Hence, because electrons are not as
tightly held (when compared to
smaller atoms) and metals want to
lose electrons, metals that are larger
in size are far more reactive.
Atomic Size and Metal Reactivity
 Francium
metal.
(Fr) is the most reactive
Atomic Size and Nonmetal
Reactivity
 With
nonmetals the greater the
tendency to gain/share electrons, the
more reactive the nonmetal is.
Atomic Size and Nonmetal
Reactivity
 Smaller
atoms have greater nuclear
charge because the outer electrons
are closer to the nucleus.
Atomic Size and Nonmetal
Reactivity
 Thus
the tendency to gain electrons
will increase.
 As a result, the reactivity of the
nonmetals should increase as you go
from left to right across the periodic
table, up to but not including the inert
(noble) gases.
Atomic Size and Nonmetal
Reactivity
 Nonmetals
that
are smaller in
size are far more
reactive.
 Fluorine is the
most reactive
nonmetal.
Group Names
 Groups
1, 2, and 13 - 18 (Group A
elements) are called representative
(main group) elements.
Group Names
 Groups
3 - 12 (Group B elements) are
called transition elements.
Group 1: The Alkali Metals
Group 1: The Alkali Metals
 Group
1 elements are called the
alkali metals and have one valence
electron.
 They form 1+ ions after losing the
one valence electron.
Group 2: The Alkaline Earth
Metals
Group 2: The Alkaline Earth
Metals
 Group
2 elements are called the
alkaline earth metals and have two
valence electrons.
 They form 2+ ions after losing the
two valence electrons.
Group 17: The Halogens
Group 17: The Halogens
 Group
17 elements are called the
halogens and have seven valence
electrons.
 They form 1- ions after gaining one
more electron.
Group 18: The Noble Gases
Group 18: The Noble Gases
 Group
18 elements are called the
noble gases and have eight valence
electrons, except for helium which
only has two.
 The noble gases with 8 valence
electrons obey the octet rule and are
generally unreactive.
Valence Electrons
Question
How many valence electrons are in
an atom of each of the following
elements?
a) Magnesium (Mg) (2)
b) Selenium (Se) (6)
c) Tin (Sn) (4)
Question
2. Match each element in Column A
with the best matching description
in Column B. Each Column A
element may match more than one
description from Column B.
Question
Column A
1. strontium
2. chromium
3. iodine
Column B
a. halogen
b. alkaline earth metal
c. representative element
d. transition element
Answers
1. strontium
b, c
2. chromium
d
3. iodine
a, c
STOP HERE
Periodic Trends
 Because
the periodic table relates
group and period numbers to
valence electrons, it’s useful in
predicting atomic structure and,
therefore, chemical properties.
Periodic Trends
Atomic Size (Atomic Radius)

The atomic radius of a chemical element
is a measure of the size of its atoms,
usually the mean or typical distance
from the nucleus to the boundary of the
surrounding cloud of electrons.
Trends in Atomic Size (Radii)
 Atomic
size is influenced by two
factors.
• Energy Level – A higher energy
level is farther away.
• Charge on nucleus - More charge
(protons) pulls electrons in closer.
Group Trend for
Atomic Radii
 As
you go down
a group, each
atom has another
energy level so
the atoms get
bigger.
H
Li
Na
K
Rb
Period Trend for Atomic Radii
 As
you go across a period, the
radius gets smaller.
 Atoms are in the same energy level,
but as you move across the chart,
atoms have a greater nuclear charge
(more protons).
 Therefore, the outermost electrons
are closer.
Period Trend for Atomic Radii
Na
Mg
Al
Si
P
S Cl Ar
Question
3. (a) State why atoms get bigger as
you go down a group on the
periodic table.
(b) State why the radius decreases
across a period.
Question
4. Choose the element from the pair
with the larger atomic radius.
a) lithium (Li) or beryllium (Be)
(lithium)
b) silicon (Si) or tin (Sn)
(tin)
Question
5. Choose the element from the pair
with the smaller atomic radius.
a) silver (Ag) or gold (Au)
(silver)
b) cesium (Cs) or barium (Ba)
(barium)
Ionic Size (Ionic Radius)
 Ionic
radius is the radius of an
atom's ion.
 When an atom gains or loses one or
more electrons, it becomes an ion.
Ionic Size (Ionic Radius)


Recall that metals tend to lose
electrons in order to achieve the
stability of a filled octet.
As a result, metals tend to form
cations which are positive ions.
Ionic Size (Ionic Radius)
A
cation has a smaller radius than
its neutral atom.
Ionic Size (Ionic Radius)


Nonmetals tend to gain electrons in
order to achieve the stability of a
filled octet.
As a result, nonmetals tend to form
anions which are negative ions.
Ionic Size (Ionic Radius)
 An
anion has a larger radius than its
neutral atom.
Question
6. Choose the element from the pair
with the smaller radius.
a) silver (Ag) or the silver ion (Ag1+)
(silver ion)
b) oxygen (O) or the oxygen ion (O2-)
(oxygen)
Question
7. For each of the following pairs,
predict which atom is larger.
a) Mg, Sr
(Sr)
b) Sr, Sn (Sr)
c) Ge, Sn (Sn)
d) Ge, Br (Ge)
e) Cr, W (W)
Question
8. For each of the following pairs,
predict which atom or ion is larger.
a) Mg, Mg2+ (Mg)
b) S, S2– (S2-)
d) Cl–, I– (I-)
e) Na+, Al3+ (Na+)
c) Ca2+, Ba2+ (Ba2+)
Ionization Energy
energy (IE) is the amount
of energy required to completely
remove an electron from a gaseous
atom.
 Removing one electron makes a 1+
ion. The energy required to do this is
called the first ionization energy.
 Ionization
Ionization Energy
What Determines
Ionization Energy (IE)
 Greater
nuclear charge (# of protons)
means greater IE.
 The shorter the distance from the
nucleus, the greater the IE.
Ionization Energy
 As
you go down a group, first IE
decreases.
 This is because the electron is farther
away, thus there is more shielding by
the core electrons from the pull of the
positive nucleus.
Ionization Energy
 All
the atoms in the same period
have the same energy level.
 They have the same shielding, but as
you move across the chart there is
an increasing nuclear charge
because of the increasing number of
protons.
 Therefore, IE generally increases
from left to right.
Question
9. (a) State why ionization energy
decreases as you go down a
group.
(b) State why ionization energy
increases across a period.
Question
10. Choose the element from the pair
with the greater ionization energy.
a) silver (Ag) or iodine (I)
(iodine)
b) oxygen (O) or selenium (Se)
(oxygen)
Question
11. Choose the element from the pair
with the smaller ionization energy.
a) chromium (Cr) or tungsten (W)
(tungsten)
b) sodium (Na) or magnesium (Mg)
(sodium)
Electronegativity
 Electronegativity
is the tendency for
an atom to attract a pair of electrons
to itself when it is chemically
combined with another element.
 Large electronegativity means the
atom pulls the electron toward it.
Electronegativity
Electronegativity
 Electronegativity
decreases down a
group.
 The farther you go down a group, the
farther the electron is away from the
nucleus. It is harder to attract extra
electrons far from the nucleus.
Electronegativity
 As
you go across a row,
electronegativity increases.
 Remember the radius of the atoms
decreases across the periodic table.
 With the smaller size, there is a
greater attraction for the nucleus to
electrons.
Question
12. (a) State why electronegativity
decreases as you go down a
group.
(b) State why electronegativity
increases across a period.
Question
13. Choose the element from the pair
with the greater electronegativity.
a) sodium (Na) or rubidium (Rb)
(sodium)
b) selenium (Se) or bromine (Br)
(bromine)
Question
14. Choose the element from the pair
with the smaller electronegativity.
a) magnesium (Mg) or calcium (Ca)
(calcium)
b) nitrogen (N) or oxygen (O)
(nitrogen)
Summary of the Periodic Trends
 Moving
Left → Right
• Atomic Radius Decreases
• Ionization Energy Increases
• Electronegativity Increases
 Moving
Top → Bottom
• Atomic Radius Increases
• Ionization Energy Decreases
• Electronegativity Decreases
Oxidation Numbers
Oxidation Numbers
 Recall
that metals lose electrons and
form cations which are positive ions,
and nonmetals gain electrons and
form negative ions called anions.
 The ion charge for an element is
called its oxidation number.