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Transcript
Ch 2.4 & 8: The Periodic Table
Atomic Radius (pm)
250
200
150
100
50
0
0
5
10
Atomic Number
15
20
I. Creation of the Periodic Table
The Person Who Organized
the Elements
Dmitri
Mendeleev
1860’s, Russia
Dmitri Mendeleev
• taught chemistry in terms of properties (St.
Petersburg University).
• Mid 1800 - molar masses of elements
were known….Wrote down the elements
in order of INCREASING MASS
• Found a pattern of repeating properties –
elements with SIMILAR CHEMICAL AND
PHYSICAL PROPERITES were GROUPED
together (there were some discrepancies)
Mendeleev’s Periodic Table
Like A Calendar
Mendeleev’s
Periodic Table of Elements
•Only 63 elements known
at time
•He left spaces on the
table for undiscovered
elements.
•He predicted the
properties of elements
that had not yet been
discovered.
Example: Eka-aluminum
He left a space for an undiscovered element
below aluminum.
He predicted the melting point to be low and
the density to be 5.9 g/cm3.
• It was later found and named Gallium with a
density of 5.91 g/cm3 and melting point of
29.7oC.
Gallium
History
Henry Moseley (1913, British)
-Re-organized elements by increasing
atomic number.
-Resolved discrepancies in Mendeleev’s
arrangement.
Periodicity & Periodic Law
250
Atomic Radius (pm)
PERIODICITY—
reoccurrence
(repetition) of
properties at
regular intervals.
200
150
100
50
0
PERIODIC LAW—
when elements are arranged by atomic number, there
is a periodic (i.e. cyclical or repeating) pattern in their
physical and chemical properties. In other words, when
elements are arranged by atomic number, elements
with similar properties occur at regular intervals
0
5
10
Atomic Number
15
20
II. Organization of the
Elements
Metals, Nonmetals, &
Metalloids
• Metals
(lustrous,
conductive,
malleable, ductile,
lose e-/ form +
CATIONS)
• Nonmetals
(dull, insulating,
brittle solids or
gases & liquids, gain e-/ form - ANIONS)
• Metalloids (both metallic & non-metal properties)
Metals
Nonmetals
Metalloids
Row
Series
Period
The periods (rows) of the periodic table
indicate the valence shell.
1
2
3
4
5
6
7
• Horizontal rows are called Periods or Series
There are 7 periods
Column
Group
Family
• Vertical columns are called Groups.
• Elements are placed in columns by
similar properties.
• Also called families
1A
• The elements in the groups in the s
and p block are called the
2A
representative elements3A 4A 5A 6A 7A
8A0
•Main Group
or
Blocks
Representative
Elements
•Transition Metals
•Inner Transition
or
Rare Earth
Metals
Chemical Reactivity
• Families
– Similar valence e- within a group result in similar
chemical properties
Valence Electrons
• The valence electrons occupy the
OUTERMOST (physically farthest away from
the nucleus) shell of an atom.
• Valence electrons play a key role in the
chemical properties of an element.
Periodicity Explained
• The valence shell fills up in a regular pattern
• The valence shell electron configuration
REPEATS
• The properties of atoms therefore repeat
when placed in order of Atomic Number
H
1
1s1
3
1s22s1
Li
Na
11
1s22s22p63s1
1s22s22p63s23p64s1
K
19
Rb
37
Cs
55
Fr
87
1s22s22p63s23p64s23d104p65s1
1s22s22p63s23p64s23d104p6 5s24d10 5p66s1
1s22s22p63s23p64s23d104p65s24d105p6
6s24f145d106p67s1
He
2
1s
1s22s22p6
1s22s22p63s23p6
2
Ne
10
Ar
18
1s22s22p63s23p64s23d104p6
Kr
1s22s22p63s23p64s23d104p65s24d105p6
Xe
1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p6
36
54
Rn
86
s- block
s1
s2
• Alkali metals all end in s1
• Alkaline earth metals all end in s2
• really have to include He but it fits
better later.
• He has the properties of the noble
gases (because the first shell is filled
when 1s2)
Transition Metals -d block
d1 d2 d3
d5
d5 d6 d7 d8 d9 d10
The p-block
p1 p 2
p3
p4
p5
p6
f - block
• inner transition elements
f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14
Group 1- Alkali Metals
• 1 valence electron (ns1), forms 1+
ion when it loses an e-.
• VERY REACTIVE WITH WATER!
Metal + H2O  H2 + MOH (a base),
reactivity increases from top to
bottom
• Note: Hydrogen, a nonmetal, is
located in the first column because
it has one valence electron.
Note:
• Sodium and Potassium are stored in oil to
keep them from reacting with oxygen and
water in the air.
• Cesium is stored in glass tubes of argon gas(
an inert gas).
Group 2- Alkaline Earth Metals
• Have 2 valence electrons.
Beryllium
• Form 2+ ions when 2 eare lost.
Magnesium
Calcium
Strontium
Barium
Radium
Group 2- Alkaline Earth Metals
Reactions with Water
• Be does not react with water.
• Mg reacts with hot water.
• Ca, Sr, Ba react easily with cold water.
• Which way along the group does reactivity
increase?
Notes:
• There is Mg in chlorophyll
C55H72O5N4Mg
• Calcium is in your bones,
chalk, limestone, toothpaste,
pearl (all as calcium
carbonate).
Transition Metals
• Groups 3-12 are the transition metals
• They have metallic properties such as high
luster, good conductivity and are relatively
non-reactive.
• They form very colorful compounds.
p-block metals (Groups 13-18)
• These are main group elements.
• The groups are sometimes named for the first
element in the group. For example, the carbon
family
• Group 17 is the halogens. They are highly
reactive and like to react with metals to form
salts. Why would they like to react with group
1 elements?
• Group 18 is the noble gases. They have filled
valence shells and are non-reactive.
PERIODIC & GROUP
TRENDS
Atomic Radius
(aka Atomic Size)
}
Radius
•First problem: Where do you
start measuring.
•The electron cloud doesn’t
have a definite edge.
•They get around this by
measuring more than 1 atom
at a time.
•Commonly measure in pm
(ex. I is 140 pm)
• Atomic Radius = half the distance between two nuclei
of a diatomic molecule.
Trends in Atomic Size
• Influenced by two factors.
– # of Shells – “Shielding”
• More shells- electrons are further away from
nucleus.
– Charge on nucleus – Zeff
• More charge (# of protons) pulls electrons in
closer.
Atomic Radius
• Atomic Radius
K
Atomic Radius (pm)
250
Na
200
Li
150
100
Ar
Ne
50
0
0
5
10
Atomic Number
15
20
GROUP TREND:
Atomic Size (atomic radii)
H
Li
Na
K
Rb
• As we go down a
group
• Each atom adds
another shell
• So the atoms logically
get bigger.
PERIODIC TREND:
Atomic Size (atomic radii)
• As you go across a period the radius gets
smaller.
1. Same number of shells (shielding).
2. More protons = increased nuclear charge.
3. Increased attraction for outermost electrons.
• So… the atoms should get smaller.
Na
Mg
Al
Si
P
S Cl Ar
Increases
DOWN
Atomic Radius
Decreases RIGHT
Atomic Radius
• Why larger going down?
– Higher shells have larger orbitals
– Shielding - core e- block or screen the attraction
between the nucleus and the valence e-
• Why smaller to the right?
– Increased nuclear charge without additional shielding
pulls e- in tighter
Shielding Effect
Valence
+
nucleus
Kernel electrons block
the attractive force of
the nucleus from the
valence electrons
-
-
-
Electron
Shield
“kernel”
electrons
Electrons
Examples
• Which atom has the larger radius?
Be
or
Ba
Ca
or
Br
Example:
Which is larger: a lithium atom or a fluorine
atom?
A lithium atom
Which is larger: an arsenic atom or a sulfur
atom?
An arsenic atom
First Ionization Energy
–The amount of energy (J) required to completely
remove
one e- from a
GASEOUS atom (a
measure of how
strongly the atom is
holding its electrons)
© 1998 LOGAL
*Ionization—removal of electrons
1st: Atom A +I1  Ion A+ + free e2nd: Ion A+ + I2  Ion A+2 + free e-
Ionization Energy
1st Ionization Energy (kJ)
• First Ionization Energy
He
2500
Ne
2000
Ar
1500
1000
Li
500
Na
K
0
0
5
10
Atomic Number
15
20
e-
e-
GROUP trend:
Ionization Energy (IE)
• As you go down a group first IE
decreases because….
• The electron is further away…because
there are more shells between the
nucleus and the e- thus….
• Shielding the e- from + nucleus.
• IT IS EASIER (takes LESS energy) to
remove an electron
Symbol IE (J)
Li
Be
B
C
N
O
F
Ne
520
900
800
1086
1402
1314
1681
2080
Periodic trends: Ionization Energy (IE)
e-
e-
• All the atoms in the same period have the
same number of shells.
• Same number of shells = same shielding, but
the number of protons increase, therefore….
• Increasing nuclear charge …helps pull e- in
tighter, therefore it is harder to remove.
• So IE generally increases from left to right.
Ionization Energy
• Why opposite of atomic radius?
– In small atoms, e- are close to the nucleus where the
attraction is stronger. SMALL atoms will require
MORE energy to remove e-; LARGE atoms will
require LESS energy to remove e-
• Why small exceptions within each period?
– Stable e- configurations (s2, s2p3) don’t want to lose
e- (will elevate the IE more than what’s expected
from the trend)
Examples
• Which atom has the higher 1st I.E.?
N
or Bi
Ba
or Ne
SUCCESSIVE Ionization Energy
Large jump in I.E. occurs when a CORE e- is
removed.
(Write in the equations below)
Mg
Core e-
1st I.E.
2nd I.E.
3rd I.E.
736 kJ
1,445 kJ
7,730 kJ
Successive Ionization Energy
Large jump in I.E. occurs when a CORE eis removed.
Al
Core e-
1st I.E.
2nd I.E.
3rd I.E.
4th I.E.
577 kJ
1,815 kJ
2,740 kJ
11,600 kJ
Electronegativity (EN)
• The relative ability of an atom IN A BOND
to attract electrons to itself (Linus Pauling)
Which has greater electronegativity in
each compound?
Electronegativities of Some Elements
Element
F
Cl
O
N
S
C
H
Na
Cs
Pauling scale
4.0
3.0
3.5
3.0
2.5
2.5
2.1
0.9
0.7
Note
• Most electronegative element is F
(EN 4.0)
• Least electronegative stable element is Cs
(EN 0.7)
PERIODIC TREND: Electronegativity (EN)
e-
e-
• From left to right atoms become smaller,
with greater nuclear charge.
e - are attracted by the + charged
nucleus.
•therefore EN will increase from left to right
GROUP TRENDS: Electronegativity (EN)
e-
• Down a group atoms become larger, and
have greater nuclear charge.
• e - are attracted by the increased
+charge but shielding also increases
…which has a greater influence!!
• therefore EN will decrease as you
go down a group or family
e-
Electronegativities
F
Cs
Probably the most important trend.
(1) Useful for explaining chemical reactivity.
(2) It is also useful for categorizing the elements.
•Metals are electropositive and lose electrons,
•nonmetals are electronegative and gain electrons,
•metalloids have intermediate electronegativities
Electronegativity & BONDING
• When two atoms have similar electronegativities,
they pull on the electrons equally and form a
NONPOLAR COVALENT BOND. Example, H-H.
• When the atoms are different, the atoms pull on
the electrons unevenly, and form either a POLAR
COVALENT BOND (if they’re not very different,
and are still sharing, but not equally) or an IONIC
BOND (if they are very different an the electron is
completely transferred). Example, HCl and NaCl
LARGE EN DIFFERENCES 
IONIC COMPOUNDS/ BONDING
(full transfer of e-)
Ions
Ion
• Atom or group of atoms
that have lost or gained
electrons & have a charge
• Charges not balanced;
differing number of protons
and electrons
Ions are formed when
atoms gain or lose
electrons. NEVER ADD
OR SUBTRACT
PROTONS!
Cation
(ca+ion)
• positively charged ion
• formed when neutral
atoms LOSE electrons
ex.
Determine
+
:
a. K
+
#p
b.
2+
Zn
c.
+3
Al
and
#e
Anion (Nion)
• Negatively charged
ion
• formed when neutral
atoms GAIN electrons
ex.
Determine
:
a. Cl
b.
-2
S
c.
3P
+
#p
and
#e
Ion charges can be
determined from the
periodic table:
metals → cations
non-metals → anions
OCTET RULE – atoms
will gain or lose e- or
bond to form a full
valence shell (usually 8
e-, except for He).
ex. Cation or anion?
Charge?:
a. Nitrogen
b. Sodium
c. Sulfur
Naming Cations
1. Group IA / 2A & Ag, Cd, Zn, &
Al: Use element name
2. d-block and p-block: Most use
Roman numerals for charge
ex. Name the following
cations:
2+
a. Mg
b.
2+
Cu
c.
+3
Fe
Naming Anions
Add “-ide” to
element name.
ex. Name the following
anions:
2a. O
b.
F
c.
-3
N
Ionic Radius
• Ionic Radius
Cations (+)
lose esmaller
Anions (–)
gain elarger
© 2002 Prentice-Hall, Inc.
Ion Sizes
+
Li,152 pm
Li + , 78 pm
3e and 3p
2e and 3 p
Forming
a cation.
• CATIONS are SMALLER than the atoms from
which they come.
• **Losing a valence electron
is
not
• just the lose of an e•
But the lose of the valence shell!!!
Ion Sizes
F, 71 pm
F - , 133 pm
9e and 9p
10 e and 9 p
Forming
an anion.
• ANIONS are LARGER than the atoms from which
they come.
• The addition of an electron increases the
repulsion in the valence shell and so size
INCREASES.
• Trends in ion sizes are the same as atom sizes.
Ionic Radii Trend
• same general trend as atomic
radii
• anions are larger than neutral
atom
• cations are smaller than neutral
atom
Trends in Ion Sizes
Figure 8.13
Examples
• Which particle has the larger radius?
S or
2S
Alor
3+
Al
Periodic & Group trends: Electron
Affinity(EA)
• The energy change associated with adding an
electron to a gaseous atom.
• e - are attracted by a + charge.
– Therefore more protons more attraction
– But remember “shielding”
• Easiest to add to group 7A.
• Gets them to full energy level.
• Increase from left to right atoms become smaller,
with greater nuclear charge.
• Decrease as we go down a group.
Periodic & Group trends: Electron
e-
Affinity(EA)
• From left to right atoms become smaller,
with greater nuclear charge.
e - are attracted by the + charged
nucleus.
•therefore EA will increase from left to right
e-
Periodic & Group trends: Electron
e-
Affinity(EA)
• Down a group atoms become larger, and
have greater nuclear charge.
• e - are attracted by the increased
+charge but shielding also increases
…which has a greater influence!!
• therefore EA will decrease as you
go down a group or family
e-
Nuclear charge increases
Shielding increases
Atomic radius increases
Ionic size increases
Ionization energy decreases
Electronegativity decreases
Summary
Shielding is constant
Atomic Radius decreases
Ionization energy increases
Electronegativity increases
Nuclear charge increases