Transcript ChemChapter11

Chapter 11 & 12
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John Newlands
• In 1864, noticed
when the
elements were
arranged in
order of
increasing
atomic mass,
their properties
repeated every
eight elements.
– THE LAW OF
OCTAVES
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Meyer & Mendeleev
• In 1869, published almost identical versions with
the elements in order of increasing atomic mass
and in columns with similar properties.
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Mendeleev
• Mendeleev is given more credit than
Meyer BECAUSE:
– He published his table first
– He better demonstrated his table
• Suggested some of the previously
measured masses were incorrect
• Left blanks for not yet discovered
elements
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Mosley
• In 1913, using X-rays, he
discovered a unique number of
protons in the nuclei of atoms for
each element.
• Today the elements are arranged
in order of increasing atomic
number
Periodic Law
• There is a periodic repetition of
chemical and physical properties
of the elements when they are
arranged in order of increasing
atomic number
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Arrangement of the Periodic Table
•
Groups/Families
– 18 vertical columns (↑↓)
– Two Labeling Systems
1. Number-and-letter system
- 1A through 8A columns (representative
elements)
- 1B through 8B short columns (transition
elements)
2. Number system, Group 1 to Group18
• Periods
– 7 horizontal rows (↔)
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GROUPS/FAMILIES
PERIODS
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Arrangements of the Periodic Table
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Non-Metals
Create a
Legend
Metalloid
Non-Metals
Metals
Metals
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•
•
•
•
•
Metals
Shiny
Solid at room temperature
Good conductors of heat and electricity
Malleable
Ductile
Group 1
Alkali Metals
Group 2
Alkaline Earth Metals
Groups 3-12
Transition Metals
Lanthanide & Actinide Groups
Inner Transition Metals or Rare Earth Metals
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Valence Electrons-electrons in an atom’s
outermost orbitals
Group #
Group Name
# Valence e1(1A)
Alkali Metals
1
2(2A)
Alkaline Earth Metals
2
3-12
Transition Metals
varies
13(3A)
Boron Group
3
14(4A)
Carbon Group
4
15(5A)
Nitrogen Group
5
16(6A)
Oxygen Group
6
17(7A)
Halogens
7
12
18(8A)
Noble Gases
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Nonmetals & Metalloids
• Nonmetals
– Dull
– Generally gases or
brittle solids at room
temperature
– Poor conductors of heat
and electricity
• Metalloids
B
Nonmetals →
Si
Ge As
Sb Te
Po At
– Elements with physical
and chemical properties
←Metals
of both metals and
nonmetals
– Rest on the “stair-step”
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Element Placement
Why are elements put into groups/families together?
Because they have similar chemical properties
Why do elements have similar chemical properties?
Because they have the same number of
valence electrons
Group 1 – Alkali Metals
Period 2
Lithium
1s22s1
[He]2s1
Period 3
Sodium
1s22s22p63s1
[Ne]3s1
Period 4
Potassium
1s22s22p63s23p64s1
[Ar]4s1
ALL ELEMENTS IN GROUP 1 (ALKALI METALS) HAVE ONE VALENCE ELECTRON 15
Dot Diagrams for Representative Elements
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Representative Elements
• s-block elements
– Groups 1&2, hydrogen & helium
– Two columns Maximum 2 electrons fills the 1 s
orbital.
– Valence electrons occupy outermost s sublevels
only.
• p-block elements
– Groups 13-18 (except helium)
– Six columns Maximum 6 electrons fills the 3 p
orbitals.
– Valence electrons include a full outermost s
sublevel and a filled or partially filled p sublevel.
Period number is equal to the principle energy level
where the valence electrons are located
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Transition Elements
• d-block elements
–Groups 3-12
–Ten columns Maximum 10
electrons fills the 5 d orbitals.
–Valence electrons include a
full outermost s sublevel and
a filled or partially filled d
sublevel.
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Inner Transition metals or
Rare Earth Metals
• f-block elements
–Fourteen columns Maximum
14 electrons fills the 7 f orbitals.
–Lanthanide & Actinide Groups
–Full or partially full outermost s
sublevel, and full or partially full
outermost f sublevel.
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Periodic Trends
• Atomic Radii: The distance from the nucleus to
the outer edge of the electron cloud
• Ionic Radii:
The distance from the nucleus to the
edge of the electron cloud of an ion.
• Ionization Energy:
The amount of energy
needed to remove an electron from an atom
• Electron Affinity:
The energy change that
occurs when an atom gains an electron.
• Electronegativity:
The ability of an atom to
attract electrons toward itself from a covalent chemical
bond.
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Atomic Radius
• Half the distance between two nuclei of
identical atoms that are chemically
bonded together
• Down the group
– atomic radius increases,
because…
• Across the period
– atomic radius decreases,
becauses….
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Practice Atomic Radius
• Which has the larger atomic radii of the
following?
B or Al
Na or Mg
F or Cl
• Which has the smaller atomic radii of
the following?
H or He
K or Cs
N or Ne
• Circle the one with the largest atomic
radius and underline the one with the
smallest.
C, Si, Ge
V, Cr, W
N, Mg, Ca
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Atomic Radius Decreases
H
He
B
Atomic Radius Increases
Na Mg
K Ca
Cs
C N
Al Si
V Cr
F Ne
Cl
Ge
W
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Ionization Energy
• The amount of energy required to remove an
electron from the atom
(how tightly an atom holds on to its electrons)
• A general term for the energy required to
remove an electron from an orbital in an
atom. Think of it also as the energy required
to make a cation.
• Down a group
– ionization energy decreases, because…
• Across a period
– ionization energy increases, because…
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Practice Ionization
Energy
• Which has the greater ionization
energy?
Ne or Ar
N or O
Sc or Ti
• Which has the smaller ionization
energy?
Al, Si, P
K, Rb, Sr
Be, Mg, Ca
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Ionization Energy Decreases
Ionization Energy Increases
Be
N
Mg
Al Si P
O
Ne
Ar
K Ca Sc Ti
Rb Sr
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Electron Affinity
• The electron affinity of an element is the energy
given off when a neutral atom in the gas phase
gains an extra electron to form a negatively
charged ion.
• The attraction to additional electrons
• A fluorine atom in the gas phase, for example,
gives off energy when it gains an electron to form
a fluoride ion.
• Down the group
– electron affinity decreases, because
• Across the period
– electron affinity increases, because
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Practice Electron Affinity
• Which has the smaller electron affinity?
Te or Xe
Ag or Au
Zn or Br
• Which has the greater electron affinity?
F, Cl, Br
B, C, Si
Ni, Cu, Au
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Electron Affinity Increases
Electron Affinity Decreases
B
C
F
Si
Cl
Ni Cu Zn
Ag
Br
Te
Xe
Au
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• Octet Rule
Ionic Radius
– Atoms tend to gain, lose, or share electrons in order to
achieve a full outer energy level (typically 8 are needed)
– Groups 1 A – 3 A, loses Valence e-( 1 to 3 e-).
– Group 4 A, share Valence e-.
– Groups 5 A – 7 A, gain electrons ( 3 to 1 e-).
• Noble Gases, Outer most shell are full. These elements
don’t gain nor lose e-, Non-reactive.
• Ion
– An atom that has an overall charge due to the gaining or
losing of electrons
– Cation, positive charges (Groups 1A – 4A)
– Anion, negative charges (Groups 5A – 7A)
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Ionic Radius Comparisons
• Metals have LOW ionization
energy and electron affinity
Atom size decreases
– They lose electrons to form
positively charged ions
– Positive charged ions are smaller
than the original atom
• Nonmetals have HIGH ionization Atom size increases
energy and electron affinity
– They gain electrons to form
negatively charged ions
– Negatively charged ions are larger
than the original atom
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Ionic Radius Practice
• Which is the smaller of the two?
Lithium ion or Lithium atom
Chlorine ion or Chlorine atom
• Underline the following that will form a positively
charged ion and circle the ones that will form a
negatively charged ion.
Mg
F
Al
Cu
Br
N
S
K
• How will the radius of each of the above change
when an ion is formed?
Mg
F
Al
Cu
Br
N
S
K
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N
Li
Mg
K
Al
S
F
Cl
Cu
Br
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Electronegativity
• The attraction an atom has for
electrons in a covalent bond.
• The ability of an atom to attract
electrons in a chemical bond.
• Down the group
– Electronegativity values decrease,
because
• Across the period
– Electronegativity values increase, because
*Noble
gases are the exception to
this rule.
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Electronegativity
Electronegativity of 0.0
to 0.3 is a non-polar
covalent bond.
Electronegativity of 0.4
to 1.9 is a polar
covalent bond.
Electronegativity of 2.0
to 4.0 is an ionic bond.
You must subtract the values of
electronegativity to determine it the
bond is covalent, polar covalent or
ionic
Covalent shares the
electrons equally
Polar is slightly
negative on one side
Ionic has electrons
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captured by one atom.
Electronegativity Practice
• Which has the greater electronegativity
value?
B or N
Si or Sn
Cr or W
• Which has the smaller electronegativity
value?
Rb, Sr, Y
Ge, In, Sn
As, Se, S
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Electronegativity Decreases
Electronegativity Increases
B
Cr
Rb Sr Y
N
Si
S
Ge As Se
In Sn
W
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How can you remember all of these trends?
• Remember the cartoon characters Mighty Mouse
and Foghorn Leghorn?
• Put Mighty Mouse in the top right hand
corner of the periodic table and Foghorn
Leghorn in the bottom left; now draw the
trends and it should make sense. Mighty
Mouse, aka fluorine, is strong
(electronegative and electron affinity) and
small (atomic radii), while Foghorn is bigger
(larger radius) and slower (less electron
affinity), etc.
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Name: ______________ Date: ________Block: ________
CUMULATIVE REVIEW, Due at the End of Class
1) Between these elements Ga, In, & Tl:
– Which has the highest ionization energy?
– Which has the smallest atomic radius?
2) Which is the smallest radius: (a) an atom of
sodium, (b) an ion of sodium, or © an atom of
potassium?
3) Between these elements (a) zinc, (b) arsenic, or
(b) bromine:
- Which has the greatest electron affinity?
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- Which has the lowest ionization energy?