Download Unit 3 Notes: Periodic Table Notes

Unit 3 Notes: Periodic Table Notes
• John Newlands proposed an organization system based on increasing
atomic mass in 1864.
• He noticed that both the chemical and physical properties repeated every 8
elements and called this the ____Law of Octaves ___________.
• In 1869 both Lothar Meyer and Dmitri Mendeleev showed a connection
between atomic mass and an element’s properties.
• Mendeleev published first, and is given credit for this.
• He also noticed a periodic pattern when elements were ordered by
increasing ___Atomic Mass _______________________________.
• By arranging elements in order of increasing atomic mass into columns,
Mendeleev created the first Periodic Table.
• This table also predicted the existence and properties of undiscovered
elements.
• After many new elements were discovered, it appeared that a number of
elements were out of order based on their _____Properties_________.
• In 1913 Henry Mosley discovered that each element contains a unique
number of ___Protons________________.
• By rearranging the elements based on _________Atomic Number___, the
problems with the Periodic Table were corrected.
• This new arrangement creates a periodic repetition of both physical and
chemical properties known as the ____Periodic Law___.
Periods are the ____Rows_____
Groups/Families are the Columns
Valence electrons across a period are
in the same energy level
There are equal numbers of valence
electrons in a group.
• When elements are arranged in order of increasing _Atomic Number_,
there is a periodic repetition of their physical and chemical properties
1
• Family (Group): ___Columns (vertical)______; tells the number of electrons
in the _Outer___ Energy level, called __Valence Electrons________ (only
for representative elements)
• Period (Series): __Rows (horizontal)____; tells the number of ____Energy
Levels__________ an atom has; the number of electrons __Increases__
across a period
• Representative Elements: Groups __1A through 8A _ (called the s and p
blocks) (Columns 1, 2, 13, 14, 15, 16, 17, and 18)
• Valence Electrons: e- in the ___outer most energy level____; farthest away
from the __nucleus (protons)___; the e- with the ___most reactive____
Energy; the e- involved with ___Bonding____ (transferred or shared)
• Metals: most of the periodic table, located to the __Left___ of the “stair-step”
Properties- good conductors of _heat_ and _Electricity_; they also are
__ Malleable___; __ Ductile____; _ High Density, BP and MP_____
• Nonmetals: to the Right of the “stair-step”, located in the upper corner of
P.T._
• Although five times more elements are metals than nonmetals, two
of the nonmetals—hydrogen and helium—make up over 99 per cent
of the observable Universe
• Properties- mostly _ Brittle __, but a few _low luster______ and _poor
conductors__; they have _ low density, low MP and BP__
• Metalloids: also called _semi-metals__, located _along_ the “stair-step”
• Properties - __ similar __ to both metals and nonmetals
•
•
•
•
•
Some metalloids are shiny (silicon), some are not (gallium)
Metalloids tend to be brittle, as are nonmetals.
Metalloids tend to have high MP and BP like metals.
Metalloids tend to have high density, like metals.
Metalloids are semiconductors of electricity – somewhere between
metals and nonmetals. This makes them good for manufacturing
computer chips.
2
Element
Lithium
Germanium
Sulfur
Symbol
Li
Ge
S
Group #
1A(1)
4A(14)
6A(16)
# of valence e-
1
4
6
Period #
2
4
3
# of E levels
2
4
3
Type of element
M
M
NM
Periodic Trends:
1. Atomic Size
- __Decreases__ from left to right across a period (smaller)
- __Increases___ from top to bottom down a group (larger)
Why?
- as you go across a period, (same __energy level__), e- are
_added_but _pulled closer to the nucleus___
- as you go down a group, you add ___energy levels___
2. Ionization Energy: the amount of E needed to _remove / lose_ an electron
- __Increases__ from left to right across a period
- __Decreases____ from top to bottom down a group
Why?
3
- as you go across a period, e- feel stronger attraction from nucleus
(protons)___,
_Energy___ to remove e-, ____Ionization___ E necessary
as you go down a group, more __Energy_, _Decreases_ to remove
outermost e- because they are further away from the Nucleus (protons)
3. Electronegativity: the tendency for an atom to __attract___ electrons;
exclude Noble Gases!
- __Increases__ from left to right across a period (except Noble Gases)
- __Decreases____ from top to bottom down a group
Why?
- as you go across a period, e- feel ___more__ attraction from nucleus
_Protons_____ to pull in more e- as you go down a group, more _shielding__ from inner e-,
__hinders the nucleus ability__ to attract more e-
4. Ionic Size:
Cations:__positive_ ions; metal atoms that ___lose__ electrons
4
- __smaller__ than corresponding neutral atom
Why?
- __fewer__ e-, so it’s _easier_ for protons to pull in remaining eAnions:__Negative___ ions; nonmetal atoms that _gain_ electrons
- ___larger____ than corresponding neutral atom
Why?
- _more_ e-, so it’s __harder_ for protons to pull in outermost eShielding:
The ability of the _inner (lower levels)_ electrons to _shield (reduce)_ the pull
of the _protons_ on the _outer (higher levels)__ electrons.
“Shielding effect”_increase_ as you add Energy levels (move down a group)
Quantum Model Notes
• Heisenberg's Uncertainty Principle- Can determine either the _velocity or the
position of an electron, cannot determine both.
• Schrödinger's Equation - Developed an equation that treated the hydrogen
atom's electron as a wave.
o Only limits the electron's energy values, does not attempt to describe the
electron's path.
• Describe probability of finding an electron in a given area of orbit.
• The Quantum Model- atomic orbitals are used to describe the possible position
of an electron.
Orbitals
• The location of an electron in an atom is described with 4 terms.
5
o Energy Level- Described by intergers. The higher the level, the more energy
an electron has to have in order to exist in that region.
o Sublevels- energy levels are divided into sublevels. The # of sublevels
contained within an energy level is equal to the integer of the energy level.
o Orbitals- Each sublevel is subdivided into orbitals. Each orbital can hold 2
electrons.
o Spin- Electrons can be spinning clockwise (+) or counterclockwise (-) within
the orbital.
Periodic Table Activity:
Complete the table on page 21 with the information found on pages 18-20. When complete color each group in
a different color in the periodic table.
The Periodic Table Notes:
Historical development of the periodic table: Highlights
• Mendeleev (1869): Put the elements into columns according to their properties. Generally ranked
elements by increasing atomic mass.
• Moseley (1911): Periodic table arranged by atomic number
Top table: Metals, nonmetals, and metalloids
• Metals: Explain the electron sea theory, and as you explain each of the properties below, discuss how
they are explained by the electron sea theory. Also make sure to explain that these are general
properties and may not be true for all metals.
o Malleable: Can be pounded into sheets.
o Ductile: Can be drawn into wires
o Good conductors of heat and electricity
o High density (usually)
o High MP and BP (usually)
o Shiny
o Hard
• Nonmetals: Explain how the bonds between the atoms are highly localized, causing each of the
properties below. Again, emphasize that these are general properties and may not be true for all
nonmetals.
o Brittle
o Poor conductors of heat and electricity
o Low density
o Low MP and BP (many are gases)!
• Metalloids: The bonding in metalloids is between that of metals and nonmetals, so metalloids have
properties of both.
o Some metalloids are shiny (silicon), some are not (gallium)
o Metalloids tend to be brittle, as are nonmetals.
o Metalloids tend to have high MP and BP like metals.
o Metalloids tend to have high density, like metals.
6
Metalloids are semiconductors of electricity – somewhere between metals and nonmetals. This
makes them good for manufacturing computer chips.
Structure of the periodic table
• Families/groups (the terms are synonymous and will be used interchangeably)
o These are elements in the same columns of the periodic table.
o Elements within families/groups tend to have similar physical and chemical properties.
o They have similar chemical and physical properties because they have similar electron
configurations.
Example:
Li = [He] 2s1, Na = [Ne] 3s1 – each has one electron in the outermost
energy level.
o Explain that s- and p-electrons in the outermost energy level are responsible for the reactions
that take place.
Valence electrons: The outermost s- and p-electrons in an atom.
Show them how to find the number of valence electrons for each atom and explain that
they are only relevant for s- and p- electrons. Do several examples.
o
Periods: Elements in the same rows of the periodic table
o Elements in the same period have valence electrons in the same energy levels as one another.
o Though you’d think this was important, it has very little effect on making the properties of the
elements within a period similar to one another.
The closer elements are to each other in the same period, the closer are their chemical
and physical properties.
• Other fun locales in the periodic table:
o Main block elements: These are the s- and p- sections of the periodic table (groups 1,2, 13-18)
o Transition elements: These are the elements in the d- and f-blocks of the periodic table.
The term “transition element”, while technically referring to the d- and f-blocks, usually
refers only to the d-block.
Technically, the d-block elements are the “outer transition elements”
Technically, the f-block elements are the “inner transition elements”
Major families in the periodic table:
(Show them examples of these elements – if available – and color each family as I discuss their properties)
• Group 1 (except for hydrogen) – Alkali metals
o Most reactive group of metals
o Flammable in air and water
o Form ions with +1 charge
o Low MP and BP (MP of Li = 181º C, Na = 98º C)
o Soft (Na can be cut with a knife)
o Low density (Li = 0.535, Na = 0.968)
• Group 2: Alkaline earth metals
o Reactive, but less so than alkali metals
o React in air and water (show Ca reacting in water)
o Form ions with +2 charge
o Low MP and BP, but higher than alkali metals (MP of Ba= 302º C, Mg = 649ºC
o Soft, but harder than alkali metals
o Low density, but higher than that of alkali metals (Ca = 1.55, Mg = 1.74).
• Groups 3-12: (Outer) transition metals
o Note: These are general properties and may vary from transition metal to transition metal!
There are many exceptions to each of these rules!
o Stable and unreactive.
o Hard
•
7
•
•
•
•
•
•
•
•
o High MP and BP (Fe = 1535º C, Ti = 1660º C).
o High density (Fe = 7.87, Ir = 22.4)
o Form ions with various positive charges (usually include +2 and several others)
o Used for high strength/hardness applications, electrical wiring, jewelry
Inner Transition Metals: Lanthanides and actinides
o Lanthanides (4f section)
Also called the rare earth metals, because they’re rare.
Usually intermediate in reactivity between alkaline earth metals and transition metals.
High MP and BP
Used in light bulbs and TV screens as phosphors.
o Actinides (5f section)
Many have high densities
Most are radioactive and manmade
Melting points vary, but usually higher than alkaline earth metals.
Reactivity varies greatly
Used for nuclear power/weapons, radiation therapy, fire alarms.
Group 13: Boron Group
Group 14: Carbon Group
Group 15: Nitrogen Group
Group 16: Oxygen Group
Group 17: Halogens
o The most highly reactive nonmetals.
o Highly volatile – F and Cl are gases, Br is a volatile liquid, and I is an easily sublimed solid.
o Strong oxidizers – they readily pull electrons from other atoms.
o Diatomic – form molecules with formula of X2
o Form ions with -1 charge
o Used in water treatment and chemical production – Cl2 was used as a chemical weapon in World
War I.
Group 18: Noble Gases
o Highly unreactive
o Used to provide the atmosphere in situations where you don’t want chemical reactions to occur
(light bulbs, glove boxes, etc).
Hydrogen – “The Weirdo”
o Has properties unlike any other element
o Diatomic – H2
o Can form either a +1 or -1 charge
o Relatively unreactive unless energy is added (under most conditions) – it can form explosive
mixtures with oxygen (as it did in the Hindenburg explosion)
8
Groups on the Periodic Table Summary Sheet:
Group
Examples
of Words
used
Alkali
Metals
Alkaline
Earth
Metals
Location on
Periodic Table
Group 1,
Group 3-12,
etc
Metals, Non-Metals,
Metalloids?
Common
Charge(s)?
Reactivity
Metal
+1
Highly
reactive,
unreactive
Interesting
Information
It can be cut
with a plastic
knife
M
+1
Y
N
Any Name in
Family 1
1
M
+2
Y
N
Any Name in
Family 2
2
M
+2
N
N
Any Name in
Family 3-12
2
2
1
2
Example:
Number of Valance
Electrons
Element’s Name
Transition
Metals
(Outer)
3-12
Inner
Transition
Metals
3 (atomic #
58-71, 90103)
M
+2
N
N
Any Name
atomic
number 58-71,
90-103
Halogens
17
NM
-1
Y
Y
Any Name in
Family 17
7
Noble
Gases
18
NM
0
N
NA
Any Name in
Family 18
8
1
M
+1
Y
NA
Hydrogen
1
Hydrogen
9
Groups
O Alkali Metals
O Alkali Earth Metals
O Boron Group
O Carbon Group
O Hydrogen
O Halogen s
O Inner Transition Metals
O Metaloids
O Nitrogen Group
O Noble Gasses
O Oxygen Group
O Transition Metals
Group 1
1.00794
H
.
1
2
Hydrogen
6.941
9.01218
Li
.
Be .
3
Lithium
22.98977
P
E
R
I
O
D
4
Beryllium
24.305
Na
Mg
11
12
Sodium
Magnesium
39.0983
40.08
3
44.9559
4
5
47.88
Periodic Table of the Elements
Atomic Mass
Mass numbers in parenthesis are those of the
most stable or most common isotope
He .
14
13
Name
14
10.81
B
.
C
8
9
58.9332
11
10
58.69
63.546
12
65.39
N
.
O
7
28.0855
17
15.9994
.
F
20.179
.
Ne
9
Oxygen
30.97376
Helium
18.998403
8
Carbon Nitrogen
26.98154
55.847
.
6
Boron
16
14.0067
5
Transition Elements
15
12.01
Metals
54.9380
2
Nonmetals
Silicon
7
51.996
4.00260
Si
Symbol
Atomic Number
6
50.9415
18
28.0855
Fluorine
32.06
35.453
10
Neon
39.948
Al
Si
P
S
Cl
Ar
13
14
15
16
17
18
Silicon
Phosphorus
Aluminum
69.72
72.59
74.9216
Sulfur
Chlorine
78.96
79.904
Argon
83.80
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
Potassium
85.4678
Calcium
87.62
Scandium
88.9059
Titanium Vanadium Chromium Manganese
91.224
92.9064
95.94
(98)
Iron
Cobalt
101.07
102.906
Nickel
Copper
106.42
107.868
Zinc
112.41
Gallium Germanium
114.82
118.71
Arsenic
121.75
Selenium
Bromine
127.60
127.60
Krypton
131.29
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
Rubidium Strontium
132.905
137.33
Yttrium
Zirconium Niobium MolybdenumTechnetium Ruthenium Rhodium Palladium
138.906
178.49
180.948
183.85
186.207
190.2
192.22
195.08
Silver
196.967
Cadmium
200.59
Indium
204.383
Tin
207.2
Antimony Tellurium
208.980
Iodine
(209)
(210)
Xenon
(222)
Cs
Ba
La
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
55
56
57
72
73
74
75
76
77
78
79
80
81
82
83
84
85
86
Lead
Bismuth
Polonium
Astatine
Cesium
(223)
Barium Lanthanum Hafnium
226.025
227.028
(261)
Tantalum
(262)
Tungsten
(263)
Rhenium
(262)
Osmium
(265)
Iridium
(266)
Platinum
(269)
Fr
Ra
Ac
Rf
Db
Sg
Bh
Hs
Mt
Ds
87
88
89
104
105
106
107
108
109
110
Francium
Radium
Actinum Rutherfordium Dubnium Seaborgium Bohrium
140.12
Lanthanoid Series
144.24
Mercury
(277?)
Uuu Uub
111
Thallium
(?)
(289?)
(289?)
Uut
Uuq
Uuh
Uuo
113
114
116
118
112
Hassium Meitnerium Dormstadtium Unununium Ununbium Ununtrium Ununquadium
(145)
150.36
151.96
157.25
158.925
Radon
(293?)
162.50
164.930
167.26
Ununhexium
Ununoctium
168.934
174.967
173.04
Ce
Pr
Nd
Pm
Sm
Eu
Gd
Tb
Dy
Ho
Er
Tm
Yb
Lu
58
59
60
61
62
63
64
65
66
67
68
69
70
71
.
Cerium
232.038
Actinoid Series
140.908
Gold
(272?)
Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium
231.036
238.029
237.048
(244)
(243)
(247)
(247)
(251)
(252)
Erbium
(257)
Thulium
(258)
Ytterbium
(259)
Lutetium
(260)
Th
Pa
U
Np
Pu
Am
Cm
Bk
Cf
Es
Fm
Md
No
Lr
90
91
92
93
94
95
96
97
98
99
100
101
102
103
. Protactinium Uranium Neptunium Plutonium Americium
Thorium
10
Curium
Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium
Orbital Diagrams
Energy Level
• Indicates relative sizes and energies of atomic orbitals. Whole numbers, ranging
from 1 to 7.
• The energy level is represented by the letter n.
Sublevels
• Number of sublevels present in each energy level is equal to the n.
• Sublevels are represented by the letter l.
• In order of increasing energy:
s<p<d<f
Orbitals
• Represented by ml
• S Sublevel- Only 1 orbital in this sublevel level.
• P Sublevel- 3 orbitals present in this sublevel.
o Each orbital can only have 2 electrons.
• D Sublevel- 5 orbitals present in this sublevel.
• F Sublevel- 7 orbitals present in this sublevel.
11
1
Total # of
Orbitals in
Energy Level
1
Total # of
Electrons in
Energy Level
2
s, p
1, 3
4
8
3
s, p, d
1, 3, 5
9
18
4
s, p, d, f
1, 3, 5, 7
16
32
Energy Level
Sublevels
Present
# of Orbitals
1
s
2
Orbital Diagrams
• An orbital diagram shows the arrangement of electrons in an atom.
• The electrons are arranged in energy levels, then sublevels, then orbitals.
Each orbital can only contain 2 electrons.
• Three rules must be followed when making an orbital diagram.
o Aufbau Principle- An electron will occupy the lowest_ energy orbital
that can receive it.
To determine which orbital will have the lowest energy, look to
the periodic table.
o Hund’s Rule- Orbitals of equal energy must each contain one
electron before electrons begin pairing.
o Pauli Exclusion Principle- If two electrons are to occupy the same
orbital, they must be spinning in opposite directions.
• Energy Levels (n) determined by the ROWS
• Sub Levels (s,p,d,f)- determined by the sections
• Orbitals - determined by the # of columns per sublevel
12
13
Orbital Diagrams
1s22s22p4
1s22s22p63s23p64s23d104p65s1
1s22s22p63s23p1
• S
1s22s22p63s23p4
• As
1s22s22p63s23p64s23d104p3
• Mn
1s22s22p63s23p64s23d5
• N
• Sc
1s22s22p3
1s22s22p63s23p64s23d1
14
Name:_________________ Date:__________ Period:______ Honor Code:__________
Electron Configuration WS
Give the COMPLETE electron configuration for the following elements:
2
2
6
2
6
1. Ar = 1s 2s 2p 3s 3p
2
2
6
2
3
2. P = 1s 2s 2p 3s 3p
2
2
6
2
6
2
6
3. Fe 1s 2s 2p 3s 3p 4s 3d
2
2
6
2
6
4. Ca = 1s 2s 2p 3s 3p 4s
2
2
2
6
2
6
2
10
2
2
6
2
6
2
5
5
5. Br = 1s 2s 2p 3s 3p 4s 3d 4p
6. Mn = 1s 2s 2p 3s 3p 4s 3d
7. U = 1s22s22p63s23p64s23d104p6 5s24d105p66s24f145d106p67s25f36d1
15
Electron Configurations and Oxidation States
• Electron configurations are shorthand for orbital diagrams. The electrons are
not shown in specific orbitals nor are they shown with their specific spins.
• Draw the orbital diagram of oxygen:
• The electron configuration should be:
1s22s22p4
• Manganese (25)
1s22s22p63s23p64s23d5
• Arsenic (33)
1s22s22p63s23p64s23d104p3
• Promethium (61)
1s22s22p63s23p64s23d104p65s24d105p66s24f4
• The Noble Gas shortcut can be used to represent the electron configuration
for atoms with many electrons. Noble gases have a full s and p and therefore
can be used to represent the inner shell electrons of larger atoms.
• For example: Write the electron configuration for Lead.
• Write the electron configuration for Xenon.
• Substitution can be used:
• Manganese (25)
Mn = [Ar] 4s23d5
• Arsenic (33)
As = [Ar] 4s23d104p3
• Promethium (61)
Pm = [Xe] 6s24f45d1
16
• Valence electrons, or outer shell electrons, can be designated by the s and p
sublevels in the highest energy levels
• Write the noble gas shortcut for Bromine
Br = [Ar]4s23d104p5
• Write only the s and p to represent the valence level.
Br = 4s24p5
• This is the Valence Configuration. Bromine has 7 valence electrons.
• Silicon
[Ne] 3s23p2
3s23p2
4 valence electrons
• Uranium
[Rn] 7s25f46d1
7s2
2 valence electrons
• Lead
[Xe] 6s24f145d106p2
6s26p2
4 valence electrons
Octet Rule and Oxidation States
• The octet rule states the electrons need __eight___ valence electrons in order to
achieve maximum stability. In order to do this, elements will gain, lose or share
electrons.
• Write the Valence configuration for oxygen
O = 2s22p4- 6 valence electrons
• Oxygen will gain 2 electrons to achieve maximum stability
-2
O = 2s22p6- 8 valence electrons
o Now, oxygen has 2 more electrons than protons and the resulting charge of
the atom will be -2
o The symbol of the ___ion____ formed is now O-2.
• Elements want to be like the Noble Gas family, so they will gain or lose electrons
to get the same configuration as a noble gas.
• When an element gains or losses an electron, it is called an __ion___.
• An ion with a positive charge is a ____cation (lost electrons)_____.
• An ion with a negative charge is an ___anion (gained electrons)___.
17
-
( 2)
18
Electron Configuration and Oxidation States Worksheet
Give the noble gas shortcut configuration for the following elements:
1. Pb
2. Eu
Eu = [Xe] 6s24f 6 5d1
3. Sn
Sn = [Kr] 5s24d105p2
4. As
As = [Ar] 4s23d104p3
Give ONLY the outer shell configuration for the following elements:
1. Ba
6s2
2. Po
6s26p4
3. S
3s23p4
4. F
2s22p5
Au
6s2
Cm
7s2
19
Quantum Number Notes
•
•
•
•
•
•
The quantum mechanical model uses three quantum numbers, n, l and ml to describe an
orbital in an atom. A fourth quantum number, ms, describes an individual electron in an
orbital.
n = principle quantum number (energy level quantum number)
o This describes the energy level and can be described as an integer from 1-7. The
larger the number, the larger the orbital. As the numbers increase, the electron will
have greater energy and will be less tightly bound by the nucleus.
l = azimuthal quantum number (sublevel quantum number)
o This describes the shape of the orbital level and can be described as an integer from
0 to n-1.
o 0 is used to describe s orbitals; 1 is used to describe p orbitals.
o 2 is used to describe d orbitals; 3 is used to describe f orbitals.
ml= magnetic quantum number (orbital quantum number)
o This describes the orientation of the orbital in space and can be described as an
integer from -l to l
o s sublevels have one orbital, therefore possible values of ml include 0 only.
o p sublevels have three orbitals, therefore possible values of ml include -1, 0, 1.
o d sublevels have five orbitals, therefore possible values of ml include -2, -1, 0, 1, 2.
o What about f?
ms= electron spin quantum number
o This describes the spin of the electron in the orbital.
o The possible values for ms are +½ and ½.
o The positive spin reflects the first electron in a specific orbital and negative spin
reflects the second electron in an orbital.
Examples:
o Give the four quantum numbers for the 8th electron in Argon:
n = 2, l = 1, ml = -1, ms = -1/2
o What is the maximum number of electrons that can have the following quantum
numbers: n = 2 and ms = -½
4 (2 electrons in the s sublevel, 4 in the p sublevel, that’s 8 electrons
total, half have +1/2 spin and half have -1/2 spin)
o Which of the following quantum numbers would NOT be allowed in an atom
n = 2, l = 2 and ml = -1
not allowed (when n=2, l cannot = 2)
n = 4, l = 2 and ml = -1
allowed
n = 3, l = 1 and ml = 0
allowed
n = 5, l = 0 and ml = 1
not allowed (when l = 0, ml can only = 0)
20
Quantum Number Practice:
Write the four quantum numbers which describe the location of the highest energy electron of the
following:
1. N #7 _____________ n = 2, l = 1, ml = 1, ms = +1/2 __
2. Ni #28_____________ n = 4, l = 2, ml = 0, ms = -1/2 __
3. Xe #54_____________ n = 5, l = 1, ml = 1, ms = -1/2 __
4. Re #75_____________ n = 6, l = 2, ml = 2, ms = +1/2 __
5. Pu #94_____________ n = 7, l = 3, ml = 1, ms = +1/2 __
6. Br #35______________ n = 4, l = 1, ml = 0, ms = -1/2 __
Give the four quantum numbers which describe the location of each of the following:
7. The 4th electron in carbon_____________ n = 2, l = 0, ml = 0, ms = -1/2 __
8. The 25th electron in Hf_____________ n = 4, l = 2, ml = 2, ms = +1/2 __
9. The 57th electron in Ho_____________ n = 6, l = 2, ml = -2, ms = +1/2 __
10. The 49th electron in Xe______________ n = 5, l = 1, ml = -1, ms = +1/2 __
Identify the element whose highest energy electron would have the following four quantum numbers:
11. 3, 1, -1, +1/2__________Al_________________________________
12. 4, 2, +1, +1/2_________Cr_______________________________________
13. 6, 1, 0, -1/2_________At_________________________________
14. 4, 3, +3, -1/2_____Not Allowed___________________________________
15. 2, 1, +1, -1/2________Ne______________________________________
Which of the following represents a permissible set of quantum numbers? (answer “yes” if permissible
and “no” if no permissible)
16. 2, 2, +1, -1/2_________NO____________________________________
17. 5, 1, 0, +1/2__________Yes______________________________________
18. 6, 3,-2, +1/2________Yes_______________________________________
19. 7, 0, 0, -1/2_________Yes_______________________________________
20. 4, 1, 8, +1/2________NO________________________________________
21
Quantum Numbers Worksheet
Rules for assigning quantum numbers:
n: can be 1, 2, 3, 4, …
Any positive whole number
l: can be 0, 1, 2, … (n-1)
Any positive whole number, up to (n-1)
ml : can be (-l), (-l +1), (-l +2), … 0, 1, … (l -1), l
Any integer, from – l to l.
ms: can be +½ or - ½
1) If n=2, what possible values does l have?
They would be 1 or 0.
2) When l is 3, how many possible values of ml are there?
There are 7 possibilities.
3) What are the quantum numbers for the 17th electron of Argon?
n=3, l= 1, ml=0, ms=-1/2
4) What are the quantum numbers for the 20th electron of Chromium?
n=4, l= 0, ml=0, ms=-1/2
5) What are the quantum numbers for the 47th electron of Iodine?
n=5, l= 2, ml=1, ms=-1/2
6) Give the quantum numbers for ALL of the electrons in Nitrogen.
e1 = n=1, l= 0, ml=0, ms=+1/2
e2 = n=1, l= 0, ml=0, ms=-1/2
e3 = n=2, l= 0, ml=0, ms=+1/2
e4 = n=2, l= 0, ml=0, ms=-1/2
e5 = n=2, l= 1, ml=-1, ms=+1/2
e6 = n=2, l= 1, ml=0, ms=+1/2
e7 = n=2, l= 1, ml=+1, ms=+1/2
7) Determine what the error is in the following quantum numbers, and explain. (Hint: 1 does not
have an error)
a. (1, 1, 0, - ½ ) l should be at least one less then n.
b. (3, 2, -1,+ ½ ) If n is 3, it can not have l as 2 because there is no d orbital in n=3.
c. (2, -1, 1, - ½ ) l can not be -1
d. (3, 3, 2 , + ½ )l can not be 3. It must be 1 less then n.
e. (2, 0, 1, - ½ ) If l is 0, ml must be 0.
f. (-1, 0, 0, + ½ ) n can not be -1.
g. (2, 1, 0, + ½ ) It is valid
22
Periodic Trends- Review Notes
• Shielding: As you go down the periodic table, the number of shells increases
which results in greater electron-electron repulsion.
o The more shells there are, the further from the nucleus the valence
electrons are.
o Therefore, more shielding means the electrons are _Less_ attracted
to the nucleus of the atom.
• Atomic Radius is defined as half the distance between adjacent nuclei of
the same element.
o As you move DOWN a group an entire energy level is added with
each new row, therefore the atomic radius __increases_(larger)_.
o As you move LEFT-TO-RIGHT across a period, a proton is added, so
the nucleus more strongly attracts the electrons of a atom, and
atomic radius __decreases (smaller)__.
• Ionic Radius is defined as half the distance between adjacent nuclei of the
same ion.
o For __cation____ an electron was lost and therefore the ionic radius
is smaller than the atomic radius.
o For __anion_____ an electron was gained and therefore the ionic
radius is larger than the atomic radius.
o As you move down a group an entire energy level is added, therefore
the ionic radius increases.
o As you move left-to-right across a period, a proton is added, so the
nucleus more strongly attracts the electrons of a atom, and ionic
radius ____decreases____.
23
However! This occurs in 2 sections. The cations form the first group,
and the anions form the second group.
• Isoelectronic Ions: Ions of different elements that contain the same number of electrons.
• Ionization energy is defined as the energy required to __remove__ the first electron from
an atom.
o As you move down a group atomic size increases, allowing electrons to be further
from the nucleus, therefore the ionization energy ___decreases_____.
o As you move left-to-right across a period, the nuclear charge increases, making it
harder to remove an electron, thus the ionization energy ______increases_____.
• Electronegativity is defined as the relative ability of an atom to attract electrons in a
____electron cloud by the nucleus________________.
o As you move down a group atomic size increases, causing available electrons to be
further from the nucleus, therefore the electronegativity ______decreases_____.
o As you move left-to-right across a period, the nuclear charge increases, making it
easier to gain an electron, thus the electronegativity __________increases_______.
• Reactivity is defined as the ability for an atom to react/combine with other atoms.
o With reactivity we must look at the metals and non-metals as two separate groups.
• Metal Reactivity- metals want to lose electrons and become cations
o As you move down a group atomic size increases, causing valence electrons to be
further from the nucleus, therefore these electron are more easily lost and reactivity
___decreases_______.
o As you move left-to-right across a period, the nuclear charge increases, making it
harder to lose electrons, thus the reactivity __increases__.
• Non-metal Reactivity- non-metals want to gain electrons and become anions
o As you move down a group atomic size increases, making it more difficult to attract
electrons, therefore reactivity ____decrease_____.
o As you move left-to-right across a period, the nuclear charge increases, making it
easier to attract electrons, thus the reactivity __increases___.
24
Periodic Table : What is the Trend?
Definition
Atomic Size
(Atomic
Radius)
Trend
Radius is defined as
half the distance
between adjacent
nuclei of the same
element.
Ability of an atom to
See above
Electronegativity attract electrons
Ionization
Energy
Energy required to
remove an e- from an
atom
See above
Metal
Having the
characteristics of a metal
Non-Metal
Having the
characteristics of a nonmetal
Shielding
This describes the
decrease in attraction
between an electron and
the nucleus in any atom
with more than one
electron shell. As more
electrons are between the
valence electrons and the
nucleus the more shielded
the outer electrons are
from the nucleus.
25
Periodic Trends Worksheet
1. Explain why a magnesium atom is smaller than both sodium AND calcium.
It is smaller than Na because it has more protons and smaller than Ca
because is has less energy levels.
2. Would you expect a Cl- ion to be larger or smaller than a Mg2+ ion? Explain
You would expect Cl- to be larger because of the electron to proton
ratio and Mg+2 now has the second energy level as its outer level.
3. Explain why the sulfide ions (S2-) is larger than a chloride ion (Cl-).
It is larger because of the electron to proton ratio. S-2 has two more
electrons than protons and Cl- only has one more.
4. Compare the ionization energy of sodium to that of potassium and
EXPLAIN.
It would require less ionization energy for K to loss an electron than
Na. K has more energy levels and the valence electrons are further
from the nucleus.
5. Explain the difference in ionization energy between lithium and beryllium.
They are the same energy level, but Be is slightly smaller so the
valence electron are closer to the nucleus so it would have a higher
ionization energy.
6. Order the following ions from largest to smallest: Ca2+, S2-, K+, Cl-. Explain
your order.
S2-, Cl-, K+, Ca2+ It is because of the electron to proton ratio.
7. Rank the following atoms/ions in each group in order of decreasing radii
and explain your ranking for each (larger to smaller).
a. I, I-
I- , I
b. K, K+
K, K+
c. Al, Al3+ Al, Al3+
8. Which element would have the greatest electron affinity: B or O? Explain.
Hint: a positive electron affinity means that the element wants to form a
negative charge.
It would be O. Because O wants to gain two electrons to achieve
noble gas configuration. While B wants to lose three electrons.
26
Part 2 - Review:
Give the Orbital Diagram for the following elements:
1. Chromium
2. Nitrogen
Give the COMPLETE electron configuration for the following elements:
2
2
6
2
6
3. Argon 1s 2s 2p 3s 3p
2
2
6
2
3
4. Phosphorous 1s 2s 2p 3s 3p
Give the Noble Gas electron configuration for the following elements:
2
5
1
5. Plutonium Pu = [Rn] 7s 5f 6d
Hg = [Xe] 6s24f145d10
6. Mercury
7. Complete the table.
Total # of
electrons
Valence
Configuration
Gain
or
Lose e
How
Many?
Ion
Symbol
New Valence
Configuration
Phosphorous
15
3s23p3
G
3
P-3
3s23p6
18
Chlorine
17
3s23p5
G
1
Cl-1
3s23p6
18
Cesium
55
6s1
L
1
Cs+1
5s25p6
54
Lithium
3
2s1
L
1
Li+1
1s2
2
Element
Give the 4 quantum numbers for the last electron of the following elements:
8. Phosphorous n=3, l=1, ml=1, ms= +1/2
9. Manganese n= 4, l=2, ml=2, ms= +1/2
10. Silver
n= 5, l=2, ml=1, ms= -1/2
11. Promethium n= 6, l=3, ml=0, ms= +1/2
12. Iodine n= 5, l=1, ml=0, ms= -1/2
27
Total # of
e
Determine if the following sets of quantum numbers would be allowed in an atom. If
not, explain why and if so, identify the corresponding atom.
1
13. n = 2, l = 1, ml = 0, ms = + Yes
2
14. n = 4, l = 0, ml = 2, ms = -
1
2
15. n = 1, l = 1, ml = 0, ms = +
1
2
No, because orbital 0 only has sub level 0
No, because l must be one less than n
Give the element with the LARGER radius, ionization energy, electronegativity and
reactivity.
IONIZATION
ELEMENTS
ATOMIC RADIUS
ELECTRONEGATIVITY
ENERGY
Sodium and
Na
Na
Al
Aluminum
Chlorine and
I
I
Cl
Iodine
Oxygen and
O
O
F
Fluorine
Magnesium
Ca
Ca
Mg
and Calcium
Circle the element / ion with the larger radius.
18. S or S216. Mg or Mg 2+
17. Sr2+ or Br-
19. Cl- or Mg2+
20. N3- or F21. B or F
For each of the following families, give their relative reactivity, the number of valence
electrons, and at least one additional piece of information (such as how they are found
in nature or what other group the generally react with).
22. Alkaline Earth Metals Very reactive, s2, most in the earth’s crust
23. Alkali Metals Very reactive, s1, they will react in air and with water
24. Halogens Very reactive, s2p5, they form salts
25. Noble Gases non reactive, s2p6, they are gases at room temperature
28
Matching (1 point each): Match the description in Column B with the correct term in
Column A. Write the letter in the blank provided. Each term matches with only one
description, so be sure to choose the best description for each term. Not all
descriptions will be used.
Column A
Column B
__A__ 26. Alkaline Earth Metal
A. located in the second column
__D__ 27. Transition Metal
B. solid or liquid mixture of two or more metals
__F__ 28. Alkali Metal
C. horizontal row of elements
__ I
D. located in columns 3-12
_ 29. Noble Gases
__K__ 30. Halogen
E. energy required to remove an e- from an atom
__C__ 31. Period
F. located in the first column
__E__ 32. Ionization Energy
G. ability of an atom to attract electrons
__H__ 33. Valence Electron
H. an electron in the outermost shell of an atom
__G__ 34. Electronegativity
I. located in column 18
__J__ 35. Group
J. vertical column of elements
K. located in column 17
_D_ 36. Elements in a family or group in the periodic table often share similar
properties because
a. They look alike.
b. They are found in the same place on Earth.
c. They have the same physical state.
d. Their atoms have the same number of electrons in their outer energy level.
_B__ 37. Groups 3-12 are commonly referred to as
a.
Alkali metals.
b.
Transition metals.
c.
Lanthanides.
d.
Actinides.
__C_ 38. Which of the following elements has the highest electronegativity?
a.
Ca
b. Cu
c. Br
d. As
_B_ 39. An atom is neutral because the number of
a.
Electrons equals the number of neutrons.
b.
Electrons equals the number of protons.
c.
Protons equals the number of neutrons.
d.
None of the above.
29