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. 1
 When elements are arranged in order of increasing _Atomic Number_,
there is a periodic repetition of their physical and chemical properties
 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____ (transferring or sharing)
 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 Melting Point and Boiling
Point__
 Metalloids: also called _semi-metals__, located _along_ the “stair-step”
 Properties - __ similar __ to both metals and nonmetals




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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.
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Valence electrons
Valence electrons the electrons that are in the highest (outermost) energy level
that level is also called the valence shell of the atom they are held most loosely
The number of valence electrons in an atom determines: The properties of the
atom The way that atom will bond chemically As a rule, the fewer electrons in the
valence shell, the more reactive the element is. When an atom has eight
electrons in the valence shell, it is stable.
Our discussion of valence electron configurations leads us to one of the cardinal
tenets of chemical bonding, the octet rule.
The octet rule states that atoms become especially stable when their valence
shells gain a full complement of valence electrons. For example, Helium (He)
and Neon (Ne) have eight outer valence electrons in their outer shells which
means it is completely filled, so they have a tendency to neither gain or lose
electrons.
Therefore, Helium and Neon, two of the so-called Noble gases or Inert gases
Group #
1
Group Name
Alkali Metals
# of valence electrons
1
2
Alkaline Earth Metals
2
3-12
Transition Metals
1 or 2
13
Boron Group
3
14
Carbon Group
4
15
Nitrogen Group
5
16
Oxygen Group
6
17
Halogens
7
18
Noble Gases
8
The number of valence electrons increases as you go across the periodic table from left to
right.
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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
ML
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 _ an electron
- __Increases__ from left to right across a period
- __Decreases____ from top to bottom down a group
Why?
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- 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, __Energy_, _Decreases_ to remove outermost ebecause 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
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- __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. 6
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. 7
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.  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 o
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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 
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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 N2 O2 F2 Cl2 Br2 I2 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) 9
Groups on the Periodic Table Summary Sheet:
Group
Examples
of Words
used
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
Example:
Number of Valance
Electrons
Element’s Name
Alkali
Metals
Alkaline
Earth
Metals
Transition
Metals
(Outer)
Inner
Transition
Metals
Halogens
Noble
Gases
M
+1
Y
N
Any Name in
Family 1
M
+2
Y
N
Any Name in
Family 2
1
2
3-12
3 (atomic #
58-71, 90103)
1
2
M
M
+2
+2
N
N
N
Any Name in
Family 3-12
N
Any Name
atomic
number 58-71,
90-103
7
8
17
NM
-1
Y
Y
Any Name in
Family 17
18
NM
0
N
NA
Any Name in
Family 18
2
2
Hydrogen
1
M
+1
Y
10
NA
Hydrogen
1
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
114.82
Germanium
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
132.905
Strontium
137.33
Yttrium
138.906
Zirconium Niobium MolybdenumTechnetium Ruthenium Rhodium Palladium
178.49
180.948
183.85
186.207
190.2
192.22
195.08
Silver
196.967
Cadmium
200.59
Tin
Indium
204.383
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
140.908
144.24
Gold
(272?)
Mercury
(277?)
Uuu Uub
111
Thallium
(?)
(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 Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium
232.038
Actinoid Series
(289?)
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
Thorium Protactinium Uranium Neptunium Plutonium Americium
11
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.
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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 13
There are two ways of representing the electron distribution among the various orbitals of an atom:
1. Electron configuration
An electron configuration consists of the symbol for the occupied subshell with a superscript indicating the
number of electrons in the subshell.
The electron configuration for sodium (atomic number 11) is
1s22s22p63s1
 The large numbers represent the energy level.
 The letters represent the sublevel.
 The superscript numbers indicate the number of electrons in the sublevel.
2. Orbital diagram
 An orbital diagram consists of a box representing each orbital and a half arrow representing each electron.
 The orbital diagram below is for sodium (atomic number 11)
Condensed Configurations
For large atoms, showing all the electrons with an electron configuration or orbital diagram can become quite
complex. Since it is the outermost electrons that are largely responsible for chemical behavior, we
can condense the electron configuration and orbital diagram to focus on those electrons.
Outer-shell electrons, those involved in chemical bonding, are called valence electrons. Those electrons below
the outer shell, inner-shell electrons, are usually referred to as core electrons.
The electron configuration and orbital diagram can be condensed by beginning with the nearest (before the atom)
noble gas symbol in brackets to represent the core electrons, then showing the valence electrons as usual.
Sodium's complete electron configuration is
1s22s22p63s1
The same electron configuration in condensed form becomes
[Ne]3s1
The complete orbital diagram for sodium is
The same orbital diagram in condensed form becomes
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Orbital Diagrams 1s22s22p4
1s22s22p63s23p64s23d104p65s1
1s22s22p63s23p1
 S 1s22s22p63s23p4
 As 1s22s22p63s23p64s23d104p3
 Mn 1s22s22p63s23p64s23d5  N 1s22s22p3
 Sc 1s22s22p63s23p64s23d1
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Name:_________________ Date:__________ Period:______ Honor Code:__________
Electron Configuration WS
Give the COMPLETE electron configuration for the following elements:
1. Ar = 1s22s22p63s23p6
2. P = 1s22s22p63s23p3
3. Fe 1s22s22p63s23p64s23d6
4. Ca = 1s22s22p63s23p64s2
5. Br = 1s22s22p63s23p64s23d104p5
6. Mn = 1s22s22p63s23p64s23d5
7. U = 1s22s22p63s23p64s23d104p6 5s24d105p66s24f145d106p67s25f36d1
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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) 1s22s22p63s23p64s23d104p65s24d105p66s24f45d1  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 17
 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 3s23p2
4 valence electrons
[Ne] 3s23p2
 Uranium 7s2
2 valence electrons
[Rn] 7s25f46d1
 Lead 6s26p2
4 valence electrons [Xe] 6s24f145d106p2
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 O‐2 = 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)___. 18
(-2)
19
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 20
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____.
21
 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___. 22
Periodic Table : What is the Trend? Atomic Size (Atomic Radius) Definition Trend Radius is defined as
half the distance
between adjacent
nuclei of the same
element. Ability of an atom to
Electronegativity attract electrons
Ionization Energy See above 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.
23
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.
24
Unit 3 Test Review:
Give the Orbital Diagram for the following elements:
1. Chromium
2. Nitrogen
Give the COMPLETE electron configuration for the following elements:
3. Argon 1s22s22p63s23p6
4. Phosphorous 1s22s22p63s23p3
Give the Noble Gas electron configuration for the following elements:
5. Plutonium Pu = [Rn] 7s2 5f5 6d1
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= 3, l=2, ml=2, ms= +1/2
10. Silver
n= 4, l=2, ml=1, ms= -1/2
11. Promethium n= 4, l=3, ml=0, ms= +1/2
12. Iodine n= 5, l=1, ml=0, ms= -1/2
25
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.
13. n = 2, l = 1, ml = 0, ms = +
14. n = 4, l = 0, ml = 2, ms = -
1
Yes
2
1
No, because orbital 0 only has sub level
2
0
1
2
15. n = 1, l = 1, ml = 0, ms = + 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
Al
Al
Aluminum
Chlorine and
I
Cl
Cl
Iodine
Oxygen and
O
F
F
Fluorine
Magnesium
Ca
Mg
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
26
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.
27