unit 3 notes: periodic table notes - loudoun county public schools · 2016-11-26 · 1 unit 3...
TRANSCRIPT
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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.
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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
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 # Group Name # of valence electrons 1 Alkali Metals 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 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
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- __smaller__ than corresponding neutral atom
Why?
- __fewer__ e-, so it’s _easier_ for protons to pull in remaining e-
Anions:__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 e-
Shielding:
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.
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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.
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o 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
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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
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)
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Groups on the Periodic Table Summary Sheet:
Group Location on
Periodic Table Metals, Non-Metals,
Metalloids? Common
Charge(s)? Reactivity
Interesting Information
Example: Number of Valance
Electrons Examples of Words
used
Group 1, Group 3-12,
etc Metal +1
Highly reactive,
unreactive
It can be cut with a plastic
knife Element’s Name
Alkali Metals
1 M +1 Y N
Any Name in Family 1
1
Alkaline Earth Metals
2 M +2 Y N
Any Name in Family 2
2
Transition Metals (Outer)
3-12
M +2 N N
Any Name in Family 3-12
2
Inner Transition
Metals
3 (atomic # 58-71, 90-
103)
M +2 N N
Any Name atomic
number 58-71, 90-103
2
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
Hydrogen
1
M +1 Y NA Hydrogen 1
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Periodic Table of the Elements 1.00794
H 1
Hydrogen
28.0855
Si 14
Silicon
4.00260
He 2
Helium
6.941
Li 3
Lithium
9.01218
Be 4
Beryllium
10.81
B 5
Boron
12.01
C 6
Carbon
14.0067
N 7
Nitrogen
15.9994
O 8
Oxygen
18.998403
F 9
Fluorine
20.179
Ne 10
Neon
22.98977
Na 11
Sodium
24.305
Mg 12
Magnesium
26.98154
Al 13
Aluminum
28.0855
Si 14
Silicon
30.97376
P 15
Phosphorus
32.06
S 16
Sulfur
35.453
Cl 17
Chlorine
39.948
Ar 18
Argon
39.0983
K 19
Potassium
40.08
Ca 20
Calcium
44.9559
Sc 21
Scandium
47.88
Ti 22
Titanium
50.9415
V 23
Vanadium
51.996
Cr 24
Chromium
54.9380
Mn 25
Manganese
55.847
Fe 26
Iron
58.9332
Co 27
Cobalt
58.69
Ni 28
Nickel
63.546
Cu 29
Copper
65.39
Zn 30
Zinc
69.72
Ga 31
Gallium
72.59
Ge 32
Germanium
74.9216
As 33
Arsenic
78.96
Se 34
Selenium
79.904
Br 35
Bromine
83.80
Kr 36
Krypton
85.4678
Rb 37
Rubidium
87.62
Sr 38
Strontium
88.9059
Y 39
Yttrium
91.224
Zr 40
Zirconium
92.9064
Nb 41
Niobium
95.94
Mo 42
Molybdenum
(98)
Tc 43
Technetium
101.07
Ru 44
Ruthenium
102.906
Rh 45
Rhodium
106.42
Pd 46
Palladium
107.868
Ag 47
Silver
112.41
Cd 48
Cadmium
114.82
In 49
Indium
118.71
Sn 50
Tin
121.75
Sb 51
Antimony
127.60
Te 52
Tellurium
127.60
I 53
Iodine
131.29
Xe 54
Xenon
132.905
Cs 55
Cesium
137.33
Ba 56
Barium
138.906
La 57
Lanthanum
178.49
Hf 72
Hafnium
180.948
Ta 73
Tantalum
183.85
W 74
Tungsten
186.207
Re 75
Rhenium
190.2
Os 76
Osmium
192.22
Ir 77
Iridium
195.08
Pt 78
Platinum
196.967
Au 79
Gold
200.59
Hg 80
Mercury
204.383
Tl 81
Thallium
207.2
Pb 82
Lead
208.980
Bi 83
Bismuth
(209)
Po 84
Polonium
(210)
At 85
Astatine
(222)
Rn 86
Radon
(223)
Fr 87
Francium
226.025
Ra 88
Radium
227.028
Ac 89
Actinum
(261)
Rf 104
Rutherfordium
(262)
Db 105
Dubnium
(263)
Sg 106
Seaborgium
(262)
Bh 107
Bohrium
(265)
Hs 108
Hassium
(266)
Mt 109
Meitnerium
(269)
Ds 110
Dormstadtium
(272?)
Uuu 111
Unununium
(277?)
Uub 112
Ununbium
(?)
Uut 113
Ununtrium
(289?)
Uuq 114
Ununquadium
(289?)
Uuh 116
Ununhexium
(293?)
Uuo 118
Ununoctium
140.12
Ce 58
Cerium
140.908
Pr 59
Praseodymium
144.24
Nd 60
Neodymium
(145)
Pm 61
Promethium
150.36
Sm 62
Samarium
151.96
Eu 63
Europium
157.25
Gd 64
Gadolinium
158.925
Tb 65
Terbium
162.50
Dy 66
Dysprosium
164.930
Ho 67
Holmium
167.26
Er 68
Erbium
168.934
Tm 69
Thulium
173.04
Yb 70
Ytterbium
174.967
Lu 71
Lutetium
232.038
Th 90
Thorium
231.036
Pa 91
Protactinium
238.029
U 92
Uranium
237.048
Np
93 Neptunium
(244)
Pu
94 Plutonium
(243)
Am
95 Americium
(247)
Cm 96
Curium
(247)
Bk 97
Berkelium
(251)
Cf 98
Californium
(252)
Es 99
Einsteinium
(257)
Fm 100
Fermium
(258)
Md 101
Mendelevium
(259)
No 102
Nobelium
(260)
Lr 103
Lawrencium
Lanthanoid Series
Actinoid Series
Transition Elements
Group 1
2
3 4 5 6 7 8 9 10 11 12
13 14 15 16 17
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Name
Atomic Mass
Symbol Atomic Number
P E R I O D
Mass numbers in parenthesis are those of the most stable or most common isotope
Nonmetals
Metals
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
. . . . . .
.
.
.
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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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Energy Level Sublevels Present
# of Orbitals Total # of Orbitals in Energy Level
Total # of Electrons in Energy Level
1 s 1 1 2
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
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
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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
S
1s22s22p63s23p4
As
1s22s22p63s23p64s23d104p3
Mn
1s22s22p63s23p64s23d5
N 1s22s22p3
Sc
1s22s22p63s23p64s23d1
1s22s22p63s23p64s23d104p65s1
1s22s22p4
1s22s22p63s23p1
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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
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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
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)___.
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(-2)
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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
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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____.
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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___.
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Periodic Table : What is the Trend?
Definition Trend
Atomic Size (Atomic Radius)
Radius is defined as half the distance between adjacent nuclei of the same element.
Electronegativity Ability of an atom to attract electrons
See above
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 non-metal
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.
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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.
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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
6. Mercury Hg = [Xe] 6s24f145d10
7. Complete the table.
Element Total # of electrons
Valence Configuration
Gain or Lose e-
How Many?
Ion Symbol
New Valence Configuration
Total # of e-
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
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
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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 = +2
1 Yes
14. n = 4, l = 0, ml = 2, ms = -2
1 No, because orbital 0 only has sub level 0
15. n = 1, l = 1, ml = 0, ms = +2
1 No, because l must be one less than n
Give the element with the LARGER radius, ionization energy, electronegativity and reactivity.
ELEMENTS ATOMIC RADIUS IONIZATION ENERGY
ELECTRONEGATIVITY
Sodium and Aluminum
Na Al Al
Chlorine and Iodine
I Cl Cl
Oxygen and Fluorine
O F F
Magnesium and Calcium
Ca Mg Mg
Circle the element / ion with the larger radius.
16. Mg or Mg 2+ 18. S or S2- 20. N3- or F-
17. Sr2+ or Br- 19. Cl- or Mg2+ 21. 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
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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 _ 29. Noble Gases D. located in columns 3-12
__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.