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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.
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.
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?
- 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
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.
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] 2s 1 , Na = [Ne] 3s 1 – 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
10
Groups on the Periodic Table Summary Sheet:
Group
Location onPeriodic Table
Metals, Non-Metals,
Metalloids?
CommonCharge(s)?
Reactivity
InterestingInformation
Example:
Number of Valance
Electrons
Examplesof Words
used
Group 1, Group 3-12,
etc
Metal
+
Highlyreactive, unreactive
It can be cutwith a plastic
knife
Element’s Name
AlkaliMetals
M
+
Y
N
Any Name in
Family 1
AlkalineEarthMetals
M
+
Y
N
Any Name in
Family 2
Transition
Metals(Outer)
M
+
N
N
Any Name inFamily 3-
Inner Transition
Metals
3 (atomic #58-71, 90-
M
+
N
N
Any Name
atomic
number 58-71,
Halogens
NM
Y
Y
Any Name in
Family 17
NobleGases
NM
N
NA
Any Name in
Family 18
Hydrogen
M
+^1
Y
NA
Hydrogen
11
P
eriodic Table of the Elements
H^1
Hydrogen
Si^14 Silicon
He^2 Helium
Li^3 Lithium
Be^4
Beryllium
B^5 Boron
C^6
Carbon
N^7
Nitrogen
O^8
Oxygen
F^9
Fluorine
Ne^10 Neon
Na^11 Sodium
Mg^12
Magnesium
Al^13
Aluminum
Si^14 Silicon
P 15
Phosphorus
S 16 Sulfur
Cl^17
Chlorine
Ar^18 Argon
K^19
Potassium
Ca^20 Calcium
Sc^21
Scandium
Ti^22
Titanium
V^23
Vanadium
Cr^24
Chromium
Mn^25
Manganese
Fe^26 Iron
Co^27 Cobalt
Ni^28 Nickel
Cu^29 Copper
Zn^30 Zinc
Ga^31 Gallium
Ge^32
Germanium
As^33 Arsenic
Se^34
Selenium
Br^35 Bromine
Kr^36 Krypton
Rb^37
Rubidium
Sr^38
Strontium
Y 39
Yttrium
Zr^40
Zirconium
Nb^41 Niobium
Mo^42
Molybdenum
(98)
Tc^43
Technetium
Ru^44
Ruthenium
Rh^45 Rhodium
Pd^46
Palladium
Ag^47 Silver
Cd^48 Cadmium
In^49 Indium
Sn^50 Tin
Sb^51
Antimony
Te^52
Tellurium
I 53
Iodine
Xe^54 Xenon
Cs^55 Cesium
Ba^56 Barium
La^57
Lanthanum
Hf^72 Hafnium
Ta^73
Tantalum
W^74
Tungsten
Re^75
Rhenium
Os^76 Osmium
Ir^77 Iridium
Pt^78
Platinum
Au^79 Gold
Hg^80 Mercury
Tl^81
Thallium
Pb^82 Lead
Bi^83
Bismuth
(209)
Po^84
Polonium
(210)
At^85 Astatine
(222)
Rn^86 Radon
(223)
Fr^87
Francium
Ra^88 Radium
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
Ce^58 Cerium
Pr^59
Praseodymium
Nd^60
Neodymium
(145)
Pm^61
Promethium
Sm^62
Samarium
Eu^63
Europium
Gd^64
Gadolinium
Tb^65 Terbium
Dy^66
Dysprosium
Ho^67 Holmium
Er^68 Erbium
Tm^69 Thulium
Yb^70
Ytterbium
Lu^71 Lutetium
Th^90 Thorium
Pa^91
Protactinium
U^92
Uranium
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
18
Name
Atomic Mass
Symbol
Atomic Number
PERIOD
Mass numbers in parenthesis are those of themost stable or most common isotope
Nonmetals
Metals
Groups
O Alkali MetalsO Alkali Earth MetalsO Boron GroupO Carbon GroupO HydrogenO Halogen sO Inner Transition MetalsO MetaloidsO Nitrogen GroupO Noble GassesO Oxygen GroupO Transition Metals
.^ .^ .^ .^ .^
.
.
.
.
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
There are two ways of representing the electron distribution among the various orbitals of an atom:
- 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 1 s^2 2 s^22 p^6 3 s^1 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 1 s^2 2 s^22 p^6 3 s^1 The same electron configuration in condensed form becomes [Ne]3 s^1 The complete orbital diagram for sodium is
The same orbital diagram in condensed form becomes
Name:_________________ Date:__________ Period:______ Honor Code:__________
Electron Configuration WS
Give the COMPLETE electron configuration for the following elements:
1. Ar = 1s^2 2s^2 2p^6 3s^2 3p^6
2. P = 1s^2 2s^2 2p^6 3s^2 3p^3
3. Fe 1s^2 2s^2 2p^6 3s 2 3p^6 4s 2 3d^6
4. Ca = 1s^2 2s^2 2p^6 3s^2 3p^6 4s 2
5. Br = 1s^2 2s^2 2p^6 3s^2 3p^6 4s 2 3d^10 4p 5
6. Mn = 1s^2 2s^2 2p^6 3s^2 3p^6 4s 2 3d^5
7. U = 1s 2 2s 2 2p^6 3s 2 3p^6 4s 2 3d 10 4p 6 5s 2 4d^10 5p 6 6s 2 4f 14 5d^10 6p^6 7s 2 5f 3 6d^1
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:
1s 2 2s 2 2p^4
Manganese (25)
1s 2 2s 2 2p^6 3s 2 3p^6 4s 2 3d^5
Arsenic (33)
1s 2 2s 2 2p^6 3s 2 3p^6 4s 2 3d^10 4p^3
Promethium (61)
1s 2 2s 2 2p^6 3s 2 3p^6 4s 2 3d^10 4p^6 5s 2 4d^10 5p^6 6s 2 4f 4 5d^1
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] 4s 2 3d^5
Arsenic (33)
As = [Ar] 4s 2 3d^10 4p^3
Promethium (61)
Pm = [Xe] 6s 2 4f 4 5d^1
Electron Configuration and Oxidation States Worksheet
Give the noble gas shortcut configuration for the following elements:
1. Pb
2. Eu
Eu = [Xe] 6s 2 4f 6 5d^1
3. Sn
Sn = [Kr] 5s 2 4d^10 5p^2
4. As
As = [Ar] 4s 2 3d^10 4p^3
Give ONLY the outer shell configuration for the following elements:
1. Ba
6s 2
2. Po
6s 2 6p^4
3. S
3s 2 3p^4
4. F 2s 2 2p^5
Au 6s 2
Cm 7s 2
