Unit 3 Notes: Periodic Table Notes - Loudoun County Public ... · Unit 3 Notes: Periodic Table...

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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

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• 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.

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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 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?

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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, 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

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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] 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

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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

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


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


7

Noble Gases

18

NM 0 N NA


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 llll.

• 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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Orbital Diagrams

• S

1s22s

22p

63s

23p

4

• As

1s22s

22p

63s

23p

64s

23d

104p

3

• Mn

1s22s

22p

63s

23p

64s

23d

5

• N

1s22s

22p

3

• Sc

1s22s

22p

63s

23p

64s

23d

1

1s22s

22p

63s

23p

64s

23d

104p

65s

1

1s22s

22p

4

1s22s

22p

63s

23p

1

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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)

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

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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]4s23d

104p

5

• Write only the s and p to represent the valence level.

Br = 4s24p

5

• 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 = 2s22p

4- 6 valence electrons

• Oxygen will gain 2 electrons to achieve maximum stability

O-2

= 2s22p

6- 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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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)

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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________________________________________

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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


electron of Chromium?

n=4, l= 0, ml=0, ms=-1/2


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

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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_____.


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_______.


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_____.


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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Part 2 - 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= 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

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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 Na Al

Chlorine and Iodine

I I Cl

Oxygen and Fluorine

O O F

Magnesium and Calcium

Ca Ca 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.

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