Atomic Structure Unit 3

http://www.unit5.org/chemistry

Atomic Structure

Unit 3

Greek Model

• Greek philosopher• Idea of ‘democracy’• Idea of ‘atomos’

– Atomos = ‘indivisible’– ‘Atom’ is derived

• No experiments to support idea

• Continuous vs. discontinuous theory of matter

Democritus’s model of atom

No protons, electrons, or neutrons

Solid and INDESTRUCTABLE

Democritus

“To understand the very large,

we must understand the very small.”

Foundations of Atomic Theory

Law of Definite Proportions

The fact that a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound.

Law of Multiple Proportions

If two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first elements is always a ratio of small whole numbers.

Law of Conservation of Mass

Mass is neither destroyed nor created during ordinary chemical reactions.

45 g H2O? g H2O

Conservation of Mass

Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 204

Highvoltage

Before reaction

electrodes

glasschamber

5.0 g H2

80 g O2

300 g (mass of chamber)+385 g total

Highvoltage

After reaction

0 g H2

40 g O2

300 g (mass of chamber)+385 g total

Law of Definite ProportionsJoseph Louis Proust (1754 – 1826)

• Each compound has a specific ratio of elements

• It is a ratio by mass • Water is always 8 grams of oxygen for

every one gram of hydrogen

The Law of Multiple Proportions

• Dalton could not use his theory to determine the elemental compositions of chemical compounds because he had no reliable scale of atomic masses.

• Dalton’s data led to a general statement known as the law of multiple proportions.

• Law states that when two elements form a series of compounds, the ratios of the masses of the second element that are present per gram of the first element can almost always be expressed as the ratios of integers.

Crookes Tube

William Crookes

Mask holder

Cathode(-)

Anode(+)

Crookes tube(Cathode ray tube)

Mask holder

Cathode Rays

• Cathode ray = electron

• Electrons have a negative charge

Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, pages 117-118

Highvoltage

cathode

source ofhigh voltage

yellow-greenfluorescence

shadow

(A) The effect of an obstruction on cathode rays

(B) The effect of an electric field on cathode rays

Highvoltage

cathode

source ofhigh voltage

positiveplate

negative plate

source oflow voltage

J.J. Thomson• He proved that atoms of

any element can be made to emit tiny negative particles.

• From this he concluded that ALL atoms must contain these negative particles.

• He knew that atoms did not have a net negative charge and so there must be balancing the negative charge.

J.J. Thomson

William Thomson (Lord Kelvin)

• In 1910 proposed the Plum Pudding model– Negative electrons

were embedded into a positively charged spherical cloud.

Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 56

Spherical cloud ofPositive charge

Electrons

Rutherford’s Apparatus

beam of alpha particles

radioactive substance

gold foil

circular ZnS - coated

fluorescent screen

Rutherford received the 1908 Nobel Prize in Chemistry for his pioneering work in nuclear chemistry.

Geiger-Muller Counter

Speaker gives“click” for

each particle

Window

Particlepath

Argon atoms

Hans Geiger

Rutherford’sGold-Leaf Experiment

Conclusions:

Atom is mostly empty space

Nucleus has (+) charge

Electrons float around nucleus

Bohr’s Model

Nucleus

Electron

Energy Levels

Bohr Model of Atom

The Bohr model of the atom, like many ideas in the history of science, was at first prompted by and later partially disproved by experimentation.

http://en.wikipedia.org/wiki/Category:Chemistry

Increasing energyof orbits

A photon is emittedwith energy E = hf

An unsatisfactory model for the hydrogen atom

According to classical physics, lightshould be emitted as the electron circles the nucleus. A loss of energywould cause the electron to be drawncloser to the nucleus and eventuallyspiral into it.

Hill, Petrucci, General Chemistry An Integrated Approach 2nd Edition, page 294

Quantum Mechanical Model

Modern atomic theory describes the electronic structure of the atom as the probability of finding electrons within certain regions of space (orbitals).

Niels Bohr &Albert Einstein

Models of the Atom

Dalton’s model (1803)

Thomson’s plum-pudding model (1897)

Rutherford’s model (1909)

Bohr’s model (1913)

Charge-cloud model (present)

Greek model(400 B.C.)

"In science, a wrong theory can be valuable and better than no theory at all."- Sir William L. Bragg

Electrons (-) charge no mass located outside the nucleus

Protons (+) charge 1 amu located inside the nucleus

Neutrons no charge 1 amu located inside the nucleus

Particles in the Atom

Discovery of the Neutron

James Chadwick bombarded beryllium-9 with alpha particles, carbon-12 atoms were formed, and neutrons were emitted.

+Be94 C12

Dorin, Demmin, Gabel, Chemistry The Study of Matter 3rd Edition, page 764 *Walter Boethe

Subatomic particles

Electron

Proton

Neutron

Name Symbol ChargeRelative mass

Actual mass (g)

1/1840

9.11 x 10-28

1.67 x 10-24

Symbols

Contain the symbol of the element, the mass number and the atomic number

X Massnumber

Atomicnumber

# protons

# protons + # neutrons mass number

Symbols

• Find the – number of protons– number of neutrons– number of electrons– Atomic number– Mass number

Symbols

Find the – number of protons– number of neutrons– number of electrons– Atomic number– Mass number

Na2311

Sodium atom

Symbols

Find the – number of protons– number of neutrons– number of electrons– Atomic number– Mass number

Na2311

Sodium ion

3545358035Br

1822184018Ar

2020204020Ca

e–n0p+MassAtomic

Ca40.08

Ar39.948

Br79.904

4 n0 2e– 1e–

Li shorthand

Bohr - Rutherford diagrams• Putting all this together, we get B-R diagrams• To draw them you must know the # of protons, neutrons,

and electrons (2,8,8,2 filling order)• Draw protons (p+), (n0) in circle (i.e. “nucleus”)• Draw electrons around in shells

Draw Be, B, Al and shorthand diagrams for O, Na

11 p+12 n°

2e– 8e– 1e–

8 p+8 n°

2e– 6e–

4 p+5 n°

5 p+6 n°

13 p+14 n°

Mass Number

• mass # = protons + neutrons

• always a whole number

• NOT on the Periodic Table!

Nucleus

Electrons

Nucleus

Neutron

Proton

Carbon-12Neutrons 6Protons 6Electrons 6

Isotopes

• Atoms of the same element with different mass numbers.

Mass #

Atomic #

• Nuclear symbol:

• Hyphen notation: carbon-12Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

3 n02e– 1e– 3 p+

4 n02e– 1e–

6Li 7Li

+Nucleus

Electrons

Nucleus

Neutron

Proton

Lithium-6Neutrons 3Protons 3Electrons 3

Nucleus

Electrons

Nucleus

Neutron

Proton

Lithium-7Neutrons 4Protons 3Electrons 3

Average Atomic Mass

• weighted average of all isotopes• on the Periodic Table• round to 2 decimal places

Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Avg.AtomicMass

= (mass)(%) + (mass)(%)

Average Atomic Mass

• EX: Calculate the avg. atomic mass of oxygen if its abundance in nature is 99.76% 16O, 0.04% 17O, and 0.20% 18O.

Avg.AtomicMass

= (16)(99.76) + (17)(0.04) + (18)(0.20)

100= 16.00

Average Atomic Mass

• EX: Find chlorine’s average atomic mass if approximately 8 of every 10 atoms are chlorine-35 and 2 are chlorine-37.

Avg.AtomicMass

= (35)(8) + (37)(2)

10= 35.40 amu

034 35 36 37

Mass spectrum of chlorine. Elemental chlorine (Cl2) contains only two isotopes: 34.97 amu (75.53%) and 36.97 (24.47%)

AAM = (34.97 amu)(0.7553) + (36.97 amu)(0.2447)

AAM = (26.412841 amu) + (9.046559 amu)

AAM = 35.4594 amu

Cl35.4594

Mass Spectrometry

Photographic plate

196 199 201 204

198 200 202

Mass spectrum of mercury vapor

Hill, Petrucci, General Chemistry An Integrated Approach 1999, page 320

Stream of positive ions

Mass Spectrum for Mercury

196 197 198 199 200 201 202 203 204

Mass number

196 199 201 204

198 200 202

Mass spectrum of mercury vapor

The percent natural abundances for mercury isotopes are:

Hg-196 0.146% Hg-198 10.02% Hg-199 16.84% Hg-200 23.13% Hg-201 13.22% Hg-202 29.80% Hg-204 6.85%

(The photographic record has been converted to a scale of relative number of atoms)

The percent natural abundances for mercury isotopes are:

Hg-196 0.146% Hg-198 10.02% Hg-199 16.84% Hg-200 23.13% Hg-201 13.22% Hg-202 29.80% Hg-204 6.85%

(0.00146)(196) + (0.1002)(198) + (0.1684)(199) + (0.2313)(200) + (0.1322)(201) + (0.2980)(202) + (0.0685)(204) = x

0.28616 + 19.8396 + 33.5116 + 46.2600 + 26.5722 + 60.1960 + 13.974 = x

x = 200.63956 amu

Hg200.59

(% "A")(mass "A") + (% "B")(mass "B") + (% "C")(mass "C") + (% "D")(mass "D") + (% "E")(mass "E") + (% F)(mass F) + (% G)(mass G) = AAM

ABCDEFG

Natural uranium, atomic weight = 238.029 g/molDensity is 19 g/cm3. Melting point 1000oC.

Two main isotopes:

U238 92

U235 92

Because isotopes are chemically identical(same electronic structure), they cannot beseparated by chemistry.

So Physics separates them by diffusion orcentrifuge (mass spectrograph is too slow)…

Separation of Isotopes

(238 amu) x (0.993) + (235 amu) x (0.007)

236.334 amu + 1.645 amu

237.979 amu

• Assume you have only two atoms of chlorine.• One atom has a mass of 35 amu (Cl-35)• The other atom has a mass of 36 amu (Cl-36)

• What is the average mass of these two isotopes?

35.5 amu

• Looking at the average atomic mass printed on the periodic table...approximately what percentage is Cl-35 and Cl-36?

55% Cl-35 and 45% Cl-36 is a good approximation

Cl35.453

Using our estimated % abundance data

55% Cl-35 and 45% Cl-36

calculate an average atomic mass for chlorine.

Cl35.453

Average Atomic Mass = (% abundance of isotope "A")(mass "A") + (% "B")(mass "B")

AAM = (% abundance of isotope Cl-35)(mass Cl-35) + (% abundance of Cl-36)(mass Cl-36)

AAM = (0.55)(35 amu) + (0.45)(36 amu)

AAM = (19.25 amu) + (16.2 amu)

AAM = 35.45 amu

Isotopes

Dalton was wrong.

Atoms of the same element can have different numbers of neutrons

different mass numbers

called isotopes

California WEB

C-12 vs. C-14

Using a periodic table and what you know about atomic number, mass, isotopes, and electrons, fill in the chart:

Element Symbol AtomicNumber

AtomicMass

# of protons

# of neutron

# of electron

charge

Potassium 39 +1

Br 45 -1

30 35 30

Atomic Number = Number of Protons

Number of Protons + Number of Neutrons = Atomic Mass

Atom (no charge) : Protons = Electrons

Ion (cation) : Protons > Electrons Ion (anion) : Electrons > Protons

Periodic Table

• Dmitri Mendeleev developed the modern periodic table.

• Argued that element properties are periodic functions of their atomic weights.

• We now know that element properties are periodic functions of their ATOMIC NUMBERS.

Atomic Mass

Magnesium has three isotopes. 78.99% magnesium 24 with a mass of 23.9850 amu, 10.00% magnesium 25 with a mass of 24.9858 amu, and the rest magnesium 26 with a mass of 25.9826 amu. What is the atomic mass of magnesium?

If not told otherwise, the mass of the isotope is

the mass number in amu.

California WEB

IsotopePercent

AbundanceMass

Mg-24 78.99 23.9850

Mg-25 10.00 24.9585

Mg-26 25.9826

24.304 amu

18.94575

2.49585

2.8606811.01

Atomic Mass

Calculate the atomic mass of copper if copper has two isotopes. 69.1% has a mass of 62.93 amu and the rest has a mass of 64.93 amu.

... )B"" )(massB"" (% )A"" )(massA"" (% (AAM) massatomic Average

Copper for amu 63.548 A.A.M.

amu 20.06337 amu 43.48463 A.A.M.

amu) .93(0.309)(64 amu) .93(0.691)(62 A.A.M.

63.548

IsotopePercent

AbundanceMass

Cu-63 69.1 62.93

Cu-65 64.93

43.48463

20.0633730.9

63.548

Given the average atomic mass of an element is 118.21 amu and it has three isotopes (“A”, “B”, and “C”):

isotope “A” has a mass of 117.93 amu and is 87.14% abundantisotope “B” has a mass of 120.12 amu and is 12.36% abundant

Find the mass of isotope “C”. Show work for credit.

Extra Credit: What is a cation?

A positively charged atom. An atom that has lost a(n) electron(s).

119.7932 amu

Protons Neutrons ElectronsMass

number

Cu-65 A = 29 B = 36 29 C = 65

Argon D = 18 E = 22 F = 18 40

Ba2+ 56 G = 81 H = 54 I = 137

Quantum Mechanics

• Heisenberg Uncertainty Principle– Impossible to know both the velocity and

position of an electron at the same time

Microscope

Electron

Werner Heisenberg~1926

Quantum Mechanics

σ3/2 Zπ

• Schrödinger Wave Equation (1926)

– finite # of solutions quantized energy levels

– defines probability of finding an electron

Erwin Schrödinger~1926

Quantum Mechanics

• Orbital (“electron cloud”)– Region in space where there is 90%

probability of finding an electron

Electron Probability vs. Distance

Distance from the Nucleus (pm)

100 150 200 2505000

Orbital

90% probability offinding the electron

Quantum Numbers

UPPER LEVEL

• Four Quantum Numbers:– Specify the “address” of each electron

in an atom

Quantum Numbers

Principal Quantum Number ( n )

Angular Momentum Quantum # ( l )

Magnetic Quantum Number ( ml )

Spin Quantum Number ( ms )

Quantum Numbers

1. Principal Quantum Number ( n )

– Energy level

– Size of the orbital

– n2 = # of orbitals in the energy level

Shapes of s, p, and d-Orbitals

s orbital

p orbitals

d orbitals

Atomic Orbitals

s, p, and d-orbitals

As orbitals:

Hold 2 electrons(outer orbitals ofGroups 1 and 2)

Bp orbitals:

Each of 3 pairs oflobes holds 2 electrons

= 6 electrons(outer orbitals of Groups 13 to 18)

Cd orbitals:

Each of 5 sets oflobes holds 2 electrons

= 10 electrons(found in elements

with atomic no. of 21and higher)

Kelter, Carr, Scott, , Chemistry: A World of Choices 1999, page 82

Quantum Numbers

px pz py

Quantum Numbers

• n = # of sublevels per level• n2 = # of orbitals per level• Sublevel sets: 1 s, 3 p, 5 d, 7 f

n = 3n = 2n = 1Principallevel

Sublevel

Orbital

s s p s p d

px py pz dxy dxz dyz dz2 dx2- y2px py pz

Maximum Capacities of Subshells and Principal Shells

n 1 2 3 4 ...n

l 0 0 1 0 1 2 0 1 2 3

Subshelldesignation s s p s p d s p d f

Orbitals insubshell 1 1 3 1 3 5 1 3 5 7

Subshellcapacity 2 2 6 2 6 10 2 6 10 14

Principal shellcapacity 2 8 18 32 ...2n2

Hill, Petrucci, General Chemistry An Integrated Approach 1999, page 320

Quantum Numbers

3. Magnetic Quantum Number ( ml )– Orientation of orbital

– Specifies the exact orbital within each sublevel

The magnetic quantum number

Third quantum is ml, the magnetic quantum number

– Value of ml describes the orientation of the region in space occupied by the electrons with respect

to an applied magnetic field

– Allowed values of ml depend on the value of l

– ml can range from –l to l in integral steps ml = l, -l + l, . . . 0 . . ., l – 1, l

– Each wave function with an allowed combination of n, l, and ml values describes an atomic orbital, a particular spatial distribution for

an electron

– For a given set of quantum numbers, each principal shell contains a fixed number of subshells, and each subshell contains a fixed number of orbitals

Principal Energy Levels 1 and 2

Quantum Numbers

4. Spin Quantum Number ( ms )– Electron spin +½ or -½

– An orbital can hold 2 electrons that spin in opposite directions.

Electron Spin: The Fourth Quantum Number

• When an electrically charged object spins, it produces a magnetic moment parallel to the axis of rotation and behaves like a magnet.

• A magnetic moment is called electron spin.

• An electron has two possible orientations in an external magnetic field, which are described by a fourth quantum number ms.

• For any electron, ms can have only two possible values, designated + (up) and – (down), indicating that the two orientations are opposite and the subscript s is for spin.

• An electron behaves like a magnet that has one of two possible orientations, aligned either with the magnetic field or against it.

Quantum Numbers

1. Principal # 2. Ang. Mom. # 3. Magnetic # 4. Spin #

energy level

sublevel (s,p,d,f)

orbital

electron

• Pauli Exclusion Principle– No two electrons in an atom can have the

same 4 quantum numbers.

– Each electron has a unique “address”:

Wolfgang Pauli

Level n 1 2 3

Sublevel l Orbital ml

Spin ms

0 0 1 0 -1 0 1 0 -1 2 1 0 -1 -2

= +1/2

= -1/2

Allowed Sets of Quantum Numbers for Electrons in Atoms

H = 1s1

He = 1s2

Li = 1s2 2s1

Be = 1s2 2s2

C = 1s2 2s2 2p2

1s 2s 2px 2py 2pz

S = 1s2 2s2 2p6

3s2 3p41s 2s 2px 2py 2pz 3s 3px 3py 3pz

THIS SLIDE IS ANIMATEDIN FILLING ORDER 2.PPT

Fe = 1s1 2s22p63s23p64s23d6

1s 2s 2px 2py 2pz 3s 3px 3py 3pz

4s 3d 3d 3d 3d

Iron has ___ electrons.26

ArbitraryEnergy Scale

4s 4p 3d

5s 5p 4d

6s 6p 5d 4f

NUCLEUS

Orbital Filling

Element 1s 2s 2px 2py 2pz 3s Configuration

Electron ConfigurationsElectron

1s22s22p63s1

1s22s22p6

1s22s22p5

1s22s22p4

1s22s22p3

1s22s22p2

1s22s1

Filling Rules for Electron Orbitals

Aufbau Principle: Electrons are added one at a time to the lowest energy orbitals available until all the electrons of the atom have been accounted for.

Pauli Exclusion Principle: An orbital can hold a maximum of two electrons.To occupy the same orbital, two electrons must spin in opposite directions.

Hund’s Rule: Electrons occupy equal-energy orbitals so that a maximum number of unpaired electrons results.

*Aufbau is German for “building up”

Spin Quantum Number, ms

The electron behaves as if it were spinning about an axis through its center.This electron spin generates a magnetic field, the direction of which dependson the direction of the spin.

Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 208

Electron aligned with magnetic field,

ms = + ½

Electron aligned against magnetic field,

ms = - ½

Energy Level Diagram of a Many-Electron Atom

4s 4p 3d

5s 5p 4d

6s 6p 5d 4f

NUCLEUS

O’Connor, Davis, MacNab, McClellan, CHEMISTRY Experiments and Principles 1982, page 177

Maximum Number of Electrons In Each SublevelMaximum Number of Electrons In Each Sublevel

Maximum Number Sublevel Number of Orbitals of Electrons

d 5 10

f 7 14

LeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World , 1996, page 146

Quantum Numbers

n shell

l subshell

ml orbital

ms electron spin

1, 2, 3, 4, ...

0, 1, 2, ... n - 1

- l ... 0 ... +l

+1/2 and - 1/2

Order in which subshells are filled with electrons

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d … 2 2 6 2 6 2 10 6 2 10

1sn = 1

Sublevels

1sn = 1

Sublevels

1s22s22p63s23p64s23d104p65s24d10…

Filling Rules for Electron Orbitals

Aufbau Principle: Electrons are added one at a time to the lowest energy orbitals available until all the electrons of the atom have been accounted for.

Pauli Exclusion Principle: An orbital can hold a maximum of two electrons.To occupy the same orbital, two electrons must spin in opposite directions.

Hund’s Rule: Electrons occupy equal-energy orbitals so that a maximum number of unpaired electrons results.

*Aufbau is German for “building up”

Energy Level Diagram of a Many-Electron Atom

4s 4p 3d

5s 5p 4d

6s 6p 5d 4f

NUCLEUS

O’Connor, Davis, MacNab, McClellan, CHEMISTRY Experiments and Principles 1982, page 177

Energy Level Diagram

4s 4p 3d

5s 5p 4d

6s 6p 5d 4f

NUCLEUS

CLICK ON ELEMENT TO FILL IN CHARTS

Fe = 1s22s22p63s23p64s23d6

H He Li C N Al Ar F Fe La

Bohr Model

Electron Configuration

Energy Level Diagram

4s 4p 3d

5s 5p 4d

6s 6p 5d 4f

NUCLEUS

CLICK ON ELEMENT TO FILL IN CHARTS

La = 1s22s22p63s23p64s23d10

4s23d104p65s24d105p66s25d1

H He Li C N Al Ar F Fe La

Bohr Model

Lanthanum

Electron Configuration

neon's electron configuration (1s22s22p6)

Shorthand Configuration

[Ne] 3s1

third energy level

one electron in the s orbital

orbital shape

Na = [1s22s22p6] 3s1 electron configuration

Shorthand Configuration

[Ar] 4s2

Electron configurationElement symbol

[Ar] 4s2 3d3

[Rn] 7s2 5f14 6d4

[He] 2s2 2p5

[Kr] 5s2 4d9

[Kr] 5s2 4d10 5p5

[Kr] 5s2 4d10 5p6

[He] 2s22p63s23p64s23d6

Fe [Ar] 4s23d6

[Kr] 5s1 4d10

General Rules

Aufbau Principle– Electrons fill the

lowest energy orbitals first.

– “Lazy Tenant Rule”

RIGHTWRONG

General Rules

• Hund’s Rule– Within a sublevel, place one electron

per orbital before pairing them.

– “Empty Bus Seat Rule”

• Shorthand Configuration

S 16e-

Valence ElectronsCore Electrons

S 16e- [Ne] 3s2 3p4

1s2 2s2 2p6 3s2 3p4

Notation

• Longhand Configuration

S32.066

d (n-1)

f (n-2) 67

Periodic Patterns

1234567

s-block1st Period

1s11st column of s-block

Periodic Patterns

• Example - Hydrogen

Periodic Patterns

• Shorthand Configuration– Core electrons:

• Go up one row and over to the Noble Gas.

– Valence electrons: • On the next row, fill in the # of e- in each sublevel.

• Full energy level

• Full sublevel (s, p, d, f)• Half-full sublevel

Stability

This fills the valenceshell and tends to givethe atom the stabilityof the inert gasses.

The Octet Rule

Atoms tend to gain, lose, or share electrons until they have eight valence electrons.

ONLY s- and p-orbitals are valence electrons.

• Electron Configuration Exceptions– Copper

EXPECT: [Ar] 4s2 3d9

ACTUALLY: [Ar] 4s1 3d10

– Copper gains stability with a full d-sublevel.

Stability

Electron Filling in Periodic Table

Zn3d10

Cr4s13d5

Cu4s13d10

2sn = 2

1sn = 1

Cr4s13d5

Cu4s13d10

Cu3d10

Stability

• Ion Formation– Atoms gain or lose electrons to become more

stable.– Isoelectronic with the Noble Gases.

Orbital Diagrams for Nickel

2s 2p 3s 3p 4s 3d1s

Excited State

VIOLATES Pauli Exclusion

VIOLATES Hund’s Rule

Ni58.6934

2 2 6 2 6 2 8

2 2 6 2 6 1 9

Write out the complete electron configuration for the following:1) An atom of nitrogen

2) An atom of silver

3) An atom of uranium (shorthand)

Fill in the orbital boxes for an atom of nickel (Ni)

2s 2p 3s 3p 4s 3d1s

Which rule states no two electrons can spin the same direction in a single orbital?

1s22s22p3

1s22s22p63s23p64s23d104p65s24d9

[Rn]7s26d15f3

Extra credit: Draw a Bohr model of a Ti4+ cation. 22+n = n

Pauli exclusion principle

Ti4+ is isoelectronic to Argon.

Answer Key

Orbitals Being Filled

3 4 5 6 7

Lanthanide series

Actinide series

Groups 8

Electron Filling in Periodic Table

Zn3d10

Cr4s13d5

Cu4s13d10

2sn = 2

1sn = 1

Cr4s13d5

Cu4s13d10

Cu3d10

Electron Configurations of First 18 Elements:Hydrogen

Lithium

Sodium

11NaMagnesium

Aluminum

Carbon

Silicon

14SiPhosphorous

Oxygen

Sulfur

Fluorine

Chlorine

Beryllium

4BeNitrogen

Helium

Electron Dot Diagrams

1A 2A 3A 4A 5A 6A 7A 8A

= valence electron

s1 s2 s2p2 s2p3 s2p4 s2p5 s2p6s2p1

1 2 13 14 15 16 17 18

First Four Energy Levels

Sublevels

3s 3p 3d

4s4p 4d 4f

Sublevel designation

Four sublevels

Three sublevels

Two sublevels

One sublevel

Principallevel 4

Principallevel 3

Principallevel 2

Principallevel 1

Principal Level 2 Divided

2s sublevel 2p sublevel

2px 2py 2pz

1sn = 1

Sublevels

Four sublevels

Three sublevels

Two sublevels

One sublevel

Principallevel 4

Principallevel 3

Principallevel 2

Principallevel 1

Metals and Nonmetals

METALS

Nonmetals

Metalloids

Isotopes of Magnesium

Atomic symbol Mg Mg Mg

Number of protons 12 12 12

Number of electrons 12 12 12

Mass number 24 25 26

Number of neutrons 12 13 14

Isotope Notation Mg-24 Mg-25 Mg-26

Timberlake, Chemistry 7th Edition, page 64

12e- 12e- 12e-

Isotopes of Hydrogen

1 p+ 1 e- 1 p+

1 n 1 e- 1 p+2 n 1 e-

Protium Deuterium Tritium

(ordinary hydrogen) (heavy hydrogen) (radioactive hydrogen)

Ralph A. Burns, Fundamentals of Chemistry 1999, page 100

H-2 H-3H-1

Isotopes of Three Common Elements

Element SymbolFractional

AbundanceAverage Atomic Mass

Carbon

Chlorine

SiliconSiSiSi

282930

27.97728.97629.974

92.21% 4.70% 3.09%

1.11%13.00313C

98.89%12 (exactly)12C

Mass (amu)

75.53%

24.47%36.96637Cl

34.96935Cl

Number

LeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World , 1996, page 110

Atomic Structure

• ATOMS– Differ by number of protons

• IONS– Differ by number of electrons

• ISOTOPES– Differ by number of neutrons

Formation of Cation

sodium atomNa

loss of one valence

electron

sodium ionNa+

11p+e-

Formation of Anion

chlorine atomCl

gain of one valence

electron

chloride ionCl1-

17p+e-

Formation of Ionic Bond

chloride ionCl1-

sodium ionNa+

11p+e-

17p+e-

Atomic Structure Unit 3

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