Modern Atomic Theory Notes

Electromagnetic radiation – energy that travels through space as waves.

Waves have three primary characteristics:

• Wavelength (- lambda) – distance between two consecutive peaks or troughs in a wave. Unit = meter

• Frequency ( = nu) – indicates how many waves pass a given point per second. Unit = Hertz (Hz)

• Speed – velocity (c = speed of light = 3 x 108 m/sec) - indicates how fast a given peak moves in a unit of time

c =

Electromagnetic radiation (light) is divided into various classes according to

wavelength.

Wave- Particle Theory – Light as waves – Light as photons (de Broglie)

Photon/quantum – packet of energy – a “particle” of electromagnetic radiation

Energy - (E – change in energy) – Unit Joules (J)

Planck’s Constant –

(h = 6.626 x 10-34 J * s)

Ephoton = hChange in Energy of a photon = (Planck’s Constant) x (frequency)

c = + Ephoton = h = Ephoton = hc

Ex: What is the wavelength of light with a frequency of 6.5 x 1014 Hz? What is the change in Energy of the photon?

Given

= 6.5 x 1014 Hz

= ? ΔE = ?

= c

= 3 x 108 m/sec

6.5 x 1014 Hz

λ = 4.6 x 10-7 m

E = hc = (6.626 x 10-34 J x s)(3 x 108 m/s)

4.6 x 10-7 m

ΔE = 4.3 x 10-19 J

Wrap – upSo with light waves, you can convert between wavelength,

frequency, and energy with two equations:

= c E = h And two constants:

c = 3 * 108 m/s h = 6.626 * 10-34 J s

In the visible part of the spectrum, different colors correspond to different frequencies, wavelengths and

energies. Blue light has a short wavelength, high frequency and high energy. Red light has a long

wavelength, low frequency, and low energy.

Learning Check

Excited State – atom with excess energy

Ground State – lowest possible energy state

Wavelengths of light carry different amounts of energy per photon

Only certain types of photons are produced (see only certain colors)

Quantized – only certain energy levels (and therefore colors) are allowed

Emission and Absorption Spectra

Emission Spectrum – bright lines on a dark background. Produced as excited electrons return to a ground state – as in flame tests.

Absorption Spectrum – dark lines in a continuous spectrum. Produced as electrons absorb energy to move into an excited state, only certain allowable transitions can be made. Energy absorbed corresponds to the increase in potential energy needed to move the electron into allowed higher energy levels. The frequencies absorbed by each substance are unique.

Intensity

Color

An Element’s Fingerprint

• When excited by heat or electricity, gases glow with characteristic colors.

• A prism can be used to spread out the light from these hot gases.

• This reveals a series of discrete lines, the element’s fingerprint.

• Chemists use these fingerprints (called spectral lines) to identify elements both in the lab and in space.

Here are some spectral lines

Learning CheckNow, try matching each of the spectra from column A with its corresponding line plot from

column B.

A B

Bohr Model – suggested that electrons move around the nucleus in circular orbits

Only Correct for Hydrogen

Wave Mechanical Model – Described by orbitals gives no information about when the electron occupies a certain point in space or how it moves *aka – Heisenberg's Uncertainty Principle

Parts of the Wave Mechanical Model

1. Principle Energy Level (n) – energy level designated by numbers 1-7.

-called principle quantum numbers

2. Sublevel – exist within each principle energy level-the energy within an energy level is slightly different-each electron in a given sublevel has the same energy-lowest sublevel = s, then p, then d, then f

1

2

3

4

5

6

7

s pd

f

Parts of the Wave Mechanical Model cont.

3. Orbital – region within a sublevel or energy level where electrons can be found

s sublevel – 1 orbitalp sublevel – 3 orbitalsd sublevel – 5 orbitalsf sublevel – 7 orbitals

- ** No more than two electrons can occupy an orbital**-an orbital can be empty, half-filled, filled

Electron Configuration – arrangement of the electrons among the various orbitals of the atom

Ex: 1s22s22p6 = Neon

Sulfur = 1s2 2s2 2p63s2 3p4

Cd = 1s2 2s2 2p63s2 3p64s2 3d10 4p65s2 4d10

Na = 1s2 2s2 2p63s1

Ne

Na

Learning CheckWrite electron configurations:

Oxygen

Chromium

Shapes of orbitalsAll s orbitals are spherical as the principle energy level increases the diameter increases.

All p orbitals are dumbbell or figure-8 shaped – all have the same size and shape within an energy level

4 of the d orbitals are 4-leaf clover shaped and the last is a figure-8 with a donut – all have the same

size and shape within an energy level

f orbitals are complicated!!!!!

Electron Spin Spin – motion that resembles earth rotating on its axis– clockwise or counterclockwise

Pauli Exclusion Principle – two electrons in the same orbital must have opposite spins

Hund’s Rule – All orbitals within a sublevel must contain at least one electron before any orbital can have two

Orbital Diagram – describes the placement of electrons in orbitals• use arrows to represent electrons with spin• line represents orbital (s=1, p=3, d=5, f=7)

____ full ____ half-full ____ empty

Orbital Diagrams

Ex: Neon = 1s__2s__ 2p__ __ __

Carbon = 1s__2s__ 2p__ __ __

Zinc = 1s__2s__ 2p__ __ __3s__ 3p__ __ __

4s__ 3d__ __ __ __ __

Gallium =1s__2s__ 2p__ __ __3s__ 3p__ __ __

4s__ 3d__ __ __ __ __ 4p__ __ __

Learning CheckDraw orbital diagrams:

Oxygen

Chromium

• Noble Gas Configuration – Shorthand configuration that substitutes a noble gas for electrons

Ex:

• Valence Electrons – Electrons in the outermost (highest) principle energy level in an atom

• Core Electrons – innermost electrons – not involved in bonding

• Valence Configuration – shows just the valence electrons

Ex:

Na = 1s22s22p63s1 or [Ne]3s1

Sn = 1s22s22p63s23p64s23d104p65s24d105p2 or [Kr]5s24d105p2

Na = 3s1 3rd Shell/1valence electron

Sn = 5s25p2 5th Shell/4 valence electrons

Na = 1s22s22p63s1 1 Valence

Sn = 1s22s22p63s23p64s23d104p65s24d105p2 4 Valence

Learning CheckWrite the noble gas configuration, valence configuration, and number of valence electrons:

Oxygen

Chromium

Periodic Table

Dimitri Mendeleev-1869- developed the first version of the periodic table.

He expressed the regularities as a periodic function of the atomic mass.

Henry Moseley- revised Mendeleev periodic table by describing regularities in physical and chemical properties as periodic functions of the atomic number

Periods – horizontal rows•Period number corresponds to the principal quantum number of valence electrons

Groups (family) – vertical columnElements with similar valence electrons configurationsGroup 1 – alkali metals – reactiveGroup 2 - alkaline earth metals – reactiveGroup 3-12 – transition metalsGroup 15 – nitrogen familyGroup 16 – oxygen family – reactiveGroup 17 – halogens – very reactiveGroup 18 – noble gases

Periodic Trends

1. Atomic Radius/Size – size of an atom

Increases – down a groupDecreases – across a period

Size of ions

Cation Ca+2/Ca Ca larger because Ca+2 lost 2 electronsAnion S-2/S S-2 larger because S-2 gained 2 electrons

2. Ionization Energy – energy required to remove an electron from an individual atom in a gas phaseM(g) M+

(g) + e-

(energy to make a positive ion)• Metals lose electrons to non-metals so

relatively low energy is needed• High ionization energy means an electron

is hard to removeDecreases – down a groupIncreases - across a period

3. Electron Affinity – Electron affinity is the energy involved when an electron is added to a gaseous atom.

• Negative values of energy mean that energy was released during the process. Atoms with negative values of electron affinity have a very strong attraction for electrons.

• Positive values of electron affinity have very little attraction for electrons.

(energy involved in negative ions)

Decreases – down a groupIncreases - across a period

4. Electronegativity is the tendency of an atom to draw electrons to itself when in a covalent bond. Consequently, the trends are the same as for electron affinity.

The atoms with the highest electronegativity are fluorine, then oxygen, then nitrogen. It is also important to know that the electronegativity of hydrogen is slightly less than that of carbon.

Decreases – down a group

Increases - across a period

5. Metallic Character

Increases – down a group

Decreases – across a periodElectronegativity

Ele

ctro

neg

ativ

ity

Learning CheckPut the following elements in order of increasing atomic radius:

a. Ge, Se, Fe, Cab. C, Pb, Sn, Si, Ge

Put the following elements in order of increasing electronegativity:

a. Ge, Se, Fe, Cab. C, Pb, Sn, Si, Ge

Notes- Chemical Bonding

Bond- force that holds groups of two or more atoms together and makes them function as a unit

bond energy- energy required to break the bond (tells the bond strength)

Ionic bonding- between ionic compounds which contain a metal and a nonmetal

• Atoms that lose electrons relatively easily react with an atom that has a high affinity for electrons

• Transfer of electronsCovalent bonding- between two nonmetals• Electrons are shared by nucleiPolar Covalent bonding- unequal sharing of electrons• positive end attracted to the negative end (delta) indicates partial charge

• electronegativity-(p. 362) relative ability of an atom in a molecule to attract shared electrons to itself

• The higher the atom’s electronegativity value, the closer the shared electrons tend to be to that atom when it forms a bonds

Increases – across a periodDecreases- down a group

Electronegativity difference

Bond type Covalent character

Ionic character

Zero (0-.4) Covalent Decreases Increases

Intermediate

(.4 – 1.4)

Polar covalent Decreases Increases

Large (<1.4) Ionic Decreases Increases

Ex. List the following in order of increasing polarity.H-H, O-H, Cl-H, S-H, F-H

H-H = 2.1 - 2.1 = 0O-H = 3.5 - 2.1 = 1.4

Cl-H = 3.0 - 2.1 = .9

S-H = 2.5 - 2.1 = .4

F-H = 4.0 - 2.1 = 1.9 H-H, S-H, Cl-H, O-H, F-H

• Dipole moment- has a center of positive charge and a center of negative charge

• Represented by an arrow

• Arrow points toward the negative charge

Chemical Formula – type of notation made with numbers and chemical symbols– indicates the composition of a compound– indicates the number of atoms in one molecule

Molecule - Bonded collection of two or more atoms of the same element or different elements

- monatomic molecule – one atom molecules

- diatomic molecule – two atom molecules (seven) MEMORIZE

Br, I, N, Cl, H, O, F

MetalsLocation: Left side of Periodic TableProperties: Ductile – drawn into wires

Malleable – hammered into sheetsMetallic Luster – shineGood Conductors of Heat and Electricity

NonmetalsLocation: Right side of Periodic TableProperties: Brittle

Lack Luster – not shinyPoor Conductors of Heat and Electricity

Semi-metalsLocation: Along Stair-stepProperties: Have properties of metals and nonmetals also called METALLOIDS Si, Ge, As, Sb, Te, Po, At

METALS

Nonmetals

Semi-metals

Molecular NomenclatureMolecular Compounds (molecules) – compounds made from two nonmetals

- electrons are shared by two atoms

Naming MolecularPrefixes: (MEMORIZE)Mono-1 tetra-4 hepta-7 deca-10di-2 penta-5 octa-8tri-3 hexa-6 non-9prefixes are used with both the first named and second named element. Exception:

mono- is not used on the first wordsecond word ends in –ideIf a two syllable prefix ends in a vowel, the vowel is dropped before the prefix is attached

to a word beginning with a vowel monooxide

N2O dihydrogen monoxide

Si8O5 tetrasulfur hexachloride

NH3 carbon monoxide

P3I10 carbon dioxide

= Dinitrogen monoxide

= Octasilicon pentoxide

= Nitrogen trihydride

= Triphosphorus deciodide

= H2O

= S4Cl6

= CO

= CO2

Writing molecular formulasTranslate prefixes

Examples:

Learning Check

Write the name:

a.C2O4

b.P2O5

Write the formula:

a. Dihydrogen monoxide

b. Phosphorus trihydride

Valence electrons are used in bonding.

• Stable elements want to achieve 8 electrons similar to the noble gases• If it’s a metal it wants to achieve the configuration for the noble gas before.• If it’s a nonmetal it wants to achieve the configuration for the noble gas

after.

1

2 3 4 5 6 7

8

2

Lewis Structure- representation of a moleculeShows how the valence electrons are arranged

among the atoms in the molecule. spz X px

py

Oxygen1s22s22p4

For an element: O•• ••• •

For a compound:

Li + [Li]+1 + [ Cl ]-1Cl

For a molecule:

F F••••••

••••••

Duet rule- only two electrons in the full shell

Octet rule- surrounded by eight electrons

Bonding pair- electrons shared with other atom

Lone pair or unshared pair- not involved in bonding

H & He

Happy Eight!!!!!

Line (-) = 2 electrons

dots (••) = 2 electrons/each dot is one electron

5 Steps for Covalently Bonded Lewis Structures1. Find the total number of valence electrons.2. Calculate the number of “needed” electrons to give each atom 8

electrons, except for H which wants 2.3. Subtract valence electrons from the “needed” electrons. This is the

number of bonding electrons.4. Divide the bonding electrons by 2, to find the number of bonds.5. Subtract the bonding electrons from the valence electrons to find

the non-bonding electrons or lone pairs.6. Choose a central atom and assemble the pieces to make all atoms

involved stable.

Ex. GeBr4

Valence = 1(4) + 4(7) = 32Needed = 1(8) + 4(8) = 40Bonding = 40 – 32 = 8Bonds = 8/2 = 4 linesLone e- = 32 – 8 = 24 dotsCentral atom = Ge

Ge

Br

Br

Br

Br

• • ••

••

• • ••

• •••••

••

• •••• •

• Single bond- involves two atoms sharing one pair

• Double bond- involves two atoms sharing two pairs

• Triple bond- involves two atoms sharing three pairs

Ex. CH4 C2H4 C2H21. 1(4) + 4(1) = 82. 1(8) + 4(2) = 163. 16 - 8 = 84. 8/2 = 4 lines5. 8 – 8 = 0 dotsCentral atom = C

1. 2(4) + 4(1) = 122. 2(8) + 4(2) = 243. 24 - 12 = 124. 12/2 = 6 lines5. 12 – 12 = 0 dotsCentral atom = C

1. 2(4) + 2(1) = 102. 2(8) + 2(2) = 203. 20 - 10 = 104. 10/2 = 5 lines5. 10 – 10 = 0 dotsCentral atom = C

C

H

H

H

H C CH

HH

HC CH H

Resonance- more than one Lewis structure can be drawn for the molecule

Ex. CO2

1. 1(4) + 2(6) = 162. 1(8) + 2(8) = 243. 24 - 16 = 84. 8/2 = 4 lines5. 16 – 8 = 8 dotsCentral atom = C

CO O••••

••••

C OO••••••

••

C OO••••••••

Exceptions to the Octet Rule1. boron and beryllium- tend to be electron

deficient– boron can hold 6 total electrons– beryllium can hold 4 total electrons

ex. BF3 BeH2 1. 1(3) + 3(7) = 242. 1(6) + 3(8) = 303. 30 - 24 = 64. 6/2 = 3 lines5. 24 – 6 = 18 dotsCentral atom = B

B

F

F F

• •••••

••••••••••

••

1. 1(2) + 2(1) = 42. 1(4) + 2(2) = 83. 8 - 4 = 44. 4/2 = 2 lines5. 4 – 4 = 0 dotsCentral atom = Be

BeH H

2. Electrons are small spinning electric charges that create magnetic fields

– Diamagnetic- substances which have paired electrons that cancel out the magnetic field

– Paramagnetic- substances the have one or more unpaired electrons that show great attraction to the magnetic field

Ex. O2 PH3

O O••••

••••

PH H

H

• •

3. Odd number of electrons– You cannot write electron dot structures that fulfill the octet

rule, when the total number of valence electrons is odd

Ex. NO

4. Expanded Octet- expand the valence shell to include more than 8 electrons

– Phosphorus and sulfur can expand to include 10 or 12 electrons

– You will know you have an expanded octet when you don’t have enough bonds for the atoms present

Ex. SF6

1. 1(5) + 1(6) = 11

No Drawing

F••

F

F

F

F••

••

F

S

••

•• ••

••••

••••••••

••••

•••••• ••

Lab – Lewis Structures

Structure (shape)

Molecular (geometric) structure- three-dimensional arrangement of the atoms in a molecule

VSEPR model- valence shell electron pair repulsion

• Lone pairs of electrons like to be as far away from each other as possible

• Double and triple bonds “act” like a single shared pair for shape.

Linear

Linear- two pairs of electrons are present around an atom– One total pair – one shared

pair– Two total pairs – two

shared pairs– Bond angle = 180

Ex. BeCl2 1. 1(2) + 2(7) = 162. 1(4) + 2(8) = 203. 20 - 16 = 44. 4/2 = 2 lines5. 16 – 4 = 12 dotsCentral atom = Be

BeCl Cl•••• • •

• • ••••

Bent

Bent – Four total pairs– Two shared pairs and

two unshared pairs– Bond angle = 104.5

Ex. H2O1. 2(1) + 1(6) = 82. 2(2) + 1(8) = 123. 12 - 8 = 44. 4/2 = 2 lines5. 8 – 4 = 4 dotsCentral atom = O

O H

H

• •••

Trigonal planarTrigonal planar- whenever

three pairs of electrons are present they should be placed at the corners of a triangle– Three total pairs– Three shared pairs– Bond angle = 120

Ex. BCl31. 1(3) + 3(7) = 242. 1(6) + 3(8) = 303. 30 - 24 = 64. 6/2 = 3 lines5. 24 – 6 = 18 dotsCentral atom = B

B

Cl

Cl Cl

• •••••

••••••••••

••

Tetrahedral

Tetrahedral– Four total pairs– Four shared pairs no

unshared pairs– Bond angle = 109.5

Ex. CCl4

Valence = 1(4) + 4(7) = 32Needed = 1(8) + 4(8) = 40Bonding = 40 – 32 = 8Bonds = 8/2 = 4 linesLone e- = 32 – 8 = 24 dotsCentral atom = C

C

Cl

Cl

Cl

Cl

• • ••

••

• • ••

• •••••

••

• •••• •

Trigonal pyramid

Trigonal pyramid– Four total pairs– Three shared pairs

and one unshared pair– Bond angle = 107

Ex. NH3

1. 1(5) + 3(1) = 82. 1(8) + 3(2) = 143. 14 - 8 = 64. 6/2 = 3 lines5. 8 – 6 = 2 dotsCentral atom = N

NH H

H

• •