Electrons and the Periodic

TableHonors Chemistry

The Periodic Table

Development of the Periodic Table

1) History of the Periodic Table – By the end of the 1700’s, scientists had identified only 30 elements (ex. Cu, Ag, Au, H2, N2, O2, C).

2) By the mid 1800’s, about 60 elements had been identified.

3) Sept 1860 – chemists assembled at the First International Congress of Chemists in Germany to settle the controversial issues such as atomic mass. Standard values set for atomic mass and improved communication for research.

Johann Dobereiner: 1817Organized the elements into

sets of three with similar properties.

He called these groups triads. The middle element is often the average of the other two.

Ex) Cl – 35.5 Br – 79.9 I – 126.9

CaAvg Sr

Ba

Cl + IAvg.

2

Triads on the Periodic Table

John Newlands: 1866• Arranged elements in

order of increasing atomic mass.

• Noticed repeating patterns in the elements’ properties every 8th element.

• Law of Octaves - properties of elements repeated every 8th element.

• There were 62 known elements at the time.

Dmitri Mendeleev: 1869• Arranged elements in

order of increasing atomic mass.

• Similar properties occurred after periods (horizontal rows) of varying lengths.

• Organized the 1st periodic table according to increasing atomic mass and put elements with similar properties in the same column.

• Periodic – repeating properties or patterns

Mendeleev’s 1st

Periodic Table

• Mendeleeve Noticed inconsistencies in the arrangement of increasing atomic mass.

• He arranged some elements out of atomic mass order to keep them together with other elements with similar properties. (Notice Te and I)

• He also left several blanks in his table. • In 1871, he correctly predicted the

existence and properties of 3 unidentified elements – Sc, Ga and Ge

• These elements were later identified and matched his predictions.

1st Periodic Law

Properties of the elements repeat periodically when the elements are arranged in increasing order

by atomic mass

Mendeleev is known as the Mendeleev is known as the

Father of ChemistryFather of Chemistry

Element 101 (Md) honors Element 101 (Md) honors MendeleevMendeleev

What’s wrong with Mendeleev’s PT?

• We know that Mendeleev's periodic table was underpinned by false reasoning. – Mendeleev believed, incorrectly, that chemical

properties were determined by atomic weight.

• In 1869 the electron itself had not been discovered until 1896 - 27 years later.

• It took 44 years for the correct explanation of the regular patterns in Mendeleev's periodic table to be found...

Henry Moseley: 19131. Fired electrons at atoms, resulting in

emission of x-rays. Each element emitted x-rays at a unique frequency. – He examined the pattern that was best

explained if the positive charge in the nucleus increased by exactly one unit from element to element.

2. Element are different from one another because their atoms have different number of positive charge (protons).

3. Analyzed data and found that the elements in the PT fit into patterns better when arranged in increasing nuclear charge, which is the Atomic Number.

The Modern Periodic

Table: When elements are arranged in

order of increasing

atomic number, their physical and chemical

properties show a periodic pattern.

Glenn Seaborg “Seaborgium” Sg #106

• Born in 1912 in Michigan, Seaborg proposed reorganizing the Periodic Table one last time as a young chemist working on the Manhattan Atomic Bomb Project during WWII.

• He suggested pulling the “f-block” elements out to the bottom of the table.

• He was the principle or co-discoverer of 10 transuranium elements.

• He was awarded the Noble prize in 1951 anddied in 1999.

Seaborgium is the exception…• After some argument between the USA and the rest

of the world, element 106 was named Seaborgium shortly before he died. This was a matter of some controversy because the International Union of Pure and Applied Chemistry, IUPAC, the body that deals with naming in chemistry, had previously ruled that elements should not be named after living people.

Atomic # 104 105 106 107

IUPACUnnilquadeum

UnqUnnilpentium

UnpUnnilhexium

UnhUnnilseptium

Uns

Agreed in 1995

Dubnium(Dubna, Russia)

Joliotium(Frederic Joliot)

Rutherfordium

(Earnest Rutherford)

Bohrium(Neils Bohr)

Agreed in 1996

Rutherfordium

Dubnium Seaborgium Bohrium

Parts of the Periodic Table

A. Horizontal Rows – PERIODS– There are 7 periods in the periodic table– Elements in a period do NOT have similar

properties.

B. Vertical Columns – GROUPS or FAMILIES– Labeled 1-18– IA-VIIIA are the Main-group or

representative elements. – Elements in a group have similar

properties. – Why?

Family Nameswrite these on your P.T.

• Hydrogen (1)• Alkali metals (1) – most

reactive metals; reactivity increases down the group

• Alkaline earth metals (2)• Boron family (13)• Carbon family (14)• Nitrogen family (15)• Oxygen or Chalcogen

family (16)• Halogens (17)• Noble gases (18) - inert

• Transition elements or metals (3-12): d-block

• Inner transition elements or metals (f-block)– Lanthanides or

lanthanide series– Actinides or actinide

series– Transuranium elements

1A 1

2A2

3A 4A 5A 6A 7A13 14 15 16 17

8A18

1 = Alkali Metals and Hydrogen Group

13 = Boron Group

18 = Noble Gas Group

17 = Halogen Group

16 = Oxygen or Chalcogen Group

15= Nitrogen Group

14 = Carbon Group

2 = Alkaline Earth Metals

Transition Metals

Lanthanide Series

Actinide Series

Inner Transition Metals

3 4 5 6 7 8 9 10 11 12

Parts of the Periodic Table

The ELECTRONHow are the

electronsarranged

around thenucleus?

To Review• Dalton Thomson Rutherford

Bohr Quantum Mechanical (Schrödinger) Model

• Bohr – electrons in a particular path have a fixed energy called energy levels– Rungs of a ladder

• Quantum Mechanical (Schrödinger) Model– Electrons better understood as WAVES– Does not tell where the electrons are located– Electrons have a certain amount of energy -

QUANTIZED

Parts of a Wave

Light as a WaveCharacteristics of a WaveA. Amplitude: Height of the wave from the

baseline. The higher the wave the greater the intensity.

B. Wavelength: (λ , “lambda”) in nanometers (1 x 10-9 m). Distance between similar points on 2 consecutive waves.

C. Frequency: (ν , “nu”) The number of waves that pass a fixed point per unit of time. Measured in cycles/second (1/s) 1 cycle/second = Hertz (Hz)

Baseline

Crest

Wavelength Amplitude

Amplitude

Wavelength Trough

3

D. Electromagnetic Radiation - a form of energy that exhibits wavelike behavior

as it travels through space- all forms of EM radiation move at the speed of

light

Speed of Light (c)

E. 3.00 x 108 m/s or 186,000 miles/sec. The relationship between wavelength

and frequency can be shown with the

following equation:

This is an indirect relationship. If λ then ν .

c = λ ν

In the US, AM stations all have longer wavelengths than the FM stations. •AM broadcast stations are licensed to operate only in the band 550-1700 kHz •FM broadcast stations are licensed to operate only in the band 88-108 MHz

Visible Light

Radar

Microwaves

Infrared

Radio/TV Ultraviolet

X-Rays

Gamma Rays

Low

Long

High

Short

Red Orange Yellow Green

Blue Violet

Energy

Energy

Low High

• Incandescent light bulbs give off most of their energy in the form of heat-carrying infrared light photons -- only about 10 percent of the light produced is in the visible spectrum. This wastes a lot of electricity. Cool light sources, such as fluorescent lamps and LEDs, don't waste a lot of energy generating heat -- they give off mostly visible light. For this reason, they are slowly edging out the old reliable light bulb.

http://home.howstuffworks.com/light-bulb2.htm

Quantum Theory

A. Planck’s Hypothesis: (Max Planck 1900)

1. Studied emission of light from hot objects

2. Observed color of light varied with temperature

3. Suggested the objects do not continuously emit E, but emit E in small specific amounts

a. Light is absorbed or emitted in a little packet or bundle called a quantum (quanta –plural).

b. Quantum = minimum amount of E that can be lost or gained by an atom

c. Energies are quantized. (Think steps not a

ramp)

e-

e-

Xe-

e-

Max Planck’s Energy Equation

4. Proposed that energy is directly proportional to frequency.

E = h Planck’s equation for each quantum

h = Plank’s constant = 6.626 x 10-34 J.s

This is a direct relationship.As energy increases, frequency

increases.

Albert Einstein

(1879 – 1955)

German Physicist

While well-known for the equation E=mc2 , Einstein’s work on the photoelectric effect resulted in being awarded the 1921 Nobel Prize in Physics.

Refers to the emission of electrons from a metal

when light shines on the metal

Observations:1. Electrons are ejected by light of

sufficient energy. Energy minimum is different for different metals.

2. The current (# of electrons emitted/s) increases with brightness of the light.

+-

Albert Einstein and the Photoelectric

Effect

my.hrw.com

Conclusions: 1. Proposed that light consists of quanta of

energy that behaves like particles. 2. Quantum of light = photon = massless

particle that carries a quantum of energy.

3. Proposed the Dual Nature of Light: its wave and particle nature.

a) Light travels through space as wavesb) Light acts as a stream of particles when it

interacts with matter.

Albert Einstein and the Photoelectric Effect

my.hrw.com

Light (Electromagnetic Radiation)

SpectroscopyDefinition: a method of studying substances

that are exposed to some sort of continuous exciting energy.

A. Emission Line Spectra: contains only certain colors or wavelengths ( ) of light.

1. Every element has its own line spectrum (fingerprint).

Continuous Spectrum – White Light

Line Spectrum – Excited Elements

White light gives off a Continuous Spectrum

a blending of every possible wavelength

Gas Discharge Tubes

• Electricity is added to the gas which causes the electrons to jump to a higher or excited state. They immediately fall back to the ground state and give off particular wavelengths of light. We see a blending of wavelengths without the spectroscopes.

Flame Tests

• used to test qualitatively for the presence of certain metals in chemical compounds. 

• the heat of the Bunsen flame excites electrons that emit visible light.

Copper(II) sulfate Lithium chloride

Potassium chloride Barium nitrate

Spectroscope• Uses a diffraction grating to diffract the

light into particular wavelengths of light.

A Line Spectra result from excited elements - as electrons of an element gain energy and

rise to an excited state they then fall back to their ground state in the same pattern

producing the same energy drop each time which we see as individual wavelengths of

light.

Atomic Spectra and the Bohr Model of Hydrogen (1913)

Neils Bohr - Danish ScientistExplained the bright-line spectrum

of hydrogenStudy: • Added E as electricity to H gas

at low pressure in a tube.• Emitted E as visible light, was

observed through a prism

Result: Hydrogen emitted 4 distinct bright lines of color, aka bright line spectrum

Electrons release energy as they fall back to a lower

energy level

Electrons absorb energy to rise to a higher or excited state and emit energy in the form of a photon of light as they fall back to their

ground states.

Path of an excited electron as it “falls” back to the Ground State

• When electrons gain energy, they jump to a higher energy level (excited state).

• Electrons are not stable at the excited state and will immediately fall back to a lower level or ground state.

• As they fall, they emit electromagnetic radiation.

• Depending on how far they fall determines the type of radiation (light) released.

Bohr Model of HydrogenConclusion: • *Unique line spectrum is due to quantized electron energies.• *Electrons are in specific orbits related to certain amounts of

energy known as stationary states. • *Orbits are related to energy levels.

• *Energy levels are identified as E1, E2, E3, … (n = 1, 2, 3, …)

• *Lowest energy level = ground state• *Electrons absorb certain amounts of energy to move to a higher

energy level farther away from the nucleus = excited state• *Electrons return to the more stable ground state and release a

photon that has energy equal to the difference in energy between the energy levels.

– from E2 to E1: Ephoton = E2 – E1 (difference in energy)

The Bohr Atom for Hydrogen -a Model

1. Successful in calculating the wavelength, frequency, & energy of hydrogen’s line spectrum.

2. Successful in calculating the energy needed to remove hydrogen’s electron

H(g) + energy H+1(g) + 1e-

Calculated ionization E = observed ionization E = 1312.1 kJ/mol

Lyman, Balmer and Paschen series of the Hydrogen Atom

• Lyman series: electrons fall to n = 1 and give off UV light.

• Balmer series: electrons fall to n = 2 and give off visible light.

• Paschen series: electrons fall to n = 3 and give off infrared light.

When electrons absorb energy they jump to a higher (excited) state.

n=2 n=3 n=4 n=5 n=6 n=7

Electrons are not stable. Radiation (light) is emitted when an electron falls back from a higher level to a lower level.

Ultraviolet Light

Visible Light

Infrared Light

Atomic Spectra

Hydrogen

Lithium

Mercury

Helium

Although Bohr’s atomic model explained the line spectra of hydrogen, it failed for heavier elements.

Limitations of the Bohr Model

a. Model could not calculate the wavelengths of observed spectra of multi-electron atoms.

b. Model could not explain the chemical behavior of atoms.

c. Bohr used classical mechanics to understand the behaviors of small particles.

d. The Bohr model is also known as the planetary, solar system, or satellite model.

Quantum Mechanical Model of the Atom

A. Louie De Broglie (1924-5) 1. Took Einstein’s idea that light can

exhibit both wave and particle properties2. Very small particles (like electrons)

display properties of waves.3. Behavior of electrons in Bohr’s quantized

orbits was similar to behavior of waves

Known: any wave confined to a space can only have specific frequenciesDe Broglie suggested electrons are waves confined to the space around the atomic nucleus. Electrons could exist only at specific frequencies which correspond to specific energies (E = h quantized E of Bohr)

French scientist

4. Experimentally proven in 1927 by diffraction of electrons by Davisson & Germer (showed diffraction of electrons by a crystal of Ni)

B. Wave-Particle Duality of Naturea. Light and electrons (very small particles like

electrons, atoms, molecules) have properties of waves and particles QUANTUM MECHANICS (based on WAVE properties)

**Large objects obey the laws of classical mechanics**

Quantum Mechanical Model

of the Atom

C. Werner Heisenberg: (1927)

1. Heisenberg’s Uncertainty Principle: states that it is impossible to determine simultaneously both the position and velocity of an electron or any other particle.

2. You cannot predict future locations of particles.

3. He found a problem with the Bohr Atom - no way to observe or measure the orbit of an electron.

D. Erwin Schrödinger Wave Equation (1926)

1. Wave nature of an electron is described by a mathematical equation.

2. Four quantum numbers in the equation are used to describe an electron’s behavior – location and energy.

3. Electron is treated as a wave with quantized energy.

4. Describes the probability of the electrons found in certain locations around the nucleus.

(1887 – 1961) Austrian Physicist

Electron DensityAn orbital is a region inwhich an electron with aparticular energy is likelyto be found. Where the density of anelectron cloud is high

there is a high probability that

iswhere the electron islocated. If the electrondensity is low then there

isa low probability.

E. Atomic Orbitals - region around the nucleus where an electron with a particular energy is likely to be found (not the same as Bohr’s orbits!)1. Orbitals have characteristic shapes, sizes,

& energies.2. Orbitals do not describe how the electron

moves.3. The drawing of an orbital represents the

3-dimentional surface within which the electron is found 90% of the time.

4. Sublevels can have 4 different shapes

s – orbital spherical

1s, 2s & 3s orbitals Superimposed on one another

Electron-Cloud Models

p-orbital – dumbbell shaped

p-orbital - dumbbell shaped

d-orbital - double dumbbell or fan blades

s,p and d orbitals

For a more complete representation and presentation of atomic orbitals go to http://winter.group.shef.ac.uk/orbitron/

x

y

z

x

y

z

x

y

z

x

y

z

x

y

z

x

y

z

x

y

z

x

y

z

x

y

z

s orbital p orbitals

d orbitals

Models of d-orbitals

f-orbital – more

complex!

f orbitals

f – orbitals (3D)

Quantum Numbers

• Each quantum number provides more specific information on the probable location of an electron.

• Each electron within an atom can be described by a unique set of 4 quantum numbers.

Quantum Numbers - Finding an address for each

electron:1. “state” Principle Quantum Number (n)

or the energy level; a. Describes the relative size of the electron cloud.b. Positive integer values (n = 1 to n = 7)

2. “city” Sublevel (l) a. Describes the shape of the electron cloud.b. The maximum number of sublevels within a

level = nc. Shapes are s, p, d,or f.d. Lowest energy = s Highest energy = f

Quantum numbers cont.

3. “street” Orbital (ml) odd # of orbitals1. Describes the orientation or direction in

spacea) s – 1 orbitalb) p – 3 orbitals (x, y, z)c) d – 5 orbitals (xy, yz, xz, x2 – y2, z2)d) f – 7 orbitals (y3 – 3yx2, 5yz2-yr2, x3-3xy2,

zx2-zy2, xyz, 5xz2-3xr2, 5z3-3zr2)

2. Orbitals within the same sublevel have the same energy are called degenerate orbitals

3. An orbital can hold a maximum of 2 electrons

4. “house” Spin (ms) 1. Describes the direction of electron spin in

an orbital. 2. The clockwise or counterclockwise motion

of electrons.3. Only electrons with opposite spins can

occupy the same orbital. 4. The opposite spin is written as+1/2 or -1/2 or or

Quantum numbers cont.

E. Electron Configurations:

1. Shorthand notation for indicating the number of electrons in each level, sublevel, and orbital.

1s2

2. Shows the distribution of electrons among the orbitals. Describes where the electrons are found & what energy they possess.

Electron Configuration Rules

1. The Aufbau Principle: electrons are added one at a time to the lowest energy orbital available.

Pauli Exclusion Principle:1. Each orbital can only hold 2 electrons.

2. The electrons must have opposite spins.

s-sublevel = max 2 electronsp-sublevel = max 6 electronsd-sublevel = max 10 electronsf-sublevel = max 14 electrons

incorrect: ↑↑↑ incorrect: ↑↑ correct: ↑↓

Hund’s Rule:• Electrons will

remain unpaired in degenerate orbitals before they pair up.

incorrect ↑↓ ↑ __

correct ↑ ↑ ↑

Electron Blocks on the

Periodic Table

Incr

easi

ng e

nerg

y

1s

2s

3s

4s

5s6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

Pauli Exclusion Principle: No more than 2 e- are put in each orbital and they must have opposite spin.Hund’s Rule: electrons spread out among equal energy orbitals in a sublevel (like charges repel)Aufbau Principle: Electrons fill lowest energy levels first (n=1)

Electron Configuration Examples:

Ex) electron configuration for Na:

1s2 2s2 2p6 3s1

Ex) orbital filling box diagram for Na:

yx z

1s 2s 2p 3s

_

Electron Dot Diagrams:

Write the symbol for the element.

Place dots around the symbol to represent the

valence s & ps & p electrons only.

Do NOT include d & f orbitals in diagram.

p orbital electrons s orbital electrons

Electron Configuration Orbital Box Diagram Electron-dot Diagram

yx z

1s 2s 2p

y yx z x

1s 2s 2p 3s 3p

z

168O

3517 Cl 1s22s22p63s23p5

12752Te 1s22s22p63s23p64s23d104p65s24

d105p4

1s22s22p4

y y y yx z x z x z x z

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p

What does the Tellurium electron-dot resemble???

Mark your Periodic Tables

1 2 13 14 15 16 17 18

Unpaired vs. Paired Electrons

Filled and Half-filled orbitals• Atoms with unpaired electrons are said to

be paramagnetic. These are weakly attracted to a magnetic field.

• Atoms with all paired electrons are said to be diamagnetic. These are weakly repelled from a magnetic field.

• ½ filled and filled orbitals have special stability

Noble Gas or Shorthand Electron Configurations

• Rb

• Se

• At

1[Kr]5s

2 10 4[Ar]4s 3d 4p

2 14 10 5[Xe]6s 4f 5d 6p

Draw the Dot Diagrams for these elements

Exceptions to the Rules

•Max stability - ½ filled and filled orbitals–Cr–Mo–Cu–Ag–Au

2 4[Ar]4s 3d 1 5[Ar]4s 3d

1 10[Ar]4s 3d2 9[Ar]4s 3d

Exceptions to the Rules

•Max stability - ½ filled and filled orbitals–Cr–Mo–Cu–Ag–Au

2 4[Ar]4s 3d 1 5[Ar]4s 3d

1 5[Kr]5s 4d

1 10[Ar]4s 3d

1 10[Kr]5s 4d

1 14 10[Xe]6s 4f 5d

2 9[Ar]4s 3d

Electron Configuration for Ions

• K

• P

• Al

• Se

1[Ar]4s • K+1

• P-3

• Al+3

• Se-2

2 2 6 2 6 11s 2s 2p 3s 3p 4s2 2 6 2 61s 2s 2p 3s 3p

[Ar]

2 2 6 2 31s 2s 2p 3s 3p

2 3[Ne]3s 3p2 2 6 2 61s 2s 2p 3s 3p

2 6[Ne]3s 3p

2 2 6 2 11s 2s 2p 3s 3p

2 1[Ne]3s 3p2 2 61s 2s 2p

[Ne]

2 2 6 2 6 2 10 41s 2s 2p 3s 3p 4s 3d 4p

2 10 4[Ar]4s 3d 4p2 2 6 2 6 2 10 61s 2s 2p 3s 3p 4s 3d 4p

2 10 6[Ar]4s 3d 4p

Excited vs. Ground State

•If an electron absorbs energy, it is in an EXCITED state

Ne: 1s22s22p53s1

•How is this different from the ground state configuration?

Ne: 1s22s22p6