atomic structure part i evolution of the atom...travels at the speed of light. 1. speed of light =...

Atomic Structure Part I

Evolution of the Atom

I. Early Theories of Matter

A. Greek Philosophers

1. Democritus and Leucippus


“Atomos” -indivisible


A. Greek Philosophers

1. Democritus and Leucippus

2. Aristotle

B. Alchemists

C. Lavoisier & Proust

Lavoisier – law of conservation of mass

C. Lavoisier & Proust

2. Proust – law of definite proportions

II Atomic Theory - Dalton

2. Proust – law of definite proportions

John Dalton

Atomic Theory

�All matter is composed of extremely small particles called atoms.

�All atoms of a given element are identical, having the same size, mass, and chemical properties. Atoms of a specific element are different from those of any other element.

Atomic Theory

�Atoms cannot be created, divided into smaller particles, or destroyed.

�Different atoms combine in simple whole-number ratios to form compounds

�In a chemical reaction, atoms are separated, combined or rearranged

Dalton’s Model/Chart

Dalton’s model

III. Modern Atomic Theory

1. Atoms have detailed sub-structure2. Atoms can be changed from one element to another, but not by chemical reactions (nuclear only)3. Atoms of the same element can have different masses, but their average mass is the same.�Periodic table was based on average masses

V. Finding Subatomic Particles

A. Radioactivity - the phenomenon of rays being produced spontaneously by unstable atomic nuclei

B. x-rays & radioactivity

B. x-rays & radioactivity

1. Wilhelm Roentgena. shot cathode rays at metal and radiation was given off that would darken photographic film

b. the rays could pass through, flesh, but not bone to expose photo platec. rays were unknown, so they were given the name "x-rays"

Visibility

2. Becquerel - obtained a sample of uranium ore (pitchblende)a. he noticed uranium also darken photographic paperb. Radioactive Decayc. Transmutation - change of one element into another

Becquerel

3. Marie Curie-noticed polonium and radium also gave off radiationa. radiation of radium is two million times greater than uraniumb. it conducted electricity in air "ionized air"

Curies

3. Marie Curie-1. ionizing radiation has sufficient

energy to change atom and molecules into ions

2. non-ionizing radiation does not, such as radio waves

C. Electrons –JJ Thomson – gas discharge tube

Electron Discovery

C. Electrons –JJ Thomson – plum pudding model

Thomson

3 Millikan’s oil drop animation of oil drop

Millikan

4 Rutherford – gold foil experiment


gold foil experiment

bombarded thin film of foil90% of particles pass with no deflectionatom was not solid, mostly empty space

4 Rutherford – gold foil experiment--some questions arose about why the electrons don't fall into the nucleus (after all, opposite charges should attract)Rutherford said that the motion of the electrons kept them from falling into the nucleus, but...

--according to the laws of physics, charged particle moving in a curved path should give off energy, and if it did, it should fall toward the nucleus


-Niels Bohr proposed a model in which orbiting electrons don't lose energy to understand this, we first have to understand the nature of light and other electromagnetic radiation

D. Subatomic Particle & the Nuclear Atom

1. Nucleuscontains positively charged protons and neutral neutrons

VERY dense! If a nucleus were the size of the dot in the exclamation point at the end of this sentence, its mass would be approximately as much as that of 70 automobiles!

D. Subatomic Particle & the Nuclear Atom2. Protons (discovered by Rutherford)

Carries a charge of 1+, equal but opposite that of an electron

Mass = 1.6726x10-24 g (1800x’s more than electron)

Referred to as atomic number

Found in nucleus

Atomic number = number of protons = number of electrons


3. Neutrons (discovered by Chadwick)

No Charge

Mass = 1.6750x10-24 g

Found in nucleus

Number of neutrons = mass number – atomic number

D. Subatomic Particle & the Nuclear Atom4. Electrons (discovered by Thomson)

Negative charge (1-)

Relative mass is 1/1840 amu (actual mass is 9.115x10-28 g).

Located in the space surrounding the nucleus

number of protons = number of electrons

What is not neutral?

What is a charged atom called? ion


Sodium? Protons?

Na = 11 protons

(+11) + (-10) = +1 charge

Na 1 +

What if it has 10 electrons?

D. Subatomic Particle & the Nuclear Atom5. Atomic number

Refers to the number of protons

Determines the elements position on the periodic table


All atoms of a particular element have the same number of protons and electrons, but neutrons may differ these types of elements are called Isotopes.

6. Isotopes and Mass Number

The sum of the number of protons and neutron is the Mass Number.

Mass number = protons + neutrons


The atomic mass of an element is the weighted average mass of the isotopes of that element.



H = 1 amu

Relative masses, assigned by Dalton

O = 16 amu



20

Mass #

Atomic #

19 K39

19 K40

19 K41

Neutrons? 21 22

93.25% = ?g 5.73025% = ?g 0.127% = ?g

Abundance of isotope?


93.25% = ?g36.3675 g

5.73025% = ?g2.2921g

0.127% = ?g.05207g

36.3675 g + 2.67873g + .05207g =

39 40 41

Electrons in AtomsI. Properties of Waves1. Definition:

Energy that exhibits wave-like (or oscillating) behavior as it travels through space

Electrons in Atoms2. Wavelength ( λλλλ) distance from peak

to peak, length of one complete wave

3. Frequency ( f)

a. number of peaks that pass at a given point each sec

b. can be called cycles per

second (peak/sec)

c. cps now called 1 Hertz (Hz)

Electrons in Atoms - Cont.4. Velocity (C = speed of light)a. distance a given peak moves in

a unit of timeb. velocity (m/s) = frequency x

wavelength

c = f x λλλλ

II. Behavior of LightA. Newton (1600) thought light consisted of particles

(beam of light is a stream of particles)

B. Maxwell (1864) thought light was a wave phenomenon.

•Calculated the velocity of the propagation of an electromagnetic wave and found it was the same for light

II. Behavior of Light1. some say light is like waves, some say

its like particles

2. modern theory says that it behaves as both "wave/particle duality"

II. Behavior of Light3. Max Planck (early 1900's) said:a. light is made up of bundles of energy called

photons (or quanta)b. the energy of each photon is proportional to

the frequency of the light (Quantum Theory)Each quantum has a specific amount of energy

• example: CONTINUOUS SPECTRUM

*** when white light is passed through a prism, it is separated into a band of colors from red � violet. It's called a continuous spectrum

c. the work of Planck & Einstein led to

E=energy,

v or f= frequency,

h=planks constants (6.6262x10-34J/sec)

E = h x νννν

III. Bright line spectrum• A. a spectrum that shows separate bright

lines, each with a specific wavelength• B. bright-line spectra occur when an

element is heated and the colored light given off is viewed through a spectroscope. Each element has a unique set of lines, characteristic of that element (like a fingerprint)

Line-Emission Spectrum

ground state

excited state

ENERGY IN PHOTON OUT

Fireworks? Hmmm…

IV. Electromagnetic Spectrum• A. visible light (like the continuous spectrum)

is only one type of radiation. All other types are not visible to the human eye.

HIGH

ENERGY

LOW

ENERGY

Electromagnetic Spectrum

LOW

ENERGY

HIGH

ENERGY

R O Y G. B I V

red orange yellow green blue indigo violet

Electromagnetic Spectrum

B. all forms of electromagnetic radiation travels at the speed of light.

1. speed of light = 3.00 x 10 8meters/sec

2. use formula:c = v x λλλλ

3. each line spectrum has a particular frequency (v). If know wavelength ( λ), we can find v using c as a constant.

C. The energy in a photon of light is directly proportional to the frequency of the light.

• 1. frequency, energy• 2. can find the energy of a single

quantum (photon) of radiation at any given frequency.

C. The energy in a photon of light is directly proportional to the frequency of the light.

• 3. proportionality constant that relates the two is called Planck's constant ( h).

• 4. formula:

E = h x f

example: a spectral line has frequency of 3.5x10 12 hertz. What is the energy of a photon of radiation of this frequency?

E = h x f (h=6.6262x10-34J/sec)

E = (3.5x1012Hz) (6.6262x10-34J • sec )

E = (2.3x10-21J)

V. Electron energy levels in Bohr's Model

A. There are certain different orbits in which an electron can travel around a nucleus.

1. each circular orbit (or shell) is at a fixed distance from the nucleus

V. Electron energy levels in Bohr's Model

2. the greater the radius of that shell, the greater the energy of the electron in that shell.

3. these electron orbits are known as energy levels

B. When electrons absorb energy firm an outside source, they jump from lower to higher energy levels.�when they fall back to their original levels , energy is emitted (light); the same amount as was absorbed.

1

23

456 �Energy of photon

depends on the difference in energy levels

�Bohr’s calculated energies matched the IR, visible, and UV lines for the H atom

C. In energy atom in its normal state, all electrons are in the lowest energy levels available (energetically stable)

VI. Atoms and Radiation

• A. When all of the lowest energy levels are occupied, the atom is in the ground state (unexcited).


• B. When electron moves to higher energy level, atom is in the excited state, and is energetically unstable.


C. Bright line spectrum of an element represents the energy levels in its atoms.

�problems with Bohr's Model:

�� only explained some of the lines in the bright line spectrum

�� really only worked for hydrogen�� need sublevels and electron cloud

model to account for all of the lines.

VII. The Modern Model of the Atom

A. Mechanics1. Classical Mechanics - Newton's Laws of

Motion (Newtonian Mechanics)

Describes the behavior of visible objects traveling at ordinary velocities.Bohr’s basis for his model, but couldn’t explain why electrons would stay at on energy level or another. When looking at H-spectral lines, noticed more one (several closely spaced).

VII. The Modern Model of the Atom2. Quantum Mechanics – (wave

mechanics)

Describes the behavior of extremely small particles traveling at velocities at or near the speed of light

a. Louis de Broglie - particles could have properties of waves

Planks quanta gave wave properties, deBroglie said electron streams are like waves of light and have properties of both particles and waves (matter behaves as waves)

•b. Schrodinger - described the behavior of electrons in terms of quantized energy changes "quantum mechanics"

Describe a wave equation used to determine the probability of finding an electron in any give place or orbital

Schrodinger’sCat

Radial Distribution CurveOrbital

c. Heisenberg - uncertainty principle

�Region of space where there is a probability of finding an electron is called an orbital

"The more precisely the POSITION is determined,the less preciselythe MOMENTUM is known"

B. Principal Energy Levels1. Energy Levels• Bohr -

High Energy(outer level)

Low Energy

1234

Principal Quantum Numbers (N) Number of electrons

2818

32

Corresponds to energy level

2. SublevelsPrincipal Quantum Numbers (N) Sublevel

Present

1 1s2 2s2p3 3s3p3d4 4s4p4d4f

• Orbital -

– Region of space where an electron is probably found

• Electron spin– An orbital can hold 2 electrons that spin in

opposite directions.

Electrons are represented by arrows

Energy Level Diagram

orbital

Place where electrons are probably found

Electrons have “up” and “down” spin

c

Aufbau PrincipleElectrons fill the

lowest energy orbitals first.

Hund’s Rule–Within a sublevel, place one e-

per orbital before pairing them.–“Empty Bus Seat Rule”

RIGHTWRONG

Rules1. No more than two electrons in

any orbital2. Electrons to be added must be

placed in unfilled orbitals of lowest energy for stable configuration

3. In sublevel second electron can't be added until each orbital in sublevel contains one electron (Hund's rule)

• Pauli Exclusion Principle

–Each orbital can hold TWO

electrons with opposite

spins.

–No two electrons in an atom can have the same 4 quantum numbers.

–Each e - has a unique “address”:

Rules1. No more than two electrons in

any orbital2. Electrons to be added must be

placed in unfilled orbitals of lowest energy for stable configuration

3. In sublevel second electron can't be added until each orbital in sublevel contains one electron (Hund's rule)

s

The s orbital

py

px

The p orbitals

pz

The d The d orbitalsorbitals

The f The f orbitalsorbitals

f

Click here for orbital viewer

C. Electron Configurations

1s2 = Helium1s22s1 = Lithium1s22s22p63s23p6

4s23d104p6

= Krypton

Energy Level

Sub Level # of Electrons

O8e-

• Orbital Diagram

• Electron Configuration

1s2 2s2 2p4

Electron Configuration Notation

1s 2s 2p

• Shorthand Configuration

S 16e-

Valence ElectronsCore Electrons

S 16e- [Ne] 3s2 3p4

1s2 2s2 2p6 3s2 3p4

Notation

• Longhand Configuration

Periodic Patterns

• Period #– energy level (subtract for d & f)

• A/B Group # – total # of valence e -

• Column within sublevel block– # of e - in sublevel

s-block1st Period

1s11st column of s-block

1 2 3 4 5 6 7

Periodic Patterns• Example - Hydrogen

1

2

3

4

5

6

7

Periodic Patterns

• Shorthand Configuration– Core e-: Go up one row and over to the

Noble Gas.– Valence e-: On the next row, fill in the #

of e - in each sublevel.

Try a few!Mg = Fe = Ru = Ir = Ca+2 =

Cl-1 =

1 2 3 4 5 6 7

Stability• Ion Formation

– Atoms gain or lose electrons to become more stable.

– Isoelectronic with the Noble Gases.

Feeling overwhelmed?

[Ar]

1 2 3 4 5 6 7

4s2 3d10 4p2

Periodic Patterns

• Example - Germanium

• Full energy level

1 2 3 4 5 6 7

• Full sublevel (s, p, d, f)

• Half-full sublevel

Stability

atomic structure part i evolution of the atom...travels at the speed of light. 1. speed of light =...

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