Atomic Theory

Atomic Theory

http://www.skanschools.org/webpages/rallen/

The Evolution of the Atomic Model

Atom Basic building block of matter

Cannot be broken down chemically

A single unit of an element

Dalton (1803)Known as the founder of the atomic theory

Dalton’s Postulates:1. All matter is composed of indivisible particles called atoms2. all atoms of a given element are identical in mass and properties. Atoms of different elements have different masses and different properties3. Compounds are formed by a combination of 2 or more atoms4. Atoms cannot be created, destroyed, or converted into other kinds of atoms during chemical reactions

1. Spherical

2. Uniform Density

Cannonball Theory/Model

J.J. Thomson (1897)

Used a cathode ray tube with charged particle field (+/-)

Cathode ray deflected by negative electrode toward positive electrode

Discovered subatomic particle called the ELECTRON

SmallNegatively charged

J.J Thomson

Plum Pudding Theory/Model

Positive “pudding”

Negative electrons embedded (just like raisin bread)

Rutherford (1909)

Conducted the GOLD FOIL EXPERIMENT where he bombarded a thin piece of gold foil with a positive stream of alpha particles

Expected virtually all alpha particles to pass straight through foil

Most passed through but some were severely deflected

Conclusion led to…

Rutherford’s Experiment

RutherfordNuclear Theory/Model

1. The atom is mostly empty space

2. At the center of the atom there is a DENSE, POSITIVE CORE called the NUCLEUS

**Provided no information about electrons other than the fact that they were located outside the nucleus**

Neils Bohr (1913)

Bohr Model or Planetary Model

1. Electrons travel AROUND the nucleus in well defined paths calls ORBITS (like planets in the solar system)

2. Electrons in different orbits possess different amounts of energy

Neils Bohr 3. Absorbing/Gaining a certain amount of energy causes electrons to jump to a higher energy level or an excited state

When excited electrons emit/lose a certain amount of energy which causes electrons to fall back to a lower energy level or the Ground State

Wave-Mechanical/Cloud Model

Modern, present day model

Electrons have distinct amounts of energy and move in areas called orbitals

Orbital= an area of high probability for finding an electron (not necessarily a circular path)

Developed after the famous discovery that energy can behave as both waves & particles

Many scientists have contribute to this theory using X-Ray diffraction

Vocabulary(of the Periodic

Table)Atomic#= the number of protons in every atom of the element (NEVER CHANGES!)

Atomic Mass= average mass of all the isotopes of an element

Element Symbol= the letter(s) used to identify an element

Subatomic

Particle

Charge

Relative Mass

Location

Symbol How to Calcula

te

Proton +1 1 amu Nucleus

or Look at atomic

#

Neutron 0 1 amu Nucleus

Mass # - Atomic

#

Electron -1 1/1836 amu

Outisde nucleus

P= e(in

neutral atoms)

**Nucleons= Protons and Neutrons (any subatomic particle found within the nucleus)

Determining Subatomic

Particles (p, n, e)# of Protons = sum of the protons and neutrons in an atom of an element (Atomic Number)

# of Neutrons = Mass # - Atomic # or Mass # - # Protons

# of Electrons = sum of the protons and neutrons in an atom of an element ( equal to # protons in a neutral atom)

Determining Subatomic

Particles (p, n, e)Atomic Number:1. Look at the element symbol and locate on the periodic table

2. same as the # of protons or the nuclear charge

Mass Number = # Protons + # Neutrons

Nuclear Charge = # of protons or the atomic #P = nuclear charge

Atoms vs. IonsVocabulary

TermDefinition Example

Neutral Atom An atom with the same number of protons and electrons

P = e (no charge indicated)

Ion Two Types

Anion and Cation

Protons and electrons are not the same

Atoms vs. Ions

Anion Cation

aNion

An atom that has GAINED one or more electrons

e > p

NEGATIVE ION

ca+ion

An atom that has LOST one or more electrons

p > e

POSITIVE ION

Ions

Isotope

Atoms of the same element with different mass numbers; same atomic number, same number of protons, different number of neutrons

Isotopes

p = p = p = e = e = e = n = n = n =

Example 1: Carbon (12, 13, 14)

Isotopes

p = p = e = e = n = n =

Example 2: Uranium (238, 240)

U-240

Practice

1. Two different isotopes of the same element must contain the same number of

2. Two different isotopes of the same element must contain a different number of

a. protons b. neutrons c. electrons*

a. protons b. neutrons c. electrons*

Practice

3. Isotopes of a given element have

a. the same mass number and a different atomic numberb. the same atomic number and a different mass numberc. the same atomic number and the same mass number

*

Calculating Atomic Mass (for any

element)Atomic mass = the weighted average of an element’s naturally occurring isotopes

% abundance of isotope 1 x (mass of isotope 1)

% abundance of isotope 2 x (mass of isotope 2)

+ % abundance of isotope 3 x (mass of isotope 3)

Average Atomic Mass of the Element

Calculating Atomic Mass

Chlorine has two naturally occurring isotopes, Cl-35 (isotopic mass 34.9689 amu) and Cl-37 (isotopic mass 36.9659 amu). In the atmosphere, 32.51% of the chlorine is Cl-37, and 67.49% is Cl-35. What is the atomic mass of atmospheric chlorine?

Step 1: Multiply the mass of each separate isotope by its percent abundance

Example 1: The exact mass of each isotope is given

CL-35 = 34.9689 amu x = CL-37 = 36.9659 amu x =

(.6749)

(.3251) 12.0176

23.6005


Step 2: Add the products of all the calculated isotopes together from step 1

23.6005

+12.0176

35.6181

***This is your average atomic mass***


The element Carbon occurs in nature as two isotopes. Calculate the average atomic mass for Carbon based on the information below

C-12 = 98.89% C-13 = 1.11%

**Since the mass numbers were not given for either isotope, use the mass number

instead**

C-12 = 12

C-13 = 13


Step 1: Multiply the mass by the percent abundance

C-12 = 12 x (.9889) =

C-13 = 13 x (.0111) =

Step 2: Add the products together

11.8668

0.1443

*These are weighted masses*

11.8668+ 0.1443

12.0111


Practice: The element Boron occurs in nature as two isotopes. Calculate the average atomic mass for Boron, using the information below.

Isotope Mass % abundance

Boron-10 10.0130 amu 19.9 %

Boron-11 11.0093 amu 80.1 %

Average Mass of Boron = 10.8104

Practice: The element Hydrogen occurs in nature as three isotopes. Calculate the average atomic mass of Hydrogen.

Isotope % abundance

Protium 99.0%

Deuterium 0.6%

Tritium 0.4%

Average Mass= 1.014

Mass Number Atomic Mass

The MASS of ONE isotope of a given

element

The AVERAGE MASS of ALL isotopes of a

given element

Electron Configurations

The dashed chain of numbers found in the lower left hand corner of an element box

Tells the number of energy levels as well as the number of electrons in each level(how the electrons are arranged around the nucleus)

http://www.lon-capa.org/~mmp/period/electron.htm

Electron Configuration

Is the representation of the arrangement of electrons distributed among the orbitals

Used to describe the orbitals of an atom in the ground state

Can be used to describe ionized atoms (cations and anions)

Many of the chemical and physical properties of elements can be correlated to their electron configuration


All electron configurations on the Periodic Table are NUETRAL ( p = e)

For IONS, add or subtract electrons from the LAST NUMBER in the electron configuration only


Orbitals

There are four types (s, p, d, and f)

They have different shapesEach orbital can hold a maximum of two electrons

P, d and f orbitals have different sublevels


Is unique to an element’s position on the periodic table

Energy level is determined by the period

Number of electrons is given by the atomic number


Orbitals on different energy levels are similar to each other but they occupy different areas in space

ex. 1s and 2s orbitals both have s orbital characteristics ( radial nodes, spherical volume, can only hold two electrons) but since they are found in different energy levels they occupy different spaces around the nucleus


Electrons fill orbitals to minimize energy

Electrons fill the principal energy levels in order of increasing energy

Electrons are getting further from the nucleus

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p


A single orbital can hold a maximum of two electrons

These electrons must have opposing spins

One electron spins up and the other spins down

Ensures they have different quantum numbersPauli Exclusion Principal


S subshell 1 orbital that can hold 2 electrons

P subshell 3 orbitals that can hold up to 6 electrons

D subshell 5 orbitals that can hold up to 10 electrons

F subshell 7 orbitals that can hold up to 14 electrons

Principal Energy Level (n)

Electron energy levels consisting of orbitals which designated s, p, d, or f.


Valence Electrons:

Electrons found in the OUTERMOST shell or orbital

Kernel Electrons:

INNER electrons (all non-valence electrons)

Orbital Notation

Orbital Notation

When filling in the orbital

Electrons fill the lowest vacancy levels first

When there’s more than one subshell at a particular energy level (ex. 3p or 4d) only one electron fills each subshell until each subshell has one electron. Then electrons start pairing in each subshell

Hund’s Rule

Every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied, and all electrons in singly occupied orbitals have the same spin.

Energy Level Diagram: Oxygen.

Oxygen is atomic number 8.

8 protons How many electrons?

8 electrons

Bohr Diagrams

A method for showing electron location in an atom/ion

All electrons must be drawn

Look up the electron configuration and the period (row) the element is in on the periodic table

Draw a circle (nucleus) write in the number of protons and neutrons

Draw the shells (energy levels) around the nucleus

Bohr Models

Add the electrons

The first shell only holds 2 electrons

Now add the rest of the electronsAdd one at a time starting on the right side and going counter clockwise

Shell (energy level) Maximum number of electrons

1 22 83 184 325 506 727 98

2n2 n=Formula for maximum # of electrons per quantum number (or energy level)

Bohr ModelsExample:

Draw the Bohr Model for Carbon

Bohr ModelsExampleDraw the Bohr Model for Oxygen

Bohr ModelsExample

Draw the Bohr Model for Na+

E- configuration 2-8

Lewis Dot Diagrams

Only illustrates valence electrons

Write the element symbol

Look up the electron configuration (use the last number in the configuration- # of valence electrons)

Use either an X or dot to represent the electrons

Place that many electrons around the symbol at (12, 3, 6 and 9)

Lewis Dot Diagram

Example: Carbon e- configuration 2-4

Practice

The number of unpaired electrons is equal to the number of BONDS that an element can form with other elements

When determining the number of bonds an element can form, arrange the valence electrons so that you have the MAXIMUM number UNPAIRED

Example: Carbon

How many bonds can carbon form?4

Ground State vs. Excited State

Ground State electrons are in the lowest energy configuration possible (the configuration found on the periodic table)

Excited State electrons are found in a higher energy configuration (any configuration not listed on the periodic table)

Ground State Excited State

EXCITED

EXCITED

EXCITED

EXCITED

EXCITED

GROUNDGROUND

GROUND

The greater the distance from the nucleus, the greater the energy of the electron

When ground state electrons absorb energy they jump to a higher energy level or an excited state

This is very unstable/temporary condition

Excited electrons fall rapidly to a lower energy level

When excited electrons fall from an excited state to a lower energy level, they release energy in the form of light

Ground ExcitedEnergy is absorbedDark line spectrum is produced

Excited Ground Energy is releasedBright line spectrum is produced

Dark Lines Absorbed

Bright LinesEmitted

Balmer Series: electrons falling from an excited state down to the second (2nd) energy level give off visible light (Bright Line Spectrum or Visible Light Spectrum)

Different elements produce different colors of light or spectra

These spectra are unique for each element Spectral lines are used to identify different elements

*

Gas A and Gas D

Atomic Theory

Documents

Transcript of Atomic Theory