UNIT 2 - ATOMIC THEORY

VOCABULARY:

Allotrope Electron Configuration Nuclear Charge

Anion Element Nucleons

Atom Excited state Nucleus

Atomic Mass Ground state Orbital

Atomic Mass unit (a.m.u.) Ion Proton

Atomic number Isotope Quantum Theory

Bohr model Kernel electron(s) Valence electron(s)

Cation Lewis Dot Diagram Wave-mechanical model

Compound Mass number

Electron Neutron

OBJECTIVES: Upon completion of the unit you will be able to do the following:

Discuss the evolution of the atomic model

Relate experimental evidence to models of the atom

Identify the subatomic particles of an atom (electron, proton, and

neutron)

Know the properties (mass, location, and charge) of subatomic

particles Determine the number of protons, electrons, and neutrons in a neutral atom

and an ion

Calculate the mass number and average atomic weight of an atom

Differentiate between an anion and a cation Distinguish between ground and excited state

Identify and define isotopes

Write electron configurations

Generate Bohr diagrams Differentiate between kernel and valence electrons

Draw Lewis Dot Diagrams for an element or an ion

THE EVOLUTION OF THE ATOMIC MODEL

Atom 1.) basic building block of matter 2.) cannot be broken down chemically 3.) a single unit of an element

Dalton (cannonball model)

FOUNDER of the atomic theory

Dalton’s Postulates:

1. All matter is composed of indivisible particles called atoms

2. All atoms of a given element are identical in mass and

properties. Atoms of different elements have

different masses and different properties 3. Compounds are formed by a combination of 2 or more atoms

4. Atoms cannot be created, destroyed, or converted into other

kinds of atoms during chemical reactions

CANNONBALL MODEL:

SPHERICAL

UNIFORM DENSITY

J.J.Thomson (plum pudding model) EXPERIMENT:

Used a CATHODE RAY TUBE with charged electrical field (+/-)

o Cathode ray deflected BY NEGATIVE electrode

TOWARD POSITIVE electrode

Discovered SUBATOMIC PARTICLE called the ELECTRON:

1. SMALL

2. NEGATIVELY CHARGED

PLUM PUDDING MODEL:

POSITIVE “PUDDING”

NEGATIVE ELECTRONS EMBEDDED (just like raisin

bread)

Rutherford (nuclear model)

EXPERIMENT:

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 (see diagram

above)

NUCLEAR MODEL:

The ATOM is MOSTLY EMPTY SPACE

At the center of the atom 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 (BOHR MODEL)

BOHR MODEL or PLANETARY MODEL:

Electrons travel AROUND the nucleus in well-defined paths called ORBITS

(like planets in a solar system) Electrons in DIFFERENT ORBITS possess DIFFERENT AMOUNTS OF ENERGY

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

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

o 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 contributed to this theory

VOCABULARY (of the Periodic Table)

Atomic Mass = AVG.

mass of all the

isotopes of an element

Atomic # = the number

of protons in EVERY

ATOM of the element

Electron configuration = number of

electrons in each energy level; add the

numbers to get the total # of electrons

Oxidation #’s = possible

charges an atom of an

element can have

Element Symbol = the

letter(s) used to

identify an element

SUBATOMIC PARTICLES

Subatomic Charge Relative Mass Location Symbol How to Calculate

Particle

p

Proton +1

1 amu

Nucleus

----------

Look at the atomic #

1

p

1

n

Neutron 0

1 amu

Nucleus

----------

Mass # - Atomic #

1

n

0

Electron -1 1/1836th

amu Outside e- or e P = e- in neutral atom

(negligible) nucleus

DETERMINING SUBATOMIC PARTICLES (p, n, e) -

DIRECTIONS: Use the information below to complete the chart that follows.

Atomic number = the # of protons in an atom of an element (p)

Mass # = sum of the protons and neutrons in an atom of an element

(p + n) = mass #

Nuclear Charge=charge w/in the nucleus; = to the # of protons or the atomic #

nuclear charge=# of p

Nucleons= any subatomic particles found w/in the nucleus

protons and neutron

Symbol # # # Atomic Mass Nuclear Nuclear

Protons Neutrons Electrons Number Number Charge Symbol

35

Cl-35 17 18 17 17 35 17 Cl 17

15 16

C-14 8

16 8 O 8

Ar-40 18 22

20 40

ATOMS (neutral) VS. IONS (charged)

Vocabulary Term Definition Example/Diagram

Neutral Atom An atom with the same

12

number of protons and

electrons 6

C

p = e

or

C

(no charge indicated)

Ion aNion

19

F

-1

An atom that has 9

GAINED one or more

electrons

the -1 indicates a

e > p negative charge with a

magnitude of one

Ca+ion

7 +1

Li

An atom that has LOST

3

one or more electrons

the +1 indicates a

p > e positive charge with a

magnitude of one

ISOTOPE = atoms of the same element with different mass #’s but the same atomic #;

(same number of protons, different number of neutrons)

p = p = p =

n = n = n =

e = e = e =

Example 2: Isotopes of Uranium (U-238, U-240)

p = p =

n = n =

e = e =

Calculating Atomic Mass (for any element):

Atomic Mass = the weighted average of ALL the element’s naturally occurring

isotopes

(% abundance of isotope in decimal form) x (mass of isotope 1)

+(% abundance of isotope in decimal form) x (mass of isotope 2)

+ (% abundance of isotope in decimal form) x (mass of isotope 3)

Average Atomic Mass of the Element

Example 1: The exact mass of each isotope is GIVEN to you.

Chlorine has two naturally occuring 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 in DECIMAL FORM (move the decimal 2 places to the left)

Cl-35 = 34.9689 x (.6749) = 23.6005

Cl-37 = 36.9659 x (.3251) = 12.0176

** These are the weighted masses **

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

23.6005

+ 12.0176

35.6181

** This is your average atomic mass **

Example 2: The exact mass for each isotope is NOT GIVEN to you.

The element Carbon occurs in nature as two isotopes. Calculate the average atomic mass

for Carbon based on the information for the isotopes below.

12C = 98.89%

13C = 1.11%

**Since the exact masses were NOT given for either of the isotopes, just use the

mass number instead. For 12

C, the mass would be 12 amu, and for 13

C, the mass would

be 13.

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

FORM (move the decimal 2 places to the left)

12

C = 12 .9889) = 11.8668

13C = 13 x (.0111) = 0.1443

** These are the weighted masses **

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

11.8668 +

0.1443

12.01

** This is your average atomic mass **

Practice 1: The element Boron occurs in nature as two isotopes. In the space

below, calculate the average atomic mass for Boron.

Isotope mass percent abundance

Boron-10 10.0130 amu 19.9%

Boron-11 11.0093 amu 80.1%

Practice 2: The element Hydrogen occurs in nature as three isotopes. In the

space below alculate the average atomic mass for Boron.

Isotope of Hydrogen Percent Abundance

1

H Protium 99.0%

1

2

H Deuterium 0.6%

1

3

H Tritium 0.4%

1

Challenge Practice 3: Chlorine has two naturally occuring isotopes, Cl-35 (isotopic

mass 34.9689 amu) and Cl-37 (isotopic mass 36.9659 amu). If chlorine has an atomic

mass of 35.4527 amu, what is the percent abundance of each isotope?

Mass Number Atomic Mass

The MASS of ONE ISOTOPE of a given The AVERAGE MASS of ALL ISOTOPES

element. of a given element

ELECTRON CONFIGURATIONS = a dashed chain of numbers found

in the LOWER LEFT CORNER of an element box (see below); tells us the number of ENERGY LEVELS as well as the number of ELECTRONS in each level (tells us how the electrons are arranged around the nucleus)

Electron configuration

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

SUBSTANCE ELECTRON CONFIGURATION

Magnesium 2-8-2

Mg+2

2-8

Bromine

Br-1

Barium

*Lead

* shortcut allows you to cut out the first two orbitals to shorten the “address”

Valence Electrons: electrons found in the OUTERMOST shell or orbital

the LAST number in the electron configuration

Kernel Electrons: INNER electrons (all non-valence electrons)

Ex: Practice 1: Practice 2:

Chlorine Nitrogen Sodium

# valence e- = 7 # valence e

- = # valence e

- =

# kernel e- = 10 # kernel e

- = # kernel e

- =

Principle Energy Level (n) = electron energy levels that contain a certain number of SUBLEVELS (s, p, d, f); each sublevel contains one a set number of ORBITALS

Maximum # of electrons in an energy level = 2n2 where n = quantum # (or period #)

Principle Energy Level (n) Maximum number of electrons (2n2)

1

2

3

4

BOHR DIAGRAMS (one method for expressing electron location in an atom or

ion); ALL ELECTRONS MUST be drawn BOHRing

1. Look up electron configuration of element at hand on

Periodic Table (if you are working with an ion,

add/subtract the proper amount of electrons from outer

shell(s) of configuration)

Example: Oxygen is 2-6

2. Draw nucleus (with a square) and notate correct amount of

protons and neutrons inside

P = 8

N = 8

3. Using rings or shells, place the proper number of ORBITS

around your nucleus.

4. Use either and “x” or a dot to represent your electrons, place the correct number of electrons in the area that would correspond to the number 12 on the face of a clock in the ORBIT CLOSEST TO THE NUCLEUS ONLY. Remember, you can have a maximum of 2

e-in the first orbit.

x x

P = 8

N = 8

5. Place one “x” or one dot at a time around your 2nd

or valence

orbit in the areas that would correspond to the numbers 12, 3, 6, and 9 on the face of a clock.

x

x x

x P = 8

x

N = 8

6. If there are any electrons remaining in the configuration, pair

them up with electrons you have already placed. You may have

no more than 2 e- in any of the 4 spots, and no more than 8 e

-

total in the 2nd

or valence orbit. In this case, we have 2 e- left

to place.

x

P = 8 x

N = 8 x

Carbon Fluorine Beryllium Al

Li Ca2+

Na+ S2-

LEWIS (ELECTRON) DOT DIAGRAMS (Only illustrates VALENCE ELECTRON CONFIGURATION)

1. Write the element’s symbol

2. Retrieve electron configuration from Periodic Table. The last number in the

configuration is the NUMBER OF VALENCE ELECTRONS

3. Arrange the valence electrons (DOTS) around the symbol using the following rules:

Only two electrons maximum per side of the symbol (therefore no

more than 8 total surrounding symbol – 8 is great!)

Always “pair” the first two If you have more than 2 valence electrons, deal them one at a time to

the other three sides until you run out

8 1 5 2 OR

8 1 2

3

4 X 6 5 X 6

3 7 4 7

Ex: Using a dot or an “x” place the valence electrons around

the symbols for carbon below in order according to each of

the models above.

C

OR

C

4. If you are working with an ION you must adjust the valence

electrons (add or subtract electrons) in the configuration

before constructing your Dot Diagram. Your final diagram

must include brackets and the charge on the ion.

Ex 1: S-2

ADD 2 e- to the 6 that S

normally has in its valence shell.

Ex 2: K+1

REMOVE 1 e- from the

valence shell of K.

*NEGATIVE IONS always end

up with EIGHT VALENCE e-

*POSITIVE IONS always end

up with ZERO VALENCE e-

QUICK NOTE:

The number of UNPAIRED VALENCE 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.

Ex: Carbon can be arranged in either of the two ways

shown on the previous page. Draw the Lewis dot

structure for carbon that gives it the maximum

number of UNPAIRED valence electrons.

How many bonds can an atom of carbon form? ____

The alternate Lewis structure for carbon with only 2

unpaired valence electrons represents the LOWEST

ENERGY STATE.

Ground State vs. Excited State

*Notice that one electron from the 2nd

orbital has moved to the 3rd

orbital

Ground State = electrons in LOWEST ENERGY CONFIGURATION possible (the configuration FOUND ON PERIODIC TABLE)

ground state electron configuration for Li is 2-1

Excited State = electrons are FOUND IN A HIGHER ENERGY CONFIGURATION (ANY configuration NOT FOUND ON PT)

excited state electron configuration for Li could be 1-2, 1-1-1

Distinguish between ground state and excited state electron configurations

below:

Bohr Electron Configuration Ground (G) or Excited (E) state?

2-1 Ground

2-0-1 Excited

1-1-1

2-7-3

2-8-2

2-8-8-2

2-8-17-6

2-8-18-8

2-6-18-1

2-5-18-32

***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 a very UNSTABLE/TEMPORARY condition

EXCITED ELECTRONS rapidly FALL BACK DOWN or DROP to a

LOWER energy level

When excited electrons fall from an excited state to lower

energy level, they release energy in the form of LIGHT.

GROUND EXCITED

Energy is ABSORBED

DARK-LINE SPECTRUM is produced

EXCITED GROUND

Energy is RELEASED

BRIGHT-LINE SPECTRUM is produced

DARK LINES show the specific

wavelengths of light being ABSORBED

by the electrons (becoming excited)

BRIGHT LINES show the specific

wavelengths of light being EMITTED

by the electrons (falling down)

Balmer Series: electrons falling from an EXCITED STATE down

to the 2ND

ENERGY LEVEL give off VISIBLE light (AKA “Bright Line SPECTRUM” or “Visible Light SPECTRUM”)

Different elements produce different colors of light or

SPECTRA.

These spectra are UNIQUE for each element (just like a

human fingerprint is unique to each person).

We use spectral lines to identify different elements.