UNIT 2 - ATOMIC STRUCTURE

ESSENTIAL QUESTIONS: What components make up an atom? How are atoms of one element different from atoms of another element? What happens when electrons in atoms absorb or release energy?

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

1.THE EVOLUTION OF THE ATOMIC MODELAtom 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 disti

nct 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

2.3.ATOMIC STRUCTURE

Atom – The smallest particle of an element that retains the chemical identity of an element; made up of protons, neutrons, and electrons4.SUBATOMIC PARTICLES

Subatomic Charge Relative Mass Location Symbol How to CalculateParticle

Proton +1

1 amu (atomic mass units)or 1 u nucleus

p=atomic numberp=e- (for a neutral atom)p

Neutron 0 1 amu nucleusn n=atomic mass-p

n=atomic mass-atomic number

Electron -1 1/1836 amuOutside nucleus e-

p=e- (for a neutral atom)

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

6.DETERMINING SUBATOMIC PARTICLES (p, n, e) .

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 nucleusprotons and neutron

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

Symbol # # # Atomic Mass Nuclear NuclearProtons Neutrons Electrons Number Number Charge Symbol

35Cl-35 17 18 17 17 35 17 Cl

17

P-31 15 16 15 15 31 15

C-14 6 8 6 6 146

O-16 8 168 8 8 16 8 O

8

Ar-40 18 22 18 1840 18

Ca-40 20 20 20 20 40 20

7.ATOMS (neutral) VS. IONS (charged)

Vocabulary Term Definition Example/Diagram

Neutral Atom An atom with the same12number of protons and

electrons6

Cp = e

or

C(no charge indicated)

Ion aNion 19F

-1

An atom that has 9GAINED one or moreelectrons

the -1 indicates ae > p negative charge with a

magnitude of one

Ca+ion 7+1

LiAn atom that has LOST 3

one or more electronsthe +1 indicates a

p > e positive charge with amagnitude of one

8.ISOTOPE = atoms of the same element with different mass

#’s but the same atomic #; (same number of protons, different number of neutrons)

p =6 p =6 p =6

n =6 n =7 n =8

e =6 e =6 e =6

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) = 23.6005Cl-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+ 12.017635.6181

** 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.

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

12C = 12 (.9889) = 11.866813C = 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 **

Isotope mass percent abundanceBoron-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 calculate the average atomic mass for Boron.

Isotope of Hydrogen Percent Abundance1H Protium 99.0%

12H Deuterium 0.6%

13H Tritium 0.4%

1

Practice 3: Chlorine has two naturally occurring 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 givenThe AVERAGE MASS of ALL NATURALLY OCCURING ISOTOPES

element. of a given element

9.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 CONFIGURATIONMagnesium 2-8-2Mg+2 2-8BromineBr-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

10. 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 = 8N = 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 = 8N = 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.

xx x

x P = 8 xN = 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.

xP = 8 xN = 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 2OR

8 1 2 34 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 C4. 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-

11. 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 Excited1-1-12-7-32-8-2

2-8-8-22-8-17-62-8-18-82-6-18-12-5-18-32

***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.