Atomic Structure

Atomic Structure-The BIG Picture

Discovery of the components of the atom and subsequent modeling of the atomic structure led to explosive advances in

chemistry, medicine, and energyATOMIC STRUCTUREChemistry

Medicine

Energy• The nature of the chemical bond• New molecule synthesis• Predictions about reactivity• Information about how reactions work• Electronics / computer development• New analytical (measuring) methods • Emergence of the field of Nuclear Chemistry

• Isotope tracers• New drugs• Cancer treatments• New cell screening methods

• Nuclear fission• Nuclear fusion• Power plants• Understanding of the nature of the sun, planets, stars, etc• Weapons

A History Lesson• Not only does history help you

become an educated person, it helps you understand current theories if you see how they developed…

A History Lesson• When you learn this stuff, try to put

yourself in the time of the person making the discovery. – What was it like back then? – How did society influence thinking? – Would you have helped to make this

discovery?– Do you think another problem of that

day was more important?

A History Lesson• People were thinking

about the ATOM at a time when indoor plumbing and electricity were not available, priorities to work on, or imminently possible!

A History Lesson

1. Thousands of years ago, a Greek philosopher Democritus speculated if a piece of matter were divided in half enough times, you’d finally find the smallest piece that could not be sub-divided any further

A History Lesson

a. This little piece of matter would have the same properties as the big piece.

b. It was called the ATOM. c. He also believed:

– matter could not be created, destroyed, or further divided

– matter is mostly empty space

A History Lesson

• So what made the atom of one type of matter different than another?

• The Greek philosophers thought SHAPE (geometry) was the key difference.

• The Greeks were great practitioners of Geometry…

• I wonder if THAT influenced their thinking on the ATOM?

A History Lesson

Let’s mark this for laterDemocritus: Shape

A History Lesson- 1800 years later…

2. Dalton’s (1766-1844) experimentation on matter led him to believe:a. All atoms were spherical in shape

but differed from one another by mass. • Mass was a big deal in Dalton’s time.• Maybe it influenced his thinking?

A History Lesson2. Dalton’s (1766-1844) experimentation on matter

led him to believe: b. All matter is composed of atomsc. All the atoms for a given element were

identical*. Atoms of a specific element are different from atoms of another element.

d. Atoms cannot be created, divided* or destroyed.

e. Atoms could combine to make compounds only in whole number ratios

f. In a chemical reaction, atoms are separated, combined or rearranged.

*later shown to be incorrect (Make sure you marked them!)

A History Lesson• These statements have come to be

known as Dalton’s Atomic Theory.• Dalton was right about the MASS part

of his theory and how they combine to make compounds

*Let’s mark this for later: Dalton: mass

A History Lesson

• Near the turn of the 20th century, evidence began to emerge that suggested charged subatomic parts made up the atom…

A History Lesson

3. Thomson’s (1897) experimentation led to his discovery of a negatively-charged subatomic particle!a. The negatively charged particle is

called the electron.b. Discovered while studying electricity

with a cathode ray tube

Cathode Ray Tube (CRT)

A History Lesson

• Thomson inferred from his data that the ATOM must be composed of a combination of (+) charged “matrix” with (–) charged particles (electrons) dispersed in it

• like raisins or plums in a puddingc. Developed the “plum pudding”

model of the atom

Thomson’s “Plum Pudding” Model• He discovered the electron (link) in 1897 before the nucleus was discovered

• Later discoveries invalidated this model

or

Progression of the Atomic Model…Discovery of the electron

J.J. Thomson in Philosophical Magazine, 1904

“... the atoms of the elements consist of a number of negatively electrified

corpuscles enclosed in a sphere of uniform positive

electrification, ... “

“Plums”

“Pudding”

• By the way, the cathode ray tube, which Thomson used to generate and study electrons evolved into the television set (or at least became the central component in TVs).

Lets mark this for later: Thomson: electron

A History Lesson- Refining the Atom

4. Millikan (1909) used Thomson’s work to determine:

a. the electron charge (-)b. the mass of an electron is 9.11x10-28

gramsc. his experiment was the oil drop

experiment

Millikan’s Oil Drop Experiment

Millikan later determined the mass and charge of the electron

A History Lesson- Refining the Atomic Model

5. Rutherford (1911) proved:a. The (+) in an atom was not spread out

but concentrated in a central location called the nucleus.

b. The volume of an atom is mostly empty space!

Gold Foil Experiment

Gold Foil Experiment: The Explanation

A History Lesson- Refining the Atomic Model

c. Rutherford’s gold foil experiment:i. He bombarded thin, gold foil with heavy,

positively (+) charged He (alpha) particles and most passed through the foil.

ii. Occasionally, the particle would bounce back (as if it were a tennis ball hitting a brick wall!)

iii. Rutherford assumed the deflected particles hit a dense, gold nucleus and could not pass because of the large mass of the gold nucleus and its (+) charge

Plum Pudding

vs.

Gold Foil Experim

ent

Gold foil animation

Rutherford’s famous gold foil experiment:• showed that the positive charge of the atom MUST be concentrated in a tiny, yet heavy volume he called the nucleus• almost ALL of the mass of the atom is in the nucleus• very light electrons surround this nucleus• the volume that an atom occupies is mostly empty space

If a nucleus were as big as you are wide, the edge of its atom (outermost electron orbital) would be over a mile away!

.About 1.25 miles

A History Lesson- Refining the Atomic Model

• *Let’s mark this for later: Rutherford: proton, nucleus

Click

A History Lesson- Refining the Atomic Model

6. Bohr (1913)- suggested electrons must move around in well-defined orbits or energy levelsa. His experiments suggested that

electrons reside at different energy levels because it took more (or less) energy to knock them loose from an atom

*Lets mark this for later: Bohr: planetary orbit of the electrons around the nucleus

7. Chadwick (1932) discovered the neutrona. The neutron is a particle with no

charge but about the same mass as a proton.

A History Lesson- Refining the Atomic Model

Chadwick’s experiment

• This history lesson gets us closer to the modern atomic model.

• We still need to understand how the electrons behave.

• BUT we’ll do that in the next unit…

II. The Subatomic ParticlesA. Protons (p+)

1.Positively-charged subatomic particle

2.Contained in the nucleus3.Confirmed by Rutherford4.Mass: 1.67 x 10-24 g (1840x

more massive than an electron!)

II. The Subatomic Particles

B. Neutron (n0)1.Not charged subatomic

particle2.Contained in the nucleus3.Discovered by Chadwick4.Mass: 1.67 x 10-24 g

(same as proton)

II. The Subatomic ParticlesC. Electrons (e-)

1.Negatively charged subatomic particle (the charge is equal and opposite to the charge of the proton)

2.Surrounding the nucleus3. Mass: 9.11 x 10-28 gram4. Tiny mass but occupies the

majority of the volume of the atom

II. The Subatomic ParticlesC. Electrons (e-)

5.Each electron has an “electronic address”• Each resides in a well defined

energy level some distance from the nucleus. The further from the nucleus, the higher the energy level.

6.Responsible for chemical bonding

III. Current Model of the Atom

• Spherically-shaped • Small, dense positively-charged

nucleus surrounded by a cloud of negatively-charged electrons

• Most of the atom is empty space• >99% of mass is in the nucleus• Very small (there are 6.5 x1021 atoms

in a drop of water)• Nucleus is held together by strong

nuclear forces

• THE STRONG NUCLEAR FORCE is the name given to the attractive force that holds protons and neutrons together in the nucleus.

• If you think it is weird that like-charged protons can get together in an atom’s nucleus without flying apart, I don’t blame you!

• Don’t worry about the theory of how this force works.

• Just remember, the STRONG NUCLEAR FORCE is able to overcome like-charge repulsion and hence atomic nuclei are quite stable.

valence shellRESPONSIBLEFOR CHEMICALBONDING

core electronsRESPONSIBLE FOR NUCLEARSHIELDING

NucleusPROTONSNEUTRONSRESPONSIBLEFOR MASS,IDENTITY OF THEELEMENT

Here’s the atom so far:

IV. The Subatomic Particles

A. Atomic Number1. Protons are responsible for the

nuclear charge2. # protons = atomic number

++

++++

++

++

NUCLEUS

++PROTONS:RESPONSIBLE FOR NUCLEARCHARGE & = ATOMIC NUMBER

IV. The Subatomic Particles

3. This is the big (or top) number shown on the periodic table

4. # of protons identifies the elementIf for some reason the number of protons changes (like a nuclear reaction), the ELEMENT CHANGES!

++

++++

++

++

++

NUCLEUS

PROTONS:RESPONSIBLE FOR NUCLEARCHARGE & = ATOMIC NUMBER

NEUTRONS

“STRONG NUCLEAR FORCE” HOLDS THE NUCLEUS TOGETHER. IT OVERCOMES THEREPULSIVE FORCE OF “LIKE CHARGES” (REMEMBER COULOMB’S LAW)

THE NUCLEUS

Practice1. Determine the number of protons in:

a. Fluorineb. Magnesium

2. Identify the element:a. 30 protonsb. 17 protonsc. 11 p+

d. 1 p+

IV. The Subatomic Particles

912

ZincChlorineSodiumHydrogen

• Remember, an atom is the smallest particle of an element that retains the identity of that element.

• When the number of protons (+) and the number of electrons (-) are the same, the atom is neutral (has no net charge).

• An ion is a charged atom or group of atoms bonded together.

• An ion can be positive or negative.

IV. The Subatomic Particles

IONS ARE CHARGED ATOMS.IF AN ATOM GAINS ELECTRONS SO THAT IT HAS MORE

ELECTRONS THAN PROTONS, IT IS A NEGATIVELY CHARGED ATOM CALLED ANANION

IF AN ATOM LOSES ELECTRONS SO THAT IT HAS FEWER ELECTRONS THAN PROTONS, IT IS A POSITIVELY CHARGED ATOM

CALLED AN

Ca+ionelectron

Atom about to become a cation Atom about to

become an anion

electron

a negative ion

-

-

-

-

Neutral BerylliumNeutral Beryllium

+

+

+

+

Beryllium IonBeryllium Ion

BeBe2+2+

-

-

-

-

-

-

-

Neutral NitrogenNeutral Nitrogen +

+

+

+

+

+

+

Nitrogen Ion Nitrogen Ion “Nitride Ion”“Nitride Ion”

-

-

-

NN3-3-

For a neutral atom: p+ = e-

Charge of an atom or ion: p+ - e- = Charge of Ion

Atoms are always neutralIons have a charge

ExampleWhat are the charges of:Lithium has 3 p+ and 3 e-

Lithium has 3 p+ and 2 e-

Oxygen has 8 p+ and 8 e-

Oxygen has 8 p+ and 10 e-

3 – 2 = +1

8 – 10 = -2

3 – 3 = 0

8 – 8 = 0

Charges = Oxidation Numbers

IV. The Subatomic ParticlesB.Mass Number

1.Both neutrons and protons are responsible for nearly all the atom’s mass

2.# protons + # neutrons = mass numberEx: An oxygen atom has 8 protons and

8 neutrons and has a mass number = __16

Practice with Atomic number and Mass number

# of protons = atomic number

The atomic number of carbonis ___.

Number of electrons will equal the number of protons for an atom with NO NET CHARGE

# of protons + # neutrons = mass number

A carbon atom with 6 protons and 6 neutronshas a mass number = ____

6

12

Ex 1:

How many p+, e- and n0 are in an atom of Neon with a mass # 22?

• Neon’s atomic number is 10 ___ p+

• Mass number = protons + neutrons22 = 10 +

neutrons12 = neutrons

• If this atom is electrically neutral, protons = electrons ___ e-

10

10

Ex 2:

Determine which element has a mass # of 23 and contains 12 n0.

Mass number = protons + neutrons 23 = protons + 12

11 = protonsp+ = atomic number the element with 11 protons is

_______________Na (sodium)

Practice

atomic #

mass # # of p+ # of no # of e- Atomic mass

symbol

7 14 7 7 7 14.007 N

9 19 9 10 9 18.998 F

19 39 19 20 19 39.098 K

27 59 27 32 27 58.933 Co

Practice

2. If 2 protons were removed from the nucleus of an oxygen atom, what nucleus remains?

O: 8 p+

- 2 p+

6 p+ ________Carbon

IV. The Subatomic Particles

C. Isotopes1. All atoms of an element must have

the same number of protons

2. BUT they may have a different number of neutrons

3. All atoms of an element must have the same atomic # but can have a different mass #.

4. These are called isotopes.

isotopes

For Isotopes

Parts of the Atom

Isotopes of Lithium

Li LiLithium – 6 Lithium - 7

Atomic # Mass # 6

373

ISOTOPES

++

++++

++

NUCLEI OF ATOMS IN THE SAME ELEMENT

++++

++++

++++

UNRAVEL

6 NEUTRONS6 PROTONS

ISOTOPE SYMBOL

C12

6

Atoms in the same element with different MASS NUMBER but identical ATOMIC NUMBER.

ISOTOPES

++

++++

++

NUCLEI OF ATOMS IN THE SAME ELEMENT

++

++++

++

++++

++++

++++

++++

++++

++++

UNRAVEL

6 NEUTRONS6 PROTONS

7 NEUTRONS6 PROTONS

ISOTOPE SYMBOL

C12

6

C13

6

XA

Z

X –ELEMENT SYMBOLA-MASS NUMBERZ-ATOMIC NUMBER

Atoms in the same element with different MASS NUMBER but identical ATOMIC NUMBER.

Weighted Averages

• The atomic mass of one atom is TINY if reported in kilograms or even grams.

• Atomic mass is reported in atomic mass units (amu).

• 1 amu is about the mass of one proton or neutron (about 1.67 x 10-24)

• Generally, atoms of a given element will have 1 isotope in high abundance and several others in much smaller numbers.

• To calculate the average atomic mass, we use a WEIGHTED AVERAGE that accounts for the abundance of each isotope.

Ex 1:

• Carbon-12 makes up 98.89% of naturally-occurring carbon. Carbon-13 makes up 1.11% of naturally occurring carbon. Use this information to determine the average atomic mass of carbon.Average atomic mass =

(12 X 0.9889) + (13 X 0.0111) = 12.0111

amu

Ex 2:

• Chromium has 4 naturally-occurring isotopes. Their abundance is as follows: Cr-50 – 4.35%, Cr-52 – 83.79%, Cr-53 – 9.50%, and Cr-54 – 2.36%. Determine the average atomic mass for chromium.

(50 X .0435) + (52 X .8379) + (53 X .0950) + (54 X .0236) =

52.0552 amu

Practice:

1. The element copper is found to contain 69.1% of copper-63 and 30.9% of copper 65. Calculate the average atomic mass of copper.

2. Gallium occurs in nature as a mixture of two isotopes. They are gallium-69 with an 60.108% abundance and gallium-71 with a 39.892% abundance. Calculate the atomic mass of gallium.

63.618 amu

69.7978 amu

3. Neon has 2 isotopes: neon-20 and neon-22. Use the information from the periodic table to determine which occurs in greater abundance.

Neon – 20

Because the atomic mass on the periodic table rounds to 20 and not 22

4. Use the table below to calculate the atomic mass of Element X and then identify it.

Isotope % Abundance

16X 99.76217X .03818X .2

(16 X .99762) + (17 X .00038) + (18 X .002) =

16.0044 amuOxygen