Chapter 4

Chapter 4: Atomic Structure

DO NOW:- Take out your lab safety sheet

with the map on the back

- Get with your lab partner and complete the rest of your lab!

John Dalton (1803) Atomic Theory

• All matter is made of tiny, indivisible,

indestructible particles called atoms.

• All atoms of a given element are identical

in size, mass and chemical properties, but

they are different from atoms of other

elements.

• Different atoms combine in simple

whole-number ratios to form compounds.

• In a chemical reaction, atoms are

separated combined or rearranged.

The Atom

• Atom = smallest particle of an element that retains its identity in a chemical reaction

– Radii of most atoms are between 5 x 10-11 m to 2 x 10-10 m

• Can see individual atoms with a scanning tunneling microscope

Chapter 4

Electrons

• Electron = negatively charged subatomic particle– Charge = -1.6 x 10-19 Coulombs

• Robert Millikan (1916) calculated mass of electron with his Oil-drop experiment– Mass of e- = 9.11 x 10-31 kg

Chapter 4

Subatomic Particles

• Electron (e-)– J.J. Thomson discovered the electron in 1897

while experimenting with cathode ray tubes

Thomson’s Experiment

Voltage source

+-

Vacuum tube

Metal Disks

Thomson’s Experiment

Voltage source

+-

● Passing an electric current makes a beam appear to move from the negative to the positive end

Thomson’s Experiment

Voltage source

+-

● Passing an electric current makes a beam appear to move from the negative to the positive end

Thomson’s Experiment

Voltage source

+-

Voltage source

Thomson’s Experiment

● By adding an electric field

+

-

Voltage source

Thomson’s Experiment

● By adding an electric field . . .

+

-

Voltage source

Thomson’s Experiment

● By adding an electric field

+

-

Voltage source

Thomson’s Experiment

● By adding an electric field

+

-

Voltage source

Thomson’s Experiment

● By adding an electric field he found that the moving pieces were negative

+

-

Thomson Plum Pudding Model

1. Atom breakable 2. Atom has structure

3. Electrons suspended in a positively charged electric field

4. Mass of atom due to

electrons

5. Atom mostly "empty" space

The Atomic Nucleus

– What he expected from the “plum

pudding” model:

• Alpha particle to deflect slightly off

gold foil

– What he experienced:

• Most particles went straight through

the gold

• Some even bounced straight back!

– Proved atom is mostly empty space, with

nucleus of protons in center

E. Rutherford – Gold Foil Experiment (1911)

Lead block

Uranium

Gold Foil

Florescent Screen

He Expected that…

• The alpha particles to pass through without changing direction very much

• The positive charges were spread out evenly. Alone they were not enough to stop the alpha particles

“alpha particles not changing direction”

He expected this because he thought the mass was evenly distributed in the atom

What he got actually got . . .

How he explained it:

+

• Atom is mostly empty

• Small dense, positive piece at center

• Alpha particles are deflected by it if they get close enough

+

Rutherford's Gold Foil Experiment

If a football stadium were the size of an atom then…

• Nucleus = smaller than a dime

• Electron = smaller than the eye on the dimes’ picture of FDR

If a football stadium were the size of an atom then…

Nuclear Model of the Atom

Positively charged nucleus-

contains protons and

neutrons

“Empty Space”- Electron cloud

containing negatively-charged

electrons

In elements, for each +1

proton there must be a -1

electron to cancel

out the electrical charge.

Atomic Number and Mass #

• Atomic # (Z) = number of protons in the nucleus

• Mass # (A)= mass of the neutrons + protons in the

nucleus (electron mass too small)

Atomic Mass Unit (amu) based on carbon-12 = 12.000 u

Na-23Sodium (Na) Mass Number = 23Atomic Number = 11 = Number of Protons in NucleusMass Number - Protons = NeutronsFor Na: 23 - 11 = 12 neutrons in nucleus

• It’s what makes carbon carbon!• It’s what makes oxygen oxygen and not carbon!

a.k.a. the number of protons

Elements on Periodic Table

10

NeNeon

20.1797

Atomic

Number

(Z)

Element

Symbol

Element

Name

Atomic

Mass

One or two letters

Two letters – 1st capital, 2nd lower case

Isotopes

• Isotope = same number of protons, different number

of neutrons in an element

• Three isotopes of hydrogen:

+

Hydrogen-1

1 proton

0 neutrons

+

Hydrogen-2

(Deuterium)

1 proton

1 neutron

+

Hydrogen-3

(Tritium)

1 proton

2 neutrons

Chapter 4

• Mass Number, Atomic Number, and Element is shown as

follows:

• For isotopes, different mass numbers are shown in the

top number.

• Another way of showing isotopes is using the element

name and mass number:

carbon-12 carbon-13 carbon-14

More Isotopes

Mass NumberAtomic Number Element Symbol

126 C 13

6 C 146 C

Determining # of Protons, Neutrons, and Electrons in an Ion or Atom

● Atomic number = number of protons● Number of neutrons =

Mass number - number of protons

● If the atom is neutrally charged: Number of

electrons = Number of protons

If the atom is an ion: Number of electrons =

Number of protons - Charge

Calculating Atomic Mass• Need to account for all isotopes of an element when

calculating atomic mass

• Weighted average of percent abundance in nature

• Example: Chlorine – Two isotopes - 35Cl (a.k.a. chlorine-35) and 37Cl (chlorine-37)– Natural Abundance: 76% 35Cl, 24% 37Cl– Atomic Mass = (76%)(35 amu) + (24%)(37 amu) = 35.453

amu