1

Electrical Charge is fundamental to our understanding of the

atom

Atoms and ElementsAtoms and Elements Electrostatics had been

discovered by the time of

Benjamin Franklin (late 1700’s)

Radioactivity

1896- Henri Becquerel:

Discovered uranium ore emitted

rays that exposed photographic

plates protected by black paper

1898- Marie/Pierre Curie: Isolated

polonium and radium which

emitted the same kind of rays

Suggested particles emitted through

atomic disintegration

Contradicted Dalton’s atomic theory

which states that atoms are indivisible

α, β and γ radiation

Radiation Movie 1 Radiation Movie 2

Components of the Atom

Electrons

Primarily discovered by J.J. Thomson in

1897 by discovery of “negatively

charged” cathode rays using Cathode

Ray Tube (Crookes tube)

Determined charge to mass ratio of the

electron.

Thomson’s method for determining the

charge to mass ratio (q/m) for the electron

Thomson developed the “plum pudding”

model of the atom

2

Robert Millikan (1909): Millikan’s Oil-drop

Experiment

Discovered the charge on the electron

(-1.602x10-19C)

Millikan’s

Experiment Movie

Eugen Goldstein (1886) – Discovered

positively charged “canal” rays

(positively charged nuclei of atoms, with

hydrogen having the smallest mass).

Protons and the Nucleus

1910- Ernest Rutherford: Gold foil

experiment

-Also experimented with alpha

irradiated gaseous elements: Alpha

deflection proportional to atomic mass

-Bombardment of nitrogen gas

produced particles consistent with that

of hydrogen (deemed a fundamental

particle)

-Developed Nuclear Model of the Atom

1919- Proton officially proclaimed

Gold Foil Movie 1 Gold Foil Movie 2

The Neutron

1932- James Chadwick: Discovered the

neutron by bombarding beryllium with

alpha particles

3

Atomic Number and Atomic Mass Atomic Mass Unit (u or amu) = 1/12 the

mass of a C-12 isotope. (aka 1dalton (Da))

About the mass of a proton or neutron (1u =

1.661x10-24g)

Isotopes

Atoms of the same element (i.e. the same

number of protons), that differ in their

number of neutrons.

(Compare to isobars (atoms with the

same mass number but different atomic

number) and isotones (atoms with the

same number of neutrons but different

number of protons)

The mass number of an isotope is not the sum

of the masses of the individual particles due to

the mass defect

(E =mc2), which is the binding energy.

Equations:

Percent abundance = (number of atoms of a

given isotope / total number of atoms of all

isotopes of that element)x100%

Atomic weight (weighted average)= (fractional

abundance of isotope 1)(mass of isotope 1) +

…

Isotopic Abundance is

Determined by mass spectroscopy

4

Mass spectroscopy of a molecule Sample Problems

a) Argon has three isotopes with 18, 20

and 22 neutrons, respectively. What are

the mass number and symbols of these

three isotopes?

b) Gallium has two isotopes:Ga-69 and

Ga-71. How many protons and neutrons

are in the nuclei of each of these

isotopes? If the abundance of Ga-69 is

60.1%, what is the abundance of Ga-71?

Answers:

36 38 40

18 18 18Ar Ar Ar

p 31 31

n 38 40

e 31 31

69

31Ga71

31Ga

Ga-71 39.9%

Sample Problems

1. There are three naturally occurring isotopes of

neon. Their percent abundances and atomic masses

are: neon-20, 90.51%, 19.99244u; neon-21, 0.27%,

20.99395u; neon-22, 9.22%, 21.99138u. Calculate

the weighted average atomic mass of neon.

2. The two naturally occurring isotopes of copper are

copper-63, mass 62.9298u, and copper-65, mass

64.9278u. What is the percent abundances of each

of the two isotopes? (The atomic mass of copper

listed on the periodic table is 63.546u)

Answers:

(.9051)(19.99244u)+ (.0027)(20.99395u) + (.0922)(21.99138u)

= 20.1794 = 20.18u

(62.9298u)(X) + (64.9278u)(1-X) = 63.546u

62.9298X + 64.9278u – 64.9278X = 63.546u

-1.998X = -1.3818u

X = .6916

Cu-63 = 69.16% Cu-65 = 30.84%

1 mole ≡ The quantity of things as there

are atoms in exactly 12.0g of the C-12

isotope.

This must be determined experimentally.

NA (Avogadro’s number) = 6.022x1023 = 1

mole

5

Mass number of an element

1) Taken in amu’s (u) = atomic mass (mass

of 1 atom)

2) Taken in grams (g) = molar mass (mass

of 1 mole of atoms)

24.30Mg If you have 24.30g of Mg you have

6.022x1023 atoms of Mg or 1mole

Determining the molar mass of a

compound

Add the molar masses of the individual

elements that the compound contains.

(multiply by MM) (divide by NA)

Grams Moles Basic Units (atoms, molecules, formula units, etc.)

(divide by MM) (multiply by NA)

Sample Problem

a) What is the mass, in grams, of 1.5 mol

of silicon?

b) What amount (moles) of sulfur is

represented by 454g? How many atoms?

c) What is the average mass (in grams) of

one sulfur atom?

Answers:

a) g Si = 1.5mol Si (28.09g/mol) = 42.135g =

42g42g42g42g

b) Mol S = 454g (1mol/32.07g) =14.1565 =

14.2mol14.2mol14.2mol14.2mol

# S atoms = 14.1565mol (6.022x1023) =

8.53x108.53x108.53x108.53x1024242424atomsatomsatomsatoms

c) Mass of 1 sulfur atom = 32.07g/mol (1 mol /

6.022x1023atom)

= 5.325x105.325x105.325x105.325x10----23232323gggg

Sample problem

The density of gold, Au, is 19.32g/cm3.

What is the volume (in cm3) of a piece of

gold that contains 2.6x1024 atoms? If the

piece of metal is a square with a

thickness of 0.10cm, what is the length (in

cm) of one side of the piece?

Answers:

2.6x1024atoms (1mol/6.022x1023atoms)

(197.0g/1mol)(1cm3 / 19.32g) = 44.02 =

44cm44cm44cm44cm3333

X2(0.10cm) = 44cm3

X = 20.982 = 21cm21cm21cm21cm

6

Dmitri Mendeleev: Russian Schoolteacher

Father of the modern periodic table (1869)

Table based on mass instead of atomic number (prior to

understanding of atomic particles)

Trends allowed for prediction of elements that had not yet

been discovered.

The Periodic Table

7

Diatomic elements

Allotropes