Chapter Menu The Structure of the Atom

Section 4.1 Early Ideas About Matter

Section 4.2 Defining the Atom

Section 4.3 How Atoms Differ

Section 4.4 Unstable Nuclei and

Radioactive Decay

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BELL RINGER! Describe the experiment proving the positive charge

in an atom resides in the nucleus

Rutherford’s gold foil experiment demonstrated that the positive charge in atoms resides in the nucleus. In this experiment, Rutherford set up an alpha particle emitter such that particles hit a very thin piece of gold foil. Alpha particles are the nucleus of the 4He isotope and have a charge of +2. In the experiment, most of the alpha particles passed directly through the gold foil but some were deflected by large angles

Niels Bohr

1885-1962

• Planetary Model 1913

– Nucleus surrounded by

orbiting electrons at

different energy levels

– Electrons have definite

orbits

• Utilized Planck’s

Quantum Energy theory

• Worked on the

Manhattan Project (US

atomic bomb)

Bohr’s Model

Bohr Model for Nitrogen

Ernst Schrödinger 1887-1961

• Quantum Mechanical

Model 1926

– Electrons are in

probability zones called

“orbitals”, not orbits and

the location cannot be

pinpointed

– Electrons are particles

and waves at the same

time

– Developed quantum

numbers based on

theories of Einstein and

Planck

Werner Heisenberg 1901-1976

Orbitals

Quantum Mechanical Theory Electron in a Hydrogen atom

The Structure of the Atom

ATOM

nucleus

energy

levels

protons

neutrons

electrons

( p+ )

( n0 )

( e- )

CHEMICAL COMPOSITION SHORTHAND

Cl 35

17

MASS

NUMBER

ATOMIC

NUMBER

NUMBER OF

PROTONS

# OF PROTONS +

# OF NEUTRONS

D. The Size of Atoms Electron

Cloud 1. Nucleus – extremely

small volume but

massive

2. Protons & neutrons

have the same mass

& are found in the

nucleus

3. Electrons are 1836

times smaller in

mass than protons

& neutrons & are

found in the

electron cloud

(large in volume)

Counting Atoms Objectives:

• Explain what isotopes are

• Define atomic number and mass

number

• Determine the number of protons,

neutrons, and electrons of a nuclide

• Define mole in terms of Avogadro’s

number, and define molar mass

A. Atomic Number

1. Represents the # of protons in an atom

Ex: Hydrogen is atomic #1 has 1 proton

2. If the atom is electrically neutral,

then # of protons (p+) = # of electrons (e-)

A = P = E

C. Using A=P=E M-A=N

M - A = N

1. How many p+, e-, and no are in

chlorine-37?

17 p+, 17 e-, 20 no

2. How many p+, e-, and no are in

bromine-80?

35 p+, 35 e-, 45 no

3. Write the nuclear symbol for carbon-13.

Worksheet

13

6 C

CHEMICAL COMPOSITION SHORTHAND

Cl 35

17

MASS

NUMBER

ATOMIC

NUMBER

NUMBER OF

PROTONS

# OF PROTONS +

# OF NEUTRONS

• EVERY CHLORINE ATOM HAS 17 PROTONS, WITHOUT EXCEPTION, –HOWEVER, NOT EVERY CHLORINE

ATOM HAS 18 NEUTRONS.

– ATOMS WITH THE SAME NUMBER OF PROTONS BUT CONTAIN DIFFERENT NUMBERS OF NEUTRONS ARE CALLED ISOTOPES.

• BECAUSE ISOTOPES OF AN ELEMENT HAVE DIFFERENT NUMBERS OF NEUTRONS THEY HAVE DIFFERENT MASS NUMBERS.

ISOTOPES

• ISOTOPES ARE CHEMICALLY ALIKE BECAUSE THEY HAVE IDENTICAL NUMBERS OF PROTONS AND ELECTRONS

– IT’S THE ELECTRONS AND PROTONS THAT ARE RESPONSIBLE FOR CHEMICAL BEHAVIOR

proton

neutron

electron

BERYLLIUM ISOTOPES

EXAMPLE OF AN

ISOTOPE

Cl 35

17 Cl 37

17

20 NEUTRONS

ATOMIC MASS

18 NEUTRONS

ATOMIC NUMBER

IONS • AN ELEMENT’S ATOMS ARE NOT

ALWAYS NEUTRAL IN CHARGE. – WHEN AN ATOM LOSES OR GAINS

ONE OR MORE OF ITS ELECTRONS IT BECOMES ION.

• AN ION THAT HAS MORE ELECTRONS THAN PROTONS HAS A NEGATIVE ELECTRICAL CHARGE

• AN ION THAT HAS FEWER ELECTRONS THAN PROTONS HAS A POSITIVE ELECTRICAL CHARGE

NOTE: IT’S THE PROTONS THAT DEFINE THE TYPE OF ATOM IT IS, BUT THE ELECTRONS DEFINE THE ATOM’S CHARGE.

SOME ATOMS GAIN ELECTRONS

O

-

- -

-

-

-

-

-

O-2

-

- -

-

-

-

-

-

- -

ATOM’S IONIC CHARGE =

# PROTONS - # ELECTRONS

ATOMS, IONS, AND ISOTOPES

ATOMS

NEUTRAL AND ARE DEFINED BY THE # OF PROTONS IN THEIR

NUCLEUS

3 p+ = Li ATOM, ETC.

IONS

HAVE AN ELECTRICAL CHARGE DETERMINED BY

# PROTONS - # ELECTRONS

N-2 = 7 p+ - 9 e- ; ETC.

ISOTOPES

TWO ATOMS WITH THE SAME # OF PROTONS, BUT DIFFERENT #’S OF

NEUTRONS OR MASSES

CALCIUM-40 & CALCIUM-44

atomic number:

--

--

To find net charge on an atom, consider

____ and ____.

mass number:

# of p+ the whole number

on Periodic Table

Does not change

determines identity

of atom

(# of p+) + (# of n0)

10

Ne 20.1797

p+ e–

(It is NOT on ―the Table.‖)

protium

Isotopes: different varieties of an element’s atoms

--

-- have diff. #’s of n0; thus, diff. masses

some are radioactive; others aren’t

Isotope

H–1

Mass p+ n0 Common Name

H–2

H–3 tritium

deuterium

1

2

3

1 0

1 1

1 2

C–12 atoms C–14 atoms

All atoms of an element react the same, chemically.

6 p+ 6 n0

stable

6 p+ 8 n0

radioactive

Proton

s

Complete

Atomic

Designation

92

Neutron

s

Electro

ns

34

11

146

45

12

92

36

10

59 27

3+ Co

37 17

1– Cl

55 7+ Mn 25

238 92

U

23 11

1+ Na

79 34

2– Se

20 18 17

30 18 25

32 24 27

Two Methods to Identify Isotopes

1. Hyphen notation - mass number is written with a hyphen after the name of the element

hydrogen-3, uranium-235

2. Nuclear symbol – shows composition of nucleus

232U

Superscript = mass number

Subscript = atomic number

Nuclide – general term for any isotope

of any element

92

Atomic mass unit – (amu) 1 amu is equal to

1/12 mass of a carbon –12 atom – used to express the masses of atoms

* carbon-12 is the standard for determining atomic mass

* depends on masses and relative abundance of the isotopes

Average atomic mass – weighted average of the atomic masses of the naturally occurring isotopes of an element

* may not equal the mass of any of its isotopes

* atomic mass listed in the periodic table is the average atomic mass

• An element’s atomic mass is

usually rounded to TWO decimal

places before it is used in a

calculation.

Average Atomic Mass (Atomic Mass, AAM)

This is the weighted average mass of all atoms of

an element, measured in amu.

For an element with

isotopes A, B, etc.:

AAM = Mass A (% A) + Mass B (% B) + …

(use the decimal form of the %;

e.g., use 0.253 for 25.3%)

% abundance

Ti has five naturally-

occurring isotopes

Lithium has two isotopes.

Li-6 atoms have mass 6.015 amu;

Li-7 atoms have mass 7.016 amu.

Li-6 makes up 7.5% of all Li atoms.

Find AAM of Li.

AAM = Mass A (% A) + Mass B (% B)

AAM = 6.015 amu (0.075) + 7.016 amu

AAM = 6.94 amu

(0.925)

AAM = 0.451 amu + 6.490 amu

Li batteries

** Decimal number on Table refers to…

molar mass (in g) OR AAM (in amu).

6.02 x 1023

atoms

1 ―average‖

atom

19

Description Net

Charge

Atomic

Number

Mass

Number

Ion

Symbol

15 p+

16 n0

18 e– 38 p+

50 n0

36 e–

18 e– 1+ 39

Te2– 128

K1+

Sr2+

P3–

88

31 15 3–

38

52 2–

2+

76 n0 52 p+

54 e–

20 n0 19 p+

a- or b-particles, g rays

Radioactive Isotopes:

Nucleus attempts to attain a lower

energy state by releasing extra

energy as __________.

e.g.,

half-life: the time needed for

½ of a radioactive

sample to decay

into stable matter

have too many or too few n0

radiation

e.g., C–14: -- half-life is 5,730 years

-- decays into stable N–14

Years

from now

0

g of N–14

present

5,730

11,460

120

60

30

15

7.5

0

Say that a 120 g

sample of C-14 is

found today.

g of C–14

present

17,190

22,920

60

90

105

112.5

= C–14

= N–14

Section 4.4 Unstable Nuclei and

Radioactive Decay

• Explain the relationship between unstable nuclei and

radioactive decay.

element: a pure substance that cannot be broken

down into simpler substances by physical or chemical

means

• Characterize alpha, beta, and gamma radiation in

terms of mass and charge.

Section 4.4 Unstable Nuclei and

Radioactive Decay (cont.)

radioactivity

radiation

nuclear reaction

radioactive decay

alpha radiation

Unstable atoms emit radiation to gain stability.

alpha particle

nuclear equation

beta radiation

beta particle

gamma rays

Radioactivity

• Nuclear reactions can change one element

into another element.

• In the late 1890s, scientists noticed some

substances spontaneously emitted radiation,

a process they called radioactivity.

• The rays and particles emitted are called

radiation.

• A reaction that involves a change in an atom's

nucleus is called a nuclear reaction.

Radioactive Decay

• Unstable nuclei lose energy by emitting

radiation in a spontaneous process called

radioactive decay.

• Unstable radioactive elements undergo

radioactive decay thus forming stable

nonradioactive elements.

Radioactive Decay (cont.)

• Alpha radiation is made up of positively

charged particles called alpha particles.

• Each alpha particle contains two protons and

two neutrons and has a 2+ charge.

Radioactive Decay (cont.)

• The figure shown below is a nuclear

equation showing the radioactive decay of

radium-226 to radon-222.

• The mass is conserved in nuclear equations.

Radioactive Decay (cont.)

• Beta radiation is radiation that has a

negative charge and emits beta particles.

• Each beta particle is an electron with a 1–

charge.

Radioactive Decay (cont.)

Radioactive Decay (cont.)

• Gamma rays are high-energy radiation

with no mass and are neutral.

• Gamma rays account for most of the energy

lost during radioactive decay.

Radioactive Decay (cont.)

• Atoms that contain too many or two few

neutrons are unstable and lose energy

through radioactive decay to form a stable

nucleus.

• Few exist in nature—most have already

decayed to stable forms.

Section 4.4 Assessment

A reaction that changes one element into

another is called what?

A. chemical reaction

B. beta radiation

C. nuclear reaction

D. physical reaction

Section 4.4 Assessment

Why are radioactive elements rare in

nature?

A. They do no occur on Earth.

B. Most have already decayed to a

stable form.

C. They take a long time to form.

D. They are too hard to detect.

• Mole (mol) – amount of a

substance that contains as many

particles as there are atoms in

exactly 12 g of carbon

* counting unit

• Avogadro’s number – 6.022 X 1023 number of particles in 1 mole of a

pure substance

• Molar mass – mass of 1 mole of a

pure substance

* numerically equal to the

average atomic mass of an

element

* must know the average atomic

mass to determine the molar mass

―When I see a cation, I see a positive ion;

ion:

anion: a (–) ion cation: a (+) ion

--

--

a charged atom

more e– than p+

formed when

atoms gain e–

--

-- more p+ than e–

formed when

atoms lose e–

a n ions negative ions. are I think that

that is, I… C ion.‖ A +