Chapter4(12-13)forwebsite.pptx

Chapter 4

4.1 Defining the Atom

•

Democritus’s Atomic

Philosophy

• _{The Greek philosopher}

Democritus

(460 BC –370 BC) was among the first to suggest the existence of atoms.

• _{Democritus reasoned that atoms}

were indivisible and indestructible.

Democritus’s Atomic Philosophy

•

_{Although Democritus’s ideas}

agreed with later scientific theory,

they did not explain chemical

behavior

•

_{They also lacked experimental}

Dalton’s Atomic Theory

–

The modern process of

discovery regarding atoms

began with John Dalton

(1766–1864), an English

chemist and schoolteacher.

–

By using experimental

methods, Dalton

transformed Democritus’s

ideas on atoms into a

4.1 Defining the Atom

•

Dalton’s Atomic Theory

indivisible particles called atoms.

• _{Atoms of the same element are}

identical. The atoms of any one element are different from those of any other element.

• Atoms combine in simple whole

number ratios to make compounds.

• In a reaction atoms are rearranged,

but they are never changed into atoms of another element as a result of a chemical reaction.

Atoms of

element A _{element B}Atoms of

4.1 Defining the Atom

• _{How was John Dalton able to study atoms even}

though he couldn’t observe them directly? What

evidence did he use to formulate his atomic theory?

• _{Dalton studied the ratios in which elements combine}

in chemical reactions. He observed that when atoms mix, they maintain their own identity unless they

4.1 Defining the Atom

•_{Atoms are very small.}

•_{A pure copper coin the size of a penny contains}

about 2.4  1022 atoms.

•_{By comparison, Earth’s population is only about}

7  109 people.

•If you could line up 100,000,000 copper atoms side

by side, they would produce a line only 1 cm long!

•Despite their small size, individual atoms are

4.2 Structure of the Nuclear Atom

•

Thomson (1897)

–

_{used the cathode ray}

experiment to

discover the electron

–

_{Established the}

charge to mass ratio

for these particles

1856-1940

Discovery of the Electron

9

In 1897, J.J. Thomson used a

cathode ray tube

4.2 Structure of the Nuclear Atom

•

_{Thomson’s model is also known as the plum}

pudding model.

•

_{Electrons are stuck into a lump of positive charge,}

similar to raisins stuck in dough.

4.2 Structure of the Nuclear Atom

•

_{Millikan (1909) – Oil-Drop}

Experiment: determined the

mass and charge of the

electron.

4.2 Structure of the Nuclear Atom

•

_{Combined with the charge/mass ratio from}

Thomson, Milikan was able to accurately

calculate the mass of a single electron.

–

_{The mass of an electron is 1/1840 the mass}

of a hydrogen nucleus.

–

_{The charge of an electron is one unit of}

negative charge.

•

_{Even with his primitive equipment, his}

4.2 Structure of the Nuclear Atom

•

_Protons

–

_{In 1886, Eugen Goldstein (1850–1930)}

observed a cathode-ray tube and found rays

traveling in the direction opposite to that of the

cathode rays.

• _{He concluded that they were composed of positive}

particles.

–

_{Such positively charged subatomic particles}

4.2 Structure of the Nuclear

Atom

•

_{Rutherford (1911):}

used the Gold Foil

experiment to discover

the nucleus and basic

structure of the atom.

– _{Electrons are outside}

the positively charged nucleus.

– _{Most of the atom is}

empty space.

Ernest Rutherford’s

Gold Foil Experiment - 1911

16

• Alpha particles are helium nuclei - The alpha

particles were fired at a thin sheet of gold foil

• Particle that hit on the detecting screen (film)

Rutherford’s problem:

17

In the following pictures, there is a target

hidden by a cloud. To figure out the shape of

the target, we shot some beams into the cloud

and recorded where the beams came out. Can

you figure out the shape of the target?

Target

#1

The Answers:

18

Rutherford (1920) – Concluded there must

be a particle in the nucleus carrying a

charge equal but opposite of the electron

and he called it a

proton

.

4.2 Structure of the Nuclear Atom

• _{Rutherford (1920) He concluded that all the positive}

charge and almost all of the mass are concentrated in a small region that has enough positive charge to account for the great deflection of some of the alpha particles.

• _{Nucleus: the tiny, dense central portion of an atom,}

composed of protons and neutrons

•

Nuclear Model:

4.2 Subatomic Particles and the nuclear atom

21



James Chadwick (1932) -

discovered the existence of a

neutrally charged particle in the

nucleus called the

neutron

.



Neutrons

are subatomic particles

with no charge but with a mass

nearly equal to that of a proton.

1.673×10-24 1 0 Nucleus n0 Neutron 1.673×10-24 1 1+ Nucleus p+ Proton 9.11×10-28 1/1840 1-Outside Nucleus e -Electron Actual Mass (g) Relative Mass (amu) Relative Electrical Charge Location Symbol Particle

Counting the Pieces

• _{Moseley (1913) – discovered each}

element has a unique positive charge in their nuclei.

– _{Each element has a different number of}

protons.

Atomic Number

• _{(# of protons determines kind of atom)}

•

_{Atomic Number}

electrons in a neutral atom

•

_{Mass Number}

protons + neutrons

• _{(NOTE: Mass number is not found on}

the periodic table.

23 12 6 C 14 6C 12 6

C

Mass Number = A

Atomic Number

24

The atomic number (Z):

tells the number of protons.

determines identity of element

4.3 Distinguishing Among Atoms

–

_{Atomic number = # protons}

•

# protons = # electrons

How many protons

does chlorine have?

17

17

How many electrons?

4.3 Distinguishing Among Atoms

Atoms of the First Ten Elements

Name Symbol Atomic

number Protons Neutrons numberMass Electrons

Hydrogen H 1 1 0 1 1

Helium He 2 2 2 4 2

Lithium Li 3 3 4 7 3

Beryllium Be 4 4 5 9 4

Boron B 5 5 6 11 5

Carbon C 6 6 6 12 6

Nitrogen N 7 7 7 14 7

Oxygen O 8 8 8 16 8

Fluorine F 9 9 10 19 9

Neon Ne 10 10 10 20 10

For each element listed in the table below,

Symbols

Contain the symbol of the element, the mass

number and the atomic number

27

X

Mass number

Atomic number

Symbols

Find the

–

number of protons

–

number of neutrons

–

number of electrons

–

Atomic number

–

Mass number

Br

80

35

= 35 = 45 = 35 = 35 = 80

http://www.chem.purdue.edu/gchelp/liquids/bromine.gif

Symbols

•

_{Find the}

–

_{number of protons}

–

_{number of neutrons}

–

_{number of electrons}

–

_{Atomic number}

–

_{Mass number}

29

F

19

9

= 10 = 9 = 9

4.3 Distinguishing Among Atoms

•

_{Look at the atoms of neon below.}

• _{All have the same number of protons (10).} • _{All have the same number of electrons (10).}

4.3 Distinguishing Among Atoms

ISOTOPES

•

Atoms of the same element with the same

number of protons, but different mass numbers

and different number of neutrons.

Isotopes

Neutrons 6 Protons 6 Electrons 6

Nucleus

Electrons

Carbon-14

Isotopes

33

Mass #

Atomic #

•

Nuclear symbol:

•

Hyphen notation:

carbon-12

Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

12

Mass Number

Mass number is the protons and neutrons in the nucleus of an isotope: Mass #= p+ + n0

We can also put the mass number after the name of the element: carbon-12

carbon-14 uranium-235

Nuclide

p

n

e

Mass #

Oxygen - ₁₀

- 33 42

- 31 15

8

8

18

Arsenic

₇₅

₃₃

Phosphorus

16

Isotopes

are atoms of the

same element

having

different masses

, due to varying numbers of

neutrons.

Isotope Protons Electrons Neutrons Nucleus

Hydrogen–1

(protium) 1 1 0

Hydrogen-2

(deuterium) 1 1 1

Hydrogen-3 (tritium)

1 1 2

An Introduction to Ions

• _{Atoms are neutral due to balanced numbers of protons}

and electrons.

An Introduction to Ions

• _{Ion notations}

– _{Charges are shown as superscripts after symbol.} – _{+ and – are used to show positive and negative}

charges

• _{Positive indicates number of electrons lost}

• _{Negative indicates number of electrons gained}

– _{For +1 and -1, the one is generally implied.} – _{Any element written without charge is neutral}

.

2 electrons

gained 2 electrons

4.3 Distinguishing Among Atoms

Calculate the p+_{, e}-_{, and n}0 _{of each.}

p+ =

e- =

n0 =

17 17 18

p+ =

e- =

n0 =

20

21 20

Iron - 56

Oxygen - 18 p+ =

e- =

n0 =

26 26 30

p+ =

e- =

n0 =

Symbols

Find the

–

number of protons

–

_{number of neutrons}

–

_{number of electrons}

–

_{Atomic number}

–

_{Mass number}

Cl

37

17

1-Chloride ion = 17

= 20 = 18 = 17

= 37

Symbols

Find the

–

number of protons

–

_{number of neutrons}

–

_{number of electrons}

–

_{Atomic number}

–

_{Mass number}

Ca

46

20

2+

Calcium ion = 20

= 26 = 18 = 20

= 46

4.3 Distinguishing Among Atoms

•

_{Because the mass of atoms is so small}

(proton = 1.67×10

g) we simplify atomic

masses by measuring them in

atomic

mass units

.

•

_{Atomic mass unit (amu)}

_{– 1/12 the mass}

of carbon-12.

4.3 Distinguishing Among Atoms

•

_{In nature, most elements occur as a mixture of}

two or more isotopes.

– _{Each isotope of an element has a fixed mass and a}

4.3 Distinguishing Among Atoms

If elements have isotopes with different atomic

masses, what is the atomic mass on the periodic

table?

•

Atomic Mass – weighted average mass of the

isotopes of an element.

•

This found by summing the mass contribution of

each isotope of the element.

There are 2 naturally occurring isotopes of

copper; copper-63 and copper-65. If copper-63

has a mass of 62.930 amu and 69.17%

abundance and copper-65 has a mass of 64.928

amu and 30.83% abundance, what is the

average atomic mass of copper?

1. First, calculate the mass contribution of each

isotope to the average atomic mass, being sure

to convert each percent to a fractional

2. Finally, the average atomic mass of the element

is the sum of the mass contributions of each

isotope.

63.55

Mass contribution = (62.930 amu)(.6917) = 43.52868 43.53 amu For copper-65:

Mass contribution = (64.928 amu)(.3083) = 20.017302 20.02 amu

amu

20.02

amu

43.53

Average Atomic Mass

•

_{EX: Calculate the average atomic mass of}

oxygen if its abundance in nature is 99.76%

2

. Finally, the average atomic mass of the

element is the sum of the mass contributions

of each isotope

.

16.003 amu Mass contribution of oxygen-16:

Mass contribution = (16.00 amu)(.9976) = 15.9616 15.96 amu Mass contribution of oxygen-17:

Mass contribution = (17.00 amu)(.0004) = .0068 .007 amu Mass contribution of oxygen-18:

Mass contribution = (18.00 amu)(.0020) = .036

4.4 Unstable Nuclei and Radioactive

Decay

Can chemical reactions change the identity of

an atom?

NO!

Why Not?

You can’t mess with the nucleus!

Does this mean the nucleus is not affected by

any reactions?

4.4 Unstable Nuclei and

Radioactive Decay

Why do they change?

Stability!

Unstable systems, like an atom with the wrong

number of neutrons or a pencil sitting on its tip,

gain stability by losing energy!

The pencil loses energy when it falls,

4.4 Unstable Nuclei and

Radioactive Decay

Types of radioactive decay

1.

Alpha decay ( )

– Mass # drops by 4 and atomic number drops by 2 – +2 charge

2.

Beta Decay( )

– Mass # stays same and atomic number increases by 1 – -1 charge

He

4 2

He

Rn

Ra

226

88





β

or

e

0 1



e

N

C

14

4.4 Unstable Nuclei and

Radioactive Decay

3.

Gamma decay ( )

– No charge

– No mass # change or atomic number change – Usually accompanies alpha or beta decay

– Most of the energy lost in radioactive decay is from

gamma decay

γ



4.4 Unstable Nuclei and

Radioactive Decay



Balancing nuclear equations

– Mass number and atomic number must be

conserved.

2 234

90

Th

He

U

4

238