Atomic Structure and Nuclear Chemistry

Chapter 4 and 18

Elements a.k.a atoms Robert Boyle first defined an element as a substance which could no longer be broken down into other substances Each element has unique properties

Many early scientists speculated how the element or atom was structured

Theory of atomic structure evolved from early thoughts to today’s atom

What do you know about the atom? Take a moment and create a concept

web about the atom. Work on your own. You have about 5 minutes. Jot down everything you can connect

to atoms.

Atoms

Timeline of the Atom

1808 - Dalton first proposed a theory on atoms

Discovery of the electron by J.J. Thompson in the late 1890’s

1910 -Lord Kelvin’s “plum pudding theory

1912 –Bohr model of the hydrogen atom

1932 – Rutherford and his coworker Chadwick identified the neutron

Mid-1920’s – wave mechanical model

Rutherford’s gold foil experiment leads to atomic nucleus and in 1919 the introduction of the proton

Democritus (460-370 B.C.) understood that if you cut a stone in two pieces, each piece contained the same material as the original stone. He also believed that you could do this an infinite number of times. He called these infinitesimally small pieces of matter atomos, meaning "indivisible.“

Who started…

John Dalton (1766 – 1844) was an English scientist who made his living as a teacher in Manchester.

Dalton’s Atomic Theory (p.88) Elements are composed of atoms All atoms of a given element are identical Atoms of a given element are different from those of

any other element Atoms of one element combine with atoms of other

elements to form compounds Law of Constant Composition: all samples of a

compound have the same proportion of the elements as in any other sample of that compound

Atoms are indivisible in a chemical process. all atoms present at the beginning of a chemical

process must also be present at the end of the process.

atoms are not created or destroyed, they must be conserved.

atoms of one element cannot be turned into atoms of another element

Atomic Structure History Discovery of the Electron

1st atomic particle identified In 1897, J.J. Thomson used a cathode ray

tube to deduce the presence of a negatively charged particle.

Cathode ray tubes pass electricity through a gas that is contained at a very low pressure. This creates a beam of negatively charged particles bent by an electric field.

Conclusions from the Study of the ElectronCathode rays have identical properties regardless of

the elemental gas used to produce them. Therefore, all elements must contain identically

charged particles (electrons). Atoms are neutral, so there must be positive particles

in the atom to balance the negative charge of the electrons

Electrons have so little mass that atoms must contain other particles that account for most of the mass

An electron is a tiny, negatively charged particle

The next step…. What is the positive charge?

Models of the atom

William Thomson’s (Lord Kelvin’s) Atomic Model

Lord Kelvin believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model (easier to think of as “chocolate chips" in chocolate chip cookie dough.)

Rutherford’s Gold Foil Experiment

Rutherford’s Gold Foil Experiment Rutherford shot α (alpha) particles at a thin

sheet of gold foil (think: bullet = alpha particles, target atoms = gold foil)

α particles are positively charged gold atoms are about 50 larger than α

particles. Particles were fired at a thin sheet of gold

foil Particles hit on the detecting screen (film)

were recorded

(a) The results that the metal foil experiment would have yielded if the plum pudding model had been correct. (b) Actual results known as Rutherford’s model.

Over 98% of the particles went straight throughAbout 2% of the particles went through but were deflected by large anglesAbout 0.01% of the particles bounced off the gold foilMost of the volume of the atom is empty space

Rutherford’s Conclusion: A Nuclear ModelThe atom contains a tiny dense center

called the nucleus the volume is about 1/10 trillionth the volume of the

atomThe nucleus is essentially the entire mass

of the atom (extremely dense)The nucleus is positively charged

the amount of positive charge of the nucleus balances the negative charge of the electrons

The electrons move around in the empty space of the atom surrounding the nucleus

Finally, the neutron.. Discovered in 1932 by Chadwick based on the

idea from Rutherford Has no charge Is located in the nucleus Mass a mass of 1 amu (actually, it’s slightly

larger than a proton but for our work the mass is the same)

So, the structure of the atom ……. The nucleus was found to be composed of two

kinds of particles Some of these particles are called protons

◦ charge = +1◦ mass is about the same as a hydrogen atom

Since protons and electrons have the same amount of charge, for the atom to be neutral there must be equal numbers of protons and electrons

The other particle is called a neutron◦ has no charge◦ has a mass slightly more than a proton

The Modern Atom We know atoms are composed of three

main atomic particles - protons, neutrons and electrons

The nucleus contains protons and neutrons

The radius of the atom is about 100,000 times larger than the radius of the nucleus

A nuclear atom viewed in cross section.

Particle Charge Mass # Location

Electron -1 0 Electron cloud

Proton +1 1 Nucleus

Neutron 0 1 Nucleus

Summary of Atomic Particles

Going beyond the electron, proton, and neutron

Describing an Atom

How many protons, neutron, and electrons does an atom have?

Element # of protons

Atomic # (Z)

Carbon 6 6Phosphorus

15 15

Gold 79 79

Atomic Structure - protonsThe number of protons in an atom of a given

element is the same as its atomic number (Z).

(Z) is found on the Periodic Table, whole # for each element

Atomic Structure - neutrons Mass number = protons + neutrons;

always a whole number.

# of Neutrons = mass number - # of protons

Atomic mass – decimal number in each element’s box on the periodic table. If you round the atomic mass of an element to the closest whole number you generally get the mass # for that element.

Atomic Structure - Electrons # of Electrons = # of protons if the atom is neutral If the chemical symbol is written with a charge,

representing an ion, the charge indicates the number of electrons that have been added or removed from the atom. If the ion has a positive charge (cation), subtract that

charge from the # of protons to get the number of electrons.

If the ion has a negative charge (anion), add that charge number to # of protons to get the number of electrons.

# of Electrons = # protons – charge Charge = # protons - # electrons

Representing atomic particles in atoms The number of each type of atomic particle (proton,

neutron, electron) is determined by using symbols. There are several different ways to write an

element: Atomic symbols Nuclear symbols

Atomic Symbols include the element symbol and a charge if any. C – neutral carbon C+4 – carbon cation C-4 – carbon anion

Charge (if any)

Cl-1 –38 Fluorine-18

Nuclear Symbols

Element symbol

Mass number (p+ + no)

Atomic number (number of p+)

U238

92

+

Mass number

Mass number

Element name

Element symbol with charge

Atom # protons # neutrons # electrons

Silver – 109

Pb-208

C-14

Isotopes atoms of an element with the same number of protons but different 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

Two isotopes of sodium.

Isotopes Examples 35

17Cl 3717Cl

H-1 H-2 H-3

Copper – 63 Copper – 65

Isotopes and Atomic Mass

AMU When we think about the mass of an atom, we

use atomic mass units (amu). A proton is 1 amu A neutron is 1 amu

Add up protons and neutrons to get the mass number (not atomic mass) Why? Most elements in nature have isotopes All these isotopes contribute to the average

atomic mass (listed on the table)

Determining Average Atomic Mass The average atomic mass seen on the periodic

table is a combination of all an element’s isotopes and their abundance.

To determine the average atomic mass for an element, you must

1. Multiply the percentage (percent abundance) of each isotope of the element by its mass number.

2. Add the products of the multiplications together.3. Divide by 100.4. Your answer should be very close to the atomic

mass of the element for that element

Average Atomic Mass Examples Find the average atomic mass of each of the

following elements from their percentages and mass numbers.

69.17% 63Cu and 30.83% 65Cu

5.85% Fe-54, 91.75% Fe-56, 2.12% Fe-57 and 0.28% Fe-58

Nuclear Reactions What you just did was write nuclear reactions.

Typical reactions are decay reactions and capture reactions.

What did you notice about the products of the reactions? The products of the reactions are isotopes of the

element. Nuclear reactions produce different particles that are

not elements. You need these particles to balance out the protons

and neutrons in the nuclei. In a nuclear reaction, the atomic number (Z) and the

mass number (A) are conserved.

Radioactive decay

Radioactive decay is a natural process. Only a handful (~200) of the known isotope nuclei (2000) do

not decay A nuclei may spontaneously kick out a particle, forming a

new element. There are four common particles:

The alpha particle - α (gold foil) The beta particle – β The gamma particle – γ The positron particle

In many cases, the process does not stop at one step, but rather a combination of steps. This is known as a decay series.

Capture Reactions

Nuclear Transformations Nuclear transformation is the changing of one element

to another (modern alchemy!!!) using larger atoms

Rutherford observed 1st transformation in 1919

Marie Curie and her husband another transformation (14 years later)

Electron capture is the “natural” nuclear transformation

Man-made elements are made by bombarding two nuclei together

The nuclear particles

Radioactivity Review Radioactivity is expressed as either a decay

process or a capture process

Decay processes can be connected as a chain

Nuclear transformations use larger atoms to create different elements than what you started with

Review of nuclear symbols Mass number is in upper left of symbol Atomic mass is in lower left of symbol

Nuclear particles are written in the nuclear format

U238

92

The alpha particle (α)

The alpha particle is actually a helium nucleus. It is the weakest of the decay particles. A alpha particle has a mass number of 4 and an

atomic number of 2. If an α particle is added to a nuclei, the mass number

will increase by 4 and the atomic number by 2. If an α particle is released from a nuclei, the mass

number will decrease by 4 and the atomic number by 2.

For example:

The Beta particle (β)

The beta particle is an electron. It has no mass. It does have a charge

Examples include:

Sometimes a nucleus will grab an electron that is close. This is called electron capture.

The Gamma Particle (γ)

This is also know as the gamma ray. This is a high energy photon of light. Picked up by specific detectors. It is the strongest (most dangerous) of the decay

particles. Example:

The Positron The positron is similar to a beta particle, but

has a positive charge. Example:

Neutron emission or decay does not change the element, only the mass

Example:

Neutron Emission

Mass # and the atomic # totals must be the same for reactants and the products.

3919K 35

17Cl + ___

20682Pb 0

-1e + ___

Balancing Nuclear Equations

Alpha decay of Cu-68

Gamma emission of Thorium-235

Positron emission of P-18

Astatine-210 releasing 3 neutrons

Electron capture with Ti-45

Writing Balanced Nuclear Equations

Nuclear Chemistry and half-life

Balancing Nuclear EquationsHalf-life

Radioactive isotopes or nuclides all decay because they are unstable, some just breakdown much faster than others

Geiger counter or scintillation counters are used to detect particles.

Half-life – amount of time for half of the original sample to decay

For two samples of the same isotope, regardless of the sample size, after one half-life, only half of the original amount of sample remains.

Half-Life and Nuclear Stability

Isotopes Half-Live Carbon – 14 5730 years Sodium – 24 15 hours Bismuth – 212 60.5 seconds Polonium – 215 0.0018 seconds Thorium – 230 75400 years Thorium – 234 24.1 days Uranium – 235 7.0 x 108 years Uranium – 238 4.46 x 109 years

Sample Half-lives

Working with half-life A material has t1/2 = 10 minutes. If you begin

with 16g, how long will it take to decay to 2 g?

Begin with 16 g, 1 half life gets you to 8 g. 2 half lives get you to 4 g. 3 half lives get you to 2 g. So, 3 x 10 minutes = 30 minutes.

A material has t1/2 = 150 years. If you begin with 100 g, how long will it take to decay to 3.125 g?

A material has t1/2 = 15 minutes. How much material is left after 75 minutes if you begin with 100 g? Calculate the number of half-lives used = 75/15 =

5 Run through 5 half lives:

After 1 – 50g After 2 – 25 g After 3 – 12.5 g After 4 – 6.25 g After 5 – 3.125 g

A material has t1/2 = 6.2 years. How much material is left after 24.8 years if you begin with 14 g?

What is the half-life of a material that decays from 16 g to 2 g over 20 minutes?

Determine the number of half lives: 16 → 8 8 → 4 4 → 2 3 half lives spent

Divide the time by the number of half-lives: 20/3 = 6.67 minutes

What is the half-life of a material that decays from 125 g to 3.9 g over 100 years?

Uses of nuclear chemistry

Uses of Nuclear Chemistry Medicine

X-rays and MRIs Chemical tracers

Energy Reactors generate electricity for homes, etc.

Destruction Bombs

Fusion – combining two smaller nuclei into one heavier, more stable nucleus.

32He + 1

1H 42He + 0

1e

Fission – splitting a large unstable nucleus into two nuclei with smaller mass numbers.

20984Po 125

52Te + 8432Ge

Fission versus Fusion

Sample fission reaction

Sample fusion reaction