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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

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elements a k a atoms
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
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.


timeline of the atom
Timeline of the Atom

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.“

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

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

Mid-1920’s – wave mechanical model

1932 – Rutherford and his coworker Chadwick identified the neutron

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
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
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 electron
Conclusions from the Study of the Electron
  • Cathode 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?
william thomson s lord kelvin s atomic model
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 experiment1
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 through

About 2% of the  particles went through but were deflected by large angles

About 0.01% of the  particles bounced off the gold foil

Most of the volume of the atom is empty space

rutherford s conclusion a nuclear model
Rutherford’s Conclusion: A Nuclear Model

The atom contains a tiny dense center called the nucleus

  • the volume is about 1/10 trillionth the volume of the atom

The 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
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
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
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.


Summary of Atomic Particles

Going beyond the electron, proton, and neutron

describing an atom
Describing an Atom

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

atomic structure protons
Atomic Structure - protons

The 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
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
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
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


Mass number

(p+ + no)

Charge (if any)

Cl-1 –38 Fluorine-18




Element symbol


Atomic number

(number of p+)

Element name

Mass number

Element symbol with charge

Mass number


atoms of an element with the same number of protons but different numbers of neutrons

isotopes examples
Isotopes Examples
  • 3517Cl 3717Cl
  • H-1 H-2 H-3

Copper – 63 Copper – 65

  • 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
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
  • Multiply the percentage (percent abundance) of each isotope of the element by its mass number.
  • Add the products of the multiplications together.
  • Divide by 100.
  • Your answer should be very close to the atomic mass of the element for that element
average atomic mass examples
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
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
  • 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.
nuclear transformations
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
radioactivity review
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
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




the alpha particle
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
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
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
  • The positron is similar to a beta particle, but has a positive charge.
  • Example:
neutron emission
Neutron Emission
  • Neutron emission or decay does not change the element, only the mass
  • Example:
balancing nuclear equations
Balancing Nuclear Equations
  • Mass # and the atomic # totals must be the same for reactants and the products.
  • 3919K  3517Cl + ___
  • 20682Pb  0-1e + ___
writing balanced nuclear equations
Writing Balanced 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
nuclear chemistry and half life

Nuclear Chemistry and half-life

Balancing Nuclear Equations


half life and nuclear stability
Half-Life and Nuclear Stability
  • 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.
sample half lives
Sample Half-lives
  • IsotopesHalf-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
working with half life
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 chemistry1
Uses of Nuclear Chemistry
  • Medicine
    • X-rays and MRIs
    • Chemical tracers
  • Energy
    • Reactors generate electricity for homes, etc.
  • Destruction
    • Bombs
fission versus fusion
Fission versus Fusion
  • Fusion – combining two smaller nuclei into one heavier, more stable nucleus.

32He + 11H  42He + 01e

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

20984Po 12552Te + 8432Ge