Ch.2

Atomic Structure, Isotopes, Ions and Moles

Understanding the Nature of Atoms

• If you cut a piece of graphite from the tip of a pencil into smaller and smaller pieces, how far could you go?

• You would eventually end up with atoms (translates to “indivisible” in greek) of pure carbon.

• You can not divide a carbon atom into smaller pieces and still have carbon

• An atom is the smallest identifiable unit of an element

• The theory that all matter is composed of atoms grew out of two primary laws

1. Law of constant composition2. Law of conservation of mass

Matter

The relative amounts of each element in a given substance are always the

same, regardless of how the substance was made.

Law of Constant Composition

• For a molecule, AB:

• mass A + mass B = total mass AB• %A + %B = 100• %A =

• Example: We analyze 1.630 g of CaS and find that it’s 0.906 g Ca. Find the mass of S? Find the mass% of Ca and S?

* 55.6% Ca, 44.4% S* This means that all pure CaS in the universe has the same composition as calculated above, regardless of how it was made or where it was found.

Law of Constant Composition

In a chemical reaction, atoms are not created or destroyed, only rearranged. The total mass of substances present before a reaction is equal to the total mass after.

Law of Conservation of Mass

• We have established that matter is comprised of atoms. But what are atoms made of?

• In the 1800’s, physicists conducted numerous experiments which revealed that the atom itself is made up of even smaller, more fundamental particles.

• The three types of sub-atomic particles that make up the atom are known as:• electrons• protons• neutrons

Atomic Structure

J.J. Thomson’s Cathode Ray Experiment (late 1800’s)

• No Electric Field• With Electric Field

Discovery of the Electron

Applying voltage to a metal cathode produces a beam of particles. The beam can be deflected by electric fields towards a positive pole. Mass of cathode plate does not change during this process. What does this mean?

• Atoms are charge neutral. If electrons reside within an atom, then an equivalent number of positive charges must also exist, appropriately named protons.

• How do all these charges coexist?

• Thomson proposed the very first theoretical model of the atom, the so-called plum pudding model (PPM) shown to the right.• Electrons reside in a sea of

uniform positive charge

Plum Pudding Model

• Ernest Rutherford sought to test the PPM using the gold foil experiment (below)

• A beam of positively charged α-particles were focused on a very thin sheet of gold

• Based on the PM model, this beam would pass right through the gold foil. In actuality, the beam was deflected at odd angles, with some α-particles bouncing directly back!!

Protons and The Nucleus

True model of the atom is a dense, positively charged, proton-loaded nucleus surrounded by a sparse electron cloud! The vast majority of an atom’s mass is contained within the nucleus.

The Atom

• Rutherford’s model was incomplete. For example, a hydrogen atom has one proton and one electron, but is only ¼th the mass of a helium atom which has two electrons and two protons.

• If all of the mass of an atom comes from its sub-atomic particles, how do we explain the unaccounted for mass?

• The answer is neutrons, particles that are equal in mass to protons, but with no electrical charge.

Neutrons

Subatomic Particles and Their Relative Masses & Charges

Particle Relative Charge Mass (amu)Proton +1 1.007

Neutron 0 1.008

Electron -1 .000548

• The number of protons in an atom is called the atomic number. An element is defined by its atomic number. (ex. only carbon has 6 protons)

• For a given element, the number of protons DOES NOT CHANGE

• In a neutral atom, the number of protons is equal to the number of electrons.

6

CCarbon

12.0107

Atomic #

Elemental Symbols

6

CCarbon

12.0107 Mass #

Elemental Symbols

• The mass number of an element is the sum of its protons and neutrons. The mass #’s listed on the periodic table are averages .

• The unit of atomic mass is the amu, or atomic mass unit, which is equal to 1/12 of the mass of a carbon atom having 6 neutrons.

• These averages are used because numerous variations of elements called isotopes exist in nature.

• Isotopes are variations of elements having different numbers of neutrons. Isotope symbols are shown below for the three isotopes of nitrogen with their % abundances in nature. The 35Cl and 37Cl isotopes have 18 and 20 neutrons, respectively.

(75.78%)

(24.22%)

mass number

atomic number

Isotopes

For any sample containing chlorine atoms, you will have the mix of isotopes shown.

Transitional PageAvg. atomic mass is obtained using the % abundance and the isotope mass.

𝐴𝑣𝑒𝑟𝑎𝑔𝑒𝑎𝑡𝑜𝑚𝑖𝑐𝑚𝑎𝑠𝑠=∑ (𝑖𝑠𝑜𝑡𝑜𝑝𝑒𝑚𝑎𝑠𝑠 ) 𝑥(𝑎𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒)

• Confirm the average mass of Cl shown on the periodic table.

Example

𝐴𝑣𝑒𝑟𝑎𝑔𝑒𝑚𝑎𝑠𝑠=(34.96885271amu ) (0.7578 )+(36.96590260)(0.2422)

𝒎𝒂𝒔𝒔𝒐𝒇 𝑪𝒍𝟏𝟕𝟑𝟓 𝒊𝒏𝒂𝒎𝒖 𝒎𝒂𝒔𝒔𝒐𝒇 𝑪𝒍𝟏𝟕

𝟑𝟕 𝒊𝒏𝒂𝒎𝒖

𝒂𝒃𝒖𝒏𝒅𝒂𝒏𝒄𝒆𝒐𝒇 𝑪𝒍𝟏𝟕𝟑𝟓 𝒊𝒏𝒂𝒎𝒖 𝒂𝒃𝒖𝒏𝒅𝒂𝒏𝒄𝒆𝒐𝒇 𝑪𝒍𝟏𝟕

𝟑𝟓 𝒊𝒏𝒂𝒎𝒖

𝐴𝑣𝑒𝑟𝑎𝑔𝑒𝑚𝑎𝑠𝑠=35.453𝑎𝑚𝑢

• Boron has two isotopes, 10B and 11B. Using the given isotope masses, determine the % abundances of each isotope.

• Hint: As decimals, the abundances must add to 1.ISOTOPE % A Mass (amu)

10.013

11.009

Group Work

• The nuclei of most naturally occurring isotopes are very stable, despite the massive repulsive forces that exist between the protons in the nucleus.

• A strong force of attraction between neutrons and protons known as the nuclear force counteracts this repulsion.

• As the number of protons increases, more neutrons are required to stabilize the atom. Stable nuclei up to atomic number 20 have equal numbers of protons and neutrons.

• For nuclei with atomic number above 20, the number of neutrons exceeds the protons to create a stable nucleus.

Proton-Neutron Ratio and Radioactivity

• Radioactive isotopes are unstable (high in energy). This instability is attributed to a neutron/proton ratio that is either too high or too low.

• To become stable, they spontaneously release particles or radiation to lower their energy.

• This release of energy is called radioactive decay.

Proton-Neutron Ratio and Radioactivity

Property α β γ

Charge 2+ 1- 0Mass 6.64 x 10-24 g 9.11 x 10-28 g 0Emitted Radiation Type

2 protons and 2 neutrons ()

High energy electron.

Pure energy (Radiation)

Penetrating Power Low. Stopped by paper. Blocked by skin.

Moderate. Stopped by aluminum foil. (10α)

High. Can penetrate several inches of lead. (10000α)

• The three most common types of radioactive decay are alpha, beta, and gamma

Radioactivity

• For example, the isotope undergoes alpha decay to decrease its n/p ratio:

• The Thorium-234 isotope then undergoes beta decay which lowers the ratio even more:

– In beta decay, a neutron is converted to a proton and an electron. This causes the proton count to increase:

• Gamma (γ) decay usually accompanies α or β decay to release residual excess energy. γ is not shown in equations.

𝑈→ h𝑇 + 𝐻𝑒24

90234

92238

h𝑇 → 𝑃𝑎+ 𝑒−10

91234

90234

𝑛→ 𝑝+ 𝑒− 10

11

01

Radioactivity

Applications of Radiochemistry: Carbon Dating

𝑵𝟕𝟏𝟒 + 𝒏→𝟎𝟏 𝑵𝟕𝟏𝟓

𝑵𝟕𝟏𝟓 → 𝑪𝟔𝟏𝟒 + 𝑯𝟏𝟏

𝑪𝟔𝟏𝟒 → 𝑵+ 𝑒−10

𝟕𝟏𝟒

Beta decay!Half life = 5700 yrs

Applications of Radiochemistry: Radiomedicine

Radioactive tracers can be linked to chemical compounds to allow doctors to monitor physiological processes. These compounds ‘glow’ upon decay via γ-decay

• Organ malfunction can be indicated if a radioisotope is taken up either too little or too much.

• Can monitor blood flow• Intestinal blockages can be detected by accumulation of

tracer.• Tumor detection

131I treatment for thyroid cancer

• Thus far, we’ve learned that each element has an exact number of protons. – For example, Hydrogen has only one proton. If you force a

second proton onto the atom, you no longer have hydrogen… you now have Helium.

• We have also learned that atoms can have variable numbers of neutrons (isotopes).

• Next, we will discuss ions.

Ions

• Ions are electrically charged atoms, resulting from the gain or loss of electrons.

• Positively charged ions are called cations. You form cations when electrons are lost

• Negatively charged ions are called anions. You form anions when electrons are gained

Ions

• A cation is named by adding the word “ion” to the end of the element name

• Anions are named by adding the suffix –ide to the end of an element

𝑳𝒊+¿ ¿

𝑵𝒂+¿ ¿

𝑴𝒈𝟐+¿ ¿

𝑨𝒍𝟑+¿¿

Lithium ion

Sodium ion

Magnesium ion

Aluminum ion

𝑪𝒍−

𝑺𝟐−

𝑶𝟐−

𝑷 𝟑−

Chloride

Sulfide

Oxide

Phosphide

Ion Nomenclature

• Fill in the missing information below

ISOTOPE P N E

2-

?? 13 14 104+ 95

Group Work

• Atoms are very tiny particles. As such, very small masses of a substance can contain extremely large numbers of atoms. For example, consider a single atom of gold, which has an average mass of 196.97 amu. The mass would be:

• The number of gold atoms in just a mere 1.0 μg of gold would be:

Moles

196.97𝑎𝑚𝑢𝑎𝑡𝑜𝑚 𝑥 1.66054 𝑥10

− 27𝑘𝑔𝑎𝑚𝑢 𝑥 10

3𝑔𝑘𝑔 𝑥1𝑎𝑡𝑜𝑚=3.2708 𝑥10− 22𝑔

!!

• As you can see, we need a unit of atoms that allows us to express large quantities.• The unit that we use is called the mole. A mole of atoms

is equal to the number of atoms in exactly 12g of .• This value, Avogadro’s number, NA, is 6.022 x 1023

• We can show that a mole of a substance has a mass equal to the atomic mass, in g. This is called the molar mass.

• 1.0079g of H has the same exact number of atoms as 196.97 g of Au !!!

Moles

• For molecules, the molar mass is the sum of the masses of each element.

• Examples: • Calculate the moles of water in a 12.5 g sample of

water.• Calculate the molecules of water in the sample.• Calculate the total number of hydrogen atoms in the

sample.

Working with Molar Masses

𝑀𝑀𝑜𝑓 𝐶𝑂2:(12. 011 𝑔𝐶𝑚𝑜𝑙 )+2 (16 𝑔𝑂𝑚𝑜𝑙 )≈44𝑔𝐶𝑂2

𝑚𝑜𝑙

𝑀𝑀𝑜𝑓 𝐶𝑎 ¿