1. Atomic Structure & Nuclear Chemistry ‘new’ book: Ch 2
Unit 2
Atomic Structure & Nuclear Chemistry
‘new’ book: Ch 2

2. Early Theories of Matter
Democritus ( B.C.) proposed & believed that
Matter was not infinitely divisible
Matter is made up of tiny particles called atomos (“uncuttable”)
Atoms could not be created, destroyed, or further divided
No protons, electrons, or neutrons
Atom was solid and indestructible
Ideas lacked experimental support because scientific testing was unknown at that time
Idea was rejected by Aristotle and thus lost for 2000 years.
Greek society was slave based; Beneath famous to work with hands, did not experiment, Greeks settled disagreements by argument
Aristotle was more famous; He won! His ideas carried through middle ages. Alchemists change lead to gold
Democritus’ model of atom
“Billiard Ball” Model
Solid and INDESTRUCTABLE

3. Development of Chemistry
Robert Boyle
Beginnings of modern chemistry were seen in the 16th and 17th centuries, where great advances were made in metallurgy, the extraction of metals from ores.
In the 17th century, Robert Boyle described the relationship between the pressure and volume of air and defined an element as a substance that cannot be broken down into two or more simpler substances by chemical means.

4. Development of Chemistry
During the 18th century, Priestley discovered oxygen gas and the process of combustion where carbon-containing materials burn vigorously in an oxygen atmosphere.
Priestley

5. Development of Chemistry
Lavoisier
In the late 18th century, Lavoisier wrote the first modern chemistry text. His most important contribution was the law of conservation of mass, which states that in any chemical reaction, the mass of the substances that react equals the mass of the products that are formed.
He is known as the father of modern chemistry.

6. Development of Chemistry
In the 19th century, John Dalton revised Democritus's ideas based upon the results of scientific research he conducted
Led to Dalton’s atomic theory
The real nature of atoms
Dalton

7. Dalton’s Atomic Theory
Elements are composed of tiny 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.
If copper atoms were placed side by side, they would form a line 1 cm long. What is this number of atoms written in scientific notation?

8. Dalton’s Atomic Theory
Atoms of different elements can physically mix together or can chemically combine with one another in simple whole-number ratios to form compounds.

9. Dalton’s Atomic Theory
Chemical reactions occur when atoms are separated, joined, or rearranged.
Atoms of one element, however, are never changed into atoms of another element as a result of a chemical reaction.
Dalton’s atomic theory is essentially correct, with four minor modifications:
1. Not all atoms of an element must have precisely the same mass.
2. Atoms of one element can be transformed into another through nuclear reactions.
3. The composition of many solid compounds are somewhat variable.
4. Under certain circumstances, some atoms can be divided
(split into smaller particles: i.e. nuclear fission).

10. Legos are Similar to AtomsH O H2 O2
H2O +
Legos can be taken apart and built into many different things.
Atoms can be rearranged into different substances.

11. Foundations of Atomic Theory
Law of Conservation of Mass
Mass is neither destroyed nor created during ordinary chemical reactions.
Law of Definite Proportions
The fact that a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound.
Lavoisier (credited with Law of Conservation of Mass).
Proust (credited with Law of Definite Proportions).
Dalton (credited with Law of Multiple Proportions).

12. Law of Definite Proportions
Joseph Louis Proust
(1754 – 1826)
Each compound has a specific ratio of elements
It is a ratio by mass
EX: Water is always 8 grams of oxygen for every one gram of hydrogen
Photo pg 100 Ihde text (Edgar Fahs SmithCollection)
Joseph Louis Proust ( ), French chemist given credit for law of definite composition.
Whether synthesized in the laboratory or obtained from various natural sources, copper carbonate always has the same composition. Analysis of this compound led Proust to formulate the law of definite proportions.

13. Law of Definite Proportions
Whether synthesized in the laboratory or obtained from various natural sources, copper (II) carbonate always has the same composition.
Analysis of this compound led Proust to formulate the law of definite proportions.
40 g of oxygen
Much of Dalton’s theory is accepted today. One important change, however, is that atoms are now known to be divisible (not known until the late 1800’s). They can be broken down into even smaller, more fundamental particles. Dozens of kinds of subatomic particles are unleashed when powerful devices known as atom smashers are used to fracture atoms. You will now learn about 3 kinds of subatomic particles.
103 g of copper (II) carbonate
53 g of copper
10 g of carbon + +

14. Structure of Atoms Scientists began to wonder what an atom was like.
Was it solid throughout with no internal structure or was it made up of smaller, subatomic particles?
It was not until the late 1800’s that evidence became available that atoms were composed of smaller parts.
Some of the details of Dalton’s atomic theory require more explanation.
Elements: As early as 1660, Robert Boyle recognized that the Greek definition of element (earth, fire, air, and water) was not correct. Boyle proposed a new definition of an element as a fundamental substance, and we now define elements as fundamental substances that cannot be broken down further by chemical means. Elements are the building blocks of the universe. They are pure substances that form the basis of all of the materials around us. Some elements can be seen in pure form, such as mercury in a thermometer; some we see mainly in chemical combination with others, such as oxygen and hydrogen in water. We now know of approximately 116 different elements. Each of the elements is given a name and a one- or two-letter abbreviation. Often this abbreviation is simply the first letter of the element; for example, hydrogen is abbreviated as H, and oxygen as O. Sometimes an element is given a two-letter abbreviation; for example, helium is He. When writing the abbreviation for an element, the first letter is always capitalized and the second letter (if there is one) is always lowercase.
Atoms: A single unit of an element is called an atom. The atom is the most basic unit of the matter that makes up everything in the world around us. Each atom retains all of the chemical and physical properties of its parent element. At the end of the nineteenth century, scientists would show that atoms were actually made up of smaller, "subatomic" pieces, which smashed the billiard-ball concept of the atom.
Compounds: Most of the materials we come into contact with are compounds, substances formed by the chemical combination of two or more atoms of the elements. A single “particle” of a compound is called a molecule. Dalton incorrectly imagined that atoms “hooked” together to form molecules. However, Dalton correctly realized that compounds have precise formulas. Water, for example, is always made up of two parts hydrogen and one part oxygen. The chemical formula of a compound is written by listing the symbols of the elements together, without any spaces between them. If a molecule contains more than one atom of an element, a number is subscripted after the symbol to show the number of atoms of that element in the molecule. Thus the formula for water is H2O, never HO or H2O2.
The idea that compounds have defined chemical formulas was first proposed in the late 1700s by the French chemist Joseph Proust. Proust performed a number of experiments and observed that no matter how he caused different elements to react with oxygen, they always reacted in defined proportions. For example, two parts of hydrogen always reacts with one part oxygen when forming water; one part mercury always reacts with one part oxygen when forming mercury calx. Dalton used Proust’s Law of Definite Proportions in developing his atomic theory.
The law also applies to multiples of the fundamental proportion, for example:
In both of these examples, the ratio of hydrogen to oxygen to water is 2 to 1 to 1. When reactants are present in excess of the fundamental proportions, some reactants will remain unchanged after the chemical reaction has occurred.
The story of the development of modern atomic theory is one in which scientists built upon the work of others to produce a more accurate explanation of the world around them. This process is common in science, and even incorrect theories can contribute to important scientific discoveries. Dalton, Priestley, and others laid the foundation of atomic theory, and many of their hypotheses are still useful. However, in the decades after their work, other scientists would show that atoms are not solid billiard balls, but complex systems of particles. Thus they would smash apart a bit of Dalton’s atomic theory in an effort to build a more complete view of the world around us.
Source:

15. Radioactivity
One of the pieces of evidence that atoms are made of smaller particles came from the work of Marie Curie ( ).
She discovered radioactivity, the spontaneous disintegration of some elements into smaller pieces.
1896, Becquerel discovered that certain minerals emitted a new form of energy.
Becquerel’s work was extended by Pierre and Marie Curie, who used the word radioactivity to describe the emission of energy rays by matter.
Rutherford, building on the Curies’ work, showed that compounds of elements emitted at least two distinct types of radiation. One was readily absorbed by matter and consisted of particles that had a positive charge and were massive compared to electrons. These particles were called α particles. Particles in the second type of radiation were called β particles and had the same charge and mass-to-charge ratio as electrons.
A third type of radiation, γ rays, was discovered later and found to be similar to a lower energy form of radiation called X -rays.
Three kinds of radiation – α particles, β particles and γ rays
1. Distinguished by the way they are deflected by an electric field and by the degree to which they penetrate matter
2. α particles and β particles are deflected in opposite directions; α particles are deflected to a much lesser extent because of
their higher mass-to-charge ratio.
3. γ rays have no charge and are not deflected by electric or magnetic fields.
4. α particles have the least penetrating power, and γ rays are able to penetrate matter readily.
Marie Curie

16. Discovery of the Electron
J.J. Thomson (1856 – 1940) performed experiments that involved passing electric current through gases at low pressure
He sealed the gases in glass tubes fitted at both ends with metal disks called electrodes
Electrodes were connected to a source of high-voltage electricity
In 1897 J. J. Thomson, an English physicist, conducted a series of experiments on cathode rays produced in cathode ray tubes. These tubes contain gas and are partially evacuated. After observing that the beam of light in the cathode ray tube is attracted to a positive charge and repelled by a negative charge he concluded that the rays consist of a stream of small, electrically negatively charged particles which have a mass over a thousand times less than that of a hydrogen atom. Thomson has discovered the electron. From this point onward, it becomes increasingly clear that atoms are not fundamental particles, but in fact are made up of smaller particles.
Thomson reasoned that since electrons could be produced from electrodes made of various types of metals, all atoms must contain electrons. Since atoms were known to be electrically neutral, Thomson further assumed that atoms also must contain some positive charge. As a result of his experiments, Thomson was able to measure the charge to mass ratio of the electron; he could not however, measure accurately the charge or mass independently.
Thomson

17. Cathode-Ray Experiment
One electrode, the anode, became positively charged
The other electrode, the cathode, became negatively charged
A glowing beam formed between the 2 electrodes (called a cathode ray)

18. Cathode-Ray Experiment
voltage
source - +
vacuum tube
metal disks

19. Cathode-Ray Experiment
voltage
source ON -
OFF +
Passing an electric current makes a beam appear
to move from the negative to the positive end

20. Cathode-Ray Experiment
voltage
source ON -
OFF + + -
By adding an electric field…
he found that the moving particles were negative.

21. Thomson’s Findings
Cathode rays are attracted to positively charged metal plates and repelled by negatively-charged plates
He proved that atoms contain tiny negative particles (electrons) and concluded that ALL atoms must contain these negative particles.
He knew that atoms were neutral in charge and deduced that there must be a positive charge within the atom.

22. Plum-Pudding Model
J.J. Thomson ( ) proposed a model of the atom with subatomic particles (1903).
This model was called the plum-pudding or raisin pudding model of the atom.
Thomson won the Nobel Prize in 1906 for characterizing the electron.
Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model.

23. Discovery of the Proton
Goldstein discovered the proton using the cathode ray tube in a similar way as did Thomson and the electron

24. Discovery of the Nucleus
Ernest Rutherford ( ) learned physics in J.J. Thomson’ lab.
Noticed that ‘alpha’ particles were sometimes deflected by something in the air.
Alpha particles are helium nuclei
Alpha particles are positively-charged
Gold-foil experiment
Ernest Rutherford received the Nobel Prize in chemistry (1908) for his work with radioactivity.
Rutherford

25. Rutherford’s Gold Foil Experiment
Alpha particles were fired at a thin sheet of gold foil
Particle-hits on the detecting screen (film) are recorded
In 1909 Rutherford undertook a series of experiments
He fired a (alpha) particles at a very thin sample of gold foil
According to the Thomson model the a particles would only
be slightly deflected

26. Florescent Screen Lead block Uranium Gold Foil
Ernest Rutherford English physicist. (1910)
Wanted to see how big atoms are.
Used radioactivity, alpha particles - positively charged pieces given off by uranium.
Shot them at gold foil which can be made a few atoms thick.
When the alpha particles hit a florescent screen, it glows.

27. What he expected…

28. he thought the mass was evenly distributed in the atom. Because- - - - -

29. -

30. What he got…

31. Expected and Actual Results of Rutherford’s Experiment
Plum-pudding atom + -
Alpha particles
Nuclear atom
Nucleus
Thomson’s model
Rutherford’s model

32. Try it Yourself!
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?

33. The Answers
Target #1
Target #2

34. Rutherford’s Findings
Most of the particles passed right through
A few particles were deflected
VERY FEW were greatly deflected
“Like howitzer shells bouncing off of tissue paper!”
Conclusions:
The atom is mostly empty space
The nucleus is small
The nucleus is dense
The nucleus is positively charged

35. Size of the Nucleus
If an atom is as large as a stadium, then the nucleus is about the size of a fly in the center of the stadium!!!

36. Nuclear Model
n +

37. Discovery of the Neutron
James Chadwick + +
Experiments showed that the helium nucleus, for instance, exerted exactly twice as much electrical force on an electron as a nucleus of hydrogen, the smallest atom, and this was explained by saying that helium had two protons to hydrogen's one. The trouble was that according to this model, helium would have two electrons and two protons, giving it precisely twice the mass of a hydrogen atom with one of each. In fact, helium has about four times the mass of hydrogen.
In 1932 James Chadwick bombarded beryllium-9 with alpha particles, carbon-12 atoms were formed, and neutrons were emitted.
Dorin, Demmin, Gabel, Chemistry The Study of Matter 3rd Edition, page 764

38. The Modern View of Atomic Structure
The atom contains:
electrons
protons: found in the nucleus, they have a positive charge equal in magnitude to the electron’s negative charge.
neutrons: found in the nucleus, virtually same mass as a proton but no charge.
In the years since Thomson and Rutherford, a great deal has been learned about atomic structure. Just an intro here… The simplest view of the atom is that it consists of a tiny nucleus (diameter of about cm) and electrons that move about the nucleus at an average distance of about 10-8 cm from it. The tiny nucleus accounts for almost all the atom’s mass. A piece of nuclear material about the size of a pea would have a mass of 250 million tons.
The chemistry of an atom mainly results from its electrons.

39. The Mass and Charge of the Electron, Proton, and Neutron
Why different atoms have different chemical properties lies in the the number and the arrangement of the electrons. Key point is that although the mass of an atom is concentrated in its nucleus, most of the volume of the atom is taken up by electron movement. The electrons are the parts that intermingle when atoms combine to form molecules. Therefore, the number of electrons possessed by a given atom greatly affects its ability to interact with other atoms. Since electrons are involved in bonding, we need a way to symbolize the number of protons, neutrons, and electrons in an atom…

40. The Chemists’ Shorthand: Nuclear Symbols39
Mass number  K
 Element Symbol 19
Atomic number 
Atomic # = Z
# neutrons = A - Z

41. Atomic Number Equal to the number of protons
Equal to the number of electrons in an atom
Determines the element!

42. Mass Number mass # = protons + neutrons always a whole number
NOT on the Periodic Table!
How do you find number of neutrons?

43. Ions Atoms that have lost or gained electrons
Cation – positive ion (lost electrons)
Example: How many electrons does Na1+ have?
Anion – negative ion (gained electrons)
Example: How many electrons does S2- have?

44. Cs 133 55 PRACTICE WITH IONS 1+ Find the number of protons
number of neutrons
number of electrons
Atomic number
Mass number
= 55
= 78 Cs
133 1+
= 54 55
= 55
= 133
Cesium ion 48

45. C C Isotopes Atoms of the same element with different mass numbers.
carbon carbon-14
Isotopes are atoms with the same number of protons but different number of neutrons 12 6 14 6

46. PRACTICE
Atomic Number & Mass Number WS

47. Calculating Relative Atomic Mass

48. Relative Atomic Mass https://www.youtube.com/watch?v=vqTg4cYwHXY
12C atom = × g (*watch above video!)
atomic mass unit (amu)
1 amu = 1/12 the mass of a 12C atom
Small masses are not easy to work with, so the atomic mass unit (amu) was developed
Instead it is more useful to compare the relative masses of atoms using a reference isotope as a standard. The isotope chosen is a carbon-12 atom. This isotope was assigned a mass of exactly 12 amu
One atomic mass unit (amu) is equal to 1/12 the mass of a carbon-12 atom
The mass of 1 amu is nearly equal to the mass of one proton or neutron
1 p = amu 1 n = amu 1 e- = amu

49. Average Atomic Mass For help: http://www. docbrown
weighted average of all isotopes
on the Periodic Table
round to 2 decimal places!
Avg.
Atomic
Mass
= (mass x relative abundance) + (mass x relative abundance) +….
% divided by 100 or part over whole

50. Average Atomic Mass Example #1Mg 12
24.31
Magnesium has 3 isotopes % Mg-24 with a mass of amu, 10.00% Mg-25 with a mass of amu, and the rest Mg-26 with a mass of amu. What is the average atomic mass of magnesium?
Isotope
Percent
Abundance
Mass
Relative Mass
Mg-24
78.99
Mg-25
10.00
Mg-26
Atomic mass is not a whole number because it is an average.
This is why their are the decimal numbers on the periodic table.
11.01
24.31 amu
If not told otherwise, the mass of the isotope is the mass number in amu.

51. Average Atomic Mass Example #2
EX: Calculate the avg. atomic mass of oxygen if its abundance in nature is 99.76% 16O, 0.04% 17O, and 0.20% 18O.
Avg.
Atomic
Mass
= amu

52. What’s the difference between mass number and average atomic mass?
Mass number- specifically about one isotope
Average atomic mass- includes the masses of all the different isotopes for that atom

53. Atomic Structure ATOMS IONS ISOTOPES Differ by number of protons
Differ by number of electrons
ISOTOPES
Differ by number of neutrons

54. Intro to Periodic Table (Book: Section 2.3)
Distinguish between:
Groups and periods-
Group # of ‘A’ elements gives valence e- (available for bonding); Can use this to determine the ‘charge’ (oxidation #) the atom needs to form a perfect ‘octet’
Period # gives the # of energy levels where electrons ‘reside’.
Metals and Nonmetals (and metalloids)
Where are they on PT?

55. Intro to Bonding (Book: Section 2.4)
Distinguish between: Molecule and Ion
Molecule: Covalent bond between atoms. Share e-. Usually between nonmetals and H. Represent with molecular formula which shows ratios using subscripts. (Ex: CH4)
Molecular Structural formula: Each bond between atom represents a shared pair of e-.
Molecules are weak electrolytes (don’t ionize in solution). Ex- sugar and water
What is a binary compound or molecule?

56. Intro to Bonding- con’t
Ionic: Usually between metal and nonmetal. Give and take of e- to form cations and anions. Compounds held together by strong intermolecular forces. Most are strong electrolytes in water. Ex of ionic binary compound: NaCl
Some ionic compounds are ternary (more than 2 atoms): Ex-NaOH. Na is 1+ and the (OH) is 1-. The OH is actually covalently bonded and carries overall 1- charge, BUT the Na+ is ionically bonded to the (OH)-.

57. Note:
You don’t have to know charges/ oxidation #’s for this test, except for the 5 polyatomic ions given on next slide

58. Polyatomic Ions
Group of covalently bonded (share e-) atoms carrying a charge. Will learn how this is done in the next unit
Start learning some common polyatomic ions and their charges. For this test- see table 2.2 in your book and know the formula name and charge of these 5: ammonium, hydroxide, carbonate, sulfate, phosphate. You will have to know them all for the next test.

59. HW- ‘New’ Book Go over Summary Problem p. 47 (omit h and i)
From p (end of Ch 2): #3, 6, 8,10,14, 16, 18, 20, 22, 26, 44, 49, 50, 52
Make sure you do fill in notes from PPt (especially on the individuals and their contributions to chemistry)
Go over all worksheets and PPt’s