1. Atomic structure
Chapter 4

2. 4.1 defining the Atom Early models of the Atom
Dalton’s Atomic theory
Democritus’s Atomic philosophy
Greek 460 B.C. – 370 B.C.
Reasoned that atoms were indivisible and indestructible
Didn’t explain chemical behavior
Lacked experimental support because Democritus approach was not based on the scientific method
The modern process of discovery began with John Dalton in
By using the scientific method Dalton transformed Democritus ideas on Atoms into a scientific theory
Studied the ratios in which elements combined in chemical reactions

3. Daltons Atomic theory (Solid Sphere)
1. All elements are composed of tiny indivisible particles called atoms
2. Atoms of the same element are identical. The atoms of any one element are different from those of any other element
3. Atoms of different elements can physically mix together or can chemically combine in a simple whole– number ratio to from compounds
4. Chemical reactions occur when atoms are separated from each other, joined, or rearranged in a different combination. Atoms of one element, however, are never changed into atoms of another element as a result of a chemical reaction.

4. Sizing up the atom
Drop of mercury – it does not matter about the size because all drops have the same properties because they all are made of the same atoms
If you were to crush a penny which is pure copper, you would still have copper. If you were to keep crushing and make the dust particles smaller, you would come across a particle of copper that could no longer be divided.
The final particle is an atom
Atoms are very small
A copper coin the size of a penny contains 2x1022 atoms
By comparison earths population is 7x109 people

5. Microscopes
Despite atoms small size you can see them with scanning electron microscope
A beam of electrons is focused on the sample
Electron microscopes are capable of much higher magnification than light microscopes

6. 4.2 Structure of the Nuclear Atom Subatomic particles
The 3 kinds of subatomic particles are the electrons, protons, and neutrons

7. Subatomic particles J. J. Thompson – (Plum Pudding Model)
electrons are negatively charged subatomic particles
Cathode ray- one electrode, anode, became positively charged the other electrode became negatively charged, the cathode
The result was a glowing beam
azk
Deflected by electrically charged particles metal plates or magnetic fields
A positively charged plate attracts the cathode ray and a negatively charged plate respells it
Opposite charges attract and like charges repel

8. Protons and neutrons
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 are called protons.
In 1932, the English physicist James Chadwick (1891–1974) confirmed the existence of yet another subatomic particle: the neutron.
Neutrons are subatomic particles with no charge but with a mass nearly equal to that of a proton.

9. Properties of Subatomic Particles Relative mass (mass of proton = 1)
Symbol
Relative charge
Relative mass (mass of proton = 1)
Actual mass (g)
Electron e– 1–
1/1840
9.11  10–28
Proton p+ 1+ 1
1.67  10–24
Neutron n0
Atoms have no net electric charge; they are electrically neutral.
Electric charges are carried by particles of matter.
Electric charges always exist in whole-number multiples of a single basic unit; that is, there are no fractions of charges.
When a given number of negatively charged particles combines with an equal number of positively charged particles, an electrically neutral particle is formed.

10. Rutherford’s Gold-Foil Experiment (The Nuclear Atom)
In the experiment, a narrow beam of alpha particles was directed at a very thin sheet of gold.
Rutherford’s results were that most alpha particles went straight through, or were slightly deflected.
What was surprising is that a small fraction of the alpha particles bounced off the gold foil at very large angles.
Some even bounced straight back toward the source.

11. The Rutherford Atomic Model
He proposed that the atom is mostly empty space.
New theory - 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.
In the nuclear atom, the protons and neutrons are located in the positively charged nucleus. The electrons are distributed around the nucleus and occupy almost all the volume of the atom.
According to this model, the nucleus is tiny and densely packed compared with the atom as a whole.
If an atom were the size of a football stadium, the nucleus would be about the size of a marble.

12. 4.3 distinguishing atoms Atomic number and mass number
Elements are different because they contain different numbers of protons.
An element’s atomic number is the number of protons in the nucleus of an atom of that element.
The atomic number identifies an element.
The atomic number gives the number of protons, which in a neutral atom equals the number of electrons.
The total number of protons and neutrons in an atom is called the mass number.

13. Atoms of the First Ten Elements
Name
Symbol
Atomic number
Protons
Neutrons
Mass number
Electrons
Hydrogen H 1
Helium He 2 4
Lithium Li 3 7
Beryllium Be 5 9
Boron B 6 11
Carbon C 12
Nitrogen N 14
Oxygen O 8 16
Fluorine F 10 19
Neon Ne 20

14. Mass number
The composition of any atom can be represented in shorthand notation using atomic number and mass number.
The atomic number is the subscript.
The mass number is the superscript
Number of neutrons= mass # - atomic #
Au may be written as gold-197.

15. How many protons, electrons, and neutrons are in each atom?
a. Be b. Ne c. Na

16. Isotopes
How do these atoms differ?

17. Isotopes All have the same number of protons (10).
All have the same number of electrons (10).
But they each have different numbers of neutrons
Neon-20, neon-21, and neon 22 are three isotopes of neon.
Isotopes are atoms that have the same number of protons but different numbers of neutrons.
Despite these differences, isotopes are chemically alike because they have identical numbers of protons and electrons, which are the subatomic particles responsible for chemical behavior.
They have different masses

18. Why are atoms with different numbers of neutrons still considered to be the same element?

19. Distinguishing mass
The mass of even the largest atom is incredibly small.
Such data about the actual masses of individual atoms can provide useful information, but in general these values are inconveniently small and impractical to work with.
Instead, it is more useful to compare the relative masses of atoms using a reference isotope as a standard.
The reference isotope chosen is carbon-12.
This isotope of carbon has been assigned a mass of exactly 12 atomic mass units.
An atomic mass unit (amu) is defined as one-twelfth of the mass of a carbon-12 atom.

20. Natural Percent Abundance of Stable Isotopes of Some Elements
Name
Symbol
Natural percent abundance
Mass (amu)
Atomic mass
Hydrogen H
99.985
0.015
negligible
1.0078
2.0141
3.0160
1.0079
Helium He
0.0001
4.0026
Carbon C
98.89
1.11
12.000
13.003
12.011
Oxygen O
99.759
0.037
0.204
15.995
16.995
17.999
15.999
Chlorine Cl
75.77
24.23
34.969
36.966
35.453

21. Percent abundance
In nature, most elements occur as a mixture of two or more isotopes.
The atomic mass of an element is a weighted average mass of the atoms in a naturally occurring sample of the element.
A weighted average mass reflects both the mass and the relative abundance of the isotopes as they occur in nature.

22. Calculating atomic mass
Carbon has two stable isotopes: carbon-12, which has a natural abundance of percent, and carbon-13, which has a natural abundance of 1.11 percent.
The mass of carbon-12 is amu; the mass of carbon-13 is amu.
The atomic mass of carbon is calculated as follows:
Atomic mass of carbon = ( amu x ) + ( amu x )
= ( amu) + (0.144 amu)
= amu

23. Practice
Element X has two naturally occurring isotopes. The isotope with a mass of amu (10X) has a relative abundance of percent. The isotope with a mass of amu (11X) has a relative abundance of percent. Calculate the atomic mass of element X.

24. Calculate the atomic mass of bromine
Calculate the atomic mass of bromine. The two isotopes have atomic masses and relative abundance of amu (50.69%) and amu (49.31%)