1. Chapter 4 Atomic Structure

2. OBJECTIVES: Understand the history of atom
Differentiate between different models of the atom
Identify the no. of protons and neutrons in a neutral atom based on atomic number Z and mass number A

3. KEY TERMS Atom Atomic mass Atomic mass unit Atomic number Cathode ray
Dalton’s atomic theory
Electron
Group
Isotopes
Mass number
Neutron
Nucleus
Period
Periodic table
Proton

4. SECTIONS 4.1 Defining the Atom 4.2 Structure of the Nuclear Atom
4.3 Distinguishing Among Atoms

5. 4.1
4.1 Defining the Atom Early Models of the Atom Democritus’s Atomic Philosophy 400 BC (430?)
“Everything is made up of a few simple parts called atomos.” Atomos means “uncuttable” in Greek. He envisioned atomos as small, solid particles of many different sizes and shapes. Democritus’s ideas were limited because they didn’t explain chemical behavior and they lacked experimental support. His ideas were rejected because Aristotle supported *the “earth, air, water, and fire” concept of matter.
Democritus believed that atoms were solid particles that are indivisible and indestructible. 5

6. ALL MATTER IS COMPOSED OF EXTREMELY SMALL PARTICLES CALLED ATOMS
ATOMS OF A GIVEN ELEMENT ARE IDENTICAL IN SIZE, MASS, AND OTHER PROPERTIES; ATOMS OF DIFFERENT ELEMENTS DIFFER IN SIZE, MASS, & OTHER PROPERTIES
ELEMENT 1
ELEMENT 2
ELEMENT 3
ELEMENT 4

7. ATOMS CANNOT BE SUBDIVIDED, CREATED, OR DESTROYED
ATOMS OF DIFFERENT ELEMENTS COMBINE IN SIMPLE WHOLE # RATIOS TO FORM CHEM COMPDS
IN CHEMICAL RXNS, ATOMS ARE COMBINED, SEPARATED, OR REARRANGED + +

8. Sizing up the Atom
The radii of most atoms fall within the range of 5x10-11m to 2 x 10-10m. Individual atoms are observable with instruments such as scanning tunneling microscopes.

9. Discovery of the Electron
In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle: the electron

10. 4.2
Subatomic Particles
Cathode Ray Tube
A cathode ray is deflected by a magnet.
A cathode ray is deflected by electrically charged plates.
In a cathode-ray tube, electrons travel as a ray from the cathode (-) to the anode (+). A television tube is a specialized type of cathode-ray tube.
Thomson concluded that a cathode ray is a stream of electrons. Electrons are parts of the atoms of all elements. 10

11. 4.2
Subatomic Particles
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. 11

12. Plum Pudding Model JJ Thomson – 1897 Most of the atom was pos. charged
Atoms had negative charges embedded in pos. charged material

13. Rutherford’s Model Ernest Rutherford – 1911
Gold Foil Experiment – Rutherford fired positively charged particles at a sheet of gold foil
Rutherford discovered the nucleus as a result of his experiment

14. The Atomic Nucleus 4.2 Rutherford’s Gold-Foil Experiment
In 1911, Rutherford and his coworkers at the University of Manchester, England, directed a narrow beam of alpha particles at a very thin sheet of gold foil.
Alpha particles scatter from the gold foil. 14

15. The Atomic Nucleus 4.2 The Rutherford Atomic Model
Ernest Rutherford concluded that the atom is mostly empty space. All the positive charge and almost all of the mass are concentrated in a small region called the nucleus.
The nucleus is the tiny central core of an atom and is composed of protons and neutrons.
In the nuclear atom, the protons and neutrons are located in the nucleus.
The electrons are distributed around the nucleus and occupy almost all the volume of the atom. 15

16. 4.3 Distinguishing Among Atoms
Just as apples come in different varieties, a chemical element can come in different “varieties” called isotopes. 16

17. The Periodic Table—A Preview
4.3
The Periodic Table—A Preview
A periodic table is an arrangement of elements in which the elements are separated into groups based on a set of repeating properties.
A periodic table allows you to easily compare the properties of one element (or a group of elements) to another element (or group of elements).
Each horizontal row of the periodic table is called a period. Within a given period, the properties of the elements vary as you move across it from element to element.
Each vertical column of the periodic table is called a group, or family. Elements within a group have similar chemical and physical properties. 17

18. The Periodic Table—A Preview
4.3
The Periodic Table—A Preview
A Period
Period goes across 18

19. 4.3
Isotopes
Isotopes are atoms that have the same number of protons but different numbers of neutrons.
Because isotopes of an element have different numbers of neutrons, they also have different mass numbers.
Despite these differences, isotopes are chemically alike because they have identical numbers of protons and electrons. 19

20. E Isotopic notation Atomic number = # p A Z
Mass number = p + n (in the nucleus)
Element symbol A E Z
Atomic number = # p
Elements are put in order of atomic number on the periodic table.
Ex: An atom of carbon with 7 neutrons:
An atom of lead with 125 neutrons:
13C 6
207Pb 82

21. Mass Number Au is the chemical symbol for gold. 4.3 How many protons,
electrons, and neutrons
does a gold atom have?
The atomic number is 79. Therefore, there are 79 protons and 79 electrons. The mass number is 197, which is the total number of protons and neutrons. Therefore, = 118 neutrons.
Au is the chemical symbol for gold. Applying Concepts How many electrons does a gold atom have? 21

22. Practice #2 (on worksheet)22

23. 4.3
Atomic Mass
Some elements and their isotopes
An atomic mass unit (amu) is defined as one twelfth of the mass of a carbon-12 atom.
It is useful to to compare the relative masses of atoms to a standard reference isotope. Carbon-12 is the standard reference isotope. Carbon-12 has a mass of exactly 12 atomic mass units. 23

24. Atomic Mass
4.3
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. 24

25. Atomic Mass
4.3
To calculate the atomic mass of an element, multiply the mass of each isotope by its natural abundance, expressed as a decimal, and then add the products.
For example, carbon has two stable isotopes:
Carbon-12, which has a natural abundance of 98.89% and a mass of amu
Carbon-13, which has a natural abundance of 1.11% and a mass of amu.
Silver is found in two isotopes with atomic masses and amu, respectively. The first isotope represents 51.82% and the second 48.18%. Determine the average atomic mass of silver.
( )(.5182)=
( )(.4818)=
= amu 25

27. Planetary Model – Bohr Model
Neils Bohr – 1913
Electrons move in definite orbits
Orbits are referred to as energy levels

29. 4.2 The Structure of the Atom
Atoms consist of subatomic particles
Protons
Positive charge
Located in the nucleus
Neutrons
No charge
Electrons
Negative Charge
Located in the electron cloud

31. Atomic Number and Mass Number
Atomic Number Z
# of protons in the nucleus of an atom
Different elements have different numbers of protons
Mass Number A
Sum of an atoms protons and neutrons
n0 = A - Z

32. Isotopes
Isotopes are atoms of the same element with different numbers of neutrons and different mass numbers.