1. Chapter 4 Atomic Structure
Ms. Wang
Lawndale High School

2. Section 4.1 - Defining the Atom
All matter is composed of particles called atoms
Atoms – the smallest particle of an element that retains it identity in a chemical reaction

3. Democritus’s Philosophy
Democritus was the first to suggest the existence of atoms being indivisible and indestructible
By using experimental methods, Dalton transformed Democritus’s ideas on atoms into a scientific theory

4. Dalton’s Atomic Theory
(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 simple whole-number ratios to form compounds.
(4.) 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.

5. Section 4.2 – The Structure of the Atom
There are three kinds of subatomic particles in an atom: electrons, protons, and neutrons
In 1897, J. J. Thomson discovered the electron (negatively charged subatomic particles) by using a cathode ray

6. Protons
Atoms were known to be electrically neutral, which meant that there had to be some positively charged matter to balance the negative charges

7. Rutherford’s Gold-Foil Experiment
Ernest Rutherford’s experiment disproved the plum pudding model of the atom and suggested that there was a positively charged nucleus (central core of an atom)

8. Conclusion of Rutherford’s Experiment
Atoms are mostly empty space, thus explaining the lack of deflection of most of the alpha particles
All the positive charge and almost all the mass of an atom are concentrated in a small region (nucleus)
Nucleus – tiny central core of an atom composed of protons and neutrons
Electrons are distributed around the nucleus and occupy almost all the volume of the atom (marble and football stadium)

9. Structure Of An Atom

10. Properties of Subatomic Particles
SYMBOL
CHARGE
Electron e- -1
Proton p+ +1
Neutron n0

11. Homework Section Assessment 4-1 #’s 4,5
Now on to the lab…The Mystery Box!!!
Each group will go their their home lab station and try to determine the shape of the object inside their box by moving the box around. After 2.5 minutes of explorations, record your observations and move to the next lab station.

12. Conclusion (guess what’s in the box)
Mystery Boxes
Purpose: To determine the shape of the object inside the box
Lab Station
Observations
Conclusion (guess what’s in the box) 1 2 3 4 5 6 7

13. Section 4.3 – Distinguishing Among Atoms
Elements are different because they contain different numbers of protons.
Atomic Number - the number of protons in the nucleus of an atom
*Remember since atoms are electrically neutral, the number of protons equals the number of electrons

14. Quick Practice… How many protons and electrons are in each atom?
1. Fluorine (atomic number = 9)
2. Calcium (atomic number = 20)
3. Aluminum (atomic number = 13)
How about these?
4. Boron
5. Neon
6. Magnesium

15. # of Neutrons = Mass # – Atomic #
Mass Number
Mass Number – the total number of protons and neutrons in an atom
Therefore…the number of neutrons in an atom is the difference between the mass number and the atomic number
# of Neutrons = Mass # – Atomic #

16. Shorthand Notation (You need to know this notation)
Mass Number Au
179 79
Atomic Symbol
Atomic Number

17. Practice Shorthand Notation…
How many protons, neutrons, and electrons are in each atom?
Atomic # Mass #
Beryllium (Be)
Neon (Ne)
Sodium (Na)
How many neutrons are in each atom?
1. Carbon-12
2. Fluorine-19
3. Sulfur 32 6 10 16

18. Isotopes
Isotopes – atoms that have the same number of protons, but different numbers of neutrons (also different mass numbers)
Write the following isotopes of oxygen:
1. Oxygen-16
2. Oxygen-17
3. Oxygen-18

19. Atomic Mass
Atomic Mass – weighted average mass of the atoms in a naturally occurring sample of the element
In order to calculate the atomic mass of an element:
(1.) Multiply the mass of each isotope by its natural abundance
(2.) Add the products together

20. Let’s practice… Calculate the atomic mass of the following element, X
The isotope 10X has a mass of amu and a relative abundance of 19.91%. The isotope 11X has a mass of amu and a relative abundance of 80.09%.
ANSWER = amu

21. More Practice…
The element copper has naturally occurring isotopes with mass numbers of 63 and 65. The relative abundance and atomic masses are 69.2% for mass = 62.93amu, and 30.8% for mass = 64.93amu. Calculate the average atomic mass of copper.
2. Calculate the atomic mass of bromine. The two isotopes of bromine have atomic masses and relative abundance of 78.92amu (50.69%) and 80.92amu (49.31%).

22. Preview of the Periodic Table
Periodic Table – an arrangement of elements in which the elements are separated into groups based on a set of properties
Period – horizontal rows of the periodic table (there are 7)
Group/Family – vertical columns of the periodic table
Elements within a group have similar chemical and physical properties

23. Chapter 4 Assessment Page 122
Homework
Chapter 4 Assessment Page 122
#’s 34 – 55, 59, 61, 64, 65, 71, 78, 81, 85, 88