1. Chapter 4
Atoms and Elements

2. Homework Assigned Problems (odd numbers only)
“Questions and Problems” to 4.65 (begins on page 96)
“Additional Questions and Problems” to (page )
“Challenge Questions” to 4.111, (page 124)

3. Elements Matter is anything with a mass and occupies space
Matter (in our world) is composed of combinations of about 100 basic substances called elements
109 elements have been discovered and isolated
88 are found in nature
21 are (synthetic) man-made
Oxygen most abundant element (by mass) on earth

4. “Element”
A pure substance that cannot be broken down into simpler substances by a chemical means
Single atom of that element
Sample of the element large enough to weigh on a scale
Generally referring to the presence of that element (compound), not necessarily in its free form
Chemical means are a reaction, an electric current, or a beam of light

5. Chemical symbols Each element has a unique symbol
One or two letter abbreviations
If two letters, the second is lower case
The letter symbol often corresponds to the name of the element
F = Fluorine
P = Phosphorous
Some symbols derived from the Latin or Greek names
Lead – Pb (plumbum)
Gold – Au (aurum)
Sodium – Na (natrium)
Two capital letters indicate two elements such as CO

6. Names of Symbols and Some Common Elements
Required: Know the name and symbol of some of the most common elements
Table 4.2 on p. 95 (know the names and symbols listed)
A periodic table will be given on each test or quiz
Required: Know how to use a periodic table to find needed information

7. Periodic Table
A chart of the elements with similar chemical properties arranged into vertical columns (groups)
Horizontal rows are called periods
The elements arranged (in rows) in order of increasing atomic mass (atomic number)
Main group elements are those in the columns labeled with numbers (1A-8A)
Transition elements are those in the columns labeled with the letter “B”

8. Main Group Elements Also called representative elements
The elements in the A-groups
First two columns (1A and 2A)
The last 6 columns (3A to 8A)
Easy to predict ionic structure

9. Transition Elements The elements in the B-groups
Middle block of elements (3B through 2B)
Includes the two groups at the bottom
Lanthanides and Actinides
Difficult to predict ionic structure

10. Classification of Elements
Certain groups of elements have their own special names due to the chemical similarity of the elements in them
Alkali Metals
Group 1A
Alkaline Earth Metals
Group 2A
Halogens
Group 7A
Noble Gases
Group 8A
Become familiar with these group names

11. The Periodic Table 1A 8A 1 2A 3A 4A 5A 6A 7A 2 3 4 5 6 7

12. Metals/Nonmetals Metals
Everything to the left of the metal/nonmetal barrier
Shiny solid, good conductor of electricity, ductile and malleable
Nonmetals
Everything to the right of the metal/nonmetal barrier
Dull appearance, not ductile or malleable not good conductors of electricity

13. Metalloids
Elements with properties intermediate between those metals and nonmetals
On the metal/nonmetal barrier
Have some physical properties of metals but some chemical properties of nonmetals
Semiconductors
Si, Ge, As, Sb, Te

14. The Atom
The smallest particle of an element that can exist and still have properties of that element
All atoms of a certain type are similar to one another and different from all other types
109 different types are known and each “type” is a different element

15. Dalton’s Atomic Theory (1808)
A set of five statements that summarize the modern scientific concept about atoms
All matter is made from small particles called atoms (109 different types)
All atoms of a given type are similar to one another and significantly different from all other types
A series of five statements that summarizes the modern day theory of the atom `

16. Dalton’s Atomic Theory
The number and arrangement of different types of atoms in a pure substance determine its identity
A chemical change is a combination, separation, or rearrangement of atoms (forms new substances)
Only whole atoms take part or result from any chemical reaction
A series of five statements that summarizes the modern day theory of the atom `

17. Dalton’s Atomic Theory
Atoms are indivisible in a chemical process (indestructible)
All atoms present at beginning are present at the end
Atoms are not created or destroyed, just rearranged
Atoms of one element cannot change into atoms of another element
Cannot turn Lead into Gold by a chemical reaction

18. Cathode Rays and Electrons
J.J. Thomson (1897) used a gas discharge tube to investigate a beam called a cathode ray
Determined that the ray was made of tiny negatively charged particles we call electrons
He determined the electrons were smaller than a hydrogen atom
Since electrons are smaller than atoms they must be parts of an atom
Atoms must be divisible
Atoms of different elements all produced these same electrons

19. Parts of an Atom: The Electron
Defined by Thompson
Tiny, negatively charged particle
Charge is -1
Very light compared to mass of atom
1/2000th the mass of a H atom
Moves very rapidly within the atom

20. Thompson’s Model of Atoms
Atoms have a structure and are divisible
Thomson reasoned that electrons must be a fraction of the entire size of the atom since their mass is much smaller that the whole atom
Thomson also reasoned since atoms are neutral, the electrons were embedded in a sphere of uniform positive charge
Thomson (1898) proposed the “Plum Pudding” model or “Raisin Muffin” model of the atom

21. Thompson’s Model of Atoms
Thomson Atomic Model (early 1900’s): Proposed a uniform, positive sphere of matter with small negative electrons attached to the surface of the sphere
This became known as the plum-pudding model
In the early 1900s plum pudding model

22. Rutherford’s Experiment
1911 Rutherford designed an experiment to test the Thompson model (“plum-pudding”) of the atom
Rutherford directed positively charged particles (alpha particles) towards a thin gold foil sheet
Rutherford expected the particles to pass straight through a uniform area of mass and positive charge

23. Rutherford’s Experiment

24. Rutherford’s Experiment
Results:
Most (alpha) particles mostly went straight through
A few particles were unexpectedly deflected from their expected (straight) path
A few deflected nearly back towards alpha particle source

25. Rutherford’s Experiment
Rutherford proposed:
A very small, dense core at the center of the atom
This dense core was called the “nucleus”
It contains most of the mass of the atom and it has a positive charge (protons)
Most of an atom is empty space filled with electrons

26. Parts of an Atom
Experimentation in the early 20th century (Thomson and Rutherford) proved atoms were not indivisible spheres
Atoms are comprised of smaller particles: Subatomic particles
More experiments led to the discovery of two more fundamental subatomic particles: Protons and neutrons
Electron: Negatively charged (1897)
Proton: Positively charged (1919)
Neutron: No electrical charge (1932)

27. Nucleus of the Atom (Rutherford Model)
A very dense, small center exists in the center of the atom called the nucleus
Volume of nucleus is about 1/10 trillionth the volume of the entire atom
Nucleus is basically the entire mass of the atom
The protons and neutrons are located in the nucleus
Most of the atom is empty space with fast-moving electrons

28. Nucleus of the Atom (Rutherford Model)
The nucleus is the center (core) of the atom
The nucleus has most of the mass of the atom
protons
neutrons
The extranuclear region
it contains all the electrons
Extranuclear
region
nucleus
Over 99.9 % of the mass is concentrated in the nucleus (called the nuclear region)

29. Nucleus of the Atom (Rutherford Model)
The nucleus is the core of the atom
Positively charged
Contains most of the mass of the atom
Within a neutral atom, there are equal numbers of protons and electrons, so atom has a net charge of zero

30. Mass of Subatomic Particles: The Proton
Has the same magnitude charge as the electron, but oppositely charged
Has a charge of +1
Weighs about 2000 times an electron
Is found in the nucleus
In a neutral atom:
Number of protons = identity of the compound
Number of protons = number of electrons

31. Mass of Subatomic Particles: The Neutron
The last of the three subatomic particles to be discovered, also located in the nucleus
The mass is about the same as a proton
Has no charge (neutral)
Variable amounts are possible in atoms of the same element
This is the basis for isotopes
Dalton reasoned the atoms of an element were identical because they had the same chemical properties
But atoms can exist can differ in a way that has minimal effect on its chemical properties

32. Mass of Subatomic Particles
The three subatomic particles have extremely small masses
Chemists base the mass of atoms on the atomic mass scale
A relative scale based on the mass of one carbon atom: amu
One amu is 1/12 the mass of one carbon atom, so the approximate mass of one proton or neutron is 1.00 amu

33. Atomic Number = number of protons in an atom
Atomic Number (Z)
All elements in periodic table arranged according to the atomic number
Equal to the number of protons in the nucleus of an atom
Determines the identity of the atom
Is also equal to the number of electrons in the neutral atom
The top number in each square in the periodic table
Atomic Number = number of protons in an atom

34. Mass Number = number of protons + number of neutrons
Mass Number (A)
The total number of protons and neutrons in an atom
Mass number is always a whole number (no decimals)
An oxygen atom has a mass number of 16 (8 protons and 8 neutrons)
Mass Number = number of protons + number of neutrons

35. Isotopes and Atomic Mass
All atoms of the same element have the same atomic number (Z)
The same element can differ in the mass number (A) due to a different number of neutrons
All Mg atoms have 12 protons, but may have 12, 13, or 14 neutrons

36. Isotopes and Atomic Mass
Atoms that have the same number of protons and electrons but different numbers of neutrons are called isotopes
Since isotopes are atoms of the same element,
They have the same atomic number
They display the same chemical properties

37. Nuclear (Isotopic) Symbols
A notation used when necessary to differentiate between isotopes
A is the mass number
Z is the atomic number
X is the chemical symbol

38. Atomic Mass
A specific element can have several mass values if it exists in isotopic forms
For example, oxygen atoms can have any one of three masses but often treated as if it has one mass
The atomic mass of an element is the mass of the “average atom” of that element
Reduces the number of masses to deal with to 109 from many hundreds
This weighted average takes into account the natural abundances and the atomic masses of the isotopes of an element

39. Atomic Mass Atomic mass is a “weighted average mass” based on:
The number of isotopes that exist for the element
The relative mass of each isotope
The percent abundance of each isotope

40. Example: Isotopes and Atomic Mass
Complete the following table:
Symbol
Name
# Protons
#Neutrons
#Electrons
Hydrogen 1 1
Fluorine 9 10 9
Copper 29 35 29
Hydrogen 1 1 1

41. end
Remaining slides (Section 4.6) will continue with chapter 10

42. Electron Energy Levels
Electrons possess energy; they are in constant motion in the large empty space of the atom
The arrangement of electrons in an atom corresponds to an electron’s energy
The electron resides outside the nucleus in one of seven fixed energy levels
Energy levels are quantized: Only certain energy values are allowed
A quantized property is one that can only have certain values
If the atom could have all possible energies, then the result would be a continuous spectrum instead of lines

43. Electron Energy Levels
Electrons of similar energy are grouped into energy levels
The major energy levels in an atom are called the principal shells symbolized by n, the principal quantum number
As n goes from 1 to 2, 3, 4, etc., the electron’s energy and distance from the nucleus increases
The maximum number of electrons in an energy level is equal to 2n2

44. Electron Arrangements
The chemical properties of an element are determined by the number and arrangement of electrons in the nucleus
The electron arrangement (configuration) is a statement of the number of electrons in each energy level
The number of electrons an atom has of various energies
The electron arrangements of the first 18 elements can be written by placing electrons in order of increasing energy

45. Valence Electrons
The electrons that reside in the highest energy level (n)
They are the furthest electrons from the nucleus
Determine the chemical properties of an element
Number at the top of each column for elements (1A-8A) equals the number of valence electrons for each element in that group

46. Electron-Dot Symbols
Consists of an element’s symbol and one dot for each valence electron placed about the symbol
Used only for main group elements (1A to 8A)
Main group elements in the same group have the same number of valence electrons
The number of valence electrons is the same as the group number
Also called Lewis structures

47. Electron-Dot Symbols

48. Electron-Dot Symbols Write the symbol of the element
Determine the number of valence electrons by the group number
Use dot or X to represent an electron
Group Number
Li• Be• •B• •C• •N• •O: :F: :Ne: • ••
Li• Li+1 :F: [:F:]-1 • ••

49. Electron-dot symbols Examples
Determine the number of valence electrons for Ba, As, and Br.
Write the electron-dot symbol for each of these elements
Ba: 2 valence electrons
As: 5 valence electrons
Br: 7 valence electrons Ba As Br

50. Ionization Energy
Energy needed to remove one electron from an atom in the gas state
It measures how tightly an atom holds its electrons
The lower the ionization energy, the easier it is to remove the electron
Metals have low ionization energies
Nonmetals have high ionization energies
Ionization Energy decreases down the group
Ionization Energy increases across the period
Left to right
Valence electron farther from nucleus

51. Ionization Energy
The more tightly an electron is held, the higher the ionization energy
The outermost electrons are the easiest to remove; as the energy level increases the farther the electron is from the nucleus
As n gets larger, removal of an electron requires less energy
Helium requires the most energy of any element due to its full (stable) energy level

52. end