1. Atoms and Elements Chapter 4
Chem 10 Spring 2008 Chp 4
Atoms and Elements Chapter 4
Tro, 2nd ed.

2. Each element has a number.
ELEMENTS AND ATOMS
All known substances on Earth and probably the universe are formed by combinations of more than 100 elements.
An element is a fundamental or elementary substance that cannot be broken down into simpler substances by chemical means.
Each element has a number.
Beginning with hydrogen as 1, the elements are numbered in order of increasing complexity.

3. ELEMENTS AND ATOMS
Most substances can be decomposed into two or more simpler substances.
Water can be decomposed into hydrogen and oxygen.
Table salt can be decomposed into sodium and chlorine.
An element cannot be decomposed into a simpler substance.
An atom is the smallest particle of an element that can exist.
An atom is the smallest unit of an element that can enter into a chemical reaction.

4. An atom is very small

5. The diameter of an atom is 0.1 to 0.5 nm.
This is 1 to 5 ten billionths of a meter.
Even smaller particles than atoms exist. These are called subatomic particles.
If the diameter of this dot is 1 mm, then 10 million hydrogen atoms would form a line across the dot.

6. Dalton’s Atomic Theory
Elements are composed of minute indivisible particles called atoms.
Atoms under special circumstances can be decomposed.
Modern research has demonstrated that atoms are composed of subatomic particles.
Atoms of the same element are alike in mass and size. Atoms of different elements have different masses and sizes. (Is this still correct?)
Chemical compounds are formed by the union of two or atoms of different elements in whole number ratios.

7. In 1897 Sir Joseph Thompson demonstrated that cathode rays:
WHAT’S IN AN ATOM?
DISCOVERING SUBATOMIC PARTICLES!
In 1897 Sir Joseph Thompson demonstrated that cathode rays:
- travel in straight lines
- are negative in charge
- are deflected by electric and magnetic fields
- produce sharp shadows
- are capable of moving a small paddle wheel

8. Subatomic Particles: Electron
This was the discovery of the fundamental unit of charge – the electron.

9. Subatomic Particles: Protons
Eugen Goldstein, a German physicist, first observed protons in 1886:
Thompson determined the proton’s characteristics.
Thompson showed that atoms contained both positive and negative charges.
This disproved the Dalton model of the atom which held that atoms were indivisible.
But the mass of the atom could not be accounted for by the mass of protons inside it. There had to be something else.

10. James Chadwick discovered the neutron in 1932.
Subatomic Particles: Neutron
James Chadwick discovered the neutron in 1932.
Its actual mass is slightly greater than the mass of a proton.

11. 4.1
1.0073
1.0087
Memorize name, symbol, relative charge, relative mass, and LOCATION (in nucleus or in orbitals outside of nucleus) of these three subatomic particles. See table 4.1 on page 90/96 in your book.

12. Radioactivity was discovered by Becquerel in 1896.
The Nuclear Atom
Radioactivity was discovered by Becquerel in 1896.
Radioactive elements spontaneously emit alpha (a) particles, beta (b) particles and gamma (g) rays from their nuclei.
By 1907, Rutherford found that a particles emitted by certain radioactive elements were helium nuclei, consisting of 2 protons and 2 neutrons.

13. The Rutherford Experiment
In 1911, Rutherford performed experiments that shot a stream of a particles at a thin piece of gold foil.
Most of the a particles passed through the foil with little or no deflection.
He found that a few were deflected at large angles and some a particles even bounced back.

14. Rutherford’s alpha particle scattering experiment.
5.5

15. The Rutherford Experiment
An electron with a mass of amu could not have deflected an a particle with a mass of ~4 amu.
Rutherford knew that like charges repel.
Rutherford concluded that each gold atom contained a positively charged mass that occupied a tiny volume. He called this mass the nucleus.

16. The Rutherford Experiment
If a positive a particle approached close enough to the positive mass it was deflected.
Most of the a particles passed through the gold foil. This led Rutherford to conclude that a gold atom was mostly empty space.
Because a particles have relatively high masses, the extent of the deflections led Rutherford to conclude that the nucleus was very heavy and dense.

17. Deflection Scattering
Deflection and scattering of a particles by positive gold nuclei.
5.5

18. General Arrangement of Subatomic Particles
Rutherford’s experiment showed that an atom had a dense, positively charged nucleus.
Chadwick’s work in 1932 demonstrated the atom contains neutrons.
Rutherford also noted that light, negatively charged electrons were present in an atom and offset the positive nuclear charge.

19. General Arrangement of Subatomic Particles
Rutherford put forward a model of the atom in which a dense, positively charged nucleus is located at the atom’s center.
The negative electrons surround the nucleus.
The nucleus contains protons and neutrons

20. 5.6

21. ATOMS AND THEIR SUBATOMIC PARTICLES
The atomic number of an element is equal to the number of protons in the nucleus of that element.
The atomic number of an atom determines which element the atom is.
Every atom with an atomic number of 1 is a hydrogen atom. Every carbon atom contains 6 protons in its nucleus.

22. atomic number 1H
Every atom with an atomic number of 1 is a hydrogen atom.
1 proton in the nucleus

23. atomic number 6C
Every atom with an atomic number of 6 is a carbon atom.
6 protons in the nucleus

24. Ions: atoms that have lost or gained electrons
Positive ions were explained by assuming that a neutral atom loses electrons.
Negative ions were explained by assuming that extra electrons can be added to atoms.

25. When one or more electrons are lost from an atom, a cation is formed.
5.4

26. When one or more electrons are added to a neutral atom, an anion is formed.
5.4

27. Atomic Structures of Ions

28. Isotopes of the Elements
Atoms of the same element can have different masses.
They always have the same number of protons, but they can have different numbers of neutrons in their nuclei.
The difference in the number of neutrons accounts for the difference in mass.
These are isotopes of the same element.

29. Isotopic Notation

30. Isotopic Notation for Carbon-12
6 protons + 6 neutrons C 6 12
6 protons

31. Isotopic Notation for Carbon-14
6 protons + 8 neutrons 14 C 6
6 protons

32. Hydrogen has three isotopes
1 proton
0 neutrons
1 proton
1 neutron
1 proton
2 neutrons

33. Examples of Isotopes
Element Protons Electrons Neutrons Symbol
Hydrogen
Hydrogen
Hydrogen
Uranium
Uranium
Chlorine
Chlorine

34. The mass of a single atom is too small to measure on a balance.
Atomic Mass
The mass of a single atom is too small to measure on a balance.
Using a mass spectrometer, the mass of the hydrogen atom was determined.
The mass of one hydrogen atom was determined to be x g.

35. A Modern Mass Spectrometer
Positive ions formed from sample.
Electrical field at slits accelerates positive ions.
Deflection of positive ions occurs at magnetic field.
From the intensity and positions of the lines on the mass spectrogram, the different isotopes and their relative amounts can be determined.
A mass spectrogram is recorded.
5.8

36. A typical reading from a mass spectrometer
A typical reading from a mass spectrometer. The two principal isotopes of copper are shown with the abundance (%) given.
5.9

37. ATOMIC MASS UNITS AND RELATIVE ATOMIC MASS
Single atomic masses are too small to weigh on a balance. To overcome this problem a system of relative atomic masses using “atomic mass units” was devised to express the masses of elements using simple numbers.
The standard to which the masses of all other atoms are compared to was chosen to be the most abundant isotope of carbon, carbon-12.
A mass of exactly 12 atomic mass units (amu) was assigned to the carbon-12 atom. An amu is defined as exactly equal to 1/12th mass of a carbon-12 atom. 1 amu = x g
Isotopes of the same element have different masses.
The listed atomic mass of an element is the average relative mass of the isotopes of that element compared to the mass of carbon-12.

38. ATOMIC MASS UNITS AND RELATIVE ATOMIC MASS
Calculating relative atomic mass:
Convert percent abundance to fraction by dividing by 100
Multiply fraction abundance times each isotopes mass
Sum up the results
Rel Atomic Mass = f1*m1 + f2*m2 + …
See example on next slide

39. Average atomic mass (amu)
To calculate the atomic mass multiply the atomic mass of each isotope by its fractional abundance and add the results.
Isotope
Isotopic mass (amu)
Abundance (%)
Average atomic mass (amu)
69.09
30.91
63.55
( amu) =
43.48 amu
( amu) =
20.07 amu
63.55 amu

40. Relationship between Mass Number and Atomic Mass
Atomic mass found on the Periodic Table is RELATIVE atomic mass, the weighted average of all the isotopes.
Each isotope of an element has a unique WHOLE number called the mass number, the sum of protons and neutrons. Mass number is NOT found on most Periodic Tables.
******************************************************************
ATOMIC MASS IS NOT THE SAME AS MASS NUMBER
REMEMBER: ATOMIC MASS IS AN AVERAGE OF THE DIFFERENT ISOTOPES’ MASSES. It will have decimal points!

41. mass number atomic number number of neutrons - = 109 - 47 = 62
Relationship Between Mass Number and Atomic Number
The mass number minus the atomic number equals the number of neutrons in the nucleus.
mass number
atomic number
mass number
atomic number
number of neutrons - =
109 - 47 = 62

42. Atomic Mass, Mass Number & Abundance of Isotopes
EXAMPLE: Chlorine has two isotopes, chlorine-35 and chlorine-37. Write down the atomic number, mass number and number of neutrons for each. Then look up the atomic mass. Which isotope is in greater abundance?
Cl-35: Z = 17, A = 35, = 18 neutrons
Cl-37: Z = 17, A = 37, = 20 neutrons
Atomic mass is amu.
Closer to 35, therefore chlorine-35 is more abundant.

43. Elements are not distributed equally by nature.
Oxygen is the most abundant element in the human body (65%).
Oxygen is the most abundant element in the crust of the earth (49.2%).
In the universe, the most abundant element is hydrogen (91%) and the second most abundant element is helium (8.75%).

44. Distribution of the common elements in nature.
3.2

45. Sources of Element Names
Greek-Color
Iodine: from the Greek iodes meaning violet.
Latin-
Property
Fluorine: from the Latin fluere meaning to flow. The fluorine containing ore fluorospar is low melting.
German-
Color
Bismuth: from the German weisse mass which means white mass.
Location
Germanium: discovered in 1866 by a German chemist.
Famous-
Scientists
Einsteinium: named for Albert Einstein.

46. These symbols have carried over from the earlier names of the elements (usually Latin).
A number of symbols appear to have no connection with the element.
Most symbols start with the same letter as the element.

48. Element Names and Symbols
Element names and symbols are in Periodic Table inside front cover.
Symbols: Learn first 36 on Periodic Table, plus Ag, Au, Pt, Hg, Sn, Pb, I, U, Ba, and Rn.
Also memorize that some elements exist as diatomic molecules : H2, O2, N2, F2, Cl2, Br2, and I2

49. N 7 THE PERIODIC TABLE OF THE ELEMENTS
The periodic table was designed by Dimitri Mendelev in 1869.
In the table each element’s symbol is placed inside of a box.
Above the symbol of the element is its atomic number. 7 N

50. The elements are arranged in order of increasing atomic number.
Elements with similar chemical properties are organized in columns called families or groups . He Ne
These elements are known as the noble gases. They are nonreactive. Ar Kr Xe Rn

51. THE PERIODIC TABLE OF THE ELEMENTS
18 columns - called groups
7 rows - called periods
Representative elements are in groups 1,2, 3A-8A (13-18)
Transition metals are in groups 3-12
Groups of elements have similar chemical properties

52. THE PERIODIC TABLE OF THE ELEMENTS
SPECIAL GROUP NAMES:
1A (1) = alkali metals
2A (2) = alkaline earth metals
8A (18) = noble gases (six)
7A (17) = halogens (four)
NOTE: H really belongs to its own group, which is why it’s shown by itself in my periodic table! It has 1 electron in outer shell, like the 1A elements, but it’s not a metal, and reacts more like group 7A.

53. The Periodic Table of Elements
Nonmetals: upper right
Metals: lower left
Metalloids touch stairs, except for Al.

54. Metals are solid at room temperature.
Mercury is an exception. At room temperature it is a liquid.
Metals are good conductors of heat and electricity.
Most elements
are metals
physical
properties of metals
Metals are malleable (they can be rolled or hammered into sheets).
Metals have high luster (they are shiny).

55. Metals have high densities
Metals are ductile (they can be drawn into wires).
Most metals have a high melting point.
Most elements
are metals
Metals have high densities

56. Chemical Properties of Metals
Metals have little tendency to combine with each other to form compounds.
Many metals readily combine with nonmetals to form ionic compounds.
They can combine with sulfur.
oxygen.
chlorine.
In nature, minerals are formed by combinations of the more reactive metals with other elements.

57. Chemical Properties of Metals
A few of the less reactive metals such as copper, silver and gold are found in the free state.
Metals can mix with each other to form alloys.
Brass is a mixture of copper and zinc.
Bronze is a mixture of copper and tin.
Steel is a mixture of carbon and iron.

58. Physical Properties of Nonmetals
Everything metals are, nonmetals are NOT!
Lack luster (they are dull)
Have relatively low melting points
Have low densities.
Poor conductors of heat and electricity
At room temperature, carbon, phosphorous, sulfur, selenium, and iodine are solids.

59. Physical State at Room Temperature
phosphorous
carbon
Solid
sulfur
selenium
iodine
NONMETALS

60. Physical State at Room Temperature
liquid
bromine
NONMETALS

61. gas nitrogen, oxygen fluorine, chlorine
Physical State at Room Temperature
gas
nitrogen, oxygen
fluorine, chlorine
helium, neon, argon, krypton, xenon, radon
NONMETALS

62. The Metalloids boron silicon germanium arsenic antimony tellurium
polonium
Metalloids have properties that are intermediate between metals and nonmetals

63. = Alkali Metals = Alkali Earth Metals = Noble Gases = Halogens
Chem 10 Spring 2008 Chp 4
= Alkali Metals
= Alkali Earth Metals
= Noble Gases
= Halogens
= Lanthanides
= Actinides
= Transition Metals
add pictures of elements from text

64. Important Groups - Hydrogen
Nonmetallic
Colorless, diatomic gas – exists as H2
Very low melting point & density
Reacts with nonmetals to form molecular compounds
HCl is acidic gas
H2O is a liquid
Reacts with metals to form hydrides
metal hydrides react with water to form H2
HX dissolves in water to form acids

65. Important Groups – IA, Alkali Metals
Hydrogen usually placed here, though it doesn’t belong
Soft, low melting points,low density
Flame tests ® Li = red, Na = yellow, K = violet
Very reactive, never find uncombined in nature
Tend to form water soluble compounds with colorless solutions
React with water to form basic (alkaline) solutions and H2
2 Na + 2 H2O ® 2 NaOH + H2
releases a lot of heat
lithium
sodium
potassium
rubidium
cesium

66. Important Groups – IIA, Alkali Earth Metals
Harder, higher melting, and denser than alkali metals
Flame tests ® Ca = red, Sr = red, Ba = yellow-green
Reactive, but less than corresponding alkali metal
Form stable, insoluble oxides from which they are normally extracted
Oxides are basic = alkaline earth
Reactivity with water to form H2, ® Be = none; Mg = steam; Ca, Sr, Ba = cold water
beryllium
magnesium
calcium
strontium
barium

67. Important Groups – VIIA, Halogens
Nonmetals
F2 & Cl2 gases; Br2 liquid; I2 solid
All diatomic
Very reactive
Cl2, Br2 react slowly with water
Br2 + H2O ® HBr + HOBr
React with metals to form ionic compounds
HX all acids
HF weak < HCl < HBr < HI
fluorine
chlorine
bromine
iodine

68. Important Groups – VIIIA, Noble Gases
All gases at room temperature
Very low melting and boiling points
Very unreactive, practically inert
Very hard to remove electron from or give an electron to

69. Review What is the atomic number of boron, B?
What is the atomic mass of silicon, Si?
How many protons does a chlorine atom have?
How many electrons does a neutral neon atom have?
Will an atom with 6 protons, 6 neutrons and 6 electrons be electrically neutral?
Will an atom with 27 protons, 32 neutrons and 27 electrons be electrically neutral?
Will a Na atom with 10 electrons be electrically neutral?
What is the mass number of fluorine?