1. Pre-IB/Pre-AP CHEMISTRY
Chapter 3 – Atoms: The Building Blocks of Matter

3. Section 1 Objectives
Be able to define and explain: law of conservation of mass, law of definite proportions, law of multiple proportions.
Be able to summarize the five essential points of Dalton’s atomic theory.
Be able to explain the relationship between Dalton’s atomic theory and the law of conservation of mass, the law of definite proportions, and the law of multiple proportions

4. Review
Matter is anything that occupies space and has mass. It’s simply what everything around us is made of.

5. Review
Mass is the amount of matter contained in an object. SI unit = kilogram

6. Review
All matter is either a pure substance or a mixture of substances.
A pure substance is made of only one type of matter. It is the same throughout.

7. Review
A mixture is made of two or more pure substances that can be separated very easily.
Substances in a mixture do not lose their identities (Ex. salt in water).

8. Pure Substances
Pure substances can be divided into two categories: elements and compounds.

9. Review
Elements are the simplest subtances of matter. Everything in the universe is made of one or more elements. There are over 100 known elements. (Letters of the alphabet) Ag

10. Review
The smallest particle of an element that still has all the properties of that element is called an atom.

11. Review
Transformation of a substance or substances into one or more new substances –
CHEMICAL reaction.

12. Story of the Atom
The ancient Greek philosophers described an atom as the smallest particle of matter which could not be divided. They imagined it to be a solid particle.
Democritus called nature’s basic particle an atom based on the Greek word meaning “indivisible”
ATOM

13. Foundations of Atomic Theory
Virtually all chemists in the late 1700s accepted the modern definition of an element as “a substance that cannot be further broken down by ordinary chemical means”.
Clear that elements combine to form compounds that have different chemical and physical properties than those elements that form them.
1790s – study of matter revolutionized by new inventions – tools allowed scientists to measure very small masses with accuracy never achieved before.
This allowed chemists to discover some very important laws.

14. Story of the Atom
The analytical balance scale was invented in the 1700’s. This tool allowed scientists to measure very small masses with accuracy never achieved before. This allowed chemists to discover some very important laws.
Democritus and Aristotle’s idea were not supported by experimental evidence and thus were remained in speculation until the eighteenth century.

15. Conservation of Mass
The Law of Conservation of Mass(Law of Conservation of Matter) states that mass(matter) is neither created nor destroyed during ordinary chemical or physical changes.
Oxygen Hydrogen  Water

16. Conservation of Mass
Soon lead to the assertion that regardless of where or how a pure chemical compound is prepared, it is composed of a fixed proportion of elements.
Oxygen Hydrogen  Water

17. Conservation of Mass

19. Definite Proportions
The Law of Definite Proportions states that a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or the source of a compound.

20. Multiple Proportions
The Law of Multiple Proportions states that if two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers.

21. Multiple Proportions

22. Story of the Atom
In 1808, John Dalton proposed an atomic theory that described atoms smallest particles of elements.

23. Story of the Atom
Dalton’s model pictured atoms as small, indivisible particles similar to the ancient Greeks. His model fit experimental evidence at the time(law of conservation of mass, law of multiple proportions, etc.).

24. Dalton’s Theory All matter is composed of atoms.
Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties.
Atoms cannot be subdivided, created, or destroyed.

25. Dalton’s Theory
Atoms of different elements combine in simple whole-number ratios to form chemical compounds.
In chemical reactions, atoms are combined, separated, or rearranged.

26. Modern Atomic Theory
By relating atoms to mass, Dalton turned Democritus’s idea into a scientific theory – one that could be tested by experiments.
Not all aspects of Dalton’s atomic theory have proven to be correct:
Today we know that atoms are divisible into even smaller particles
We know that a given element can have atoms with different masses (isotopes – element with the same number of PROTONS but a different number of NEUTRONS)
Atomic theory has been MODIFIED to fit new data/information
Aspects that have NOT Changed:
All matter is composed of atoms
Atoms of any element differ in properties from atoms of another element

27. Section 2 Objectives
Be able to define: atom, proton, neutron, electron, electron cloud, nuclear forces.
Be able to describe the properties of cathode rays that led to the discovery of the electron.
Be able to summarize the experiments of Rutherford that led to the discovery of the atomic nucleus.
Be able to list the properties of protons, neutrons, and electrons.

28. Story of the Atom
Although John Dalton thought that atoms were indivisible, investigators in the late 1800s proved otherwise. It became clear that atoms are actually composed of several basic types of smaller particles and that the number and arrangement of those particles determined the atom’s chemical properties.

29. Story of the Atom Atoms:
Today’s definition: the smallest particle of an element that retains the chemical properties of that element.
Atoms are put together a certain way and are made up of subatomic particles
They have a core we call the nucleus – very small region located at the center of every atom
Nucleus of every atom is made up of at least one POSITIVELY charge particle (proton) and usually one or more NEUTRAL particle (neutron)
Surrounding the nucleus is a region occupied by NEGATIVELY charged particles (electrons)
This region is very large compared to the nucleus

30. Story of the Atom
John Thomson demonstrated, in 1897, that atoms contained smaller, negatively charged particles. They were later called electrons. Atoms were no longer viewed as solid particles with no parts.

33. Story of the Atom Experiments revealed:
Cathode rays were deflected by a magnetic field in the same manner as a wire carrying electric current (negative charge)
The rays were deflected away from a negatively charged object
Thomson was able to measure the ratio of the negative charge of cathode-ray particles to their mass – found this ratio was always the SAME
These negatively charged “objects” were later called electrons. Atoms were no longer viewed as solid particles with no parts.

34. Story of the Atom
Thomson imagined the atom to be like a plum pudding. Electrons were like the scattered plums and the pudding was thought of as positively charged matter(all atoms are neutral).

35. Story of the Atom
In 1911, Ernest Rutherford’s experiments showed that atoms had a central core called a nucleus.

37. Rutherford’s Experiment

38. Story of the Atom
Rutherford suggested that negative electrons orbited a positively charged nucleus. He could not explain how this occurred.

39. Nucleus
Atoms are put together a certain way. They have a core we call the nucleus.
All atomic nuclei are made of TWO kinds of particles – protons and neutrons.
Except Hydrogen nucleus

40. Protons Inside the nucleus we find protons and neutrons.
Protons are positively charged particles, equal in magnitude to the negative charge of the electron

41. Neutrons Neutrons are neutral, they have no charge.
Atoms are neutral because they contain equal number of protons and electrons
Hydrogen has a single proton nucleus with a single electron moving about it.

42. Question
What forces are holding the nucleus together? Gravitational? Electrostatic? Magnetic?

43. Nuclear Forces
Generally – particles that have the SAME electric charge will repel one another.
When two protons are extremely close to each other, there is a strong ATTRACTION between them – they help stabilize the nucleus

44. Nuclear Forces
Nuclear forces are short range forces of attraction that exist between protons and neutrons in the nucleus.
Neutron – proton
Neutron – neutron
Proton – proton

45. Size of Atoms
Atoms are extremely small. Atomic radii are on the range from about 40 to 270 picometers(pm).
1 pm = m
1 pm is to 1 cm as 1 cm is to 103 km(about 600 mi)

46. Size of Atoms
Region occupied by the electrons as an electron cloud – a cloud of NEGATIVE charge
Radius of an atom is the distance from the center of the nucleus to the outer portion of this electron cloud

47. Section 3 Objectives
Be able to define: atomic number, isotope, mass number, nuclide, atomic mass unit(amu), average atomic mass, mole, Avogadro’s number, molar mass.
Be able to describe how the terms atomic number and mass number apply to isotopes.

48. Section 3 Objectives
Be able to determine the number of protons and neutrons in a nuclide when given the identity of the nuclide.
Be able to state how the terms mole, Avogadro’s number, and molar mass are related.

49. Section 3 Objectives
Be able to solve problems involving mass in grams, amount of moles, and number of atoms of an element using conversion factors(dimensional analysis).

50. Atomic Number
All atoms are composed of the same basic particles, but not all atoms are the same. Atoms of different elements have different numbers of protons & atoms of the same element have the same number of protons.

51. Atomic Number
The atomic number (Z) of an atom indicates the number of protons in its nucleus.
The atomic number also tells us the number of electrons in a neutral atom.

52. Atomic Number
On the periodic table you will notice that an element’s atomic number is located above its symbol. The elements are placed on the periodic table by increasing atomic number. The atomic number identifies an element.

53. Atomic Number

54. Isotopes
Isotopes - Atoms of an element that have the same number of protons (atomic number) & different mass numbers. They contain different numbers of neutrons.

55. Isotopes
The simplest atom is hydrogen – all hydrogen atoms have only one proton, but hydrogen atoms can have different numbers of neutrons. Three types of hydrogen atoms are known:
Protium - most common – accounts for % of the hydrogen atoms found on Earth; the nucleus has one proton and one ELECTRON moving about it.
Deuterium - accounts for % of Earth’s hydrogen atoms; nucleus has one proton and one NEUTRON

56. Los Alamos National Laboratory
Radioisotopes
The isotopes of some atoms are unstable and undergo nuclear changes. They are called radioisotopes.
Los Alamos National Laboratory

57. Mass Number
The mass number (A) of an atom indicates the number of protons and neutrons in its nucleus.

58. Naming Isotopes 12 6 14
Mass Number = A
At. Number = Z C
Isotopes are designated by giving the name of the element followed by its mass number (Carbon-12, Hydrogen-2).

59. Naming Isotopes
Hyphen Notation: The mass number is written with a hyphen after the name of the element
Example: hydrogen – 3
Nuclear Symbol: shows the composition of a nucleus
Example: uranium – 235 can be written as 𝑈
Superscript indicates the mass number (protons + neutrons)
Subscript indicates the atomic number (number of protons)
Nuclide – a general term for a specific isotope of an element

60. Relative Atomic Mass
As we learned in Chapter 2, scientists use standards of measurement that are constant and are the same everywhere. In order to set up a relative scale of atomic mass, one atom as been arbitrarily chosen as the standard and assigned a mass value – all other masses are expressed in relation to this defined standard.

61. Relative Atomic Mass
One atomic mass unit (1 amu) is exactly 1/12 the mass of a carbon-12 atom. The atomic mass of any other atom is determined by comparing it with the mass of the carbon-12 atom.
Isotopes of an element may occur naturally, or they may be made in the lab (artificial isotopes). Although isotopes have different masses, they do no differ significantly in their chemical behavior.

62. Relative Atomic Mass Mass of electron: .0005486 amu
Mass of proton: amu
Mass of neutron: amu

63. Atomic Mass Unit
Most elements occur naturally as mixtures of isotopes – the percentage of each isotope in the naturally occurring element on Earth is nearly always the same, no matter where the element is found. Average atomic mass – the weighted average of the atomic masses of the naturally occurring isotopes of an element.

64. Example 1
Example: suppose you have a box containing two sizes of marbles – if 25% of the marbles have masses of 2.00g and 75% have masses of 3.00g each, how is the weighted average calculated? 2 different ways:
1. Count the number of total marbles, calculate the total mass of the mixture and divide by the total number of marbles. If we have 100 marbles, the calculations would be as follows:
25 marbles (25% of 100) x 2.00g = 50g
75 marbles x 3.00g = 225g

65. Example 1 Adding these masses gives the total mass of the marbles:
50g + 225g = 275g
Dividing the total mass by 100 gives an average mass of the marbles = 2.75g
2. (Simpler way) multiply the mass of each marble by the decimal fraction representing its percentage in the mixture. Then add the products.
25% = % = .75
(2.00g x .25) + (3.00g x .75) = 2.75g

66. Calculating Average Atomic Mass
The average atomic mass of an element depends on both the mass and relative abundance of each of the element’s isotopes.
( x amu) = amu
( x amu) = amu
63.55 amu

67. Quantities of Particles
Often chemists are interested in knowing the amounts of atoms or molecules they have in a sample of a substance.

68. Counting Particles
Counting units for certain kinds of objects vary depending on the size and the use they serve.

69. Mole The SI unit for the amount of a substance is called the mole.
A mole of a substance is the amount of that substance that contains the same number of particles as exactly 12 g of carbon-12.

70. Mole
It has been determined that a mole of any substance contains x 1023particles, just like a dozen always contains 12 items.
6.022 x 1023 is called the Avogadro number. So, a mole contains the Avogadro number of particles.

71. Mole of Atoms
A mole of atoms would be the Avogadro number of atoms.

72. Molar Mass
Another definition of a mole is the mass of a substance that contains the Avogadro number of particles. This is called the molar mass of a substance.

73. Mass of a Mole
You wouldn’t expect a dozen golf balls to have the same mass as a dozen baseballs. Likewise, the mass of a mole of a substance will vary.

74. Molar Mass
The mass in grams of a mole of any substance is called its molar mass.

75. Molar Mass
The mass in grams of a mole of any element is numerically equal to its atomic mass expressed in grams.
1 mole of Cu atoms
6.022 x 1023 Cu atoms
63.5 g of Cu

76. Molar Mass
The mass in grams of a mole of any molecular substance is numerically equal to its molecular mass expressed in grams.
1 mole of H2O molecules
6.022 x 1023 H2O molecules
18.0 g of H2O