1. Atoms: The Building Blocks of Matter
Chapter 3

2. The Atom: From Philosophical Idea to Scientific Theory
When you crush a lump of sugar, you can see that it is made up of many smaller pieces of sugar. You may grind these particles into a very fine powder, but each tiny piece is still sugar. Now suppose you dissolve the sugar in water. The tiny particles seem to disappear completely. Even if you look at the sugar-water solution through a powerful microscope you cannot see any sugar particles. Yet if you were to taste the solution, you’d know that the sugar is still there. Observations like these led early philosophers to ponder the fundamental nature of matter. Is it continuous and infinitely divisible, or is it divisible only until a basic, invisible particle that cannot be divided further is reached?
The Atom: From Philosophical Idea to Scientific Theory

3. The particle theory of matter was supported as early as 400 b. c
The particle theory of matter was supported as early as 400 b.c. by certain Greek thinkers, such as Democritus. He called nature’s basic particle an atom, based on the Greek word meaning “indivisible.” Aristotle was part of the generation that succeeded Democritus. His ideas had a lasting impact on Western civilization, and he did not believe in atoms. He thought that all matter was continuous, and his opinion was accepted for nearly 2000 years. Neither the view of Aristotle nor that of Democritus was supported by experimental evidence, so each remained speculation until the eighteenth century. Then scientists began to gather evidence favoring the atomic theory of matter.

4. Foundations of Atomic Theory
Nearly all chemists in late 1700s accepted the definition of an element as a substance that cannot be broken down further
Knew about chemical reactions
Great disagreement as to whether elements always combine in the same ratio when forming a specific compound

5. Law of Conservation of Mass
With the help of improved balances, investigators could accurately measure the masses of the elements and compounds they were studying
This lead to discovery of several basic laws
Law of conservation of mass  states that mass is neither destroyed nor created during ordinary chemical reactions or physical changes

6. Law of Definite Proportions
law of definite proportions  A chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound
Each of the salt crystals shown here contains exactly
39.34% sodium and 60.66% chlorine by mass.

7. Law of Multiple Proportions
Two elements sometimes combine to form more than one compound
For example, the elements carbon and oxygen form two compounds, carbon dioxide and carbon monoxide
Consider samples of each of these compounds, each containing 1.0 g of carbon
In carbon dioxide, 2.66 g of oxygen combine with 1.0 g of carbon
In carbon monoxide, 1.33 g of oxygen combine with 1.0 g of carbon
The ratio of the masses of oxygen in these two compounds is exactly to 1.33, or 2 to 1

8. 1808 John Dalton
Proposed an explanation for the law of conservation of mass, the law of definite proportions, and the law of multiple proportions
He reasoned that elements were composed of atoms and that only whole numbers of atoms can combine to form compounds
His theory can be summed up by the following statements

9. Dalton’s Atomic Theory
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, and other properties.
Atoms cannot be divided, created or destroyed.
Atoms of different elements combine in simple whole-number ratios to form chemical compounds.
In chemical reactions, atoms are combined, separated, or rearranged.

10. Modern Atomic Theory
Dalton turned Democritus’s idea into a scientific theory which was testable
Not all parts of his theory have been proven correct
Ex. We know atoms are divisible into even smaller particles
We know an element can have atoms with different masses

11. The Structure of the Atom
Although john dalton thought atoms were indivisible, investigators in the late 1800s proved otherwise. As scientific advances allowed a deeper exploration of matter, it became clear that atoms are actually composed of several basic types of smaller particles and that the number and arrangement of these particles within an atom determine that atom’s chemical properties. Today we define an atom as the smallest particle of an element that retains the chemical properties of that element.
The Structure of the Atom

12. All atoms consist of two regions
Nucleus  very small region located near the center of an atom
In the nucleus there is at least one positively charged particle called the proton
Usually at least one neutral particle called the neutron
Surrounding the nucleus is a region occupied by negatively charged particles called electrons

13. Discovery of the Electron
Resulted from investigations into the relationship between electricity and matter
Late 1800s, many experiments were performed  electric current was passed through different gases at low pressures
These experiments were carried out in glass tubes known as cathode-ray tubes

14. Cathode Rays and Electrons
Investigators noticed that when current was passed through a cathode-ray tube, the opposite end of the tube glowed
Hypothesized that the glow was caused by a stream of particles, which they called a cathode ray
The ray traveled from the cathode to the anode when current was passed through the tube

15. Observations
1. cathode rays deflected by magnetic field in same way as wire carrying electric current (known to have negative charge)
2. rays deflected away from negatively charged object

16. Observations led to the hypothesis that the particles that compose cathode rays are negatively charged
Strongly supported by a series of experiments carried out in 1897 by the English physicist Joseph John Thomson

17. He was able to measure the ratio of the charge of cathode-ray particles to their mass
He found that this ratio was always the same, regardless of the metal used to make the cathode or the nature of the gas inside the cathode-ray tube
Thomson concluded that all cathode rays are composed of identical negatively charged particles, which were later named electrons

18. Charge and Mass of the Electron
Confirmed that the electron carries a negative electric charge
Because cathode rays have identical properties regardless of the element used to produce them, it was concluded that electrons are present in atoms of all elements
Cathode-ray experiments provided evidence that atoms are divisible and that one of the atom’s basic constituents is the negatively charged electron

19. Thomson’s experiment revealed that the electron has a very large charge for its tiny mass
Mass of the electron is about one two-thousandth the mass of the simplest type of hydrogen atom (the smallest atom known)
Since then found that the electron has a mass of × 10−31 kg, or 1/1837 the mass of the hydrogen atom

20. Based on information about electrons, two inferences made about atomic structure
b/c atoms are neutral they must have positive charge to balance negative electrons
b/c electrons have very little mass, atoms must have some other particles that make up most of the mass

21. Thomson’s Atom Plum pudding model (based on English dessert)
Negative electrons spread evenly through positive charge of the rest of the atom
Like seeds in a watermelon

22. Discovery of Atomic Nucleus
1911 by New Zealander Ernest Rutherford and his associates Hans Geiger and Ernest Marsden
Bombarded a thin, gold foil with fast-moving alpha particles (positively charged particles with about four times the mass of a hydrogen atom)
Assume mass and charge were uniformly distributed throughout atoms of gold foil (from Thomson’s model of the atom)
Expected alpha particles to pass through with only slight deflection

23. What Really Happened…
Most particles passed with only slight deflection
However, 1/8,000 were found to have a wide deflection
Rutherford explained later it was “as if you have fired a 15-inch artillery shell at a piece of tissue paper and it came back and hit you.”

25. Explanation
After 2 years, Rutherford finally came up with an explanation
The rebounded alpha particles must have experienced some powerful force within the atom
The source of this force must occupy a very small amount of space because so few of the total number of alpha particles had been affected by it
The force must be caused by a very densely packed bundle of matter with a positive electric charge
Rutherford called this positive bundle of matter the nucleus

26. Rutherford had discovered that the volume of a nucleus was very small compared with the total volume of an atom
If the nucleus were the size of a marble, then the size of the atom would be about the size of a football field
But where were the electrons?
Rutherford suggested that the electrons surrounded the positively charged nucleus like planets around the sun
He could not explain, however, what kept the electrons in motion around the nucleus

27. Rutherford’s Atom

28. Composition of Atomic Nucleus
Except hydrogen, all atomic nuclei made of two kinds of particles
Protons
Neutrons
Protons = positive
Neutrons = neutral
Electrons = negative
Atoms are electrically neutral, so number of protons and electrons IS ALWAYS THE SAME

29. The nuclei of atoms of different elements differ in the number of protons they contain and therefore in the amount of positive charge they possess
So the number of protons in an atom’s nucleus determines that atom’s identity

30. Forces in the Nucleus
Usually, particles that have the same electric charge repel one another
Would expect a nucleus with more than one proton to be unstable
When two protons are extremely close to each other, there is a strong attraction between them
Nuclear forces  short-range proton-neutron, proton-proton, and neutron-neutron forces that hold the nuclear particles together

31. The Sizes of atoms
Area occupied by electrons is electron cloud – cloud of negative charge
Radius of atom is distance from center of nucleus to outer portion of cloud
Unit – picometer (10-12 m)
Atomic radii range from pm
Very high densities – 2 x 108 tons/cm3

32. Counting Atoms
Consider Neon, Ne, the gas used in many illuminated signs. Neon is a minor part of the atmosphere. In fact, dry air contains only about 0.002% Ne. And yet there are about 5 x 1017 atoms of neon present in each breath you inhale. In most experiments, atoms are too small to be measured individually. Chemists can analyze atoms quantitatively, however, by knowing fundamental properties of the atoms of each element. In this section, you will be introduced to some of the basic properties of atoms. You will then discover how to use this information to count the number of atoms of an element in a sample with a known mass. You will also become familiar the the mole, a special unit used by chemists to express amounts of particles, such as atoms and molecules.

33. Atomic Number 3 Li Lithium 6.941 [He]2s1
Atoms of different elements have different numbers of protons
Atomic number (Z)  number of protons in the nucleus of each atom of that element
Shown on periodic table
Atomic number identifies an element 3 Li
Lithium
6.941
[He]2s1

35. Isotopes
Hydrogen and other atoms can contain different numbers of neutrons
Isotope  atoms of same element that have different masses

37. Isotopes of Hydrogen

38. Mass Number
Mass number  total number of protons and neutrons in the nucleus of an isotope

39. Naming Isotopes
Isotopes are usually identified by specifying their mass number
There are two methods for specifying isotopes
Mass number is written with a hyphen after the name of the element
Tritium is written as hydrogen-3
Refer to this method as hyphen notation
Shows the composition of a nucleus as the isotope’s nuclear symbol
Uranium-235 is written as 23592U
The superscript indicates the mass number and the subscript indicates the atomic number
Nuclide – general term for specific isotope of an element

40. Sample Problem
How many protons, electrons, and neutrons are there in an atom of chlorine-37?
atomic number = number of protons = number of electrons
mass number = number of neutrons + number of protons
mass number of chlorine-37 − atomic number of chlorine = number of neutrons in chlorine-37
mass number − atomic number = 37 (protons plus neutrons) − 17 protons = 20 neutrons

41. Practice Problems
How many protons, electrons, and neutrons are in an atom of bromine- 80?
Answer 35 protons, 35 electrons, 45 neutrons
2. Write the nuclear symbol for carbon-13.
Answer 136C
3. Write the hyphen notation for the element that contains 15 electrons and 15 neutrons.
Answer phosphorus-30

42. Relative Atomic Masses
Because atoms are so small, scientists use a standard to control the units of atomic mass  carbon-12
Randomly assigned a mass of exactly 12 atomic mass units (amu)
1 amu is exactly 1/12 the mass of a carbon-12 atom
Atomic mass of any atom determined by comparing it with mass of C-12
Ex. H  1/12 C-12, so 1 amu

43. Divide by 100 – average marble mass = 2.75g
Average Atomic Masses
Average atomic mass  the weighted average of the atomic masses of the naturally occurring isotopes of an element
Ex. You have a box containing 2 size of marbles
25% are 2.00g each
75% are 3.00g each
25 x 2.00g = 50g 75 x 3.00g = 225g
50g + 225g = 275g
Divide by 100 – average marble mass = 2.75g

44. Method 1 25 x 2.00g = 50g 75 x 3.00g = 225g 50g + 225g = 275g Divide by 100 – average marble mass = 2.75g Method 2 25% = % = 0.75 (2.00g x 0.25) + (3.00g x 0.75) = 2.75g

45. Calculating Average Atomic Mass
Copper 69.17% Cu-63 – amu 30.83% Cu-65 – amu ( x ) + ( x ) = amu

46. Relating Mass to Numbers of Atoms
The relative atomic mass scale makes it possible to know how many atoms of an element are present in a sample of the element with a measurable mass
Three very important concepts provide the basis for relating masses in grams to numbers of atoms
The mole
Avogadro’s number
Molar mass

47. The Mole SI unit for an amount of substance (like 1 dozen = 12)
Mole (mol)  amount of a substance that contains as many particles as there are atoms in exactly 12g of C-12

48. Avogadro’s Number
Avogadro’s number  the number of particles in exactly one mole of a pure substance
6.022 x 1023
How big is that?
If 5 billion people worked to count the atoms in one mole of an element, and if each person counted continuously at a rate of one atom per second, it would take about 4 million years for all the atoms to be counted

49. Molar Mass Molar mass  the mass of one mole of a pure substance
Written in unit g/mol
Found on periodic table (atomic mass)
Ex. Molar mass of H = g/mol

50. Example
How many grams of helium are there in 2 moles of helium? 2.00 mol He x = ? g He 2.00 mol He x 4.00 g He = 8.00 g He 1 mole He

51. Practice Problems
What is the mass in grams of 3.50 mol of the element copper, Cu? 222 g Cu What is the mass in grams of 2.25 mol of the element iron, Fe? 126 g Fe What is the mass in grams of mol of the element potassium, K? 14.7 g K What is the mass in grams of mol of the element sodium, Na? g Na What is the mass in grams of 16.3 mol of the element nickel, Ni? 957 g Ni

52. 1. A chemist produced 11. 9 g of aluminum, Al
1. A chemist produced 11.9 g of aluminum, Al. How many moles of aluminum were produced? mol Al 2. How many moles of calcium, Ca, are in 5.00 g of calcium? mol Ca 3. How many moles of gold,Au, are in 3.60 × 10−10 g of gold? 1.83 × 10−12 mol Au

53. Conversions with Avogadro’s Number
How many moles of silver, Ag, are in 3.01 x 1023 atoms of silver?
Given: 3.01 × 1023 atoms of Ag
Unknown: amount of Ag in moles
Ag atoms × = moles Ag
3.01x1023atomsAg x =0.500 mol Ag

54. Practice problems
1. How many moles of lead, Pb, are in 1.50 × 1012 atoms of lead? 2.49 × 10−12 mol Pb 2. How many moles of tin, Sn, are in 2500 atoms of tin? 4.2 × 10−21 mol Sn 3. How many atoms of aluminum, Al, are in 2.75 mol of aluminum? 1.66 × 1024 atoms Al

55. 1. What is the mass in grams of 7. 5 × 1015 atoms of nickel, Ni. 7
1. What is the mass in grams of 7.5 × 1015 atoms of nickel, Ni? 7.3 × 10−7 g Ni 2. How many atoms of sulfur, S, are in 4.00 g of sulfur? 7.51 × 1022 atoms S 3. What mass of gold,Au, contains the same number of atoms as 9.0 g of aluminum,Al? 66 g Au