1. Chemistry: Atoms First Julia Burdge & Jason Overby
Chapter 2
Atoms and the
Periodic Table
Kent L. McCorkle
Cosumnes River College
Sacramento, CA
Copyright (c) The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Copyright (c) The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

2. Subatomic Particles and Atomic Structure
2.2
In the late 1800’s, many scientists were doing research involving radiation, the emission and transmission of energy in the form of waves.
They commonly used a cathode ray tube, which consists of two metal plates sealed inside a glass tube from which most of the air has been evacuated.

3. Subatomic Particles and Atomic Structure
When metal plates are connected to a high-voltage source, the negatively charged plate, or cathode, emits an invisible ray.
The cathode ray is drawn to the anode where it passes through a small hole.
Although invisible, the path is revealed when the ray strikes a phosphor-coated surface producing a bright light.

4. Subatomic Particles and Atomic Structure
Researches discovered that like charges repel each other, and opposite charges attract one another.
J. J. Thomson ( ) noted the rays were repelled by a plate bearing a negative charge, and attracted to a plate bearing a positive charge.

5. Subatomic Particles and Atomic Structure
This prompted him to propose the rays were actually a stream of negatively charged particles.
These negatively charged particles are called electrons.
By varying the electric field and measuring the degree of deflection of cathode rays, Thomson determined the charge-to-mass ratio of electrons to be 1.76×108 C/g. (C is coulomb, the derived SI unit of electric charge.)

6. Subatomic Particles and Atomic Structure
R. A. Millikan ( ) determined the charge on an electron by examining the motion of tiny oil drops.
The charge was determined to be ×10-19 C.

7. Subatomic Particles and Atomic Structure
Knowing the charge, he was then able to use Thomson’s charge-to-mass ratio to determine the mass of an electron.

8. Subatomic Particles and Atomic Structure
Wilhelm Rontgen ( ) discovered X-rays. They were not deflected by magnetic or electric fields, so they could not consist of charged particles.
Antoine Becquerel ( ) discovered radioactivity, the spontaneous emission of radiation.
Radioactive substances, such as uranium, can produce three types of radiation.

9. Subatomic Particles and Atomic Structure
Alpha (α) rays consist of positively charged particles, called α particles.
Beta (β) rays, or β particles, are electrons so they are deflected away from the negatively charged plate.
Gamma (γ) rays, like X-rays, have no charge and are unaffected by external electric or magnetic fields.

10. Subatomic Particles and Atomic Structure
Ernest Rutherford used α particles to prove the structure of atoms.
The majority of particles penetrated the gold foil undeflected.
Sometimes, α particles were deflected at a large angle.
Sometimes, α particles bounced back in the direction from which they had come.

11. Subatomic Particles and Atomic Structure
Rutherford proposed a new model for the atom:
Positive charge is concentrated in the nucleus.
The nucleus accounts for most of an atom’s mass and is an extremely dense central core within the atom.
A typical atomic radius is about 100 pm
A typical nucleus has a radius of about 5×10–3 pm
1 pm = 1×10–12 m

12. Atoms First
2.1
An atom is the smallest quantity of matter that still retains the properties of matter.
An element is a substance that cannot be broken down into two or more simpler substances by any means.
Examples: gold, oxygen, helium
Atoms can also be divided smaller and smaller until eventually only a single atom remains. Dividing it any smaller would give pieces that are no longer an atom.
A DVD collection can be separated into smaller numbers until you have just one DVD left. But a single DVD cannot be separated into smaller pieces that are still DVDs.

13. What are atoms made of?
Protons: positively charged particles that are housed in the nucleus
of an atom and have significant mass
Neutrons: neutral particles that are housed in the nucleus. They act to
hold the protons in place since like charges repel each
other. Neutrons have significant mass
Electrons have negligible mass, have a negative charge and are allowed
to roam freely in the electron cloud so they take up significant
volume in the atom
Taylor 2010

14. Atomic Number, Mass Number, and Isotopes
2.3
All atoms can be identified by the number of protons and neutrons they contain.
The atomic number (Z) is the number of protons in the nucleus.
Atoms are neutral, so it’s also the number of electrons.
Protons determine the identity of an element. For example, nitrogen’s atomic number is 7, so every nitrogen has 7 protons.
The mass number (A) is the total number of protons and neutrons.
Protons and neutrons are collectively referred to as nucleons.
Mass number
(number of protons + neutrons)
Elemental symbol
Atomic number
(number of protons)

15. Atomic Number, Mass Number, and Isotopes
Most elements have two or more isotopes, atoms that have the same atomic number (Z) but different mass numbers (A).
1 proton
0 neutrons
1 proton
1 neutron
1 proton
2 neutrons
Isotopes of the same element exhibit similar chemical properties, forming the same types of compounds and displaying similar reactivities.

16. Worked Example 2.1
Determine the numbers of protons, neutrons, and electrons in each of the following species: (a) Cl, (b) Cl, (c) K, and (d) carbon-14. 35 17 37 41
Strategy Recall the superscript denotes the mass number (A) and the subscript denotes the atomic number (Z). If no subscript is shown, the atomic number can be deduced from the elemental symbol or name. Atoms are neutral so the number of electrons equals the number of protons.
Solution
Z = 17, so 17 protons
A = 35, so = 18 neutrons
# of electrons = # of protons, so 17 electrons
Element is chlorine again, so Z must be 17; 17 protons
A = 37, so = 20 neutrons
17 protons, so 17 electrons

17. Worked Example 2.1 (cont.)
Determine the numbers of protons, neutrons, and electrons in each of the following species: (a) Cl, (b) Cl, (c) K, and (d) carbon-14. 35 17 37 41
Strategy Recall the superscript denotes the mass number (A) and the subscript denotes the atomic number (Z). If no subscript is shown, the atomic number can be deduced from the elemental symbol or name. Atoms are neutral so the number of electrons equals the number of protons.
Solution
(c) Potassium’s atomic number is 19, so 19 protons
A = 41, so = 22 neutrons
# of electrons = # of protons, so 19 electrons

18. Worked Example 2.1 (cont.)
Determine the numbers of protons, neutrons, and electrons in each of the following species: (a) Cl, (b) Cl, (c) K, and (d) carbon-14. 35 17 37 41
Strategy Recall the superscript denotes the mass number (A) and the subscript denotes the atomic number (Z). If no subscript is shown, the atomic number can be deduced from the elemental symbol or name. Atoms are neutral so the number of electrons equals the number of protons.
Solution
(d) Carbon-14 can also be represented as C
Carbon’s atomic number is 6, so 6 protons
A = 14, so = 8 neutrons
6 protons, so 6 electrons 14
Think About It Verify that the number of protons and the number of neutrons for each example sum to the mass number that is given. In part (a), there are 17 protons and 18 neutrons, which sum to give a mass number of 35, the value given in the problem. In part (b), 17 protons + 20 neutrons = 37. In part (c), protons + 22 neutrons = 41. In part (d), 6 protons + 8 neutrons = 14.

19. Do You Understand Isotopes?
How many protons, neutrons, and electrons are in C 14 6 ?
6 protons, 8 (14 - 6) neutrons, 6 electrons
How many protons, neutrons, and electrons are in C 11 6 ?
6 protons, 5 (11 - 6) neutrons, 6 electrons
2.3

20. Ions Atoms are neutral—meaning that the number of protons
is equal to the number of electrons
If an atom loses or gains electrons the atom is
No longer neutral but has a charge.

21. cation – ion with a positive charge
An ion is an atom, or group of atoms, that has a net positive or negative charge.
cation – ion with a positive charge
If a neutral atom loses one or more electrons
it becomes a cation. Na
11 protons
11 electrons
Na+
11 protons
10 electrons
anion – ion with a negative charge
If a neutral atom gains one or more electrons
it becomes an anion.
Cl-
17 protons
18 electrons Cl
17 protons
17 electrons

22. A monatomic ion contains only one atom
Na+, Cl-, Ca2+, O2-, Al3+, N3-
A polyatomic ion contains more than one atom
OH-, CN-, NH4+, NO3-

23. Nuclear Symbol of Ions X A C Z X = element symbol A = atomic number
Z = mass number
C = charge of Ion
Number of Protons = Z
Numbers of Neutrons = A – Z
Number of electrons = Z - C

25. Do You Understand Ions? Al Se
How many protons, neutrons and electrons are in ? Al 27 13 3+
Protons = 13
Neutrons = = 14
Electrons = 13 – 3 = 10
How many protons, neutrons and electrons are in ? Se 78 34 2-
Protons = 34
Neutrons = 78 – 34 = 44
Electrons = 34 – (-2) = 36
Taylor 2012

26. Atomic mass is the mass of an atom in atomic mass units (amu)
Micro World
atoms & molecules
Macro World
grams
Atomic mass is the mass of an atom in atomic mass units (amu)
By definition:
1 atom 12C “weighs” 12 amu
On this scale
1H = amu
16O = amu
3.1

27. 1 amu = 1/12 the mass of a carbon-12 atom
2.4
Average Atomic Mass
Atomic mass is the mass of an atom in atomic mass units (amu).
1 amu = 1/12 the mass of a carbon-12 atom
The average atomic mass on the periodic table represents the average mass of the naturally occurring mixture of isotopes.
Average mass (C) = (0.9893)( amu) + (0.0107)( amu)
Isotope
Isotopic mass (amu)
Natural abundance (%)
12C
98.93
13C
1.07
= amu

28. Worked Example 2.2
Oxygen is the most abundant element in both Earth’s crust and the human body. The atomic masses of its three stable isotopes, O ( percent), O ( percent), O (0.205 percent), are , , and amu, respectively. Calculate the average atomic mass of oxygen using the relative abundances given in parentheses. 16 8 17 8 18 8
Strategy Each isotope contributes to the average atomic mass based on its relative abundance. Multiplying the mass of each isotope by its fractional abundance (percent value divided by 100) will give its contribution to the average atomic mass.
Solution
( )( amu) + ( )( amu) + ( )( amu)
= amu
Think About It The average atomic mass should be closest to the atomic mass of the most abundant isotope (in this case, oxygen-16) and, to four significant figures, should be the same number that appears in the periodic table on the inside front cover of your textbook (in this case, amu).

29. Average atomic mass of lithium:
Natural lithium is:
7.42% 6Li (6.015 amu)
92.58% 7Li (7.016 amu)
Average atomic mass of lithium:
(7.42 x 6.015) + (92.58 x 7.016)
100
= amu
3.1

30. Average atomic mass (6.941)

31. Dozen = 12 Names associated with an amount Pair = 2
Can you think of any more?????

32. The Mole….. The mole (mol) is the amount of a substance that
contains as many elementary entities as there
are atoms in exactly grams of 12C
1 mol = NA = x 1023
Avogadro’s number (NA)

33. Molar mass is the mass of 1 mole of in grams marbles atoms
eggs
shoes
Molar mass is the mass of 1 mole of in grams
marbles
atoms
1 mole 12C atoms = x 1023 atoms = g
1 12C atom = amu
1 mole 12C atoms = g 12C
1 mole lithium atoms = g of Li
For any element
atomic mass (amu) = molar mass (grams/mol)
3.2

34. 2.6
The Mole and Molar Mass
The mole is defined as the amount of a substance that contains as many elementary entities as there are atoms in exactly 12 g of carbon-12.
This experimentally determined number is called Avogadro’s number (NA).
We normally round this to 6.022×1023.
1 mole = 6.022×1023, just like 1 dozen = 12 or 1 gross = 144.
NA = x 1023

35. Worked Example 2.3
Calcium is the most abundant metal in the human body. A typical human body contains roughly 30 moles of calcium. Determine (a) the number of Ca atoms in moles of calcium and (b) the number of moles of calcium in a sample containing 1.00×1020 Ca atoms.
Strategy Use Avogardo’s constant, 1 mole = 6.022×1023, to convert from moles to atoms and from atoms to moles.
Solution
30.00 mol Ca ×
1.00×1020 Ca atoms ×
6.022×1023 Ca atoms
1 mol Ca
= 1.807×1025 Ca atoms
1 mol Ca
6.022×1023 Ca atoms
= 1.66×10-4 mol Ca
Think About It Make sure that units cancel properly in each solution and that the result makes sense. In part (a), for example, the number of moles (30) is greater than one, so the number of atoms is greater than Avogadro’s number. In part (b), the number of atoms (1×1020) is less than Avogadro’s number, so there is less than a mole of substance.

36. Salt (Sodium Chloride)
The Mole
One mole each of some familiar substances:
Helium (in balloon)
Water
Sugar (Sucrose)
Aluminum
Copper
Salt (Sodium Chloride)

37. Molar Mass
The molar mass of a substance is the mass in grams of one mole of the substance.
By definition, the mass of a mole of carbon-12 is exactly 12 g.
Mass of 1 carbon-12 atom: exactly 12 amu
Mass of 1 mole of carbon-12: exactly 12 g
Although molar mass specifies the mass of one mole, making the units (g), we usually express molar masses in units of grams per mole (g/mol) to facilitate cancellation of units in calculations.

38. Worked Example 2.4
Determine (a) the number of moles of C in g of carbon, (b) the number of moles of He in g of helium, and (c) the number of moles of Na in g of sodium.
Think About It Always double-check unit cancellations in problems such as these–errors are common when molar mass is used as a conversion factor. Also make sure that the results make sense. For example, in the case of part (c), a mass smaller than the molar mass corresponds to less than a mole.
Strategy Molar mass of an element is numerically equal to its average atomic mass. Use the molar mass for each element to convert from mass to moles.
Setup (a) The molar mass of carbon is g/mol. (b) The molar mass of helium is g/mol. (c) The molar mass of sodium is g/mol.
Solution
25.00 g C ×
10.50 g He ×
15.75 g Na ×
1 mol C
12.01 g C
= mol C
1 mol He
4.003 g He
= mol He
1 mol Na
22.99 g Na
= mol Na

39. Interconverting Mass, Moles, and Number of Atoms
Molar mass is the conversion factor from mass to moles, and vice versa.
Avogadro’s constant converts from moles to atoms.

40. Worked Example 2.5
Determine (a) the number of C atoms in g of carbon, and (b) the mass of helium that contains 6.89×1018 He atoms.
Strategy Use the conversions depicted in the previous slide to convert (a) from grams to moles to atoms and (b) from atoms to moles to grams.
Think About It A ballpark estimate can help you prevent common errors. For example, the mass in part (a) is smaller than the molar mass of carbon. Therefore, you should expect a number of atoms smaller than Avogadro’s number. Likewise, the number of atoms in part (b) is smaller than Avogadro’s number. Therefore, you should expect a mass of helium smaller than the molar mass of helium.
Setup (a) The molar mass of carbon is g/mol. (b) The molar mass of helium is g/mol. NA = 6.022×1023
Solution
0.515 g C ×
6.89×1018 He atoms ×
1 mol C
12.01 g C
6.022 ×1023 C atoms
1 mol C ×
= 2.58×1022 C atoms
1 mol He
6.022 ×1023 He atoms
4.003 g He
1 mol He ×
= 4.58×10-5 g He

41. How many atoms are in 0.551 g of potassium (K) ? 1 mol K = 39.10 g K
1 mol K = x 1023 atoms K
1 mol K
39.10 g K x x
6.022 x 1023 atoms K
1 mol K =
0.551 g K
8.49 x 1021 atoms K

42. What is the mass of 1.21  1020 atoms of sulfur (in mg) ?