# study guide chapter 3

```STUDY GUIDE SECTION 3-1
Dalton’s Atomic Theory
1. All matter is composed of indivisible, extremely small particles called atoms.
2. Atoms are indestructible and unchangeable. Chemical reactions involve the
combination of atoms, not the destruction of atoms.
3. All atoms of a given element are identical. Atoms of different elements have different
properties.
4. Compounds are formed when atoms of different elements combine chemically. When
elements react to form compounds, they react in defined, whole number ratios.
5. In chemical reactions, atoms are combined, separated, or rearranged.
Modern Atomic Theory
1. All matter is made up of very small particles called atoms.
2. Atoms of the same element have the same chemical properties, while atoms of different
elements have different chemical properties.
3. While individual atoms of a given element may not all have the same mass (due to
differences in nuclear structure), any natural sample of the element will have a definite
average mass that is characteristic of the element and different from that of any other
element.
4. Compounds are formed when atoms of two or more elements unite, with each atom
losing its characteristic properties as a result of this combination.
5. Atoms are not subdivided in physical or chemical reactions.
Law of Conservation of Mass: The total mass of a system remains the same whether the
elements are combined, separated, or rearranged
A
+
1 mass unit
AB 
B
AB
+ 3 mass units

A
B
+
4 mass units
4 mass units1 mass unit + 3 mass units
AB
4 mass units

+
AC
+
5 mass units
AA
+
BC
2 mass units
+
7 mass units
LAW OF DEFINITE COMPOSITION: 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.
LAW OF MULTIPLE PROPORTIONS: If two or more different compounds are composed off the
same two elements, the masses of the second element combined with a certain mass of the first element
can be expressed as ratios of small whole numbers
CO
A
+ B
1unit + 3 units 4units
CO2
A +2BAB2
carbon monoxide
 AB
carbon dioxide
3-2 Structure of the Atom
Atom: smallest particles of an element that can exist either alone or in combination with other
atoms.
Atomic Structure:
Nucleus
1. contains neutrons & protons
2. positively charged dense central portion of the atom
3. contains nearly all the atom’s mass, but takes up a very little volume of the atom—nearly
99.9% of mass is in the atom’s nucleus
Neutron
 electrically neutral subatomic particles in nucleus
 the mass of a neutron is about the same as the mass of a proton (1.675 x 10-24g)
 relative charge 0
 mass number 1
 relative mass 1.00866 u
 symbols, no, 1n
0
Proton
 positively charged subatomic particles in nucleus that have a charge equal in magnitude
to the negative charge of the electron
 mass is 1.673 x 10-24 g
 mass is 1836 times greater than an electron
 electric charge +1
 mass number 1
 relative mass 1.007276 u
 symbol 1H , p+
1
Electron
 surround the nucleus
 negatively charged subatomic particles
 very large area compared to the nucleus
 relative electric charge –1
 mass number 1
 relative mass 0.000 548 6 u
 actual mass 9.109 x 10-28 g
 symbol e-, 0e
-1
DISCOVERY OF SUBATOMIC PARTICLES:
J.J. Thompson (Crooke’s Tube—Cathode Ray Tube)
 subjected cathode rays to magnetic and electrical fields and measured their deflection.
 discovered that these rays were electrons.
 determined the ratio of the electrical charge (q) to mass (m) q/m
Anode: A positive electrode; the electrode where oxidation occurs (lose an electron become
positive)
Cathode: A negative electrode; the electrode where reduction occurs (gain an electron become
negative).
Cathode Ray: The beam of electrons in a gas discharge tube.
Millikan: Oil Drop Experment
 confirmed Thompson’s ideas
 determined the charge on an electron (-1 or 1.5921 x 10-19 Coulomb)
 determined the mass of an electron (0.000549amu or 1/1850 amu or 9.109 x 10-28g)
 discovered the neutron
 charge q=0
 mass=1amu=1.675 x 10-24 g
Rutherford GOLD FOIL EXPERIMENT
 determined the atom consisted of a small nucleus
 nucleus consists of some protons and some neutral mass
 nucleus surrounded by orbiting electrons
ISOTOPES OF Hydrogen
Isotope: Atoms of the same element that differ in mass; they have the same number of protons,
but a different number of neutrons
Protium: nucleus consists of one proton only, has one electron (99.985% of Hydrogen is
protium)
Deuterium: has one proton and one neutron in the nucleus and one electron
Tritium: radioactive, contains one proton, two neutrons, and one electron
NUCLEAR FORCES: short-range proton-proton, neutron-proton, neutron-neutron forces that
hold the nucleus together
METHODS TO DESIGNATE ISOTOPES:
Hyphen notation: protium Hydrogen-1
Deuterium Hydrogen-2
Tritium Hydrogen-3
Nuclear symbol:
protium
1H
1
Deuterium
1H
2
Tritium
1H
3
3-3 WEIGHING and COUNTING ATOMS
Atomic number (Z): the number of protons in the nucleus of each atom of that element
3
Li
Lithium
6.941
[He]2s1
Example: lithium—atomic number is 3
How many protons in Lithium? 3
THE NUMBER OF PROTONS DETERMINES THE IDENTITY OF THE ELEMENT
Neutral atoms: no overall charge—electrically neutral Same number of protons and
electrons
EXAMPLES:
lithium—atomic number is 3
How many electrons in Lithium? 3
How many electrons in H? 1
How many electrons in Beryllium? 4
What element has 47 electrons? Silver
MASS NUMBER: the total number of protons and neutrons in the nucleus of an atom
Example: lithium—mass number is 7
How many neutrons in Lithium? 4
How many neutrons in H? 0
How many neutrons in Beryllium? 5
Isotopes: atoms of the same element that have different masses (due to a different number of
neutrons)
Nuclide: a general term for any isotope of any element
PROBLEM: How many protons, electrons, and neutrons are there in an atom of chlorine37?
RELATIVE ATOMIC MASSES:
 Relative means in relation to
 more convenient to use
 standard of measurement—Carbon-12
 A single nuclide C-12 is assigned a mass of EXACTLY 12 atomic mass units.
 Atomic mass unit symbol u
 ONE ATOMIC MASS UNIT, 1u, is EXACTLY 1/12 the mass of a carbon-12 atom, or
1.6605402 x 10-24 gram.
 The mass of an atom expressed in atomic mass units is called the ATOMIC MASS of the
atom.
 Hydrogen has an atomic mass about 1/12th of a C-12, or 1 amu, 1u
AVERAGE ATOMIC MASS: is the weighted average of the atomic masses of the naturally
occurring isotopes of an element
Calculation of a weighted average
Example: Copper
Cu-63 69.17%
62.939598 u
Cu-65 30.83%
64.927793u
.69l7 x 62.939598 u = 43.54 u
.3083 x 64.927793 u = 20.01 u
total
=63.55 u
Individual atomic mass-Exact mass in atomic mass units of one atom (different for isotopes of
same element)
Average atomic mass—Average mass in atomic mass units of all natural isotopes on an element
Mass Number—Total number of protons and neutrons in a nucleus of an isotope (it is the
average atomic mass rounded off to the nearest whole number)
Atomic Number—Number of protons in an element
The mole: is the amount of a substance that contains the same number of particles as the
number of atoms in exactly 12-g of carbon-12. The abbreviation for mole is mol
Avogadro’s number: 6.022 137 x 1023 particles in one mole of a pure substance
In terms of units 6.022 137 x 1023 particles/mole
Molar mass: is the mass in grams of one mole of a pure substance.
EXAMPLES:
How many grams of helium contain Avogadro’s number of helium atoms? 4.00 g
CONVERSION FACTORS: THIS IS A MOLMASS problem
If you have 2.00 mol of He, how many grams do you have?
4.00 g He/1.00 mol He x 2.00 mol He = 8.00 g He
What is the mass in grams of 3.50 mol of the element copper?
3.50 mol Cu x
63.5 grams Cu
=
222 g Cu
mol Cu
This is a massmole problem
EXAMPLE: A chemist produced 11.9 grams of aluminum. How many moles of aluminum
have been produced?
11.9 g of Al x one mol Al = 0.441 mol Al
27.0 g of Al
Atoms  moles
How many moles of silver are in 3.01 x 1023 atoms?
3.01 x 1023 atom Ag x
one mol Ag = .500 mol Ag
6.023 x 1023 atoms Ag
atomsmolesmass
SAMPLE PROBLEM: What is the mass in grams of 1.20 x 108 atoms of copper?
1.20x108Cu atoms x 1mol Cu x 63.5 g Cu =1.27x10-14gCu
6.023x1023Cu atoms mol Cu
ATOMIC MASS (or weight) (symbol A) the number of
protons plus the number of neutrons in the
(average mass is in the
periodic table)
# of Neutrons (n) = A - Z
nucleus of an atom.
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