Atomic Structure NotesName ______
I. Atom History
A. John Dalton (1766 – 1844)
1. He revisited and revised Democritus’ theory.
2. Dalton’s Theory (1803)
____a. All matter is composed of tiny indivisible particles called ______.
____b. All atoms of an element are ______.
____c. Atoms cannot be ______, ______, or ______.
____d. Atoms of different elements combine chemically to form ______. Ex.
____e. Atoms of 1 element ______into a new element during chemical
reactions. (Alchemists of the time were trying to change Pb into Au.)
B. JJ Thomson – 1897
1. Experiment - ______
a. To determine mass of single cathode ray particle.
b. Found it to be ______charged w/ mass less than hydrogen atom.
2. Discovered the ______.
3. He believed the atom was a big ______“blob” with little
negatively charged ______floating around in it.
- Called his model the “______” model.
- Which part of Dalton’s theory did Thomson disprove?
______
C. Ernest Rutherford – 1911
1. Experiment - ______
a. Thought alpha particles (+) would pass right through ______.
b. Some particles actually bounced back or were greatly deflected.
2. Discovered the ______.
3. Concluded that the atom:
a. was mostly ______through which electrons move
b. contained a ______, dense, ______charged area (______)
in its center
D. James Chadwick – 1932
1. Problem – Rutherford’s model couldn’t account for the total mass of atom. It was
always heavier than predicted.
2. He proposed existence of ______.
a. has mass nearly equal to the mass of a ______
b. carries no charge (______)
E. Niels Bohr – early 1900’s
1. Problem – Since oppositely charged particles attract each other, why didn’t the positive
nucleus draw the negative electrons into it, causing the atom to collapse?
2. He proposed that the electrons are arranged in 7circular ______
around the nucleus
II. Subatomic Particles and their Properties
A. The Pieces
Particle / Discoverer / Location / Charge / Symbol / Relative Mass / Actual MassB. The Whole
1. Inside nucleus
a. Composed of ______& ______
b. Overall ______charge
c. Very dense (Extremely small % of total volume of atom, BUT 99.97% of its ______)
2. Outside nucleus
a. 99.9% of atom is this empty space through which the ______travel.
b. Overall ______charge
C. How they fit together
1. Electrons are held within the atom due to their attraction to the nucleus.
2. Neutrons help to stabilize the nucleus
a. they act as “pillows” to separate the positive protons
b. too many neutrons can make the atom radioactive
3. Since all atoms are electrically neutral, ______= ______
III. How Atoms Differ:
A. Atomic Number
1. Henry Moseley (1912) – discovered that atoms of 2 different elements contain different
numbers of ______.
2. Atomic Number – The number of ______in an element
a. How many protons does a gold atom have?______An aluminum atom?______
b. How many electrons are in a gold atom?______An aluminum atom?______
B. Isotopes – Atoms of the ______element with different number of ______.
1. Disproved Dalton’s Theory (Dalton said all atoms of the same element were alike.)
2. All atoms of the same element MUST have same # of ______, but not necessarily
the same # of neutrons.
3. Example – Potassium (There are 3 different kinds of K atoms.)
Potassium – 3993.25% / Potassium – 40
6.73% / Potassium – 41
0.12%
Protons
Neutrons
Electrons
4. Two ways to write the symbol for an isotope:
potassium – 39 or
5. Mass Number = ______+ ______
(The atomic # identifies the element, the mass number identifies the isotope.)
6. Isotope Similarities
a. Chemically and physically alike (# of electrons determines the chemical behavior)
b. Same # of ______and ______(atoms)
7. Isotope Differences
a. Different # of ______
b. Different ______
c. Different ______
8. Examples: Fill in the chart for the atoms listed below:
Element / Atomic # / Mass # / # Protons / # Neutrons / # Electrons / SymbolNeon / 22
Calcium / 26
8 / 17
26 / 31
64 Zn
124 / 80
C. Atomic Mass – The weighted average mass of the ______of that element.
1. Expressed in AMU’s
2. AMU = atomic mass unit (1amu = 1/12 the mass of a carbon – 12 atom)
a. proton = 1 amub. neutron = 1 amuc. electron = 1/1840 amu
3. Usually not a whole number since it is a weighted ______of the isotopes.
4. Average Atomic Mass = (% abundance)(mass) + (% abundance)(mass) + ...
Isotope 1 Isotope 2
a. Ex1There are 2 isotopes of chlorine:
35Cl (75.77% abundance) and 37Cl (24.23%abundance)
b. Ex2 There are 2 isotopes of copper.
Copper-63 has an abundance of 69.1% and Copper-65 has an abundance of 30.9%.
Calculate copper’s average atomic mass.
IV. Ions
A. Ion – An atom or group of atoms that have a ______or ______charge.
They are formed when an atom ______or ______electrons.
****Elements want to have a ______(______), to become ______.
1. Cation – An atom with a ______charge.
a. Metals tend to form positive ions by losing electrons
b. Ex: Sodium-23 atomvsSodium-23 ion
2. Anion – An atom with a ______charge.
a. Non-metals tend to form negative ions by gaining electrons.
b. Ex: Fluorine-19 atomvs.Fluoride-35 ion
Draw a Bohr model of the following elements:
HeliumBoron Oxygen
3. Summary:p+ = e- (______)p+ > e- (______) p+ < e- (______)
4. Ex: Fill in the chart below:
Name / p+ / no / e- / Atomic # / Mass # / Charge / Atom, cation, or anion / symbolBoron
atom / 6
20 / 21 / +2
17 / 35 / -1
Chapter 5: The Periodic Table
I: Periodic Table arrangement:
- Elements ordered by ______
- Elements with similar ______fall into ______
A. Groups or ______
- Designated by numbers ______
- ______
- Group number (ones) represents # of valence electrons – ______
- Similar ______and ______
Which of the following elements act like carbon? Li, Na, Al, Si, P, Ga, Ge, Cl F
B. Periods
- Designated by numbers ______
- ______
- Number of ______
*** You are more like your ______than you are like your ______
C. Blocks of Elements – designated by ______
D. Zig-Zag Line
- Separates ______from ______(label them)
Metals
- ______
- ______
- ______
Non-metals
- ______
- ______
- ______
Metalloids
- found along and to either side of the zig-zag line, except ______
- some properties of ______and some properties of ______
V. Periodic Table Groups
Group 1: ______-
Group 2: ______-
Group 3-12: ______-
Group 17: ______-
Group 18: ______-
VI. Periodic Table Trends
A. Oxidation Number
Ion – an atom (or group of combined atoms) that ______because of the ______Ex:
Group # / Valence e- / Oxidation Number / Ion Ex.Metals
1
2
13
14
Non-
metals
15
16
17
18
Oxidation Number - the ______on an _____
atoms ______or ______electrons to get ____ electrons in their ______(octet rule – atoms always want a ______valence shell).
B. Atomic Radii –
Ex: Li / C has the larger atomic radius
Ex: Ba / Mg has the larger atomic radius
Which element has the largest atomic radius?
C. Ionic Radii – distance from the center of the nucleus to the outer edge of the ion
Cations - ______electrons, less e- means protons are now pulling fewer e- resulting is a ______radius.
Anions - ______electrons, more e- means protons are now pulling more e- and also more e- are repelling each other pushing them out. Results in a ______radius.
D. Ionization Energy - the energy required to ______from an atom (kJ/mol)
- ______
- ______
Ex: Ca / Br has a higher IE
Ex: Na / Cs has a higher IE
Which element has the highest IE?
IE for Period 2 Elements (kJ/mol)
Period 2Li / Be / B / C / N / O / F / Ne
520 / 900 / 800 / 1090 / 1400 / 1310 / 1680 / 2080
E. Electronegativity – ______, a measure of an atom’s ability to ______
Ex: Be / O has a higher electronegativity
Ex: N / Sb has a higher electronegativity
What is the most electronegative element?
F. Reactivity – refers to how likely an atom is ______, depends on ...
-how easily electrons can be ______
-how easily electrons can be ______
What is the most reactive metal?
What is the most reactive nonmetal?
Recap:
a)Largest Atomic Radii =
b)Largest Ionic Radii =Smallest Ionic Radii =
c)Highest Ionization Energy =
d)Highest Electronegativity =
1