Download I. Atoms

I. Atoms
• Atom:
The smallest particle of an element
Electronically neutral: no + or - charge
• Dalton’s Atomic Theory: developed ~ 1800; based on experiments
1. Each element is composed of tiny *indivisible* particles called
atoms.
2. All atoms of a given element are *identical.* Atoms of different
elements are unique.
3. Atoms of different elements combine in simple whole number
ratios to form compounds (ex. 2:3)
4. Chemical reactions occur when atoms are separated, joined, or
rearranged. However, atoms of one element cannot change into
atoms of another element.
• Size of an atom: smaller than can be seen with the naked eye or a
standard microscope; about 1 x 10-10 m across (angstrom)
A pure copper coin the size of a penny contains about 2.4 x 1022
atoms of copper. By comparison, Earth’s population is about
6 x 109 people. If you could line up 100 000 000 copper atoms side
by side, they would produce a line only 1 centimeter long (about the
width of your pinky).
Iodine atoms (green) on a
platinum background
Japanese characters for
“atom” written w/ atoms
II. Structure of the Nuclear Atom
• Electron:
─ Charge is ─ 1
─ Located OUTSIDE atom’s nucleus
─ Mass is 0
─ Symbol is e─
*Discovered by JJ Thompson in 1897
• The cathode ray experiment (1897): proved that electrons existed
and that they have a negative charge
When a negative plate was placed
near tube, the ray bent AWAY
When a positive plate was placed
near tube, the ray bent TOWARDS
• Proton:
─ Charge is + 1
─ Located INSIDE atom’s nucleus
─ Mass is 1
─ Symbol is p+
• Neutron:
─ Charge is 0
─ Located INSIDE atom’s nucleus
─ Mass is 1
─ Symbol is n0
• Nucleus the tiny core of the atom containing all of its mass; has
a positive charge
•Scientists used to think the atom looked like this.
•It was called the “plum pudding model” of the atom
• Rutherford’s Gold Foil Experiment (1911): discovered the
“nuclear model” of the atom: proved that p+ & n0 are in atom’s
nucleus & the rest of the atom is mostly empty space
• Atomic mass unit (amu): the mass unit of an individual atom or
subatomic particle (p+, e-, n0)
Summing It Up:
Subatomic
particle
Location
in Atom
Electron
Outside
nucleus
Inside
nucleus
Proton
Neutron
Inside
nucleus
Charge
Mass
-1
0
+1
1
0
1
eNucleus
p+
n0
Note: The size of the nucleus in the atom on the right is exaggerated!
The nucleus is so tiny in comparison to the rest of the atom that if
the atom were the size of a football stadium, the nucleus would be
the size of a marble
III. Distinguishing Between Atoms
• Atomic number: ─ equal to the number of protons in an atom
─ in a neutral atom: # protons = # electrons
How to find on the periodic table: the WHOLE number
The Golden Rule: Every element has a different atomic
number. The number of protons identifies an element.
• Mass number:
─ equal to the number of protons + neutrons
─ # neutrons = Mass number – Atomic number
How to find on the periodic table: ROUND decimal to a WHOLE number
• Atomic mass: ─ weighted average of all of the isotopes of an element
calculated by using the formula on pg. 4
How to find on the periodic table:
The ENTIRE DECIMAL VALUE
SUMMING IT UP:
A
ZX
Format for representing element data:
Mass #
Atomic #
A
Z
X
Symbol
12
6 C
14
7 N
• Isotope: atoms of the same element with different numbers of
neutrons and different masses; # p+ and e- are the same
Example:
12
C
6
13
C
6
14
C
6
p+
6
6
6
e-
6
6
6
n0
6
7
8
HOW TO CALCULATE ATOMIC MASS FROM ISOTOPE DATA:
Formula: Atomic Mass = (Abundance x Mass) + (Abundance x Mass) +…
Must be written as a decimal , not %
(move decimal point 2 spaces left)
Note: The more abundant an isotope, the closer its mass will be to the
average atomic mass
SAMPLE PROBLEM 1: Element X has two natural isotopes:
Isotope 1: mass of 10.012 amu and relative abundance of 19.91 %
Isotope 2: mass of 11.009 amu and relative abundance of 80.09 %
Calculate the atomic mass of this element.
Atomic Mass = (Abundance x Mass) + (Abundance x Mass)
Atomic Mass = (0.1991 x 10.012 amu) + (.8009 x 11.009 amu)
Atomic Mass =
10.81 amu
SAMPLE PROBLEM 2: Mg has three isoptopes:
24Mg=
mass of 23.985 amu, 78.99% abundance
25Mg=
mass of 24.986 amu, 10.00% abundance
26Mg=
mass of 25.983 amu, 11.01% abundance
Calculate the atomic mass of this element.
Atomic Mass =
(0.7899 x 23.985 amu) + (0.1000 x 24.986 amu) + (0.1101 x 25.983 amu)
Atomic Mass = 24.31 amu
SAMPLE PROBLEM 3:
Carbon has three isotopes. Their masses are 12, 13, and 14 amu.
The average atomic mass of carbon is 12.011 amu.
Without doing a calculation, which isotope of carbon is most abundant and why?
The one w/ a mass of 12
12 is closest to the average mass of 12.011
IV. The Periodic Table: Organizing the Elements
• The Periodic Table:
an arrangement of the elements according
to similarities in their properties
• Dmitri Mendeleev (mid-1800’s):
▪ 1st to arrange elements in a logical, systematic way
▪ Arranged elements by MASS
• Henry Moseley (1913):
▪ Arranged elements by ATOMIC #
▪
Made the periodic table look like it does today
• Periodic Law: when elements are arranged by increasing atomic
number, their properties repeat
• Period: HORIZONTAL ROWS; numbered 1-7 from top to bottom
• Group:
▪ VERTICAL COLUMNS; numbered 1-8 from left to right
▪ Also called FAMILIES
▪ Elements in the same group have similar properties!
Group Numbers
Period Numbers
1
1
2
3
4
5
6
7
2
3 4
5
6
7
8
• THREE CLASSES OF ELEMENTS AND THEIR LOCATIONS:
1.) METALS all elements to LEFT of staircase EXCEPT H
2.) NONMETALS all elements to RIGHT of staircase AND H
3.) METALLOIDS all elements along the staircase EXCEPT Al
Transition Metals
Lanthanides
Actinides
A. Group 1: Alkali metals
B. Group 2: Alkaline earth metals
C. Group 7: Halogens
D. Group 8:
Noble gases
E. Transition Metals
F. Metalloids
G. Lanthanide/
Actinide
Transition Metals