Download U2notes2015

Unit 2 Atomic Structure
Atomic Theory
Dalton’s Postulates (1808)
1.
An element is composed of extremely small particles called atoms.
2.
Atoms of a given element have the same size, mass, and properties whereas
atoms of different elements have different size, mass and properties.
3.
Atoms cannot be subdivided, created or destroyed
4.
Atoms of two or more elements combine in small whole number ratios to form
compounds – supports law of definite composition and law of multiple
proportion
5.
In a chemical reaction, atoms combine, separate or rearrange.- supports Law
of conservation
 Dalton’s Model of Atom: a solid sphere
ATOM – defined as the smallest particle of an element that can enter into a chemical
reaction
Basic Laws of Chemistry explained the postulates of the Atomic Theory
1. Lavoisier’s Law of conservation of mass: mass can neither be created or
destroyed but only changed in form, the mass
of reactants equal the mass of products
2. Proust’s Law of definite proportion(1799):
a compound always contains the
same elements in the same proportions
by mass even if we look at different
samples
3. Dalton’s Law of multiple proportions:
when 2 elements form a series of
compounds, the ratios of the masses of
nd
the 2 element that combine with 1 gram of the
first element can always be reduced to small whole
numbers.
Part 2 and 3 of Dalton’s theory has been modified after
experimental evidence proved them incorrect; Modern
Atomic Theory states
 #3 Atoms were discovered divisible (made up of subatomic particles) between
1850-1900; Atoms cannot be created, or destroyed in ordinary chemical
reactions. However, they CAN in nuclear reactions.
 #2 Isotopes (atoms of the same element but with different masses) were
discovered; atoms of an element have a average mass which is specific to that
element
Components of the atom
 If the Houston Astrodome was an atom, a marble placed in the stadium would be
the size of the nucleus
 Most of the mass of the atom is in the nucleus
 Most of the atom is empty space
Electron Cloud
The Electron: negatively charged particles located in the outer regions of the atom
 JJ.Thomson discovered its charge using a cathode ray tube
 relative charge of –1
 Millikan discovered its actual mass
 Actual mass = 9.109 x 10-28 g relative mass 1 amu;; approximately 2000 times
smaller than the proton or neutron
 JJ Thomson formulated the PLUM PUDDING MODEL. (the positively charged
pudding had negatively charged electrons(plums)
embedded in it)
Nucleus
 Discovered by Rutherford in the Gold Foil
experiment (1910)
 alpha particles (+ helium nuclei) were shot through a piece of gold foil. It was
assumed that the rays would come out on the other end but some were deflected
and even bounced back causing the belief that there was a small, dense,
positively charged area in the atom
 formulated the “Rutherford
Model”
The Proton
 relative charge +1
 relative mass 1.00728 amu or actual mass of 1.6726 x 10-24 g
The Neutron- discovered by Chadwick(1932)
 relative charge : neutral
 relative mass 1.00867 amu or 1.6749 x 10-24 g
Atomic Mass
 Atomic masses are relative, representing the mass of an atom of 1 element
compared to the mass of another. We used carbon-12 as the standard and
assign it a mass of exactly 12 amu.
 is the average of all the naturally occurring isotopes; expressed as amu; depends
on the number of isotopes and percent abundance
 Equation to calculate Average Atomic Mass (AAM)is:
average atomic mass of Y = (atomic massY1) x % Y1 + (atomic massY2) x % Y2 + …
Example 1: There are two isotopes of carbon 12C with a mass of 12.00000
amu(98.892%), and 13C with a mass of 13.00335 amu (1.108%). What is the average
atomic mass of carbon?
Example 2: There are two isotopes of nitrogen,one with an atomic mass of 14.0031amu
and one with a mass of 15.0001 amu. What is the percent abundance of each?
Experimental determination of Relative Mass
 instrument used is a mass spectrometer
 gaseous atoms or molecules at very low pressure are ionized by removing one or
more electrons. The cations formed are accelerated toward a magnetic field,
deflecting them from a straight path. The extent of deflection is inversely
proportional to the mass of the ion. By measuring the voltages required to bring
the 2 ions of different mass to the same point will determine its relative mass
 Analyze Mass Spectrographs: see additional handouts
Atomic number- the number of protons in a nucleus; symbol is Z
 in a neutral atom, the number of protons is equal to the number
of electrons
Mass number is the number of protons and neutrons in a nucleus; symbol is A
Calculating neutrons: Mnemonic: MAN; mass number (A) minus the atomic number
(Z) equals the number of neutrons in an atom
Isotopes: atoms that contain the same number of protons but a different number of
neutrons; they have different masses
 represented by a nuclear symbol(isotopic symbol)
A
X
Z
X = Element Symbol
Z = Number of Protons
A = Number of Protons + Neutrons
Example 3:
Write the isotopic symbol for the 3 isotopes of silicon in which there are 14, 15, and 16
neutrons.
Example 4:
How many protons, neutrons, and electrons are in
197Au
?
Example 5:
How many neutrons are in the 3 isotopes of magnesium-24, magnesium-25, and
magnesium-26?
Example 6: Write the symbol for the atom that has an atomic number of 9 and a mass
number of 19. How many electrons and neutrons does this atom have?