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1.1.1 The Changing Atom
You will be able to…
• Describe how the model of the atom has changed over the
years, and how it continues to do so.
• Understand that scientific knowledge is always evolving.
• Describe how new theories are accepted by scientists.
Atoms and their fundamental particles
• Chemistry is the study of matter and the
changes it undergoes
• Matter is defined as being composed of
atoms (first proposed by Democritus in
400BC)
• Dalton’s Atomic Theory (1808) – All matter
consists of atoms which cannot be created,
destroyed or split
Elements and atoms
• Elements are substances composed
of only one type of atom
• The composition of atoms was
established by a series of experiments
• The ‘plum pudding’ model had
suggested that atoms were made up of
positive ‘sponge’ material with negative
electrons distributed throughout like
currants
The composition of atoms
• However, this was all changed in 1909, when Rutherford’s
students, Geiger and Marsden, performed an experiment
where they fired alpha particles* at a sheet of gold foil and
found that some immediately bounced back!
• This led to Rutherford proposing that atoms are made up of a
positive nucleus with the electrons at some distance away
The Geiger-Marsden Experiment
The radioactive source produces alpha particles which are collimated
into a beam and directed at a gold foil. The alpha particles scatter off the
foil and are detected by a flash of light when they hit the deflecting
screen.
The Rutherford model of the atom
•
Atoms are made up
of a positive nucleus
•
Electrons adopt
‘planetary orbits’
around the nucleus
•
Atoms are mainly
space!
1.1.2 Atomic Structure
You will be able to…
•
Describe protons, neutrons and electrons
•
Describe the distribution of mass and charge in an atom
•
Describe the contribution of protons and neutrons to the nucleus in
terms of atomic number and mass number
•
Explain the term isotopes
•
Deduce the atomic structure in atoms and ions
What is atomic number?
What are the particles in this helium atom?
neutron
electron
helium
nucleus
proton
Atomic number = the number of protons in an atom
What is the atomic number of helium?
Protons, Neutrons, Electrons
particle
relative
mass
charge
proton
1.0
1+
neutron
1.0
0
electron
1/2000
1-
Atoms of elements contain equal numbers of protons
and electrons.
How does this affect the charge of such atoms?
Atoms of elements have no charge, they are neutral.
What is the charge on a fluorine atom?
19
F
9
9 protons
charge = +9
9 electrons
charge = - 9
10 neutrons
charge = 0
total charge = 0
How are the number of electrons and atomic number
related in a neutral atom? They are the same.
Isotopes
Ions
Ions are atoms that have either extra electrons
added or electrons removed. e.g.
Atoms
Protons
Electrons
Lose 1 electron
+1 ion
Gain 1 electron
Protons
Electrons
So in ions the
number of
electrons no
longer equals the
number
of protons
-1 ion
Protons
Electrons
1.1.3 Atomic Masses
You will be able to…
• Explain why 12C is used as the standard measurement of relative
mass.
•
Define the terms relative isotopic mass and relative atomic mass.
•
Calculate relative atomic masses.
•
Work out relative molecular masses and relative formula masses.
What is an element?
• A substance that consists of atoms having the same chemical
properties
• A substance that cannot be broken down into simpler
components using chemical techniques
e.g
Hydrogen
Helium
Oxygen
Copper
Gold
H
He
O
Cu
Au
What is an atom?
• The smallest particle of an element that has the chemical
properties of that element
• Atoms are the building blocks of elements
What is a compound?
• A substance consisting of atoms of two or more elements in a
definite ratio i.e a specific combination of atoms
• A specific combination of elements that can be broken down
by chemical techniques
• Examples
–
–
–
–
Sodium Chloride
Water
Ethanol
Viagra
NaCl
H 2O
C2H6O
C22H30N6O4 S
What is a Molecule?
• A substance containing two or more atoms
• Examples
–
–
–
–
Hydrogen gas
Nitrogen gas
Ethanol
Viagra
H2
N2
C2H6O
C22H30N6O4 S
Diatomic
What is an ion?
• When an atom loses or gains an electron it is called an ion
e.g
Sodium ready loses an electron to make a cation
Na  Na+ + eChlorine readily accepts an electron to make an anion
Cl + e-  Cl-
What is an ion?
• Ionic compounds are compounds resulting from a reaction
between ions
e.g Sodium Chloride (Salt)
Na+ + Cl-  NaCl
As the compound contains equal numbers of anions and cations
then the compounds themselves have no overall charge
Relative Atomic Mass
• Atoms are amazingly small. In order
to get a gram of hydrogen you would
need to count out around
602,204,500,000,000,000,000,000
atoms
• So instead, their masses are
compared with the mass of an atom
of the carbon-12 isotope
• On this scale one atom of carbon
weighs exactly 12 units
Why carbon-12?
• Before 1961 we used oxygen as the standard for
atomic masses
• But chemists were using naturally occurring
oxygen, which has a mixture of the isotopes oxygen16, oxygen-17 and oxygen-18.
• Physicists had chosen the single isotope oxygen-16.
• So we had two different sets of atomic masses
• In 1961 chemists and physicists agreed to
compromise on carbon-12
• So if one atom of the carbon-12 isotope weighs
exactly 12 units
• An atom of the commonest isotope of magnesium
weighs twice as much as that and so is said to have
a relative isotopic mass of 24
• An atom of the commonest isotope of hydrogen
weighs only one twelfth that of the carbon-12
isotope, so have a relative isotopic mass of 1
• The basic unit on this scale is therefore 1/12 of the
mass of a 12C atom.
Relative Molecular Mass Mr
• This is the average mass of a single molecule and is
sometimes termed RMM (Mr)
• To calculate Mr add up the relative atomic masses of the
different atoms in the molecule
H2
H2O
HCl
C6H12O6
Mr = 1 x 2 = 2
Mr = (1 x 2) + 16 = 18
Mr = 1 + 35.5 = 36.5
Mr = (12 x 16) + (1 x 12) + (16 x 6) = 180
Key definitions:
• Relative isotopic mass: The mass of an atom
of an isotope compared with one-twelfth of
the mass of an atom of carbon-12
• Relative atomic mass, Ar: The weighted
mean mass of an atom compared with onetwelfth of the mass of an atom of carbon-12
Isotopes and relative atomic mass
JFe worked example
A sample of bromine contains 53% bromine-79 and
47% of bromine-81. Determine the relative atomic
mass of bromine.
Calculate Ar (Br)
Questions!