Chapter 2
Atoms and Elements
atom- smallest identifiable unit of an element
element- a substance that cannot be broken
down into simpler substances
-there are about 91 naturally occurring elements
-scientists have made over 20 elements
Modern Atomic Theory and Laws That Led to It
1) Law of conservation of mass- matter is
neither created nor destroyed
-when a reaction is carried out, the mass of the
reactants must equal the mass of the products
ex- 2Na + Cℓ2  2NaCℓ
7.7g 11.9g
19.6g
19.6g
2) Law of Definite Proportions- all samples of a
given compound have the same proportions
of their constituent elements
ex- 18.0g H20 results in 16.0g O2 and 2.0g H2
mass ratio = 16.0g O2/2.0g H2 = 8 O2:1 H2
-this is true for all samples of water
Page 49 Ex 2.1 and For practice 2.1
3) Law of Multiple Proportions- when two
elements (A and B) form two different
compounds, the mass of element B that
combines with 1g of element A can be
expressed as a ratio of small whole numbers
Page 50
Ex- 2.2
For practice 2.2
John Dalton’s Atomic Theory
1. All elements are composed of atoms
2. All atoms of the same element have the same
properties and are identical
3. Atoms combine in simple whole number
ratios with other atoms to form compounds
4. Atoms of one element cannot be changed into
atoms of another element
-Scientists began to think that atoms were
composed of smaller particles
Discovery of Electron
-J.J. Thomson (1856-1940)
-used a cathode ray tube
-believed that the cathode ray was composed of
tiny particles with an electrical charge
-these are electrons
-rays travel from – charged electrode (cathode)
to + electrode (anode)
Discovery of Charge of Electron
-Robert Millikan (1868-1953)
-said electrons have a negative charge
Plum Pudding Model
-it was known that – charged particles attract +
charged particles
-it was also known that atoms were neutral
-so there must be a + charged particle in the atom
-Thomson proposed that – charged electrons
were small particles held within a + charged
sphere
-protons discovered by Eugen Goldstein (1850 –
1930)
Ernest Rutherford (1871-1937)
-performed experiments that allowed him to
conclude that the atom must have a positive mass
in a much smaller space than proposed
Nuclear Theory
1) most of atom’s mass and all of its + charge are in
a small core called the nucleus
2) most of volume is empty space through which
negative particles move
3) there are as many negatively charged particles as
there are positive (protons) so the atom is neutral
-There were still some parts of atom missing
James Chadwick (1891-1974)
-discovered that missing mass was neutrons
Subatomic Particles and the Atom
protons (p+) – positively charged, found in
nucleus, mass = 1.67 x 10-24g
neutrons (n0) – neutrally charged, no charge,
found in the nucleus, mass = 1.67 x 10-24g
electrons (e-) – negatively charged, found
surrounding nucleus in clouds or energy levels,
mass = 9.11 x 10-28g
**most of mass in the nucleus of an atom
atomic number- # of protons in an element,
defines the element, the smaller # on the
periodic table
-since atoms are neutral:
# of p+ = # of e-so the atomic # also tells you the # of eHow many protons in the following?
argon
uranium
iron
lithium
18
92
26
3
*each element also has that # of e-
mass number- sum of the # of p+ and n0, larger #
on the periodic table rounded to a whole #
How do you find the # of n0?
# n0 = mass # - atomic #
How many neutrons in the following?
argon
uranium
iron
lithium
22
146
30
4
Can be written this way:
Ar-40
U-238
Fe-56
Li-7
atomic mass- larger # on the table not rounded,
average mass of all the isotopes of that element
isotopes- atoms with the same # of p+ and e- but
different # of n0
exNe-20
Ne-21
Ne-22
p+ = 10
p+ = 10
p+ = 10
e- = 10
e- = 10
e- = 10
n0 = 10
n0 = 11
n0 = 12
Page 59 Ex 2.3 and For Practice 2.3
natural abundance- the relative % of an isotope
with respect to other isotopes of the same
element
-back to Ne isotope
Which isotope of Ne would have the highest
relative abundance?
Ne-20, because it is closest to the atomic mass
-Relative Abundance problems
Modern Periodic Table
-grew from the work of Dmitri Mendeleev
-arranged elements according to increasing mass
-called periodic law- when elements are
arranged in order of increasing mass/atomic #,
similar properties will recur
periods- horizontal rows
groups/families- vertical columns, indentified by
a # and a letter
Groups of the Periodic Table
Alkali Metals- Group 1A elements
Alkaline Earth Metals- Group 2A elements
Transition Metals- Group B elements, bridge
Inner Transition Metals- two rows at the bottom
non-metals- upper and lower right side
Metalloids- staircase- starting at boron
Halogens- Group 7A
Noble Gases- Group 8A or Group 0
-also called inert gases- they do not react to
form compounds
Main group or Representative elements
-Groups 1A to 8A minus the Transition metals
Properties of Metals
1) conduct electricity
2) ductile (can be drawn into wires)
3) shiny/lustrous
4) malleable (can be pounded into sheets)
-Most elements are solids
Gases = H, O, N, F, Cℓ, He Ne, Ar, Kr, Xe, Rn
Liquids = Hg, Br (Ga, Fr, Cs)