Download Atomic mass

Matter
All things are composed of matter.
 Kinetic Theory of Matter – All matter
is composed of tiny particles that are
in constant motion.

Phases of matter
1. Solid
a. definite shape and volume
b. can be crystallized
c. closely packed slow moving particles
2. Liquid
a. definite volume but takes the shape of the
container
b. closely packed particles that flow past each
other
c. cannot be compressed
3. Gas
a. takes the shape and volume of its container
b. can be compressed
c. fast moving particles separated by space
4. Plasma
a. fastest moving particles
b. high temperature ionized gas
c. composes 90% of the matter in the universe
Changes of phase
1. evaporation – liquid to a gas below
boiling point
2. condensation – gas to a liquid
3. boiling – liquid to a gas at boiling point
4. melting – solid to liquid at melting point
5. freezing – liquid to solid at freezing point
6. sublimation – solid to gas or reverse
7. Each of these changes of phase results in
the gain or loss of energy
Physical properties
– Can be observed without changing the
substance
a. melting/freezing point
b. boiling point
c. color
d. odor
e. ductility
f. malleability
g. solubility
h. conductivity
Physical changes
– Does not change the composition to the
matter
a. change in state
b. change in shape
c. change in size
Chemical properties
– Seen when the substance changes into
another substance
a. reactivity
b. combustibility
Chemical change
– Change that produces a new substance.
Recognized when the following occurs
a. gas production
b. formation of a precipitate
c. burning
d. rusting
Classification of matter
1. substance – same throughout, cannot
be separated by physical means
a. compound – more than one type of
atoms
b. element – only one type of atom
2. mixture – not the same throughout, can be
separated by physical means
a. homogeneous – same throughout
1. solutions – particles evenly distributed
2. Ex: salt water, iced tea
b. heterogeneous – not the same throughout
1. suspension – settles out
Ex: muddy water
2. colloid – never settles out
Ex: milk, smoke
Methods of Separating Mixtures
1. physically – using hand to sort
2. filtration – using a membrane (filter) to catch
large particles
3. magnetism – using a magnet to pick-up
materials
4. evaporation – solid remains after liquid
evaporates
5. fractional distillation – heating a mixture of
liquids with different boiling points
and collecting the vapors released at each
boiling point
Atomic Theory Scientist
A. 400 B.C. – Democritus
B. Aristotle
C. early 1800’s – John Dalton
D. 1897 – J.J. Thompson
E. 1912-1913 - Lord Ernest Rutherford
John Dalton (1807)

Billiard Ball Model

atom is a
uniform,
solid sphere
J. J. Thomson (1903)

Cathode Ray Tube
Experiments


Discovered Electrons


beam of negative particles
negative particles within
the atom
Plum-pudding Model
J. J. Thomson (1903)
Plum-pudding Model

positive sphere
(pudding) with
negative electrons
(plums) dispersed
throughout
Ernest Rutherford (1911)

Gold Foil Experiment

Discovered the nucleus


dense, positive charge in
the center of the atom
Nuclear Model
Ernest Rutherford (1911)

Nuclear Model

dense, positive nucleus surrounded
by negative electrons
Niels Bohr (1913)

Bright-Line Spectrum


Energy Levels


tried to explain presence
of specific colors in
hydrogen’s spectrum
electrons can only exist in
specific energy states
Planetary Model
Niels Bohr (1913)
Bright-line spectrum

Planetary Model

electrons move in circular
orbits within specific
energy levels
1912-1913 - Lord Ernest Rutherford
1. The Gold Experiment – A sheet of gold foil
was bombarded with positively charged
subatomic particles.
a. Most particle passed through.
b. A few particle were deflected at large
angles.
c. About 1 in 8000 bounced back in almost
the opposite direction.
Rutherford’s Experiment

http://glencoe.mcgrawhill.com/sites/0078807220/student_view0/
chapter28/concepts_in_motion.html
Elements
A. Have their own characteristics
1. physical properties
2. chemical properties
B. Unique name and symbol
1. universally accepted
2. symbol is 1-3 letter abbreviation
a. 1st letter is always capitalized
b. 2nd and 3rd letters are always lower case
Make-up of atoms
1. nucleus
a. protons – positive with 1 amu of mass – ID’s
element
b. neutron – no charge with 1 amu of mass – ID’s
isotope
2. electron clouds
a. area where you are most likely to find electrons
b. electrons – negative particles with extremely
small mass – ID’s ion
D. Atomic number (Z) – number of protons
E. Atomic mass (A) – number of protons plus
the number of neutrons
# of neutrons = A – Z
F. Electrically neutral atoms have equal
numbers of protons and electrons
G. Ions
1. positive - # of proton > # of electrons
2. negative - # of electron > # of protons
H. Isotopes – atoms of the same element with
different numbers of neutrons.
History of the periodic table
1869 Dimitri Mendeleev – Russian chemist,
proposed the use of atomic weight to
classify elements. His table was more
widely accepted than Meyer’s table. He
used the properties of the elements to
group them. He had blank spaces, that
he predicted what the element would be
like.
E. 1913 Henry Mosely – English scientist,
proposed the use of atomic numbers to
classify elements as a modification of
Mendeleev’s Table. This cleared up
inconsistencies in Mendeleev’s Table.
Using the Periodic Table
A. groups or families are in columns (18)
B. periods are in rows (7)
C. stair steps separate metals and
nonmetals
Metals
1. to the left and below the stair steps
2. properties
a. shiny surface
b. good conductors
c. give up electrons to form compounds
d. usually have 3 or less electrons in outer shell.
e. can be ductile, malleable or brittle
Metal Families
a. alkali metals – Group 1 except hydrogen
1. most reactive
2. exist only as compounds in nature
3. good conductors, shiny, ductile, and malleable
4. filling s orbital
b. alkaline earth metals – Group 2
1. second most reactive group of metals
2. exist only as compounds in nature
3. filling s orbital
c. transition metals – Groups 3-12
1. most commonly known metals
2. contains the best conductors:
Copper, Silver, and Gold
3. contains the natural magnets:
Iron, Nickel, and Cobalt
4. contains the most unreactive metals:
copper, silver, and gold
5. contains a liquid metal: mercury
6. filling d orbitals
d. rare earth metals
1. 2 series
i. lanthanide
ii. actinide
2. filling f orbitals
4. Reactivity for metals increases form right
to left and top to bottom.
Nonmetals
1. to the right and above the stair steps
2. Properties
a. dull surface
b. poor conductors
c. share or gain electrons to form
compounds
d. usually have 3 or more electrons in
outer shell
3. families
a. hydrogen
i. diatomic molecule in nature (H2)
ii. highly reactive
iii. can gain or lose electrons
b. halogens – Group 17
i. most reactive of nonmetals – needs one
electron to be stable
ii. range from gases to liquids to solids
c. noble gases – Group 18
i. most stable – outer orbital is filled
ii. gases at room temperature
iii. originally called inert gases because
they were thought not to react
Mixed Families
a. Boron family – Group 13
i. one metalloid – boron (B)
ii. four metals
iii. three metals used for semiconductors –
Ga, In, and Tl
iv. contains the most abundant element in
the earth’s crust – Al
b. Carbon family – Group 14
i. one nonmetal – carbon (C)
ii. two metalloids – Si and Ge
iii. two metals – Sn and Pb
c. Nitrogen family – Group 15
i. two nonmetals – N and P
ii. two metalloids – As and Sb
iii. one metal – Bi
d. Oxygen family – Group 16
i. three nonmetals – O, S, Se
ii. two metalloids – Te and Po
5. Reactivity increases from left to right and
bottom to top.
Metalloids
1. along the stair steps
2. properties of both metals and nonmetals
3. There are eight – B, Si, Ge, As, Sb, Po,
Te, and At
4. Boron is considered the least reactive
nonmetal because it has three
electrons in its outer shell and needs
five more to be stable
Chemical Bonding
A. Metallic bonding – metal atoms closely
attracted to each other allow their
electrons to flow from one
atom to another.
B. Ionic bonding
1. Ion – atom or group of atoms with a
charge ( + or - )
2. Attraction between ions of different
charges (+ or - )
3. Weak bond that is easily broken.
4. Can be form by the exchange of electrons
a. one atom gains electrons
b. one atom loses electrons
C. Covalent Bonding
1. Strong bond
2. Sharing of electrons by two atoms.
Writing and naming compounds
A. Naming binary compound
1. Write the name of the first element or
polyatomic ion.
2. Write the name of the second element
changing the ending to –ide, or write the
name of the polyatomic ion.
B. Writing formulas
1. Cancel change method
a. Write the element symbol for the first
element or polyatomic ion
b. Write the element symbol for the second
element or polyatomic ion
c. Use subscripts to indicate how many of
each type of ion is needed to cancel the
charges
d. Use parentheses if more than one of a
polyatomic ion is needed.
EX: sodium chloride
Na
Cl
+
+1 + (-1) = 0
Charges cancelled
Formula is NaCl
EX: aluminum chloride
Al
Cl
+3
+3 + (-1) = 2
Charges are not cancelled.
+3 + (3) (-1) = 0
Charges are cancelled.
Formula is AlCl3.
2. Cross over method
a. Write the formula for each ion
b. If the charges are opposite but
equal in value, drop the charges and
write the formulas together.
EX: lithium chloride
Li+
ClFormula LiCl
c. If the charges are not equal in value, drop
the sign for the charge and use the
number from the charge as the subscript
for the opposite ion.
EX: aluminum chloride
Al+3
Cl↓←←←←↑
Al3
Cl
↓→→→→↑
AlCl3
formula
d. If a metal has a Roman numerals in
parentheses after its name, the metal can form
more than one ion and the Roman numeral
indicates the metals positive charge.
EX: iron (III) sulfide
Fe+3
S-2
↓←←←↑
Fe3
S2
↓→→→↑
Fe2S3
formula
e. When polyatomic ions are part of the formula,
use parenthesis if there is more than one
polyatomic ion.
EX: mercury (II) phosphate
Hg+2
PO4-3
↓←←←←↑
Hg 2
PO4 3
↓→→→→↑
Hg 3
PO4 2
Hg3(PO4)2 formula
EX: hydronium phosphate
H3O+
PO4-3
↓→→→↑
H3O
PO4 3
↓→→→↑
H3O 3
PO4
(H3O)3PO4
←
formula
C. Naming compounds containing metals
with more than one oxidation state.
1. Oxidation state refers to the charge for
the ion (the number of electrons that can
be gained or lost)
2. Determine the charge of the negative
ion.
EX: Fe2O3
Oxygen is always -2
3. Multiply the negative charge by the subscript on
the negative ion.
EX: Fe2O3
Charge is -2. Subscript is 3.
-2 x 3 = -6
4. Drop the negative sign and divide by the
subscript of the positive ion.
EX: Fe2O3
Total negative charge is -6.
Positive
subscript is 2
6 ÷ 2 = 3
5. Write the name of the metal and in
parentheses using a Roman numeral write
the positive charge (oxidation state) just
calculated.
EX: Fe2O3
iron (III)
6. Then write the name of the negative ion
EX: Fe2O3
iron (III) oxide
7. Roman numerals
1 I
6
VI
2 II
7
VII
3 III
8
VIII
4 IV
9
IX
5 V
10
X
8. Naming negative ions
a. Drop the ending on the element name
and add –ide.
b. hydrogen
nitrogen
oxygen
sulfur
carbon
phosphorus
HN-3
O-2
S-2
C-4
P-3
hydride
nitride
oxide
sulfide
carbide
phosphide
chlorine
fluorine
iodine
bromine
arsenic
silicon
selenium
astatine
tellurium
Cl FIBr As -3
Si -4
Se -2
At Te-2
chloride
fluoride
iodide
bromide
arsenide
silicide
selenide
astatide
telluride
VIII. Lewis Structures – Electron Dot
Diagrams
A. Show the placement of valance
electrons
B. Used to explain bonding angles and
crystal structures
C. Writing electron dot diagrams
1. Determine the number of valance electrons the
element has
a. Group 1 has 1
b. Group 2 has 2
c. Group 13 has 3
d. Group 14 has 4
e. Group 15 has 5
f. Group 16 has 6
g. Group 17 has 7
h. Group 18 has 8, except He which has 2
i. For the transition metals and rare earth metals
refer to their electron configuration. The
number of valance electrons is always the last
number.
2. Write the element symbol
3. Start on one side of the element symbol,
place one dot on each side before pairing
to equal the total number of valance
electrons.
EX: hydrogen
Group 1 element → has 1 valance electron
EX: phosphorus
Group 15 element → has 5 valance electrons
IX. Formula mass – The sum of the masses of
all atoms making up a molecule.
A. Calculating formula mass
1. Determine what elements are present
a. determine how many atoms of each element are
present
b. look up the average atomic mass for the element
and round to the nearest .5 amu.
2. Calculate the total mass of each element present by
multiplying the number of atoms of an element by
the element’s average atomic mass.
3. Calculate the total mass of the molecule by adding
the total masses of all the elements in the molecule.
B. Examples
1. FeO
Element
# of atoms
Average atomic
mass
Total mass for
element
Formula Mass
Fe
1
O
1
55.847
15.999
55.847
71.846
15.999
(Fe)(1)
+
(O)(1) = FeO
(55.847) (1) + (15.999)(1) = 71.846
Element # of atoms
Fe
1
x
O
1
x
Formula mass
mass
t. mass
55.847 = 55.847
15.999 = +15.999
= 71.846
2. NaOH
Element
# of atoms
Average atomic
mass
Total mass for
element
Formula mass
Na
1
O
1
H
1
22.990
15.999
1.0079
22.990
39.9969
15.999
1.0079
(Na)(1)
+ (O)(1)
+ (H)(1) = NaOH
(22.990)(1) + (15.999)(1) + (1.0079)(1) = 39.9969
Element # of atoms
Na
1
x
O
1
x
H
1
x
Formula mass
mass
t. mass
22.99 =22.99
15.999 = 15.999
1.0079= +1.0079
= 39.9969
3. AlPO4
Element
Al
Number of atoms
1
Average atomic
mass
26.983
Total mass of
element
26.982
Formula mass
121.952
P
1
O
4
30.974
15.999
30.974
63.996
(Al)(1)
+
(P)(1)
+
(O)(4) = AlPO4
(26.983)(1) + (30.974)(1) + (15.999)(4) = 121.952
Element # of atoms
Al
1
x
P
1
x
O
4
x
Formula mass
mass
t. mass
26.983 = 26.983
30.974 = 30.974
15.999 = +63.996
= 121.952
4. Hg3(PO4)2
Element
Number of atoms
Average atomic
mass
Total mass of
element
Formula mass
Hg
3
P
2
O
8
200.59 30.974 15.999
601.77 61.948 127.992
791.71
(Hg)(3) + (P)(2) + (O)(8)
= Hg3(PO4)2
(200.59)(3) + (30.974)(2) + (15.999)(8) = 791.71
Element # of
Hg
P
O
Formula mass
atoms
3
2
8
mass
t.mass
200.59 = 601.77
30.974 = 61.948
15.999 = 127.992
791.71