Laboratory measurements are
made by reading all digits on the
instrument, and estimating one
digit.
• All the digits of a measurement
that you are sure of (markings on
the instrument) plus one estimated
digit.
• Used to express limitations in a
measurement.
Rules for significant digits.
• Nonzero digits are always significant.
– Example: 5.67
• All final zeros used after the decimal point
are significant.
– Example: 5.60
• Zeros between two other significant digits
are always significant.
– Example: 5006, 5.006
• Zeros used solely for spacing the decimal
point are not significant.
– Example: 56,100 , 0.566 , .0561
Significant zeros
If a zero is used only to place the
decimal, it is NOT significant.
Examples:
• Desk measures 32.10 cm
– What is the last marking on the instrument?
– How many significant figures in this number.
• Counting numbers
are exact whole numbers.
– 30 people in the room
– $20.05 dollars in my pocket
Reading a Measurement
A) 720.29g
B) 722.9 g
C.) 722.90 g D.) 723 g
Reading a Measurement
A.) 200 g
B.) 200.5 g
C.) 200.55 g
D.) 20.5g
Reading a Measurement
A.) 76.0 ml
B.) 76 ml
C.) 77 ml
D.) 75.00 ml
Reading a Measurement
A.) 47 ml
B.) 47.00 ml
C.) 47.0 ml
D.) 44 ml
Reading a Measurement
A.) 3.44 ml
B.) 3.4 ml
C.) 3.50 ml
D.) 3.5 ml
E.) 3.48 ml
Lec Sup #1
How many significant figures are in
the measurement 405cm?
A) 0
B) 1
C) 2
D) 3
E) 4
Lec Sup #2
How many significant figures are in
the measurement 3.208 cm?
A) 0
B) 1
C) 2
D) 3
E) 4
Lec Sup #3
How many significant figures are in
the measurement 3000 cm?
A) 0
B) 1
C) 2
D) 3
E) 4
Lec Sup #4
How many significant figures are in
the measurement .0045 cm?
A) 0
B) 1
C) 2
D) 3
E) 4
Lec Sup #5
How many significant figures are in
the measurement .0450 cm?
A) 0
B) 1
C) 2
D) 3
E) 4
Rules for calculating with
significant figures.
Addition or subtraction
Your final answer may contain no
more places after the decimal than
your least known quantity.
(Round the answer so that it has the same
number of decimal places as the measurement
having the fewest decimal places.)
Example:
42.253 mL
125.6 mL
1.75 mL
169.603 mL
*Answer can only have as many
places to the right of the decimal as
that of the known quantity with the
least:
= 169.6 mL
Multiplication and division
Your final answer may have no more
total Significant digits than your
known quantity with the least number
of significant figures.
Example:
62 cm x 33.03 cm = 2047.86 cm2
*only good to 2 figs
2.0 x 10
3 cm2
Rules for rounding off
Look at digit to the right of digit to be
rounded. IF:
• Greater than or equal to 5 round up, less
than 5 leave.
Lec Sup #6
Round to three significant figures.
8.7257cm
= A) 8.7300
B) 8.7200
C) 8.72
D) 8.73
E.) 8.73 X 104
Lec Sup #7
Round to three significant figures.
125.699cm
=
A) 126.00
B) 125
C) 126
D) 125.00
E.) 130.
Lec Sup #8
Round to three significant figures.
124,292
= A) 124
B) 124,000
C) 120,000
D) 1.20 X 104
E.) 125,000
Lec Sup #9 Example:
If 4.383 g of oxygen
combine with 0.0023 g of carbon,
what’s the mass of the resulting
compound?
A)4.4 g
4.383 g
0.0023 g
4.3853 g
B)4.38 g
C)4.385 g
D)4.3853 g
E)4.39 g
Lec Sup #10 Example:
If a line of 1.0 x 108 water
molecules is 1.00 inches long,
what is the average diameter, in
millimeters of a water molecule?
1.00 inch
2.54 cm 1 m
1000 mm
1.0 x 108 molecules 1 inch 100 cm 1 m
= 2.54 x 10-7
= 2.5 x 10-7
Lec Sup #11
A student places 28.70 g of iron,
0.3807 oz of aluminum, and 0.00389 lb of
copper in a beaker that weighs 138 g.
What is the total mass in grams of the
beaker and its contents?
A)179.27 g
Note: conversions will count in sig
B) 179.3 g
fig analysis if the conversion is not
C) 179 g
exact. For instance metric to english
(1lb is approx 454 g). But not in the
D)180 g
case of exact conversions (1 lb is
E) 200 g
exactly 16 oz
#11
A student places 28.70 g of iron,
0.3807 oz of aluminum, and 0.00389 lb of
copper in a beaker that weighs 138 g.
What is the total mass in grams of the
beaker and its contents?
0.3807 oz 1.00 lb 454 g = 10.8 gAl 28.70 g
16 oz 1.00 lb
10.8 g
1.77
g
0.00389 lb 454 g
= 1.77 gCu
138 g
1.00 lb
179.27 = 179 g
Lec Sup #12
A girl needs to reflux a mixture
for 9.85 hours. How long must
the mixture reflux in minutes?
9.85 h 60 min
1h
A)= 591 min
B)= 600 min
C)= 590 min
Lec Sup #12
A girl needs to reflux a mixture
for 9.85 hours. How long must
the mixture reflux in minutes?
9.85 h 60 min
1h
A) Exact conversions do
not dictate the
number of significant
figures in answer
since infinitely
significant.
Lec Sup #13
A group of chemistry students are
instructed to measure a 0.75 m
length of magnesium ribbon and a
.100m length of ribbon. How long
will the total ribbon be in mm?
A).850mm
If multiple operations,
calc sig figs through
B)850mm
order of operations to see
2
C)8.50 X 10 mm
how answer should be
reported.
A group of chemistry students are
instructed to measure a 0.75 m
length of magnesium ribbon and a
.100m length of ribbon. How long
will the total ribbon be in mm?
0.75 m + .100m = .85m
.85m
1000 mm
= 850 mm
1m
Accuracy & Precision
Accuracy:
how close a measurement is to the true
value of the quantity that was measured.
Precision:
(1) how closely two or more measurements
of the same quantity agree with one another
as well as (2) the degree of exactness.
• Chemistry:
–Study of matter and
changes in matter
• Matter
–Anything that occupies
space and has mass
• Pure Substance:
–Form of matter that has a
constant composition and
distinct properties.
• Mixture:
–Combination of 2 or more
substances in which the
substances retain their
distinct identities.
Types of Mixtures
• Homogeneous:
–Composition of mixture is the
same throughout. (solutions)
• Heterogeneous:
–Composition is not uniform
(suspensions)
Types of Pure Substances
• Element:
– Substance which cannot be separated
into simpler substances by chemical
means.
• Compound:
Na, K, Cl
– Composed of atoms of 2 or more
elements chemically fixed in definite
proportions.
C6H12O6
Classification of Matter
All Matter
Mixtures
Solutions
Heterogeneous
Mixtures
Pure Substances
Elements
Compounds
Example 14:
Classify Dry Ice as one of the
following
(a) an element,
(b) a compound,
(c) or a mixture:
Example14:
Classify gasoline as one of the
following
(a) an element,
(b) a compound,
(c) or a mixture:
Example 14:
Classify air as one of the
following
(a) an element,
(b) a compound,
(c) or a mixture:
Example 14:
Classify blood as one of the
following
(a) an element,
(b) a compound,
(c) or a mixture:
Example 14:
Classify methane as one of the
following
(a) an element,
(b) a compound,
(c) or a mixture:
Example 14:
Classify Iodine as one of the
following
(a) an element,
(b) a compound,
(c) or a mixture:
Similar properties on the
periodic table
• Families(groups)/Periods
• Oxidation numbers
• Polyatomic ions
• metals, nonmetals, and
metalloids
• Chemical
• Physical
• Extensive
• Intensive
• Macroscopic
• Microscopic
Lec Supplement 15: The following
are properties of the element
silicon; classify them as
(a) physical intensive,
(b) physical extensive,
(c) or chemical properties:
_______ Melting point, 1410oC
Lec Supplement 15: The following
are properties of the element
silicon; classify them as
(a) physical intensive,
(b) physical extensive,
(c) or chemical properties:
_______ Reacts with fluorine to form
silicon tetrafluoride
Lec Supplement 15: The following
are properties of the element
silicon; classify them as
(a) physical intensive,
(b) physical extensive,
(c) or chemical properties:
_______ Gray
Lec Supplement 15: The following
are properties of the element
silicon; classify them as
(a) physical intensive,
(b) physical extensive,
(c) or chemical properties:
_______ Not affected by most acids
Matter Changes
• Physical change: does not change the
identity (composition) of substance
• Chemical change: changes chemical
composition (identity of substance)
• both involve energy transformation
• endothermic: absorb energy
• exothermic: release energy
Lec Sup 15b.)
Classify each of the following as a
chemical or physical change.
Boiling Water
A) Chemical Change
B) Physical Change
Tearing Clothes
A) Chemical Change
B) Physical Change
Lighting a Match
A) Chemical Change
B) Physical Change
A Rusting Nail
A) Chemical Change
B) Physical Change
Melting Ice Cream
A) Chemical Change
B) Physical Change
Metabolizing Food for Energy
A) Chemical Change
B) Physical Change
Finger Nails Growing
A) Chemical Change
B) Physical Change
Frying an Egg
A) Chemical Change
B) Physical Change
Sweat evaporating from your
forehead
A) Chemical Change
B) Physical Change
A balloon pops due to the
expansion of air
A) Chemical Change
B) Physical Change
Dissolving of Toluene in Carbon
Tetra Chloride
A) Chemical Change
B) Physical Change
Terms to know:
• Group/Period
• Metals/nonmetals/metalloids
• Alkali metals, Alkaline earth
metals, Chalcogens, Halogens,
and Noble gases (rare gases)
• Molecule:
–2 or more atoms tightly bound
together.
• Molecular Compound:
–Compounds made up of
molecules & contain more than 1
type of atom.
Molecule- A neutral group of atoms
that act as a unit.
a) In molecules atoms always
combine in simple whole
number ratios, therefore their
proportions by mass must
always be the same.
Molecular Compounds
compounds composed of molecules
• low melting and boiling points
• many exist as gases or liquids at
room temperature
• Most are composed of 2 or
more nonmetallic elements
Ionic Compounds
composed of cations(+) and anions(-)
• Difference in charges holds them
together.
• ions are arranged in orderly 3dimensional pattern.
• Each cation is between two or
more anions.
• Most are crystalline solids at room
temperature.
• Usually formed from a
metallic and nonmetallic
element.
Ionic Charges of Elements
• You should know the table in your
handout.(grp. 1A,2A,3A,5A,6A,7A)
• Also remember:
1+
2+
Ni Ag Cd 2+ Zn 2+
• Transition metals and group 4A
and 5A metals have multiple ionic
charges (oxidation numbers)
Formula writing and compound
naming summary
1.) Classify the compound as:
ionic or molecular
2.) If the compound is ionic :
CATION
first name
ANION
last name
Naming binary Ionic compounds
• Cation(no prefix) then the anion
root plus ide.
• HCl (contains cation hydrogen
therefore ionic rules)
LS: Name the following
1. CaCl2
2. NaF
3. BaO
A.)Two naming systems for
transition metals
1.) Stock System
Uses roman numerals
i.e. Copper(II) = Cu+2
2.) Classic system
Uses suffixes ous and ic.
ous = lower charge; ic = higher
Name the following using the
classic naming system
4. FeO
5. Fe2O3
6. CuCl2
Name the following using the
stock naming system
7. CuF2
8. FeCl3
Polyatomic IonsTightly bound groups of atoms that
behave as a unit and carry a Charge.
• During reactions, these atoms
usually stay together.
• Notice some have ite / ate pairs
• ite indicates one less oxygen than
the ate.
Naming Ternary compounds
• Contain atoms of 3 different elements-one
is typically a metal
• Usually have 1 or more polyatomic ions.
• Name using Ionic procedures.
However the ending of the name of the
polyatomic ion is not changed.
• Magnesium chloride MgCl2
How do you know the subscript?
+2
Mg
-1
Cl
• Formula Unit- lowest whole
number ratio of ions in an ionic
compound. (also called empirical
formula)
Magnesium Chloride (MgCl2)
1:2 Ratio Mg to Cl
Molecule must be neutral so balance
oxidation states
3.) If the compound is molecular,
classify as an acid or nonacid.
4.) If a nonacid,
Prefix + element Prefix + root name + -ide
first name
last name
5.) If an acid,
classify as nonoxyacid or oxyacid.
Naming binary molecular
- Element on left preceded by a prefix of
how many atoms.
-Element on the right preceded by prefix
of how many atoms and ide at the end.
P2O5
1=mono etc...
Name the following
9. N2O5
10. PCl3
11. N2O4
-When there is 1 atom of the first
element, omit the prefix mono.
-Binary molecular compounds
containing hydrogen as the
first element are named by
the ionic system.
6.) If it is a nonoxoacid,
hydro + root name + ic acid
HCl, HI, HCN, H2S
Hydrochloric acid, Hydroiodic acid,
Hydrocyanic acid, hydrosulfuric acid.
7.) If it is an oxoacid,
polyatomic anion root + -ic or -ous acid
• ate ion = acid with ic
/ ite ion = ous
• H2SO4,, HNO2, H3PO4
Name the following acids
12. H2SO4
13. HNO3
14. HI
#1
• Write the chemical formula for Copper (II)
Phosphate
a.) CuPO4
b.) Cu2PO4
c.) Cu2(PO4)3
d.) Cu3(PO4)2
e.) CuP
#2
• Write the name for Fe(CN)3
a.) Iron Cyanide
b.) Iron (I) cyanide (III)
c.) Iron (III) Cyanide
d.) Iron tricyanide
e.) Iron Carbon Nitride
#3
• The name for Na2S
a.) sodium sulfide
b.) disodium sulfide
c.) disodium monosulfide
d.) sodium sulfate
e.) Sodium sulfite
#4
• The name for FeO
a.) Iron Oxide
b.) diiron trioxide
c.) Iron (II) Oxide
d.) Iron (III) Oxide
#5
• The name for CuCO3
a.) Copper carbide
b.) Copper (I) Carbonate
c.) Copper Carbonate
d.) Copper (II) Carbonate
e.) Copper monocarbonate
#6
• The formula for nitrogen dioxide
a.) N2O
b.) NO2
c.) NO
d.) N(I)O(II)
#7
• Name C3H8
•
•
•
•
•
A.) Carbon Hydride
B.) Carbon (IV) Hydride
C.) tricarbon octahydride
D.) propane
E.) butane
#8
•
•
•
•
•
•
Name HC2H3O2
A.) Hydrogen Acetate
B.) Hydrogen carbon oxide
C.) ????
D.) Acetous acid
E.) acetic acid
#9
•
•
•
•
•
•
Name IF5
A.) Iodine Fluoride
B.) Iodine (I) Fluoride
C.) Iodine pentafluoride
D.) Monoiodine pentafluoride
E.) iodine pentaflourine
#10
•
•
•
•
•
•
Name Na2SO4•10H2O
A.) Sodium (II) sulfate water
B.) Sodium (I) sulfate water
C.) Sodium sulfite
D.) Disodium sulfate decahydrate
E.) Sodium sulfate decahydrate
Terms You Should Know
•atomic number (Z)
•mass number (A)
•isotopes
A
• X
Z
•atomic mass
•average atomic mass
Particles in the Atom
Atoms consist of protons, neutrons and electrons
Particles in the Atom
Electrons
(-) charge
no mass
located outside the nucleus
1 amu
located inside the nucleus
1 amu
located inside the nucleus
Protons
(+) charge
Neutrons
no charge
Structure of the Atom
There are two regions
The nucleus
• With protons and neutrons
– Positive charge
– Almost all the mass
Electron cloud
– Most of the volume of an atom
– The region where the electron can be found
Counting the Pieces
Atomic Number (Z) = number of protons in the
nucleus of an atom, always the same for a given
element
Mass Number (A) = total number of nucleons in an
atom (protons+neutrons)
Mass #
12
Atomic #
6
C
Isotopes
Atoms of the same element with a different mass
number.
Same # protons, different number of neutrons
C-12
C-13
C-14
Isotope of Magnesium
12e-
12e12p+
12n0
Atomic symbol
24
12
Mg
12e-
12p+
13n0
12p+
14n0
25
12
26
12
Mg
Mg
Number of protons
12
12
12
Number of electrons
12
12
12
Mass number
24
25
26
Number of neutrons
12
13
14
Mg-24
Mg-25
Mg-26
Isotope Notation
Isotope of Chlorine
•
•
•
•
•
•
Atomic #
Mass #
# protons
# neutrons
# electrons
Shorthand name
37
17
Cl
Mass # - Atomic # = # neutrons
(protons & neutrons)
(protons)
1. How many electrons, protons, neutrons
are in an atom of phosphorus with mass
number 31?
2. How many electrons, protons, neutrons
are in an atom of thorium with mass
number 232?
Ions
Cations
Na+
Mg2+
Al3+
Anions
ClS2N3-
Ions
An atom that has become charged by either gaining
or losing electrons
Cation- positive ions, loses electrons
Anion - negative ions, gains electrons
125
53
I
1–
Formation of a Cation
sodium atom
Na
sodium ion
Na+
ee-
e-
e-
e-
e-
ee-
e-
11p+
ee-
loss of
one valence
electron
e-
e-
11p+
e-
e-
e-
e-
e-
e-
e-
e-
Formation of an Anion
e-
chlorine atom
Cl
e-
e-
egain of
one valence
electron
e-
e-
chloride ion
Cl1-
e-
e-
e-
eee-
e-
e-
ee-
e-
17p+
17p+
e-
e-
e-
e-
ee-
e-
e-
ee-
e-
e-
e-
e-
e-
ee-
e-
e-
Description
Net
Charge
Atomic
Number
Mass
Number
Ion
Symbol
128
Te2–
15 p+
16 n0
18 e–
38 p+
50 n0
36 e–
1+
18 e–
39
Parts of the atom Summary
• # Protons - determine the identity of the
atom, atomic number (Z)
• # Neutrons - determine the particular isotope
of that element, mass number (A)
• # Electrons - determine the charge on the
atom
Average Atomic Mass
Atomic mass - the average mass of the isotopes in their
relative abundance
Not all isotopes of an element are present in the same
amount, so we use a weighted average
Avg.
(mass)(%) + (mass)(%)
Atomic =
100
Mass
Average Atomic Mass
• EX: Calculate the avg. atomic mass of oxygen if its
abundance in nature is 99.76% O-16, 0.04% O-17, and
0.20% O-18.
Avg.
(16)(99.76) + (17)(0.04) + (18)(0.20)
16.00
=
=
Atomic
amu
100
Mass
Average Atomic Mass
• Ex: Find chlorine’s average atomic mass if
approximately 8 of every 10 atoms are Cl-35 and
2 are Cl-37
Avg.
(35)(8) + (37)(2) 35.40 amu
Atomic =
=
Mass
10
:) Atom Jokes :)
A neutron walks into a restaurant and orders a couple
of drinks. As she is about to leave, she asks the waiter
how much she owes. The waiter replies, “For you,
No Charge!!!”
Two atoms are walking down the street.
One atom says to the other, “Hey! I think I lost an electron!”
The other says, “Are you sure??”
“Yes, I’m positive!”
IV.)Chemical Formulas
• Molecular formulashows the kinds and number of
atoms present in a molecule of a
molecular compound.
Ethanol - (C2H6O)
1 Molecule of ethanol contains
2 carbon, 6 hydrogen, 1 oxygen atom
Ions
• Ions are formed by gaining or
losing electrons.
• In general,
–Metal atoms tend to lose e–Nonmetallic atoms tend to
gain e-
Molecular formula:
Indicates the actual numbers and types of
atoms in a molecule.
Empirical Formula:
elements present in simplest whole
number ratios of their atoms.
H2O2
HO Empirical formula
Molecular formula
empirical formula:
the formula for a compound
containing the smallest ratio of the
atoms.
molecular formula:
the "true" formula of the compound,
contains the actual number of atoms
of each element present in one
molecule of the compound.
EXAMPLES:
Molecular
Formula
Empirical
Formula
P5O10
PO2
N2O4
NO2
H2O2
HO
Structure of Ionic Compounds
• Ions arrange in 3-dimensional
structures. (see page 54)
• We write an empirical formula:
NaCl
• All we can do for ionic compounds
is write empirical formulas
Structural Formula
Shows which elements are
attached to which in a
molecule.
– Types of Models
• Ball and Stick models
• Space-filling models
• molar mass
• chemical formulas
• empirical formula
• molecular formula
• diatomic molecules
• polyatomic molecules
• allotropes