Inorganic
Physical
Organic
Analytical
Biochemistry
Mass :
quantity of matter
Matter
Solid
Liquid
Gas
Melting
Heat
Solid
Liquid
Cool
Solidification
Evaporation
Heat
Liquid
Vapor
Cool
Condensation
Physical state and Changes in
Matter
 Heat



Solid
Vapor
Cooling
Sublimation
Physical state and Changes in
Matter
 Heat



Ice
Water
Cool
HETEROGENEOUS
MIXTURE
HOMOGENEOUS
SUBSTANCES
SOLUTIONS
PURE
SUBSTANCES
Homogeneous mixture
of variable composition.
Can be separated into
Homogeneous matter of
fixed composition
COMPOUNDS
Composed of 2 or more
elements.
Can be separated into
ELEMENTS
Heterogeneous and
Homogeneous
Solutions, Pure Substance
and Compounds
Mass
 A mass of an object pertains to the
quantity of the matter that object
contains.
A physical property that every
Manager possesses is a mass.
The amount of mass in a pizza will
never change when the object is
moved from place to place.
A physical property
that is related to mass is weight
The weight of a chef may change
if it is moved to Uranus because
weight is determined by gravity.
Atoms are the basic building
blocks of all the chalk around you.
It is the smallest particle of
matter that can enter into
chemical combinations with other
particles.
A smallest particle of an element or
compound that can have a stable
independent existence.
Atoms make up molecules. Molecules
make up a hairy eagle.
Elements are pure
substances, made from one
type of atom. Soda can be
broken down into many
elements but nitrogen can
not be broken down.
Name
Sodium
Potassium
Gold
Silver
Iron
Symbol
Na
K
Au
Ag
Fe
Latin
name
natrium
kalium
aurum
argentum
ferrum
Gold, silver, copper,
and iron are
examples of metals.
A gold diamond is
shiny because of its
metal properties.
Gold conducts heat and
electricity. Nickel can be
hammered into thin sheets
without breaking. Platinum
can be pulled into wire.
The helium in my Christmas
balloon is a nonmetal. The
Oxygen in the air is not
shiny because of its
nonmetal properties.
A dog cannot conduct
electricity. A snap dragon
cannot be hammered into
thin sheets. A snicker
cannot be pulled into wire
because they are not
metals.
Metalloids have properties of both
metals and nonmetals.
Silicon is a metalloid that can be
found in many materials such as the
sand on Lake Tahoe the glass in a
vase and certain plastics that make
up a favorite toy, car.
Iron is abundant easy to shape when
heated and relatively strong.
Chemical Property ability of a substance to
undergo chemical change
• Composition of matter always changes
Another term for Chemical change
• One or more substance change into
one or more new substance during
chemical reaction
Reactant a substance present at the
start of the reaction
Product substance produced in the
reaction
•
How can you tell whether a
chemical change has taken
place?
 transfer in energy
 change in color
 production of gas
 formation of a precipitate
An atom or a group of
atoms that has acquired
electric charge by gaining
or losing one more electron
• Cathode
• Anode
• Anion
• Cation
•
• Any
physical change
or chemical reaction,
mass is conserved.
• Mass is neither
created nor destroyed.
•A
given compound always shows a
fixed proportion.
• A chemical compound always
contains the same elements in the
same percent by mass.
• When two elements combine to form
a given compound, they always do so
in a fixed proportion.
Trial Mass of C (g)
Mass of O2 (g)
Mass of CO2
(g)
1
2.00
5.34
7.34
2
15.00
40.05
55.05
3
5.00
13.36
18.36
Finding the % of Carbon and Oxygen
% C = mass C x 100
% O = mass of O x 100
mass of CO2 27.2%
mass of CO2
72.8%
• When
two elements combine to form more
than one compound, the masses of one
element which combine with a fixed mass of
the other element are in a ratio of small
whole numbers such as 2:1, 1:1, 2:3, etc.
Example
C
D
1st Compound
2.276
0.792 0.348
2nd
1.422
0.948 0.667
A. Mass fixed at C
therefore the formulas of the two compounds are
C
CD
1
= 1
D
0.348
0.348
CD2
1
0.667 = 2
0.348
Folder at the desktop : New Bio lectures
Find the File name: introduction to Biology
page 61 (Scientific Measurements)
• Encounter
very large or very small
numbers.
Examples:
A single gram of hydrogen, contains
approximately 602 000 000 000 hydrogen
atoms 6.02 x 10 ?
The mass of an atom gold is 0.000 000
000 000 327 gram. 3.27 x 10 ?
A given number is written as the
product of two numbers:
 a coefficient
 a 10 raised to a power
Accuracy how close a measurement
to the True value
Precision series of measurement
Accuracy
Precision
Correct value
repeated
measurements
Accepted value: true value
Experimental value: measured in lab
Formula
Error: experimental value – accepted value
Percent error:
_____error_______
accepted value
x 100
Include all the digits that are known,
plus a last digit that is estimated.
Measurements
must
always
be
reported to the correct number of
significant figures because calculated
answers often depend on the number
of significant figures in the values used
in the calculation.
1. Every nonzero digit in a reported measurement is
assumed to be significant. Ex. 24.7 meters, 0.743
meters and 714 meters each has 3 significant
measurement.
2. Zeros appearing between nonzero digits are
significant. Examples 7003 meters and 40.79 metes
have 4 s.f.
3. Left zeros appearing in front of nonzero digits are
not significant. They are just a placeholder. Ex.
0.000 099 meters has 2 s.f. you will write them as
7.1 x 10 -³
4. Zeros at the end of a number and to the right
of a decimal point are always significant. Ex.
43.00 meters, 1.010 meters have 4 s.f.
5. Zeros at the right most end of a measurement
that lie to the left of an understood decimal
point are not significant if they serve as
placeholders to show the magnitude of the
number. Example 7000 meters and 27210
meters have 1 and 4 s.f respectively.
6. The numbers are all in s.f. if it is exact
amount/count for ex. 23 students or 60 mins= 1
hour.
 24.7
74.3
512 meters
 7.003
1.505
87.29
 0.0071
0.043
0.000 0044
 9.000
43.00
1.010
 300
7000
27210
Calculate the sum of the three
measurements. Give the answer to the
correct number of significant figures.
12.52 meters + 349.0m + 8.24m
Answer: 369.8 or 3.69 x 102 meters
2.10 meters x 0.70 meter = 1.47
(meter)2
Answer: 1.47 (meter)2 = 1.5 meters 2
• Basic unit of length or linear measure is meter
METRIC UNITS OF LENGTH
Kilometer (km)
1 km = 103 m
Length of 5 city
blocks
Meter (m)
Base unit
Height of
doorknob from the
floor
Decimeter (dm)
101 dm
Diameter of large
orange
Centimeter (cm)
102 cm
Width of shirt
button
Millimeter (mm)
103 mm
Thickness of dime
Micrometer (um)
106 um
Diameter of
bacterial cell
Nanometer (nm)
109 nm
Thickness of RNA
Volume is the space occupied by any sample of
matter.
• Unit being use cubic meter (m3)
Unit
Relationship
Example
Liter (L)
Base unit
Quart of milk = L
Milliliter (mL)
103 mL + 1 L
20 drops of water =
1 mL
Cubic centimeter
(cm3)
1 cm3 =1 mL
Cube of sugar = 1
cm3
Microliter (uL)
106 uL = 1 L
Crystal of table salt
= 1uL
Kilogram (kg) is the basic unit of mass
Platform balance to measure mass of an object
Metric Units of Mass
Kilogram
(kg)
103 g
Small textbook
Gram (g)
10-3 kg
Dollar bill
Milligram
(mg)
103mg = 1
g
Ten grains of salt
Microgram 106 ug = 1g Particle of baking
(ug)
powder
• When
you hold a glass of hot water the transfer of
heat.
• Almost all substances expand with an increase in
temperature and contract as the temperature
decreases. (very important exception is water)
•Celsius was named after to Anders Celsius a
Swedish astronomer.
• Celsius scale sets freezing point of water at 0
degree and the boiling temperature is 100 degree C.
• Kelvin, named after to Lord Kelvin a Scottish
physicist and mathematician
• freezing point 273.15 and the boiling point 373.15
degree C
°F = 9 °C + 32
5
°C = 5 (°F – 32)
9
K = °C + 273
° C= K - 273
Normal human body temperature is 37 °C.
What is the temperature in Kelvin?
Given:
37 °C
Unknown: Kelvin
Formula : K = °C + 273
Solution: K = 37 °C + 273
Answer: K= 310
Correct! It lies between
273K up to 373K
Convert 14 °F to °C and Kelvin
Given: 14 °F
Unknown: °C and Kelvin
Formula: °C = 5 (°F – 32)
9
K = °C + 273
Solution:
Anwers: -10 °C
and 263 K
• Energy
is the capacity to do work or to
produce heat.
• Joule (J), named after the English physicist
James Prescott Joule and the Calorie (cal) are
common units of energy.
• One calorie is the quantity of heat that raises
the temperature of 1 g of pure water by 1 °C
Formula
1J = 0.2390
1 cal = 4.184 J
Calculate the quantity of heat in joules required to
raise the temperature of 135 g of water from 11 °C
heat to 41 °C.
Given : 135 g of water
11 to 41 °C
Formula:
Heat required = mass x specific heat x temperature
change
1 cal = 4.184 J/ g °C
Solution:
135g x 4.184 J x (41-11 °C)
g °C
= 1.7 x 104
• Are
ratio of equivalent measurements.
• Useful in solving problems in which a
given measurement is multiplied by a
conversion factor, the numerical value is
generally changed, but the actual size of
the quantity measured remains the same.
Example:
I meter = 10 decimeters = 100 centimeters
= 1000 millimeters
Express 750 dg to g
Given:
mass : 750 dg
1g = 10 dg
or
1g
10 dg
Solution:
750 dg x 1g
10 dg
Answer: 75 g
What is 0.073 cm in micrometers?
Given:
0.073 cm = 7.3 x 10 -2 cm
10 2 = 1 m
1m = 10 6 um
Unknown: um
Formula:
cm
meters
micrometers
Solution:
7.3 x 10 -2 cm x 1 m x 10 6 um
10 2
1m
Answer: 7.3 x 10 2 um
• Mass per unit volume of a substance
• Ratio of the mass of an object to its
volume.
• Is an intensive property that depends
only on the composition of a substance,
not on the size of a sample.
• Formula:
Density =
mass
volume
• Corn oil and corn syrup
Material
Density at
20°C (g/cm3)
Material
Density at
20°C
Corn oil
0.9222
Helium
0.166
Corn syrup
1.35 – 1.38
Oxygen
1.33
Table sugar
1.59
Carbon
Dioxide
1.83
Gold
19.3
Ammonia
0.718
Example :
A copper penny has a mass of 3.1 g and a volume of
0.35 cm 3. What is the density of copper?
Given:
Mass: 3.1 g
volume= 0.35 cm3
Unknown: density= ?g/cm3
Formula:
Density = mass = 3.1 g
volume 0.35 cm3
= 8.8571 g/cm3
= 8.9 g/cm3 (rounded off to two
significant figures)
Density of a substance generally
decreases as its temperature
increase
•
Atom is the smallest
particle of an element that
retains its identity in a
chemical reaction.
Democritus (460 B.C.-370
B.C.) is a Greek philosopher
was among the first to
suggest the existence of
atom.
• He believed that atoms
were indivisible and
An English chemist and school teacher
responsible for the modern process of
discovery regarding atoms.
• By
using
experimental
methods,
he
transformed Democraticus’s ideas on atoms
into a scientific theory.
 All
elements are composed of tiny
indivisible particles called atoms.
 Atoms of the same element are identical.
 Atoms of different elements can physically
mix together or can chemically combime in
simple
whole-number
ratios
to
form
compounds.
 Chemical reactions occur when atoms are
separated, joined, or rearranged.
One important change in Dalton’s atomic
theory is that atoms are now known to be
divisible. They can be broken down into
even smaller, more fundamental particles
called subatomic.
Three kinds of Subatomic Particles:
• Electrons
• Protons
• Neutrons
•
•


•
•
ELECTRONS
Negatively charged subatomic
particles.
Thomson performed
experiments that involved
passing
electric
current
through gases at low
pressure.
Travels from cathode (-) to
anode (+)
Thomson
examine two ways
that a cathode ray can be
deflected by using magnet and by
using electrically charged plates.
• A positively charged plate attracts the
cathode ray, while negatively charged
plate repels it.
•Thomson knew that opposite charges
attract and like charges repel, so he
hypothesized that a cathode ray is a
stream of negatively charged particles
moving at high speed.
• He called these particles corpuscles,
later named electrons. He concluded
that electrons must be parts of the
atoms of the elements.
• US physicist Robert Millikan carried
out experiments to find the quantity of
charged carried by an electron.
• He is the one responsible for charge
and mass.
Positively
charged
subatomic
particles.
• Example is a hydrogen atom (lightest
kind of atom) loses an electron, what
is left?
• Eugen
Goldstein
(1850-1930)
a
German Physicist observed a cathoderay-tube and found rays travelling in
the direction opposite of that cathode
rays.
• He called that canal rays and
concluded that they were composed of
positive particles
•
.
• No charge but with a mass nearly
equal to that of a proton
• James Chadwick (1891-1974)
English Physicist
confirmed
existence
an
its
Particle
Symbo Relative Relative
l
Charge mass
(mass of
proton= 1)
Actual
mass
(g)
Electron
e -
1-
1/1840
9.1 x 10
28
Proton
p+
1+
1
1.67 x 10
-24
Neutron
no
0
1
1.67 x 10
-24
-
• He concluded that all the positive charge
and almost all the mass are concentrated
in a small region that has enough positive
charge to account.
• He called this region as Nucleus.
• He said that a nucleus is a tiny central
core of an tom and is composed of proton
and neutrons.
• Rutherford atomic model is known as the
nuclear atom.
• In nuclear atom, the protons and
electrons are located in the nucleus.
• While the Electrons are distributed
around the nucleus and occupy almost all
of the volume of atom.
• of an element is the number of protons in the nucleus of an atom
of that element.
• Elements are different because they contain different number of
protons.
Name
Symbol
Atomic #
Protons
Neutron
Mass #
# of
Electrons
Hydrogen
H
1
1
0
1
1
Helium
He
2
2
2
4
2
Lithium
Li
3
3
4
7
3
Beryllium
Be
4
4
5
9
4
Boron
B
5
5
6
11
5
Carbon
C
6
6
6
12
6
Nitrogen
N
7
7
7
14
7
Oxygen
O
8
8
8
16
8
Fluorine
F
9
9
10
19
9
Neon
Ne
10
10
10
20
10
• Total number of protons and neutrons in an atom
• Example a helium atom has 2 protons and 2
neutrons so its mass is 4.
• The number of neutrons in an atom is the
difference between the mass number and atomic
number.
• Number of neutron = mass number – atomic
number
How many protons, electrons and neutrons are in
each atom?
Atomic number
Mass Number
Beryllium (Be)
4
9
Neon (Ne)
10
20
Sodium
11
23
 are atoms that have the same
number of protons but different
neutrons.
 Because isotopes of an element
have different numbers of neutrons,
they also have different mass
numbers.
 Have an identical numbers of
protons and electrons
• Hydrogen
has a mass number of 1
and is called hydrogen -1
• second isotope has one neutron and
a mass number of 2 or a hydrogen -2
or deuterium.
• third isotope has 2 neutrons and a
mass number of 3, or hydrogen -3 or
tritium.
• Remember mass number superscript;
atomic number subscript
Example is Carbon -12, This isotope of a carbon
was assigned a mass exactly of 12 atomic mass
units.
• AMU is defined as one-twelfth of the mass of a
carbon -12 atom. Using these units, a helium -4
atom with a mass of 4.0026 amu, has about
one-third the mass of a carbon -12.
• While a nickel -60 atom has about 5 times the
mass of a carbon -12 atom.
• Atomic Mass of an element is a weighted
average mass of the atoms in a naturally
occurring sample of the element.
Name
Hydrogen
Helium
Symbol
Natural
Percent
Abundance
Mass (amu)
₁¹H
99.985
1.0078
₁²H
0.015
2.0141
³He
2
4He
2
0.0001
3.0160
Average
atomic mass
1.0079
4.0026
99.9999
4.0026
Calculate the atomic mass of Hydrogen
(To calculate: multiply the mass of each isotope by
its natural abundance, express as a decimal, and
then add the products.)
Given:
1H
2H
Mass (amu) = 1.0078
Mass (amu) = 2.0141
Nat’l % A = 99.985
Nat’l % A = 0.015
Formula: Average atomic mass =
(amu) (Nat’l %) + (amu) (Nat’l %)
Isotope = 10 X
Mass (amu) = 10.012
Natural percent abundance = 19.91% = 0.1991
AMU = ?
Isotope = 11 X
Mass (amu) = 11.009
Natural percent abundance = 80.09% = 0.8009
AMU = ?
10.012 amu x 0.1991 =
11.009 amu x 0.8009 =
Answer
=
1.993 amu
8.817 amu
10.810 amu
Given:
Isotope 1
1 Helium
Mass (amu) = 3.0160 amu
Nat’l % A = 0.0001
Isotope 2
2 Helium
Mass (amu) = 4.0026 amu
Nat’l % A = 99.9999
Isotope 1
63 Copper
Mass (amu) = 62.93 amu
Nat’l % A = 69.2%
Isotope 2
65 Copper
Mass (amu) = 64.93 amu
Nat’l % A = 30.8%
Name
#
Carbon
1
2
1
2
1
2
1
2
1
2
Nitrogen
Oxygen
Sulfur
Chlorine
Natural Percent
Abundance
98.89
1.11
99.63
0.37
99.759
0.037
95.002
0.76
75.77
24.23
Mass
(amu)
12.000
13.003
14.003
15.000
15.995
16.995
31.972
32.971
34.969
36.966
Name
#
Natural Percent
Abundance
Mass
(amu)
Bromine
1
2
50.69%
49.31%
78.92
80.92
Boron
1
2
20.0 %
80.0%
10.01
11.01
Lithium
6
7
7.5%
92.5%
6.015122
7.016003
Iron
54
56
5.845%
91.754%
53.9396
55.9349
Copper
63
65
69.17%
30.83%
62.9296
64.92779
• An arrangement of elements in which the elements are
separated into groups based on a set of repeating
properties.
• Allows you to easily compare the rpoperties of one
element (or group of elements) to another element.
•Notice that the elements are listed in order of increasing
atomic number, from left to right and top to bottom.
•Each horizontal row of the periodic table is called a
PERIOD.
•Each vertical row of the periodic table is called a GROUP.
Electrons in Atom
Niels Bohr a young Danish Physicist, said that Rutherford’s model
need to be improved.
 He proposed that an electron is found only in specific circular
paths,
or orbits, around the nucleus..
Energy levels these are the fixed energies within the electron
Quantum of energy is the amount of energy required to move an
electron from one energy level to another energy level. The energy
of one electron is said to be quantitized.
 It
determines the allowed energies
an
electron can have how likely it is to find the
electron in various locations around the
nucleus.
 Erwin Schrodinger an Australian physicist
used new results to devise and solve a
mathematical equation describing the behavior
of the electron in a hydrogen atom.
 It is the probability of finding an electron at
various location around nucleus.
 is often thought of as a region of space in
which there is a high probability of finding an
electron.
 The energy levels of electrons in the quantum
numbers (n).
 Each energy sublevel corresponds to an
orbital of a different shape, which describes
where the electron is likely to be found.
Principal Energy
Level
Number of
Sublevels
Type of Sublevel
1 S (1 orbital)
n=1
n=2
1
2
2 s (1 orbital), 2p (3
orbitals)
3 s (1 orbital), 3p (3
n=3
n=4
3
4
orbitals),
3d (5 orbitals)
4 s (1 orbital), 4 p (3
orbitals),
4 d (5 orbitals), 4 f (7
Electron Configuration electrons are arranged
in various orbitals around the nuclei atoms
There are 3 rules to tell you how to find the electron
configurations of atoms:
Aufbaf Principle
Pauli Exclusion Principle
Hund’s Rule
 Electrons occupy the orbitals of lowest energy first.
Table 3-6a - Orbital and Electron
Capacity for the Four Named
Sublevels
Maximum
Sublevel # of orbitals number of
electrons
s
1
2
p
3
6
d
5
10
f
7
14
According to
Pauli, an atomic
orbital may
describe at most
two electrons.

Ex. Either 1 or
2 electrons can
occupy s or p
orbitals
 states that electron occupy orbitals of the same energy in a
way that makes the number of electrons with the same spin of
direction as alarge as possible.
Classical mechanics adequately describes the
motions of bodies much larger than atoms, while
quantum mechanics describes the motions of
subatomic particles and atoms as waves.
Periodic comes from the Greek roots peri
meaning “around” and hodos, meaning “path”.
In a periodic table, properties repeat from left to
right across each period.
The Greek word metron means “measure”
What does perimeter mean?
A Russian chemist and a
teacher published a table of
the elements.
Mendeleev arranged the
elements in the periodic
table in order of increasing
atomic mass.
Henry Moseley 1913, British
determined atomic number
for each known element.
Elements are arranged in order of increasing
atomic number.
Periodic Law When elements are arranged in
order of increasing atomic number, there is a
periodic repetition of their physical and chemical
properties.
Sodium printed in black because it is solid in
room temperature.
Symbol for gases are in red.
Symbol for two elements that are liquids at room
temperature, mercury and bromine are color blue.
Are used to distinguish groups of elements.
Two shades of gold are used for metals in Group
1A and 2A.
Group 1A alkali metals
Group 2A alkaline earth metals
“alkali” Arabic word al aqali means the ‘ashes’
Wood ashes are rich in the alkali metals sodium
and potassium.
Nonmetal of Group 7A (nonmetals) are called
halogens, comes from the Greek word hals,
meaning “salt”, and a Latin word genesis, means
“to be born”
There is a general class of compounds called salts,
which include the compound called table salt.
Chlorine, Bromine and Iodine the most common
halogens, can be prepared from the salts.
Electrons play a key role in determining the properties of
elements. So there should be a connection between an
element’s electron configuration and its location in the
periodic table.
Elements can be sorted into noble gases, representative
elements, transition metals
These are the elements in Group 8A.
These nonmetals are sometimes called the inert gases
because they rarely take part in a reaction.
Helium (He)
1s2
Neon (Ne)
1s 2
Argon (Ar)
1s 2 2s 2 2p6 3s2
Krypton (Kr)
1s 2 2s 2 2p6 3s2 3p6
3d10 4s2 4p6
2s 2
2p6
3p6
Because they display a wide range of physical and
chemical properties. Some are metals… Most of tem are
solids, but a few are gases at room temperature, and one,
bromine, is a liquid.
Lithium (Li)
1s 2
2s 1
Sodium (Na)
1s 2 2s 2 2p6 3s1
Potassium (K)
1s 2 2s 2
4s1
2p6 3s2 3p6
In atoms of Carbon, Silicon, Germanium in Group 4A,
there are four electrons in the highest occupied energy
level.
Carbon (C)
1s 2 2s 2 2p2
Silicon (Si)
1s 2 2s 2
Germanium (Ge)
1s 2 2s 2 2p6 3s2 3p6
3d10
4s2 4p2
2p6 3s2 3p2
It means the Elements in the B groups, which
provide a connection between the two sets of
representative elements.
Two types of transition elements
1. Transitional metal
2. Inner Transitional metal
These are the Group B elements that are
usually displayed in the main body of a
periodic table. Ex. Copper, silver, gold, and
iron.
In atoms here, the highest occupied s
sublevel and a nearby d sublevel contain
electrons. These elements are characterized
by the presence of electrons in d orbitals.
It appear below the main body of the periodic
table. In atoms of an inner transition metal, the
highest occupied s sublevel and a nearby f
sublevel generally contain electrons.
The inner transition metals are characterized by
f orbitals that contain electrons.
It is the energy required to remove an electron
from an atom.
First ionization energy tends to decrease from
top to bottom within a group and increase from
left to right across a period.
During reactions between metal and nonmetals, metal
atoms tend to lose electrons and nonmetal atom tend to
gain electrons. The transfer has a predictable affect on
the size of the ions that form.
Cations are always smaller than the atoms from which
they form. Anions are always larger than atoms from
which they form.
It is the ability of an atom of an element to attract
electrons when the atom is in a compound. Scientists use
such factors such as ionization energy to calculate values
for electronegativity.
Linus Pauling won a Nobel Prize in Chemistry for his
work on chemical bonds. He was the first to define
electronegativity.
Electronegativity values decrease from top to bottom
within a group.
For representative elements, the values tend to increase
from left to right across a period.
Metals at the far left of the far left of the periodic table
have low values. By contrast, nonmetals at the far right
(excluding noble gases) have high values.
Example:
Least is Cesium, has least tendency to attract electrons.
Most electronegative is Flourine, has strong tendency to
attract electrons.
Atomic size decreases
Ionization energy increases
Electronegativity increases
Nuclear charge increases
Shielding is constant
Atomic size increases
Ionic size increases
Ionization energy decreases
Electronegativity decreases
Nuclear charge increases
Shielding increases
Are the electrons in the highest occupied energy level of
an element’s atoms. The number of valence electrons
largely determines the chemical properties of an
element.
The number of valence electron is related to the group
numbers in the periodic table.
To find the number of the valence electrons in an atom
of a representative element, simply look at its group
number.
Except for the noble gases (Group 8A); wherein Helium
has 2 valence electron.
Are usually only the electrons used in chemical bonds.
Therefore as a general rule, only the valence electrons
are shown in electron dot structures.
Electron Dot Structures are diagrams that show
valence electrons as dots.
Octet Greek word okto meaning eight. Like in the
electron.
Gilbert Lewis used this fact to explain why atoms form
certain kinds of ions and molecules. He called this octet
rule.
In forming compounds, atoms tend to achieve the
electron configuration of a noble gas. Just like the
electrons in highest energy level ns2 np6.
Atoms of the metallic elements tend to lose
their valence electrons, leaving a complete
octet in the next-lowest energy level.
Atoms of some nonmetallic elements tend to
gain electrons or to share electrons with
another nonmetallic element to achieve a
complete octet.
Atom is electrically neutral = number of protons and
electrons.
Therefore, an ion forms when an atom or group of
atoms loses or gains electrons.
An atom’s loss of valence electrons produces a cation, or
a positively charged ion.
Example: sodium atom forms a sodium cation.
Similar in in their names but different chemically.
Na
-e –
1s2 2s2 2p6 3s1
Sodium Atom
Na + 1s2 2s2 2p6
Sodium ion
Na .
Na +
Ionization Na.
Na +
Neon
atom
..
: Ne :
..
+
e-
Magnesium Group 2A
Mg
Mg 2+
+
2e
Atom of Iron / Fe may lose 2 or 3 electrons = Fe 2+ or
Fe 3+
There is an exception like silver atom would have to lose
11 electrons. They don’t have noble-gas electron
configuration.
Anion is an atom or a group of atoms with a negative
charge.
The gain of negatively charged electrons by a neutral
atom produces an anion.
The name of anion typically ends in –ide.
Chlorine atom (Cl) forms a chloride ion (Cl -)
Oxygen atom (O) forms an oxide ion (O2-)
Because they have relatively full valence shells, atoms of
nonmetallic elements attain noble-gas electron
configuration more easily by gaining electrons than by
losing them.
Ex. Chlorine belongs to Group 7A (the halogen family)
+e
Cl 1s2 2s2 2p6 3s2 3p5
Cl - 1s2 2s2 2p6 3s2 3p6
Notice that it has the same electron configuration as the
noble gas argon.
Chlorine atom
Chloride ion
Argon atom
7dots
8 dots with negative
8 dots
The ions that are produced when atoms of chlorine and
other halogens gain electrons are called halide ions.
All halogen atoms have seven valence electron to achieve
the electron configuration of noble gases. (F, Cl, Br, & I)
Oxygen atoms attain the electron configuration of neon
by gaining two electrons.
Oxygen atom
Oxide ion
Neon atom
..
..
..
: O.
:O: 2:Ne:
.
..
..
O
+
2e
O2-
Compounds composed of cations and anions. Ionic
compounds are usually composed of metal cations and
nonmetal anions.
Anions and cations have opposite charges and attract
one another by means of electrostatic forces. The
electrostatic forces that hol ions together in ionic
compounds are called Ionic bonds.
Ex. Sodium cations and chloride anions
It shows the kinds and numbers of atoms in the smallest
representative unit of a substance.
Chemical formula is the lowest whole-number ratio of
ions in an ionic compound. Ex. Is NaCl one Na+ to each
Cl- )
Another is Magnesium chloride contains Magnesium
cations (Mg2+) and chloride anions (Cl-)
 Its ratio 1:2 so the formula unit is MgCl2.

crystalline solids at room temperature
 generally have high melting points
 can conduct an electric current when melted or
dissolved in water.
The coordination number of an ion is the number of ions
of opposite charge that surround the ion in a crystal.
The valence electrons of metal atoms can be modeled as
a sea of electrons. That is the valence electrons are
mobile and can drift freely from one part of the metal
to another.
Metallic bonds consist of the attraction of the freefloating valence electrons for the positively charged
metal ions.
Are mixture composed of two or more elements, at least
one of which is a metal. Ex. Brass, an alloy of copper
and zinc.
Alloys are important because their properties are often
superior to those of their component elements.
Covalent Bonding
Covalent bond is the atoms held together by sharing
electrons.
Molecule is a neutral group of atoms joined together by
covalent bonds.
Diatomic molecule is a molecule consisting of two atoms.
Ex. Oxygen molecule
Molecular compound is a compound composed of
molecule Ex. Water and Carbon monoxide (CO).
 compounds tend to have relatively lower melting
and boiling points than ionic compounds.
A molecular formula is the chemical formula of a
molecular compound.
 It shows how many atoms of each element a
molecule contains.
 it reflects the actual number of atoms in each
molecule.
 it describes molecules consisting of 1 element
 It doesn’t tell you about a molecules structure.
In forming covalent bonds, electron sharing
usually occurs so that atoms attain the electron
configuration of noble gases.
Two atoms held together by sharing pair of electrons
are joined by a single covalent bond
Ex.
H.
+
hydrogen
atom
.H
hydrogen
atom
H:H
hydrogen
molecule
Name
Chemical Properties and Uses
Formula
Fluorine
F2
Greenish yellow reactive toxic gas.
Compounds of this is halogen, are added
to drinking water and toothpaste.
Chlorine CL2
Greenish yellow reactive toxic gas.
Chlorine is a halogen used in household
bleaching agents.
Bromine
Br2
Iodine
I2
Dense red-brown liquid with pungent
odor. Compounds of bromine, a halogen
are used in the photographic emulsions
Dense gray-black solid that produces
purple vapors; a halogen. A solution of
iodine in alcohol is used as an antiseptic.
Name
Chemical
Formula
Properties and Uses
Hydrogen
H2
Colorless,
Hydrogen
element.
odorless, tasteless gas.
is the lightest known
Nitrogen
N2
Oxygen
O2
Colorless, odorless, tasteless gas. Air is
almost 80% nitrogen by volume.
Colorless, odorless, tasteless gas that is
vital for life. Air is about 20% oxygen
by volume.
Name
Chemical
Formula
Properties and Uses
Hydrog H2O2
en
peroxid
e
Colorless, unstable liquid when pure. It is
used as rocket fuel. A 3% solution is used
as a bleach and antiseptic.
Sulfur SO2
dioxide
Oxides of sulfur are produced in
combustion of petroleum products and
coal. They are major air pollutants in
industrial areas. Oxides of sulfur can lead
to respiratory problems.
Oxides of nitrogen are major air pollutants
produced by the combustion of fossil fuels
in automobile engines. They irritate the
Nitric NO
oxide/
dioxide
Name
Chemical
Formula
Nitrous
oxide
N2O
Hydroge HCN
n cyanide
Hydroge
n
fluoride/
chloride
HF
HCl
Properties and Uses
Colorless, sweet-smelling gas. It is used as
an anesthetic commonly called laughing
gas.
Colorless, toxic gas with the smell of
almonds.
Two hydrogen in halides, all extremely
soluble in water. Hydrogen chloride, a
colorless gas with pungent odor, readily
dissolves in water to give a solution called
hydrochloric acid.
Monatomic Ions consists of a single atom with a
positive or negative charge resulting from the loss
or gain of one or more valence electrons,
respectively.
Symbol
Cu+
Cu2+
Fe2+
Stock name
Copper(I) ion
Copper(II) ion
Iron(II) ion
Classical name
Cuprous ion
Cupric ion
Ferrous ion
Fe3+
Iron(III) ion
Ferric ion
Composed of more than one atom
- ite
SO3 2-
sulfite
- ate
SO4 2- sulfate
NO2 -
nitrite
NO3 – nitrate
CIO2 -
chlorate
CIO3 - chlorate
Intermolecular attractions are weaker than either ionic or
covalent bond
 These attractions are responsible for determining
whether a molecular compound is a gas, a liquid, or a
solid at a given temperature.
Van der Waals Forces named after the Dutch chemist
Johannes van Der Waals. It has two weakest attractions
between molecules.

Dipole interactions occur when
molecules are attracted to one another.
negative attracts positive
polar
Dispersion Forces weakest of all molecular interactions
and caused by the motion of electrons.
They occur even between non-polar molecules.
Naming and Writing
Formulas for Ionic
Compounds
Antoine-Laurent Lavoisler a French chemist, he
determined the composition of many compounds
in his experiments to show how chemical
compounds form.
1. Write the symbol for the cation first then followed
by the anion.
2. Write the charge for each ion in each compound.
3. Balance the formula using appropriate subscripts
Binary Ionic Compound is composed of two elements
and can be either ionic or molecular.
To name any binary ionic compound, place the cation
name first, followed by the anion name.
Ex.
 cesium oxide
CS2O
 copper(I) oxide
CuO
 copper(II) oxide
Cu2O
 stannous fluoride
SnF2
 stannic sulfide
SnS2




potassium chloride
KCl
calcium bromide
CaBr2
iron III oxide
Fe2O3
calcium sulfide
CaS
Write the formulas for these binary ionic
compounds:
 copper sulfide
 potassium nitride
Cu2+ and S2+
2
Cu
2-
S
Ans 1 (2+) + 1 (2-) = 0
K+ and N32-
1+
K
N
Ans 3 (1+) + 1 (3-) = 0
Naming and Writing
Formulas for Acids and
Bases
Acids are a group of ionic compounds with unique
properties.
 a compound that contains one or more
hydrogen atoms and produces hydrogen ions
(H+) when dissolved in water.
 chemical formula of acid in general form is
HnX; wherein n is monatomic or polyatomic
anion.
When the name of the anion (X) ends in -ide, the
acid name begins with the prefix hydro-. The stem of
the anion has the suffix –ic and is followed by the
word acid.
2. When the anion name end in –ite, the acid name is
the stem of the anion with the suffix –ous, followed
by the word acid.
3. When the anion name ends in –ate, the acid name is
the stem of the anion with the suffix -ic followed by
the word acid.
1.
Anion
Ending
Example
Acid Name
Example
Chloride, Cl-
Hydro-(stem)-ic Hydrochloric
acid
acid
- ite
Sulfite, SO3
2-
(stem)-ous acid
- ate
Nitrate, NO3- (stem)- ic acid
- ide
Sulfurous
acid
nitric acid
Name
Formula
Hydrochloric acid
HCl
Sulfuric Acid
H2SO4
Nitric Acid
HNO3
Acetic Acid
CH3COOH
Phosphoric Acid
H3PO4
Carbonic Acid
H2CO3
Base is an ionic compound that produces hydroxide ions
when dissolved in water.
Bases are named in the same way as other ionic
compounds. The name of the cation is followed by the
name of the anion.
Ex. sodium hydroxide NaOH
aluminum hydroxide (Al 3+) (OH-) then balance
Al(OH3)
Practicing Skills: Naming
Chemical Compounds
Yes
(HNO3, nitric acid)
Q = H?
Compound is binary
name ends in ide
No
No
No
>2
Elements?
Compound is an acid.
Yes
Compound contains a
polyatomic ion ;name
generally ends in ite or ate
Yes
Q = Metal?
Compound is binary molecular
use prefixes in the name
(N2O3, dinitrogen trioxide)
Yes
Name the ions
(BaS, barium sulfide)
Q = Group
A?
No
Name the ions
Name the ions
(Li2CO3, lithium
Roman Num w/ cation
carbonate)
(CuSO 4, copper(II) sulfate
Q = Group
A?
No
Name the ions; use a Roman
Roman numeral w/ cation.
(FeCl2, iron(II) chloride)
 An -ide ending generally indicates a binary
compound
 An –ite or -ate ending means a polyatomic ion that
includes oxygen is in the formula
 Prefixes in a name genrally indicate that the
compound is molecular
 A Roman numeral after the name of a cation shows
the ionic charge of the cation.
Contains prefixes? Yes
No
Molecular compound
Uses prefixes to write formula
Ionic compound
identify symbols
Group A elements
Roman Numerals
Use table 9.1 for
for charges
Give charges
for cations
Balance
charges
Polyatomic Ions
Use table 9.3
for charges
Uses criscross method
parenthesis for any multiple
polyatomic ions
The Mole: A Measurement
of Matter
What is the mass of 90 averaged-sized apples if one
dozen of apples has a mass of 2.0 kg?
Given:
Number of apples= 90 apples
12 apples = 1 dozen apples
1 dozen apples = 2.0 kg apples
Unknown:
Mass of 90 apples = ? Kg
Formula:
Number of apples
dozen of apples
mass of apple
Solution:
Mass of apples = 90 apples x 1 dozen x 2.0 kg
12 apples
1 dozen
Answer:
= 15 kg apples
The mass of 90 average-sized apples is 15 kg.
 chemist
use this as a unit in a specified number of
particles.
 1 mole = 6.02 x 10 23 SI unit for measuring the amount
of a substance.
 it is called Avogadro’s number ( Italian scientist
Amedeo Avogadro di Quaregna.
 he helped clarify the difference between atoms and
molecules.
 representative particle refers to the species present in
a substance, usually atoms, molecules, or formula units.
 representative particle of most element is atom
How many moles of magnesium is 1.25 x 1023 atoms of
magnesium?
Given:
Number of atoms = 1.25 x 1023 atoms Mg
1 mol Mg = 6.02 x 10 23 atoms Mg
Unknown :
Moles= ? Mol Mg
The desired conversion is atoms
moles
Solution:
The conversion factor is
1 mol Mg
6.02 x 10 23 atoms Mg
Multiplying atoms of Mg by the conversion factor gives
the answer.
Moles = 1.25 x 10 23 atoms Mg x 1 mol Mg
6.02 x 10 23 atoms Mg
Answer: 2.08 x 10-1
0.208 mol Mg
Ex. (CO2) has 3 atoms but I mole only so it its , thus a
mole of Carbon dioxide is three times Avogadro’s
number of atoms
Sample Problem;
Propane is gas used for cooking and heating. How many
atoms are in 2.12 mol of propane (C3H8)?
Given:
# of moles = 2.12 mol C3H8
1 mol C3H8 = 6.02 x 10 23 molecules C3H8
I molecule C3H8 = 11 atoms
The desired conversion is:
moles
molecules
atoms
Solution:
1st conversion factor 6.02 x 10 23 molecules C3H8
1 mol C3H8
2nd conversion factor 11 atoms
1 molecule C3H8
Multiply the moles of C3H8 by the proper conversion
factors:
2.12 mol C3H8 x 6.02 x 10 23 molecules C3H8 x 11 atoms
1 mol C3H8
1molecule
C3H8
Answer :
= 1.4039 x 1025 atoms
= 1.40 x 1025 atoms
Writing Chemical
Equations
 One more substance (the reactants) change into one or
more new substances (the products).
 Example is the chemical changes occur when bread is
baked
 ex. Rusting of iron (iron reacts with oxygen to produce
iron(III) oxide (rust).
 use a word equation:
 iron + oxygen
iron (III) oxide
 ex. Hydrogen peroxide (formation of bubbles of
 oxygen to a wound); the gas is he chemical change
 two new substance are produced: oxygen, gas and
 therefore, Hydrogen peroxide decomposes to form
water and oxygen gas.

Hydrogen peroxide
water + oxygen
 ex. Methane gas (for cooking) major component for
natural gas
 in writing equation urning burning a substance
requires oxygen
 methane + oxygen
carbon dioxide + water
 representation of chemical reaction
The formulas of the reactants are connected by
an arrow with the formulas of the products.
 chemical equation for rusting:
 Fe + O2
Fe 2O3
 Catalyst is a substance that speeds up the
reaction but is not used up in the reaction.
Catalyst is neither a reactant or a product so
its
formula is written above the arrow.
 ex. Compound manganese(IV) oxide (MnO2)
catalyzes the decomposition of an aqueous
solution of
hydrogen peroxide (H2O2) to produce water
and oxygen

H2O2 MnO2
H 2O + O 2
+
Used to separate two reactants or two products
“Yields,” Separates reactants from products
(aq)
heat
Pt
Designates an aqueous solution, the substance
is dissolves in water, placed after the formula
Indicates that heat is supplied ti the reaction
A formula written above or below the yield sign
indicates its use as a catalyst (ex. platinum
 bicycle; wherein
Frame + wheel + handlebar + pedal
 will become
F+W+H+P
FW2HP 2
 C + O2
CO 2
bicycle

Hydrogen and oxygen react to form water. The reaction
releases enough energy to launch a rocket.
 H 2 + O2
2H2O
Reactants
4 hydrogen atoms
2 oxygen atoms
Products
4 hydrogen atoms
2 oxygen atoms
Aluminum is a good choice for outdoor furniture
because it reacts with oxygen in the air to form a thin
protective coat of aluminum oxide. Balance the equation
for this reaction.
Al + O2
Al2O3
Answer
8 Al + 6O2
4Al2O3
Classifying
Reactions
1.
2.
3.
4.
5.
Combination
Decomposition
Single- Rplacement
Double Replacement
Combustion
 is
a chemical change in which two or
more substances react to form a single new
substance.
 ex. 2Mg + O2
2MgO
 Cu + S
Cu2S
 2Cu + S
Cu2S
 a single compound breaks down into two or more
simpler products.
 involve only one reactant aand two or more products
 requires energy in the form heat, light, or electricity
electricity

H 2O
H2
+ O2
electricity

2H 2O
2H2
+ O2
 in which one element replaces a second element in a
compound.
 ex. Zn + Cu (NO3)2
Cu + Zn (NO3)2
 Bromine is more active than iodine, so this reaction
occurs

Br2 + NaCl
NaBr + I2
 But Bromine is less active than chlorine, so this
reaction does not occur.

Br2 + NaCl
No reaction

Cl2 + 2NaBr
2NaCl + Br

Cl NaBr
NaCl
a chemical change involving an exchange of positive
ions between two compounds.
 also called double-displacement reaction
 take place in aqueous solution and often produce a
precipitate, a gas, or a molecular compound such as
water.
 ex. One of the product is a gas. Poisonous hydrogen
cyanide gas is produced when aqueous sodium cyanide
is mixed with sulfuric acid.
 2NaCn + H2SO4
2HCN + Na2SO4

 a chemical change in which an element or compound
reacts with oxygen.
 often produce energy in the form of heat and light.
 always involves oxygen as a reactant
 gasoline is a mixture of hydrocarbons that can be
approximately represented by the formula C8H18

2 C8H18 + 25O2
16CO2
+ 18H2 O
Stoichiometry
 Gr. Stoikheloin, means elements and metron,
meaning measure
 is the calculation of amounts of substances involved
in chemical reactions.
 just like a bookeeping, accountants can track income,
expenditures, and profits for small business by
tallying each in dollars and cents
Chemists can track reactants and products in a
reaction.
 Allows chemist to tally the amounts of reactants and
products using ratio of moles.
In a five-day workweek. Tiny Tyke is scheduled to make
640 tricycles. How many wheels should be in the plant on
Monday morning to make these tricycles.
Given
number of tricycles = 640 tricycles = 640 FW3HP2
F + 3W + H + 2P
FW3HP2
Unknown
number of wheels = ? Wheels
Solution
3W
1 FW3HP2
FW3HP2
3W
 the
desired unit is W. Multiply the number of
tricycles by the conversion factor.
640 FW3HP2 x
3W
1 FW3HP2
= 1920 W