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
³₁H
negligible
3.0160
³He
2
4He
2
0.0001
3.0160
Average
atomic mass
1.0079
4.0026
99.9999
4.0026
Calculate the atomic mass of Helium
(To calculate: multiply the mass of each
isotope by its natural abundance, express
as a decimal, and then add the products.)
AMU of He = (3.0160 x 0.0001) + (4.0026 x
99.999)
=
Isotope = 10 X
Mass # = 10.012
Relative abundance = 19.91% = 0.1991
AMU = ?
Isotope = 11 X
Mass # = 11.009
Relative 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
• 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.