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SCIENTIFIC METHOD
You have completed kindergarten at [North Hills Preparatory]
and are awarded this diploma in recognition of your accomplishments.
How to memorize?
 Angry
 Hippos
 Eat
 Dry
 Cereal
What does it mean?
 Angry – Ask a testable question
 Hippos – Form a reasonable hypothesis
 Eat - Experiment
 Dry – Data Analysis
 Cereal – Form a conclusion
SCIENTIFIC
METHOD
By: Keiji Chan and Nael Alami
The Basics
The Scientific Method is the logical
process to solving problems by:
1. Observing
2. Collecting data
3. Creating hypotheses
4. Testing hypotheses
5. Formulating theories backed up by
data
• The scientific method can also be a
continuous cycle
Observing
& Collecting
Data
• Also known as research
Quantitative Data
• Numerical
• Measured
• Example: There are 7 birds outside.
Qualitative Data
• Describes things
• Not in number form
• Example: The birds are blue.
• Data does not have to be only numbers, but
descriptive ideas too.
Formulating
& Testing
Hypotheses
• A statement that provokes further
investigation and can be proven right or
wrong
• Can be written in an “if…then” format, a
form of cause and effect statement
• Test one variable, preferably
– Ensures more definite results
• Must be tested multiple times for accuracy
• Based off of some sort of observation or
past evidence
• Even if proven incorrect, a hypothesis
should be revised, not rejected
Controls &
Variables
Controls
• Factors not tampered with during an
experiment
• Examples:
– Moisture, light exposure, most
environmental conditions
Independent Variable
•
The factor being changed for the purpose of
the experiment
Dependent Variable
•
Factor being measured
•
Depends on independent variable
Models &
Theories
Models
• Explanation of the process and connections of
an occurrence
– Can come in many forms (ex: diagram)
– Small scale version of something that
proves a theory
Theories
• Basic overview of the facts supporting a
conclusion
• Can be represented in models
• Can foresee the results of an experiment
SCIENTIFIC METHOD
JOSUE AND FARLEY
HYPOTHESIS
• IS A EDUCATED GUESS
• SERVES AS BASIS FOR MAKING PREDICATIONS
• ALSO CARRYING OUT FURTHER EXPERIMENTS(LEADING TO THEM).
CONTROLS
• EXPERIMENT DESIGNED TO MINIMIZE THE EFFECTS OF VARIABLES OTHER THAN THE
INDEPENDENT VARIABLE.
• IN OTHER WORDS, DOING AN EXPERIMENT THE SAME BUT WITHOUT VARIABLES THAT YOU'RE
TESTING FOR, SO YOU CAN COMPARE RESULTS.
• E.G: SEE HOW WELL CAFFEINE WILL AFFECT YOUR BODY. THE ONE WITHOUT CAFFEINE IN
THE THEIR BODY IS THE CONTROL.
• WANTS TO REMAIN THE SAME
VARIABLES
• INDEPENDENT VARIABLE
-IT'S MANIPULATED OR IT IS CHANGED BY THE SCIENTIST.
• DEPENDENT VARIABLE
-IT IS OBSERVED OR MEASURED IN THE EXPERIMENT.
• A VARIABLE IS ANY ITEM OR FACTOR OR CONDITION THAT CAN BE CONTROLLED OR
CHANGED IN THE EXPIREMENT.
THEORIES
• THEORY- A GROUP OF PROPORTIONS CREATED TO EXPLAIN A GROUP OF FACTS OR
PHENOMENA.
• THEORIES IN THE EXPERIMENT ARE PREDICTIONS OF WHAT IS THOUGHT TO BE THE OUTCOME.
• A THEORY IS CONSIDERED TO BE CORRECT TO BE TRUE/ USEFUL IF IT CAN PREDICT THE
RESULTS OF NEW EXPERIMENTS.
The Scientific Method
Thomas Arnett
Steps
• Ask a Testable Question
• Make a good hypothesis
• Do the experiment
• Analyze Data
• Make a conclusion
• Angry Hippos Eat Dried Cereal
Question
• Needs to be testable
• Not a yes or no question
• Should be what your experiment is based on
Formulating a hypothesis
• The hypothesis is a short statement of what you will be doing and
what you think will happen
• Should be testable as well
• Does not need to be an if…then statement
Experimenting
• The experiment needs to be easily replicable
• You will need:
• Procedures
• Materials
• In addition to those, you need to observe data
• Data can be qualitative or quantitative
Data Analysis
• This is where you review your data, do your math, and eventually
begin to conclude
• Should explain what your observations and data mean and why they
are significant
• Needs to analyze everything recorded
Conclusion
• Conclude if your hypothesis was correct or not
• Explain why this information is relevant
• Should use your data to back up your concluding statements
MEASUREMENTS
MEASUREMENTS, SIGNIFICANT
FIGURES, SCIENTIFIC NOTATION,
PRECISION AND ACCURACY
By Manasi Taduri & Sampreeti Bingi
METRIC UNIT MEASUREMENTS
• The IS (International System of units) is used all around the world. Except America.
• It was created in 1799 and was used temporarily to replace the previously existing
system and then replaced it for good since it as easier to understand.
Base
Measurement
Symbol
Customary Equivalent
Meter
Length
m
Feet
Gram
Mass
g
Ounces
Liter
Volume
L
Fluid ounces
Second
Time
s
second
PREFIXES
LESS THAN THE BASE
• Deci – 10^-1 x base; 1/10 base units
• Centi – 10^-2 x base; 1/100 base units
• Milli - 10^3 x base; 1/1,000 base units
• Micro - 10^6 x base; 1/1,000,000 base units
• Nano - 10^-9 x base; 1/1,000,000,000 base units
• Pico - 10^-12 x base; 1/1,000,000,000,000 base units
PREFIXES CONT.
GREATER THAN THE BASE
• Deca – 10^1 x base; 10 base units
• Hecto – 10^2 x base; 100 base units
• Kilo – 10^3 x base; 1,000 base units
• Mega - 10^6 x base; 1,000,000
• Giga – 10^9 x base; 1,000,000,000
• Tera – 10^12 x base; 1,000,000,000,000
SIGNIFICANT FIGURES
• All number 1 through 9 are significant
• Zeros are NEVER significant when at the beginning of a number; they are only place
holders.
• Example: 0.000748 only has three sig figs.
• Zeros in the middle of the number are ALWAYS significant.
• Example: 9087 has four sig figs
• Zeros are only significant at the end of a number if there is a decimal point.
• Examples: 7850 has three sig figs. 8490.0 has five sig figs. 0.09840 has four sig figs.
SCIENTIFIC NOTATIONS
• Scientific notation is a way to easily convert very large numbers or very small numbers to something more
manageable.
• For numbers less than zero, you need to move the decimal to the right until you hit the SECOND non-zero
number.
• For example, if you have .000000312, you would move the decimal 7 numbers to the RIGHT so you end up
with 3.12 and then you add the power of 10 which is the number of spaces you moved (negative exponent if
you moved right) so you would get 3.21 x 10^-7.
• For numbers larger than zero, you do the same thing with one difference. You move the decimal to the left
until your decimal is behind the first non-zero digit in the number.
• Example: If you have 456000000. then you move the decimal 8 places to the LEFT so you get 4.56 and then
you add the power of ten and it would be 10^8 so your final answer would be 4.56 x 10^8. :)
ACCURACY AND PRECISION
• Accuracy: the degree to which the result of a measurement, calculation, or
specification conforms to the correct value of the correct value or a standard
• This means that the values are close to the intended target but not close to each
other.
• Precision: refinement in a measurement, calculation, or specification, especially as
represented by the number of digits given
MEASUREMENT
By: Alea Badayos &
Serena Myoung
BASE UNITS
Meter units
Quantity
Abbreviation
Customary units
Meter
Length
m
foot
Gram
Mass
g
Ounces
Liter
Volume
L
Fluid ounces
Second
Time
s
seconds
PREFIXES
Prefix
Abbv.
Exponential form
Tera
T
10^12 x base
giga
G
10^9 x base
mega
M
10^6 x base
kilo
K
10^3 x base
Hecto
h
10^2 x base
Deca
D
10^1 x base
BASE
(m/L/g/s)
Deci
d
10^-1 x base
Senti
c
10^-2 x base
Milli
m
10^-3 x base
Micro
µ
10^-6 x base
Nano
n
10^-9 x base
pico
p
10^-12 x base
UNIT CONVERSIONS
Prefix  Base Units
1. Replace prefix with corresponding integer power of ten.
2. Simplify value (move decimal the correct number of places in the correct
direction).
EX: 354.84 milliliter  354.84 x 10 -3 L
“10 -3 ” means the decimal moves 3 places to the LEFT
354.84 milliliter s = 0.35484 liters
Negative power of ten = decimal is moved to the left
Positive power of ten = decimal is moved to the right
UNIT CONVERSION
Base Units  Prefix
1. Divide the measurements by its corresponding integer of power of ten.
2. Add prefix name to the unit.
If multiplication is preferred, instead of division …
1. Apply -1 exponent to power of 10 to move denominator to numerator, so it is
easier to solve..
2. Multiply measurement by new power of ten.
EX: meters  kilometers (kilo= 10 3 )
1609 meters/10 3  1609 x 10 -3  1 .609 kilometers
Negative exponent =
move the decimal to
the left.
UNIT CONVERSION
Prefix  Prefix
1. Multiply the number by its own integer power of 10.
2. Divide the second prefix’s integer power of 10 last.
EX: 818 mL  µL; m=10 -3 , µ=10 -6
818 x (10 -3 /10 -6 ) x 10 6  10 3  818 x 10 3 = 818,000µL
Add the exponents
ACCURACY & PRECISION
 Accuracy: the closeness of the measurements to the accepted value (your “bulls
eye”) of the quantity measured.
 Precision: closeness of the set of measurements of the same quantity; the numbers
are close to eachother.
SIGNIFICANT FIGURES (SIG FIGS)
 Digits that carry meaning contributing to the exactness of the measurements.
 Consist of all digits known with certainty plus one final digit that is estimated.
 Rules of sig for 0:
 Beginning of the number: NEVER
 Middle of the number: ALWAYS
 End of the number: SOMETIMES
- When the zero has non-zero digits before it (e.g.0.650): YES
- When there’s a decimal present after the zero (e.g. 650.0): YES
- When there’s no decimal present after the zero (e.g. 650): NO
SCIENTIFIC NOTATION
 Large numbers written in an easy format and is rounded to the hundredth place:
#.## x 10 n
 When the exponent is positive, the decimal place is moved to the right, making the
number greater than one.
 When the exponent is negative, the decimal place is moved the left, making the
number less than one.
EX: 0.0000076392  7.64 x 10 -6
4315.09  4.32 x 10 3
0.00054925 x 10 4  5.4925 x 10 4
Measurements
By Waris Sandhu and John Henry Nwachukwu
International System of Units
•
•
•
System of measurements based on prefixes changing the amount of the base
Bases include grams, liters, meters
Prefixes change the amount of base in integers of 10
Base Units
• Grams measure mass (g)
• Meters measure distance (m)
• Seconds measure time (s)
• Liters measure volume (L)
Prefixes
• Pico= 10-12
• Nano=10-9
• Micro=10-6
• Milli=10-3
• Centi=10-2
• Deci=10-1
To get the value of a base
with a prefix multiply the
base by the values.
Prefixes Contd.
• Tera= 1012
• Giga= 109
• Mega= 106
• Kilo= 103
• Hecto= 102
• Deca= 101
Prefix Abbreviations
• Pico (p)
• Nano (n)
• Micro (µ)
• Mega (M)
• Giga (G)
Prefix Abbreviations Cont.
• Kilo (k)
• Hecto (h)
• Deca (D)
• Deci (d)
• Centi (c)
• Milli (m)
Measuring with Metric Rulers
• Measurements on a ruler should be estimated one digit larger that the
certain measurement. (if the measurement is .4 centimeters, it should be
written as .40 cm)
Graduated Cylinder
• Meniscus- A curve at the top of a liquid
caused by surface tension
• Measuring should be done from the bottom
of the meniscus
Accuracy and Precision
ACCURACY
• How close you are to the true value.
Precision
• How close your measurements are to each other.
Sig Figs
• Significant Figures- Digits that contribute exactness towards the
measurement.
• Measurement where all digits are certain except for the last which is
rounded or estimated.
Counting Sig Figs
• All non-zero numbers are sig figs
• Zeroes at the beginning of a number are never sig figs
• Zeroes in the middle of the number are always sig figs
• Zeroes at the end of a number are sometimes sig figs
•
•
When there is a decimal present it is a sig fig
When there is no decimal it is not a sig fig
Counting Sig Figs contd.
• Whole quantities, Constants, and Conversion Factors have infinite sig figs.
• To find sig figs after adding and subtracting, the larger of the smallest place
value is used to determine what the place value of the final sig fig will be in
the result.
• To find sig figs after multiplying or dividing, the total sig figs in the product
is equal to the least number of total sig figs of the numbers that are being
divided or multiplied.
Scientific Notation
• Used for writing really small or large numbers into forms that are easier to
read
• Positive exponent, number is greater than one
• Negative exponent, number is less than one
• Number should always be to the hundredths place
Unit Conversions
• To convert to base unit, replace the prefix with its corresponding integer
power of ten. Then, simplify the value by moving the decimal the
appropriate number of times in the correct direction.
• 400.00 millimeters > .4 meters
Converting from the Base Unit
• To convert from base to prefix-unit, divide by the corresponding integer
power of ten and add the correct prefix name to the unit
• 1700 meters/ 103 > 1.7 kilometers
ATOMS
Atoms
AT O M - T H E BA S I C U N I T O F A C H E M I C A L E L E M E N T
Early Atomic Theory
1766-1844
Aristotle
Democritus
 Believed in the four elements of
 Believed matter could only be
air, water and fire
 Believed matter could be
divided infinitely without
changing it’s properties
divided until you got to the
smallest particle (atom)
 Atoms join with out
Plato
 Believed the world was created
from four elements (air, earth,
fire, water)
Modern Atomic Theory
1913
J.J. Thompson
 Discovered the
Electron
 Plum Pudding Model
(1904)
 Identified the
negativity charged
electron in the cathode
ray tube in 1897
Plum Pudding
Model
According to this model, the
atom is a sphere of positive
charge, and negatively charged
electrons are embedded in it to
balance the total positive charge.
Cathode Ray Tube
A cathode ray tube (CRT) is a
specialized vacuum tube in
which images are produced
when an electron beam strikes a
phosphorescent surface.
Ernest Rutherford
 Proposed the nuclear atom as
the result of the gold-foil
experiment in 1911
 Concluded that the atom is
mostly empty space
Gold-Foil Experiment
Robert Millikan
 A famous physicist
 The oil drop experiment was the
most famous scientific work of
his career
 Goal: to quantify the charge of
an electron
Law of conservation of mass
“Mass is neither destroyed nor
created during ordinary
chemical reactions or physical
changes.”
Law of definite proportions
“The fact that a chemical compound contains
the same elements in exactly the same
proportions by mass regardless of the size
of the sample or source of the compound”
Law of multiple proportions
“When two elements combine to
form more than one compound,
the mass of one element, which
combines with a fixed mass of
the other element, will always be
ratios of whole numbers.”
ATOMIC THEORY
BY Z.C P.S
ATOMIC THEORY
JOKES
▸ funny jokes
Thats Punny.
-Parker S.S.
ATOMIC THEORY
▸ Definition: any theory in which matter is regarded as
consisting of atoms.
‣ T current theoretical model of the atom involves a dense nucleus surrounded
by a probabilistic "cloud" of electrons. In chemistry and physics, atomic theory
is a scientific theory of the nature of matter, which states that matter is
composed of discrete units called atoms.
ATOMIC THEORY
HISTORY LESSON
‣ over the centuries there have been multiple versions of Atomic Theory.
Early Atomic Theory, Early Modern Atomic Theory, Dalton’s Atomic Theory,
and Modern Atomic Theory.
‣ The ancient Greek philosopher Aristotle he postulated that all mater
was made of earth, water, air and fire
‣ in 1789, Antoine Lavoisier formulated the total mass of reactants and
products in a chemical reaction remains constant
ATOMIC THEORY
FOUNDERS OF THE ATOMIC THEORY
▸ Democritus who predicted that everything is composed of
atoms and empty void.
▸ Plato believed that the world was created from four elements
▸ Aristotle argued that the elements were not made of atoms
▸ Antoine Lavoisier formulated that total mass of reactants and
products in a chemical reaction remains constant
▸ John Dalton created a Atomic Theory that helped support
Democritus’s prediction
▸ Nelis Bohr’s model of the atom is important because it
introduced the concept of the quantum in explaining atomic
properties.
ATOMIC THEORY
LAWS OF THE ATOMIC THEORY
▸ law of conservation of Mass ,1789
Mass is neither destroyed nor created during ordinary chemical
reactions or physical changes.
▸ law of Definite Proportions, 1799
A chemical compound contains the same elements in exactly the same
proportions by mass regardless of the size.
▸ law of multiple proportions, 1808
if two or more different compounds of the same two elements, then the
ratio of the masses of the second element combined with the first
element is always a ration of small whole numbers.
THE DIFFERENT ATOMIC THEORIES
▸ Early Atomic Theory: Proposed that everything is composed of atoms and empty voids.
▸ Early Modern Atomic Theory: Is the total mass of reactants and products in a chemical reaction remains constant.
▸ Dalton’s Atomic Theory: Dalton’s is the first complete attempt to describe all matter in terms of atoms and their
properties.
▸ Modern Atomic Theory: is a theory that all matter is composed of tiny particles called atoms.
FUN FACTS:
▸ “Atoms comes from Greek word “Atomos,”meaning “indivisible.”
▸ In 1897, J.J. Thompson created the Cathrode Ray Tube that discovered the
Electron
▸ In 1909, Ernest Rutherford performs the gold foil experiment to discover the
nucleus.
▸ In 1932, James Chadwick discovered the Neutron
BIBLIOGRAPHY:
▸ https://www.reference.com
▸
www.iun.edu/~cpanhd/.../modern-atomic-theory/modern-atomic-theory.
▸ https://www.khanacademy.org/.../atomic-structure.../daltons-atomic-theory-version-
Atoms
By Hrishikesh and Tarun
Early Atom Theory
Democritus
➔
Early 5th century BCE
➔
Proposed everything is made of atoms and empty void
➔
Atom comes from the word, Atomus meaning indivisible
Plato
➔
Late 5th century BCE
➔
Believed the world was made from 4 elements: Water, Earth, Fire, and Air
Aristotle
➔
Approx: 330BCE
➔
Argued against atoms, saying that the elements were not made of atoms
➔
Also believed change happened through transformation instead of rearrangement
Modern Atomic Theory
Further discoveries led to 2 postulates of Dalton’s atomic theory
being discredited.
➔ Atoms can be split into 3 subatomic particle- the electron (1e), neutron (neutral) and the proton (+1e).
➔ Atoms of same element can have different masses as
proved by the discovery of Isotopes
Law of Conservation of Mass
In 1789
➔ Mass cannot be destroyed nor created in
chemical reactions
➔ Every chemical reaction has an equal mass
for reactants and products.
Formulated by Antoine Lavoisier
Law of Definite Proportions
In 1799
➔ A chemical compound contains the same
elements in the exact same proportion by
mass, regardless of the source or size of
the sample
Proved by Joseph Louis Proust
Law of Multiple Proportions
In 1808
➔ States that when two elements combine to form more
than one compound, the mass of one element, which
combines with a fixed mass of the other element, will
always be ratios of whole numbers.
➔ Applies only to compounds composed of the same
elements
Proposed by John Dalton
Subatomic Particles
Electron:
➔
Extremely small negatively charged particle that orbits the nucleus
➔
AMU ≈ 0
➔
Discovered by J.J. Thompson in the cathode ray experiment
Charge: -1e
Proton
➔
Positively charged particle
➔
AMU = 1
➔
DIscovered by Ernest B. Rutherford
Charge: +1e
Neutron
➔
Particle roughly the same size as a proton with no charge
➔
AMU =1
Charge: 0 (neutral)
Isotopes
➔ Concept first proposed by F. Soddy.
➔ Atoms of the same element that contain different masses
due to different number of neutrons
➔ Unstable isotopes are radioactive.
➔ Isotopes can be represented in 2 ways: helium-4 or ⁴He (
With 4 being the mass number of that respective isotope).
Average Atomic Mass
➔ Calculated by summing the mass of the element’s isotopes, each multiplied by its natural abundance on
Earth.
➔ Average atomic mass = f1M1 + f2M2 + ... + fnMn where f is the fraction representing the natural
abundance of the isotope and M is the mass number (weight) of the isotope.
Example:
Average atomic mass of chlorine
1st isotope of chlorine= 35 amu, 75.77%, 2nd isotope= 37 amu, 24.23%
(0.7577 * 35 amu) + (0.2423 * 37 amu) = 35.483 amu
Bibliography:
http://www.ducksters.com/science/chemistry/isotopes_hydrogen.gif
http://images.tutorvista.com/cms/images/44/Atom.png
https://upload.wikimedia.org/wikipedia/commons/thumb/5/56/Law_of_Definite_Proportions.svg/2000pxLaw_of_Definite_Proportions.svg.png
http://study.com/cimages/multimages/16/nitrogenoxygencompound2.png
Subatomic Particles,
Isotopes, Avg. Atomic Mass
By: Tejas Saboo, Nick Patel, and Mariana Zollinger
Subatomic Particles
Protons
Located: in the nucleus
Charge: Positive (+)
Mass: 1 amu
Discoverer: Ernest Rutherford
Neutrons
Located: in the nucleus
Charge: No Charge (0)
Mass: 1 amu
Discoverer: James Chadwick
Subatomic Particles
Electron
Located: in the electron cloud
Charge: Negative (-)
Mass: (basically) 0 amu
Discoverer: J. J. Thompson
Isotopes
● Atoms of the same element have different masses,resulting in them having different mass numbers
-isotopes can be written as the element name followed by the mass number
The mass of the atom changes if the number of neutrons change
Some isotopes are more stable than others and are commonly found in nature
Average Atomic Mass
Atomic Mass Unit (AMU) is exactly 1/12 the mass of a Carbon-12 Atom
Average Atomic Mass= Average of the atomic masses of the different isotopes of a single element (based
on their percent of natural abundance)
Natural Abundance= Abundance of isotopes of an element naturally found on a planet
Calculating the average atomic mass of an element is just like calculating a weighted average!
HOW TO CALCULATE YOUR CHEM GRADE
CATEGORY
GRADE
X WEIGHT
=
CONTRIBUTION TO
AVG.
Minor Work
100
X 20%
=
20
Minor Lab
70
X 30%
=
21
=
45
=
86
In calculating the Average Atomic Mass of an element:
Test
90
X 50%
Category= Isotope
Grade= Atomic Mass (AMU)
Weight= Natural Abundance
TOTAL
Subatomic Particles,
Isotopes, and Average
Atomic Mass
By: Kumar Kudumula and Zain Abbas
The Electron
 Discovered by Joseph John Thomson a English physicist
through the use of the cathode ray tube experiment .
 Conclusions:
 Cathode Rays are negatively charged
 Cathode ray particles are 2000 times lighter than
the hydrogen Ion
 Atoms must have positive and negative charges
Atomic Model after the Discovery of the
Electron
The Plum Pudding Model
 Visualizes Thomson's last conclusion that atoms are positive
spheres with electrons scattered throughout
Electrons Cont.
 Robert A. Millikan preformed the oil drop experiment in
1909.
 Drop of oil moves up or down based on the voltage of the
metal plate
 Conclusion:
 The charge of an electron is (-1e) or -1 times the
elementary charge which is 1.602 x 10-19 Coulombs
Discovery of the Nucleus
 Ernest Rutherford preformed the gold foil experiment in
1909.
 Fired a beam of positively charged particles at a sheet of
gold foil
 Most particles went through the foil but 1 in 20,000 were
deflected back
 Rutherford’s Conclusions:
 The positive charge of an atom is concentrated in one
place
 Atoms have a dense, positively charged center
Rutherford's Conclusions Cont.
 Electrons are concentrated in a cloud around the nucleus
 Mass of an element comes from the nucleus
 Different elements have differently sized nuclei
 Electrons are made of positively charged particles, which are known as
protons
 Protons have a mass of 1 amu and have a charge of +1 time the
elementary charge
Rutherford’s Model of the Atom
 A central Nucleus made up of positively charged protons
 Electrons existing in a cloud around the nucleus
The Neutron
 James Chadwick shoots Beryllium with alpha particles from
Polonium.
 Radioactive particles with no charge are emitted from the
nucleus.
 Chadwick’s Conclusions:
 The uncharged particles are now known as Neutrons
 Neutrons have a mass of 1 amu
 Neutrons stabilize the nucleus
Isotopes
 Atoms of the same element that have different masses
 Can be written as the element name followed by the mass
number
 Ex: Titanium with a mass of 58 Titanium-58
 Mass of an atom changes if the total amount of neutrons
varies between each atom
Average Atomic Mass
 Definition: The weighted average of the atomic masses of the naturally
occurring isotopes of an element.
 Atomic Mass:
 An atomic mass unit is 1/12 of the mass of a carbon-12 atom
 The atomic mass of any nuclide is determined by comparing it with the
mass of a Carbon-12 atom.
Finding Average Atomic Mass
 In order to find the average atomic mass of an element you need both the
atomic masses of its isotopes and their natural abundance.
 Finding average atomic mass of Silver:
 Step 1: Find the masses of each isotope
 Step 2:Find the abundance of each isotope
 Step 3: Turn the percentages of each isotopes into decimals
 Step 4: Multiply the atomic mass of each isotopes with their respective
isotope
Finding Average Atomic Mass of Ag
Natural
Abundance
Contribution
to the
Average
Isotope
Atomic Mass
Ag-107
106.950 Amu
X
0.5186
=
54.441 amu
Ag-109
108.904 Amu
X
0.4814
=
52.4267 amu
Finally add both numbers together to get your average atomic
mass which is 107.868 atomic mass units.
SUBATOMIC PARTICLES, ISOTOPES,
AND AVERAGE ATOMIC MASS
SHEREBANOO & NIBITIKA
SUBATOMIC PARTICLES
• PARTICLES THAT ARE SMALLER THAN AN ATOM.
• ALSO KNOWN AS ELEMENTARY PARTICLE.
• INCLUDES: PROTONS, NEUTRONS, AND ELECTRONS.
ELECTRONS
• SUBATOMIC PARTICLES THAT HAVE A NEGATIVE CHARGE.
• REPRESENTED BY THE - SIGN
• FOUND IN THE ELECTRON CLOUD, OUTSIDE THE NUCLEUS.
• DISCOVERED BY JJ THOMSON THROUGH THE CATHODE RAY TUBE EXPERIMENT.
• THOMSON CONCLUDED:
1.
PARTICLES IN THE CATHODE RAY WERE 2000 TIMES SMALLER THAN THE LIGHTEST PARTICLE AT THE
TIME, WHICH WAS THE HYDROGEN ION.
2.
ATOMS ARE POSITIVE SPHERES WITH ELECTRONS SCATTERED THROUGHOUT.
PLUM PUDDING MODEL
PROTONS
• SUBATOMIC PARTICLES THAT HAVE A POSITIVE CHARGE.
• REPRESENTED BY THE + SIGN.
• FOUND IN THE NUCLEUS OF ATOM.
• DISCOVERED BY ERNEST RUTHERFORD THROUGH THE GOLD FOIL EXPERIMENT.
• RUTHERFORD’S DISCOVERIES:
1.
POSITIVE CHARGE IS FOCUSED ON ONE PART OF THE ATOM.
2.
THAT ONE PART IS CALLED THE NUCLEUS.
3.
ELECTRONS ARE IN A CLOUD AROUND THE NUCLEUS.
RUTHERFORD’S MODEL OF THE ATOM
NEUTRONS
• SUBATOMIC PARTICLES THAT HAVE A NEUTRAL CHARGE.
• FOUND IN THE NUCLEUS.
• FOUNDED BY JAMES CHADWICK IN 1932.
• CHADWICK BOMBARDED BERYLLIUM WITH ALPHA PARTICLES.
• HE DISCOVERED:
1.
NEUTRONS STABILIZE THE NUCLEUS AND KEEP PROTONS FROM REPELLING EACH OTHER.
2.
NEUTRONS HAVE A MASS OF ABOUT 1 AMU. (AN ATOMIC MASS UNIT)
ISOTOPES
• ATOMS OF THE SAME ELEMENT THAT HAVE DIFFERENT MASSES, THEREFORE DIFFERENT MASS
NUMBERS.
• SAME NUMBER OF PROTONS BUT DIFFERENT NUMBER OF NEUTRONS.
• IDENTICAL CHEMICAL PROPERTIES (REACTIVITY, FLAMMABILITY, ETC.) BUT DIFFERENT NUCLEAR
PROPERTIES (RADIOACTIVITY, ETC.)
AVERAGE ATOMIC MASS
• THE WEIGHTED AVERAGE OF THE ATOMIC MASSES OF THE NATURALLY OCCURRING ISOTOPES
OF AN ELEMENT.
• NATURAL ABUNDANCE:
THE PERCENTAGE OF THE ISOTOPE OF ALL NUCLIDES FOR THAT GIVEN ELEMENT.
NUCLIDE IS A GENERAL TERM FOR ANY ISOTOPE OF ANY ELEMENT.
FOUND ON THE PERIODIC TABLE: THE ATOMIC MASS
Atoms
WITH YA GIRLS LEI AND GEN
What is an atom?
 Atoms are the #basic units of matter.
 They make up everything. >.> those liars…
 They’re made up of THREE particles: Electrons. Protons, and Neutrons
Pre- Info
 Protons’ Mass≈Neutron’s Mass
 One of either>1800 electrons
Atoms…
 Have equal protons and electrons
 Adding a proton to an atom makes a new element, whereas adding a neutron
makes an isotope
PROTONS
 Positively Charged
 Found in Nucleus
 Ernest Rutherford
 Number of Protons=Atomic # determines what element it is and chemical
behavior
 EX. H:1, C:6, O:8
Electrons
 Negatively charged
 Electrically attracted to protons
Electron configuration: orbital description of the locations of the electrons in an
unlit/unexcited atom –helps chemists predict boiling point, stability, and
conductivity
Ex. Be:1s^2,2s^2, also written as Be:{HE}2s^2
SPDF
Electron Configuration
Electron configuration: orbital description of the locations of
the electrons in an unlit/unexcited atom –helps chemists
predict boiling point, stability, and conductivity
Ex. Be:1s^2,2s^2, also written as Be:{HE}2s^2
SPDF: SIMPDF
Interactive.. Activity…
 Cu
 He
 Ar
 Mg
 Se
Neutrons
•
•
•
•
Uncharged particles
Slightly larger than a proton
James Chadwick in 1932
Aren’t named or look like Jimmy
ISOTOPES
 any of two or more forms of a chemical element, having the same number of
protons in the nucleus, or the same atomic number, but having different
numbers of neutrons in the nucleus, or different atomic weights. There are 275
isotopes of the 81 stable elements, in addition to over 800 radioactive isotopes,
and every element has known isotopic forms. Isotopes of a single element
possess almost identical properties.
 # of neutrons determines the isotope of that element
Average Atomic Mass
To find aam…
 Make note of the isotopes and their percentages
 The final calculation is a weighted average meaning it roughly the mass of the
isotope times the relative abundance
 Add the results together
AAM Activity
Bibliography
 http://www.livescience.com/37206-atom-definition.html
 http://www.dictionary.com/browse/isotope
ATOMIC STRUCTURES
ATOMIC
STRUCTURES
SECTION: 4-1, 4-2, 5-3, 6-2
By: Macy and Kayleen
Lights and
Waves
• Electromagnetic Radiation is a form of energy that
exhibits wavelike behavior as it travels through space
• All forms of electromagnetic radiation travel at a
constant speed: (c= 3.00×108 m/s)
• Wavelength (lambada λ) is the distance between
corresponding points on adjacent waves
• Frequency (v) is the number of waves that pass a
given point in a specific time usually one second
• Planck’s constant (h= 6.62607004 × 10-34 m2 kg / s) is
the relationsip between quantuma nd frequency
• C= λv
E= hc/ λ
E=hv
Electromagnetic
Spectrum
Visible Light: 400 to 700
• Used to describe location of electrons in electron cloud
Quantum
Numbers
• Four quantum numbers= complete set
• Every electron has unqique set of numbers
• Quantum numbers: n, ℓ, mℓ, ms
• n= any whole number but 0
• ℓ= any number lower than n
• mℓ = every number in between the negative value of ℓ
and ℓ plus the negative value of ℓ and ℓ
• ms = 1/2 or -1/2
Quantum
Numbers and
Orbitals
• Orbitals: three dimensional region around the nucleus
that indicates the probable location of an electron
• Each l value corresponds with a different orbital shape
• Each orbital shape can hold different amounts of
electrons
• S= sphere (1)
• P= dumbbell (3)
• D= Clover (5)
L Value
0
1
2
3
4
5
6
Orbital
S
Max # of e 2
P
6
D
10
F
14
G
18
H
22
I
26
Lewis Dot
Diagrams
• Diagrams that show the bonding between
atoms of a molecule and the lone pairs of
electrons that may exist in the molecule
• Created with the valence electrons of an
element
Electron
Configuration
• Is the arrangement of electrons in an atom
• An electron always occupies the lowest energy orbital
that can receive it (Aufbu principle)
• No 2 electrons in same atom can have same set (Pauli
exclusion principle)
• Orbitals of equal energy are occupied by one electron
before occupied by a second (Hund’s rule)
Valence
Electrons
• Electrons lost, gained, or shared in chemical
compound formation
• These are electrons in the highest occupied
energy level (max: 8)
• Number of valence electrons can be
determined by electron configuration
• Inner shell electrons: not in the highest
energy level
Ions
• An ion is a charged atom or molecule
• Atoms form ions to satisfy octet rule
• Will decrease to 0 or increase to 8 depending
on what is easier
• Ex. 2 valence electrons will lose 2 electrons to
decrease to 0
• Ex. 6 valence electrons will increase by 4 to
get to 8
• Cation=0
Anion=8
Atomic Structures
Light
• Visible Light- a kind of electromagnetic radiation, which is a form
of energy that exhibits wavelike behavior as it travels through
space.
• All forms of electromagnetic radiation move at constant speeds of
3.0 x 10^8 m/s
• Wavelength: distance between corresponding points on adjacent
waves.
• Frequency: the number of waves that pass a given point in a
specific time
Electromagnetic Spectrum
Quantum Numbers
• Photoelectric effect: refers to the emission of electrons from a
metal when light shines on the metal
• Max Planck studied the photoelectric effect, and concluded that a
hot object does not continuously emit energy. Instead, it emits
energy in small, specific amounts called quanta.
• Quantum: the minimum number of energy that can be lost or
gained by an atom.
• Planck’s constant (referred to as h): 6.626 x 10^-34 J/s
Quantum Models
• Quantum Theory: describes mathematically the wave properties of electrons
and other very small particles.
• Quantum numbers: specify the properties of atomic orbitals and the properties
of electrons in orbitals
• The first three quantum numbers result from the Schrodinger equation, and
indicate the main energy level, shape, and orientation of an orbital
• Principal quantum number (n): indicates the main energy level
• Values begin at n=1 continues with whole numbers
• As value of n increases, the average distance from the nucleus increases
• Angular quantum number (l): indicates the shape of an orbital
• When L equals 0: the orbital is s, 1=p, 2=d, 3=f
• S orbitals are spherical, p orbitals have dumbbell shapes, and d orbitals have
clover shapes
Atomic structures
Electron Configuration
• Electron configuration: the arrangement of electrons in an atom
• First, orbital energy levels must be determined
• Orbital notation: orbitals are indicated by lines, and the name written
beneath. Each electron in an orbital is indicated by a half-arrow. If there
are two, one must be going up and the other down.
• Electron configuration notation: eliminates the lines and arrows of orbital
notation. Instead, a superscript is added to the orbital name to represent
the number of electrons.
Examples
Chart for electron configuration
Rules for electron configuration
• Aufbau principle: an electron occupies the lowest-energy orbital
that can receive it.
• Pauli exclusion principle: no two electrons in the same atom can
have the same set of four quantum numbers. Thus, an orbital can
hold two electrons of opposite spin.
• Hund’s rule: orbitals of equal energy are each occupied by one
electron before any orbital is occupied by a second electron, and
all electrons in singly occupied orbitals must have the same spin.
Lewis dot diagrams
• Valence electrons: the electrons available to be lost gained, or shared in
the formation of chemical compounds
• Since valence electrons are outer-shell electrons, their number can be
determined. For example, s2p4 would have 6 valence electrons.
• Electron-dot notation: an electron configuration notation in which only
the valence electrons of an atom of a particular element are shown,
indicated by dots placed around the element’s symbol.
• Example: If you complete the electron configuration notation for
fluorine, you will find that there are 7 valence electrons. In addition,
you can find the element’s group on the periodic table if it is in the A or
B block.
Lewis dot diagrams
Ions
• Ion: a charged atom
• Ions have either negative or positive charges through the gain or
loss of electrons.
• Cation: atom with a positive charge
• Anion: Atom with a negative charge
ATOMIC STRUTURE
(WOOHOO!)
BY: BITHIAH AND SHRAYA!!!
WHAT IS AN ATOM?
(To refresh your memory because if you don’t know by now, yikes!)
1) Atom: An atom is the basic (and smallest) unit of matter and is
the “defining” structure of an element.
2) The atom was first proposed by Democritus (lucky atom!) in early
5th century BCE, where Democritus suggested that everything is
composed of atoms & empty void.
3) Atoms consist of three subatomic particles, protons, neutrons,
and electrons.
4) In their natural state, atoms are electrically neutral.
WHAT IS THE ATOM? (CONT.)
1)
2)
3)
The body of an atom is made up of two different parts, the nucleus and
the electron cloud.
Nucleus: Is located in the center of the atom. It consists of both protons
and neutrons and is makes up most of the atom’s mass.
Electron Cloud: Surrounds the nucleus and obtains most of the atom’s
space.
Nucleus
Electron Cloud
IDENTIFICATION
1)
2)
Atoms are identified by their atomic number (which is the
amount of protons the atom contains)
An atom’s mass number is the total of protons and
neutrons in the nucleus, each is worth one amu.
(AMU=atomic mass unit).
IONS
1) Ion: An atom with a net electric charge, due to the losing or
gaining of one (or more) electrons.
2) Cations: A positively charged atom, due to losing electrons.
3) Anions: A negatively charged atom, due to gaining electrons,
QUANTUM NUMBERS
1. Quantum numbers describe the location of an electron in the cloud.
2. Each electron is made up of 4 quantum numbers.
3. Every electron has a unique set of quantum numbers, like an id, zipcode, or
finger print.
VALENCE ELECTRONS
1. An electron located in the outer shell (valence shell) that can be shared or
transferred to another atom
2. The typically reside in the highest energy level.
3. When two different atoms interact, the valence electrons determine how an atom
will react in a chemical reaction by coming into contact with other atoms.
ELECTRON CONFIGURATION
1) The dispersal of electrons in atomic/molecular orbitals.
2) Electrons move independently throughout the orbital.
3) From configuration, an atom's reactivity for corrosion can be
determined.
- Aufbau Principle: Electrons that orbit one or more atoms fill the lowest energy levels before filling
the highest.
- Pauli Exclusion Principle: Two electrons in the same atom can NOT have the same set of four
quantum numbers.
- Hund's Rule: Orbitals with the same amount of energy have only one electron before another
orbital can be occupied by a second electron. All electrons in a singly occupied orbital have the
same spin.
1.
2.
3.
ELECTRON CONFIGURATION NOTATION
Easy way to present the electrons.
A subscript represents the amount of electrons in a sub-level.
Ex. Neon, has 10 electrons, so its electron notation, 1s^2 2s^2 2p^6.
ORBITAL NOTATION
1.
Orbitals are represented by arrows with the orbital name underneath.
2.
Half Arrow = One electron
ATOM JOKE!
Atomic
Structur
es
Atomic Structures
Atomic structure is the concept that atoms are centralized with a positively charged nucleus (contained
with protons and neutrons) with negatively charged electrons surrounding it.
Electromagnetic Spectrum
Types of Wavelengths
Longest to Shortest:
Radio waves
Microwaves
Infrared
Optical
Ultraviolet
X-rays
Gamma Rays
Wavelength and Frequency
Wavelength (λ): The distance between two identical points on adjacent waves
Frequency (ν) (nu): The number of waves that pass an indicated point in a given time period.
Wavelength and frequency are inversely related to each other through the speed of light.
c=λν
Calculating Wavelength(l), Frequency(n), and Enegry (E)
E= hc/λ where
C=λV where
V is frequency in hertz, 1/s or s-1
E is Energy in Joules (J)
λ is wavelength
λ is wavelength
C=3.0 x 108 m/s
h=6.626 x 10-34 J s
C= 3.0 x 108 m/s
Example Question
1. A certain microwave has a wavelength of 0.032 meters. Calculate the frequency of this microwave.
Answer: 3.00 * 10^8 / .032 = 9.375 * 10^9
Atomic Structure
Atomic Number: the number of protons in an atom
Atomic Mass: the number of protons and neutrons in an atom
Number of electrons= number of protons
Valence Electrons
Valence electrons are electrons in an atom, mostly in the highest energy level, that are able to
gained, lost, or shared in the formation of a chemical bond.
The number of valence electrons can be indicated from the period of an atom in the periodic table or
its electron configuration.
Maximum number of valence electrons is 8.
Lewis Dot Diagram
Gilbert N. Lewis created the Lewis dot diagram in 1916.
The Lewis Dot Diagram is a diagram in which the valence electrons of an atom are indicated by the use of
dots circling around the symbol.
Example Question:
1) Draw the Lewis dot diagram for carbon:
Electron Configuration
The elements are represented numerically by the number of electrons in their shells and number of shells
Ions
Atoms or groups of atoms with a positive or negative charge.
Cations:
-
positive charge
-
usually metals
- fromed when an electron is taken away from an atom
Anions:
.
-
negative charge
-
usually nonmetals
-
formed when adding an electron to an atom
Bibliography
"A Brief Tutorial on Drawing Lewis Dot Structures." Lewisdots.html. N.p., n.d. Web. 13 Dec. 2016.
<http://web.chem.ucla.edu/~harding/lewisdots.html>.
"A Brief Tutorial on Drawing Lewis Dot Structures." Lewisdots.html. N.p., n.d. Web. 13 Dec. 2016.
<http://web.chem.ucla.edu/~harding/lewisdots.html>.
"The Definition of Valence Electron." Dictionary.com. N.p., n.d. Web. 13 Dec. 2016. <http://www.dictionary.com/browse/valenceelectron>.
"The Electromagnetic Spectrum: Home Page." The Electromagnetic Spectrum: Home Page. N.p., n.d. Web. 13 Dec. 2016.
<http://www.darvill.clara.net/emag/>.
NASA. NASA, n.d. Web. 13 Dec. 2016. <http://imagine.gsfc.nasa.gov/science/toolbox/emspectrum1.html>.
NASA. NASA, n.d. Web. 13 Dec. 2016. <http://science.hq.nasa.gov/kids/imagers/ems/waves4.html>.
N.p., n.d. Web. <http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch8/>.
Bibliography
"Calculations between Wavelength, Frequency and Energy - Probs 1 - 10." Calculations between Wavelength, Frequency and Energy - Probs
1 - 10. N.p., n.d. Web. 13 Dec. 2016.
"Quantum Numbers." Chemistry LibreTexts. N.p., 21 July 2016. Web. 13 Dec. 2016.
"Wavelegnth, Frequency and Energy Calculations." Wavelegnth, Frequency and Energy Calculations. N.p., n.d. Web. 13 Dec. 2016.
Atomic Structure
Zainab, Julian, Sritha 3B
Light and Waves
• Electromagnetic radiation can be seen in terms of energy, wavelength,
or frequency.
• Frequency is measured in hertz
• Wavelength is measured in meters
• Energy is measured in electron volts.
Electromagnetic Spectrum
• The electromagnetic spectrum is the
range of all types of EM radiation.
• Radiation is energy that travels and
spreads out as it goes
• Types of EM radiation that make up
the electromagnetic spectrum
are visible light, microwaves, infrared
light, ultraviolet light, Xrays and gamma-rays.
Quantum Numbers
• Used to describe
the location of an
electron within the
electron cloud.
• Four quantum
numbers make up a
complete set for
each individual
electron.
Lewis Dot Diagram
• A Lewis dot diagram is a method of
writing the chemical symbol of an
element by surrounding it with dots
to indicate the number of valence
electrons.
• Valence electrons are found in an
atom's outer shell and are the ones
involved in chemical reactions.
• Atoms tend to gain, lose or share
electrons in order to acquire a full
set of valence electrons, to satisfy
the octet rule.
Electron Configuration
.
• An atom's electron configuration is
a numeric representation of its
electron orbitals.
• Eliminates lines and arrows from
orbital notation.
• Number of electrons in a sublevel
is shown by adding to the orbital
name.
Valence Electrons
• The valence electrons are the number of electrons an atom must lose or gain
to have a full outer shell, to satisfy the Octet rule.
• Highest occupied energy level
Ions
• Ions are charged atoms that exist when atoms have an unequal number of
protons and electrons.
Atomic Structures
Created by Edward Vasquez and Syed Husaini
Electron Configuration
• Definition
The organization of electrons within the orbitals of an
the way the orbitals are filled out
Aufbau principle: Electrons enter the lower energy orbitals to the
orbitals.
S,P,D, and F are all orbitals. Starting from S the lowest energy orbital
highest energy orbital.
Pauli Exclusion Principal: There are no two electrons that can be the
Hund’s Rule: Orbitals with the same energy is singly occupied before
orbital is twice as occupied.. Electrons in singly occupied orbitals
amount of spin.
Electron Configuration
Sub- shells
# of orbitals
# of Electrons
S
1
2
P
3
6
D
5
10
F
7
14
• Purpose: When atoms contact or interact with each
other, the valence shell (outermost layer) is the part of
each that will interact first. And so an atom is most
reactive when it’s valence shell is not basically filled.
And so an atom is most stable when it’s valence shell is
full. Electron configurations can help predict various
ways elements can react.
Ions
• Definition - An atom/molecule with an electric charge due to the loss or
gain of one or more electrons.
- cation: positively charged atom (lost electrons, metals)
– anion: negatively charged atom (gained electrons, non
metals)
• Each proton carries a +1e charge
• Each electron carries a -1e charge
Ions - Trends
• As you go down a group (periodic table), the number of electrons
in the valence shell stays the same
• As you go across a period (left to right) the number of electrons
in the valence shell of the main group elements increases one by
one.
- This ignores the transition metals as they require different
process, which includes the utilization of roman numerals, to
determine their charge.
Ions – The Formation
• Atoms form ions to satisfy the octet rule
• Chemical compounds form so that each atom (by gaining, losing, or
sharing electrons) has an octet of electrons in its highest occupied
energy level.
- The number of valence electrons will either decrease to “0” or
increase to 8, whichever requires less energy
e.g.
* an atom with 2 valence electrons will lose 2 electrons to
reach 0.
* an atom with 6 valence electrons will gain 2 electrons to
reach 8.
Lewis Dot Structure
• Lewis dot structure is a quick and easy diagram that shows the valence
electrons in an element.
- The nucleus of the element is represented by its symbol.
- The valence electrons are represented by dots placed around the
symbol in pairs.
• This diagram uses the Octet rule
- Octet rule says that atoms like to have full outer shells of only
eight electrons.
- Atoms will lose or gain valence electrons to make their outer
shells full with eight electrons, and they do this by bonding with other
atoms.
Lewis Dot Diagram(s)
Lewis Dot Diagram(s)
Chemical Bond(s)
Valence Electrons
• The electron of an atom, which is found in the valence shell. that can be
transferred to or shared with another atom.
- They can be transferred/shared with another atom
• The number of electron shells with electrons is the same as the period
number.
• To find the Valence electron of an atom
- Identify the group of the atom, based on the group discover the valence
electrons.
Valence Electrons
Exp.
If an element is in group 1A then it has one
Valence electron.
Exp. If an element is in group 1A then it
has two Valence electrons.
Light and Waves
• Lambda (λ) : the actual distance of
wavelenghts
• Frequency (ν): the number of times
a point on a wave passes another
fixed point in one second
• Equations
- c=λν
-E=hν
- E = h c/ λ
- (J) = (J·s)
(m/s)/ λ
Light and waves
• Quantum- least amount of energy that an atom could gain or lose,
a distinct packet of energy.
• Photon- quantum energy that each particle carries
Electromagnetic spectrum
• forms of energy around us.
• Deals with how charged particles have interactions with each other
• Consists of wavelengths and
frequencies.
• Divided into seven different
regions
(radio waves, microwaves, infrared,
visible light, ultraviolet,
X-rays and gamma rays)
Quantum numbers
• Quantum numbers are utilized to determine and discuss the
location of an electron in the electron cloud.
MATTER
MATTER
BY: JOSE MARTINEZ AND SARRAJ AL-HILAL
WHAT IS MATTER…
• Matter is anything that has mass and takes up space
WHAT IS CHEMISTRY?
• The branch of science that deals with the identification of the substances of which
matter is composed; the investigation of their properties and the ways in which
they interact, combine, and change; and use of these processes to form new
substances
PHASES OF MATTER…
SOLID
• It is a definite shape
• Has definite volume
• Solid turns to liquid which shows it is melting
• Solid turns to gas which shows it is sublimation
LIQUID
• It is an indefinite shape, it will take shape of any container
• It has definite volume
• When a liquid turns to a solid, it is freezing
• When a liquid turns to a gas it is boiling then evaporates
GAS
• It is an indefinite shape, it will fill a container
• It has indefinite volume and will fill container
• When a gas turns into a liquid it is condensation
• When a gas turns to solid it is deposition
PLASMA
• It is a high temperature physical state of matter in which atoms loose their
electrons
• Lighting
• Fluorescent bulbs
EXTENSIVE PROPERTIES
A property that changes when the size of the sample changes.
ex of extensive properties: mass,volume,length,and total charge.
INTENSIVE PROPERTIES
An Intensive property does not change when you take away some of
the sample
Physical Property
A characteristic that can be observed or measured without changing the
composition of the sample
-can be described mixtures as well as pure substances
Physical Change
A usually reversible change in the physical properties of a substance, as size or
shape
Ex of physical change: Melting a sugar cube because the substance is still sugar
Chemical Properties
A property or characteristic of a substance that is observed during a
reaction in which the chemical composition or identify of the
substance is changed.
Chemical Change
Any change that results in the formation of new chemical substances. At the
molecular level, chemical change involves making or breaking of bonds between
atoms.
Ex. of chemical change: Iron Rusting
Classification of Matter
Matter can be classified into two basic categories. Matter is either a mixture or it is
a pure substance. We can classify mixtures into two categories-homogenous
mixtures and heterogeneous mixtures. We can also classify pure substances into
two categories- elements and compounds.
Matter
By: Ezra Aguilar and Joseph Kwon
What is chemistry?
Chemistry- The study of the composition,
structure, and properties of matter and the
changes it undergoes.
Phases of matter
Phase Changes
Extensive Properties
Depends on the amount of matter that is present
Mass
Volume
Amount of energy
Intensive Properties
Does not depend an amount of matter that is present
Melting point
Boiling point
Density
Conductivity
Physical Properties
● Characteristic can be observed/measured
○ Using the senses
Examples:
Color
Smell
Freezing point
Boiling point
Physical Change
Change in a substance that does not change its identity
Examples:
phase/state
Breaking apart or lessening quantity
Is always reversible
Chemical Properties
A substance’s ability to undergo change in identity
Creates a new substance
How a substance reacts
Combustibility
Reactivity with other chemicals or substances like water or oxygen
Chemical Change
Change where one or more substances change into a different substance
Examples:
Burning
Chemical reaction
Cooking and baking
Classification of Matter
Matter can be found in two forms
Pure Substance- has a fixed composition
All parts have the same characteristics
Can be either elements or compounds
Mixture- physical blend of two or more kinds of matter where each have their own properties
Homogeneous-uniform composition (kool-aid)
Heterogeneous-not uniform composition
Solutions are mixtures
M AT T E R & C H A N G E
BY: SARAH, JOSEPH, & GABBY
DEFINITION REVIEW
Chemistry : is the study of the structure, composition, and properties of
matter and the changes it undergoes
Matter: anything that has mass and takes up space
Mass: a measure of the amount of matter
PHASES OF MATTER
SOLID
definite shape & volume
solid ⟶ liquid
– melting
 solid ⟶ gas
– sublimation
PHASES OF MATTER
LIQUID
indefinite shape
– takes shape of container
 definite volume
 liquid ⟶ solid
– freezing
 liquid ⟶ gas
– boiling
PHASES OF MATTER
GAS
indefinite shape
– fills container
indefinite volume
– fills container
 gas ⟶ liquid
– condensation
 gas ⟶ solid
– deposition
PHASES OF MATTER
PLASMA
high temperature physical
state of matter in which
atoms lose their electrons
-lightning , fluorescent bulbs
PROPERTIES & CHANGES IN MATTER
EXTENSIVE PROPERTIES
depend on amount of matter that is present
– mass
– volume
– amount of energy
INTENSIVE PROPERTIES
do not depend on amount of matter that is
present
– melting point
– boiling point
– density
– conductivity
PROPERTIES & CHANGES IN MATTER
PHYSICAL PROPERTIES
• characteristic that can be observed or measured without changing the
identity of the substance
• observed using the senses
PHYSICAL CHANGES
• a change in a substance that does not involve a change in the identity of
• the substance
– changes of state / phase changes
– breaking apart matter, lessening quantity
PROPERTIES & CHANGES IN MATTER
CHEMICAL PROPERTIES
 relates to a substance’s ability to undergo changes that transform it into a different substance
 how the substance reacts
– combustibility
– reactivity with water, oxygen, etcetera
CHEMICAL CHANGES
 a change in which one or more substances are converted into different substances
– burning a substance
– chemical reactions
– cooking / baking
PERIODIC TABLE
PERIODIC TABLE
BY: LIANN SEBASTIAN AND ALIYA SAYANI
HISTORY AND DEVELOPMENT
• Simple list made in 1789 by Antoine Lavoisier
• 33 elements
• First table published in 1869 by Dmitri Mendeleeve
• Included more than 60 elements
• Ordered by increasing atomic weight
• Modern periodic table created in 1945
• Included 118 elements
ORGANIZATION AND LAYOUT
• Modern periodic table:
• Arranged by atomic mass across rows
• Arranged by physical and chemical properties
• Horizontal rows= periods, numbered 1-7
• Columns= groups, numbered 1-18
• Can be split into A and B blocks
•
•
•
Metals
TYPES OF ELEMENTS
•
Left side of the periodic table
•
Solid at room temperature
•
Silver-gray luster
•
Malleable and ductile
Nonmetals
•
Located on right side
•
Solid or gas at room temperature (except bromine)
•
Poor conductor of heat and electricity
•
Brittle, no common color
Metalloids
•
Has characteristics of both metals and nonmetals
•
Between metals and nonmetals
•
Semi-conductors
•
Boron, silicon, germanium, arsenic, antimony, tellurium
CHEMICAL FAMILIES
•
•
•
•
Alkali Metals
•
Group 1 or 1A
•
Most reactive, especially with nonmetals
Alkaline Earth Metals
•
Group 2 or 2A
•
Reactive, especially with water
Transition Metals
•
Group 3 or 3A
•
Less reactive than alkali/alkaline earth metals
Rare Earth Metals
•
Lanthanide series- period 6 below table
•
Actinide series- period 7 below table
CHEMICAL FAMILIES CONT.
•
Other Metals
•
Group 13- aluminum, gallium, indium, thallium
•
Group 14- tin, lead
•
Group 15- bismuth
•
Group 16-polonium
•
Halogens
•
Group 17 or 7A
•
Most reactive nonmetals, especially with metals
•
Noble gases
•
Group 18 or 8A
•
Unreactive, colorless, odorless
PERIODIC TRENDS
• Atomic Radius- measure of the size of an elements atoms
• Across a period: decreases
• Down a group: increases
• Ionic Radius- distance between the nucleus and the electron in the outermost shell of an ion
• Metals: always cations
• Nonmetals: always anions
• Ionization Energy- amount of energy needed to remove the valence electron to form a cation
• Across a period: increases
• Down a group: decreases
• Electron Affinity- a neutral atom’s likelihood of gaining an electron
• Across a period: decreases
• Down a group: increases
• Electronegativity- measure of the ability of an atom to attract electrons
• Across a period: increases
• Down a group: decreases
NOBLE GAS NOTATION
• Shorthand for electron configuration
• Beginning of electron configuration is represented by the symbol of the noble gas in the period above, in
brackets
• Remaining portion of electron configuration is normal for desired element
•
•
Ex:
•
Ne: 1s22s22p6
•
Al: 1s22s22p63s23p1
Noble Gas Notation= Al:[Ne] 3s23p1
The Periodic Table
By: Jacob Knackstedt and Austin McCrae
Antoine Lavoisier
First official periodic table published.
Antoine described elements as a material which could not be
decomposed into a simpler form.
Contained 33 elements out of our know 118.
It was made in to columns.
It was organized by metals, gases, earths, and nonmetals
Dimitri Mendeleev
•
•
•
•
•
Second published periodic table.
It was published in the year 1869.
Contained 60 plus elements it looked like a jigsaw puzzle.
Elements were ordered by their atomic weight.
Set in columns.
Modern periodic table
•
•
•
•
•
•
Created in 1945
Total of 118 known elements
Rows are arranged by atomic mass
Columns arranged by elements properties
Horizontal rows 1-7 called periods.
Columns called groups organized 1-18.
Separated as a and b block.
Element types: Metals
•
•
•
•
•
•
Good conductors of both heat and
electricity.
Most are solid at a room temp.
Silver and gray-white in appearance.
Malleable
Ductile
On the left side of the periodic table
Alkali metals
•
•
•
•
•
•
1a elements
Silver
Relatively soft
Most reactive out of all elements
Not found in nature as a free element
Only 1 valence electron
Alkaline earth metals
•
•
•
•
2a elements
Silver colored
Reactive, but nit as reactive as Alkali
Stronger than Alkali metals.
Transition metals
•
•
•
•
B block elements
High luster value
Not as reactive as Alkali or Alkaline
Mercury exists as liquid in room
temp.
Rare earth metals
•
•
•
•
•
•
2 types
Lanthanides
Period 6 and below
Actinides
Period 7 and below
All are radioactive elements
Other metals
•
•
•
•
3a: aluminum, gallium, thallium, indium
4a: lead, tin
5a:bismuth
6a:polonium
Nonmetals
•
•
•
•
Not a good conductor of electricity
and heat.
Can be gases or solid at room
temperature.
Very delicate
Don’t share a common color
between all of them
Halogens
•
•
•
•
•
•
•
7a
Are highly reactive with metals
Most reactive out of the nonmetals
Liquids in this group are Bromine.
Gases in this group are Chlorine and
Fluorine.
Solids in the group are Astatine and
Iodine.
Are toxic elements.
Noble gases
•
•
•
•
•
8a
Gases
Most of the time are unreactive
Odorless and colorless
Maximum amount od electrons in
valence shell
Metalloids
•
•
•
•
They have some of the attributes of
both metals and nonmetals
Found in between the metals and
the nonmetals
Semiconductors
Less malleable than the metal
elements
Periodic trends
Noble gas notation
•
When doing electron configuration notation write the noble gas from
the period above then continue writing notation.
•
Phosphorus:
•
Chlorine:
[Ne]3s 2 3p 3
[Ne]3s 2 3p 5
Bibliography
•
•
•
•
•
•
•
http://www.rsc.org/education/teachers/resources/periodictable/pre16/develop/mendeleev.htm
http://www.chemicalelements.com/groups/alkali.html
http://www.chemicalelements.com/groups/rareearth.html
http://www.chemicalelements.com/groups/nonmetals.html
http://www.chemicalelements.com/groups/noblegases.html
file:///C:/Users/Jacob%20Knackstedt/Downloads/Periodic_Table_of_the_Elements.pdf
Notes
Periodic Table
By: Puja Nayak & Rawan Layth
HISTORY AND DEVELOPMENT
• Dmitiri Mendeleev published the first periodic table in
1869
•
•
•
•
•
More than 60 elements
Rotated
Order by increasing atomic weight
New columns with repeating characteristics
Modern table made in 1945
Development
•
•
•
•
•
118 elements
Arranged by atomic mass across rows
Arranged by physical and chemical qualities in columns
Horizontal rows (periods) 1-7
Vertical columns (families) 1-18 A block and B block
About the Table
TYPES OF ELEMENTS
• an element that is a good conductor of heat and
electricity
• left “side” of the periodic table
• solid at room temperature
• most have silver or gray-white luster
• malleable: can be hammered or rolled into thin sheets
• ductile: can be drawn into a thin wire
METALS
▪ group 1 or 1A
▪ silver in appearance
▪ soft enough to cut with a knife
▪ react vigorously with nonmetals
ALKALI METALS
▪ group 2 or 2A
▪ harder, denser, stronger than alkali metals
▪ silver colored
▪ react with water
▪ reactive metals
ALKALI EARTH METALS
▪ group 3 – group 12 elements
– B block elements
▪ good conductors
▪ high luster
▪ less reactive
▪ mercury exists as a liquid at room temperature
TRANSITION METALS
▪ lanthanide series
– period 6 below table
– most are found in the Earth’s crust
▪ actinide series
– period 7 below table
▪ these elements sometimes categorized as transition
metals
RARE EARTH METALS
▪ from group 13 or 3A elements
– aluminum, gallium, indium, thallium
▪ from group 14 or 4A elements
– tin, lead
▪ from group 15 or 5A elements
– bismuth
▪ from group 16 or 6A elements
– polonium
OTHER METALS
NONMETALS
▪ an element that is a poor conductor of heat and electricity
▪ right “side” of the periodic table
▪ either solid or gas at room temperature exception of
bromine (liquid at room temperature)
▪ brittle
▪ no common color
NONMETALS
▪ group 17 or 7A elements
▪ most reactive nonmetals
▪ toxic
HALOGENS
▪ group 18 or 8A elements
▪ gases at room temperature
▪ odorless and colorless
▪ generally unreactive
NOBLE GASES
▪ an element that has some characteristics of metals and
some characteristics of nonmetals
▪ found between the metals and nonmetals (“stair step”)
▪ solid at room temperature
▪ less malleable than metals, not as brittle as nonmetals
▪ semiconductors
▪ boron, silicon, germanium, arsenic, antimony, tellurium
METALLOID
Periodic Trends
• Shorter way to write electron configuration
• Beginning of electron configuration is represented by a
symbol of the noble gas in the period above the wanted
element.
• The rest is written as normal
• Example: Al: [Ne]3s 23p 1
NOBLE GAS NOTATION
ORBITAL
• # of electrons (in a-block) can be determined by its
location on the periodic table
• Al found in group 3A= 3 valence electrons
VALENCE ELECTRONS
• Can be determined
from valence
electrons
• Add or remove to
satisfy octet rule
Valence
Electrons
Electrons
gained/lost
Ion charge
1
1 lost
1+
2
2 lost
2+
3
3 lost
3+
4
4 lost/4 gained
4+ 4-
5
3 gained
3-
6
2 gained
2-
7
1 gained
1-
8
none
none
IONIC CHARGE
• Half distance of nuclei of atoms bonded together
Down a group
Across a period
Increasing
positive charge
of nucleus
ATOMIC RADIUS
Caused when
electrons go to
higher energy
levels further
away from the
nucleus.
Anion
Cation
Removing valence
electrons=smaller electron
cloud
IONIC RADIUS
Electron cloud
spread out because
electrons not drawn
towards nucleus
• Energy needed to remove an electron from the atom
(neutral)
• Each ionization level increases
Across a period
Down a group
Metals lose
electrons easily,
nonmetals do not
IONIZATION ENERGY
Electrons are
further from
nucleus so easier
to remove
• Energy change when an electron is gained (Neutral
atom)
Across a period
Energy released by
positive value NOT
FAVORABLE
Down a group
Energy released
by negative value
FAVORABLE
Nonmetals easily gain
electrons while metals
don’t
ELECTRON AFFINITY
Hard to add electrons
to large atoms
• The ability for an atom to attract electrons (chemical
compounds)
Down a group
Across a period
Larger radius so
electrons feel less pull
ELECTRONEGATIVITY
Electrons
attracted to
nonmetals
PERIODIC TABLE
BY: MELLENY , QAYLAH, AHAN
HISTORY AND DEVELOPMENT
• ANTOINE LAVOISIER- made the first list of 33
elements
• JOHANN W. DOBEREINER- found 3 elemts
with similar characteristics
- Lithium, sodium, and potassium made one triad
• JOHN NEWLANDS- arranged all elments
known at the time according to incressing
amotic mass
• LOTHAR MEYER- arranged 56 elemts into one
table according to their characteristics
• DMITRI MENDELEEVGrouped elements by similar chemical
properties and arranged elements by
incressing mass
• HENRY MOSELY- arranged elements by
increasing proton number
ORGNIZATION AND LAYOUT
• 18 vertical columns called families or groups
• 7 horizontal rows called periods or series
• Split into A nad B block
• Arranged by atomic mass across rows
TYPES OF ELEMENTS
•
Metals – good conducter of heat and electricity
left side of the periodic table
solid at room temp.
•
Non Metals – poor conductors of heat and electricity
On the right side of the periodic table
Solids or gasses at room temp.
•
Metalloids- some characteristics of nonmetals and others of metals
Found between metals and nonmetals (stair step)
Solid at room temp.
CHEMICAL FAMILIES
• ALKALI METALS ( Group 1)
Most reactive metals
• ALKALINE EARTH METALS (Group 2)
Less reactive than group one
• TRANSITION METALS ( Group 3-12)
Good conductors of electricity
• Halogens( Group 17)
Most reactive nonmetals
• Noble Gases (Group 18)
Rarely react with other elements because they have a full valence shell
PERIODIC TRENDS
• Atomic Radius- size of the atom
Increases down and left
• Ionic Radius- Adding an electron makes atom bigger
Taking away an electron makes it smaller
• Ionization Energy – energy requred to remove an electron from the atom
increases up and right
• Electron Affinity- how much an atom wants to gain
Increases up and right
• Electronegativity- ability of atom to hold electrons tightly
Increases up and right
NOBEL GAS NOTATION
• Short cut method for electron arrangement
• EXAMPLE::
• Na 1s^2 2s^2 2p^6 3s^1
• -Use Neon
• - Use neon’s configuration as [Ne]
• - Use in Na: [Ne] 3s^1
COVALENT
COMPOUND
STRUCTURES
Covalent Compound
Structures
LEWIS STRUCTURES, VSEPR THEORY, AND INTERMOLECULAR FORCES
CAROLINA ROMERO MADELEINE ROOK
Lewis Structure

Formula which represents a nucleus and inner
shell electrons using covalent bonds to show
unshared electrons
 Covalent
Bond: The sharing of an electron pair
between 2 atoms that result in a chemical bond.
 Lone
bond.
Pair: A pair (2) electrons not involved in a
Lewis Structures
Step 1
Ex: PCl3
Count the number of valence electrons in each of the elements in the
compound
-1
Phosphorus atom: 5 valence electrons 1(5)=5 é
-3
Chlorine atoms: 7 valence electrons 3(7)=21 é
Add the numbers of totaled electrons
- 21é+5é=26é
Lewis Structures
Step 2
Put the least electronegative element in the center, then all the other
elements surrounding it. Connect the outer elements to the center
element.
Ex: PCl3
Subtract original é from the
number on bonds used
Cl
P
Cl
Cl
each bond is 2 electrons
26é-6é=20é
Lewis Structures
Step 3
Fill all the outside elements with the rest of the electrons.
*One bond is 2é
20é
Cl
P
Cl
Cl
Lewis Structures
Exceptions
More Than:
-Some atoms have more
than 8
-Non-metals in period 3
or below
-Add a lone pair to central
atom if electrons remain
Ex: NO2-1 O
14é- 2é
N
O
Less Than:
-less than 8
-duet (2): Hydrogen
-quartet (4): Beryllium
-sextet (6): Boron
Ex: CH3Cl
6é
H
Cl C
H
H
VSEPR Theory
Stands for Valence Shell Electron Pair Repulsion
3 dimensional arrangements of all domains (EDG)
3 dimensional arrangement of bonded pairs (MDG)
Angle between bonded pair domains (Bond Angle)
Intermolecular Forces
forces between molecules (covalently bonded compound)
London Dispersion Model (weakest)
- found in every molecule (constant motion of electrons)
Intermolecular Forces
Dipole-Dipole
-Created by equal or opposite charges with short distances of separation
+ I
- Cl
Cl I+
Intermolecular Forces
Hydrogen Bonding (strongest)
-Highly electronegative atom bonded with Hydrogen.
-Creates a very strong dipole
+
H
H
-
H+
O
H
+
O
-
MOLECULAR STRUCTURES
BY AKUL, JAYA, JOSH
VSEPR THEORY
Valence
Shell
Electron
Pair
Repulsion
VSEPR theory is a model used to predict the
geometry of specific molecules by the
number of electron pairs around the central
atom
The basic idea of VSEPR theory is that the
valance electrons will repel each other and
go into a specific geometric formation that
reduces the amount of repulsion that they
undergo
VSEPR THEORY
Electron
Domain
Geometry
The number of lone pairs and bonded
pairs around a particular atom in a
molecule.
Molecular
Domain
The number of lone pairs around a
particular atom in a molecule.
Geometry
INTERMOLECULAR FORCES
Forces of attraction between
molecules
(covalently bonded
compounds)
 London-dispersion forces
 Dipole-Dipole
 Hydrogen bonding
LONDON-DISPERSION FORCES
Weakest of the three forces
-Cause by constant motion of electrons
Everything has Londondispersion Forces
DIPOLE DIPOLE
 Attracted forces between the positive end and the negative end of a
polar molecule
 Polar and Non Polar
NO DIPOLE
 Carbon - Hydrogen bond
 The same element
HYDROGEN BONDING
Strongest of the three forces
-results from hydrogens bonded to a very
electronegative atom
-creates strong dipole
-no permanent attraction
There are three ways a
hydrogen atom can be
formed
-Have fon
Fluorine
Oxygen
Nitrogen
QUIZ
1. What does it mean to have FON?
2. Name two intermolecular forces?
3. What is VSEPR theory?
CHEMICAL FORMULAS
AND NAMES
CHEMICAL
FORMULAS AND
NAMES
SOHAN GADE & MARIJ AHMED
3A CHEMISTRY
OVERVIEW
• Ionic Bonds / compounds
• Covalent bonds / compounds
• Polyatomic ions
• Nomenclature
COVALENT BONDS
• Bonds in which electrons are shared between atoms of different elements
• These are formed with two non-metals
• Example: H2O, CO2
• The least electronegative element is named first
• To name the compound you take the first element, look at how many atoms there
are of that compound, and put an appropriate prefix (e.g. tetra for four), if it is
one then there is no prefix
• Take the second element, and do the same process and add the prefix, and then
add the suffix, -ide
IONIC BONDS
• Ionic Bonds are formed when two elements either give up or gain electrons.
• Metals will form cations(lose electrons) and nonmetals will form anions
(gain electrons).
• This differs from covalent bonding, because the atoms lose or gain the
electrons, rather than sharing them
• The atoms involved will usually try to follow the octet rule (having a
full/stable eight electrons in the outer shell)
• Examples : NaCl (Table Salt), 𝐶𝑎𝐶𝑙2 (Calcium Chloride)
POLYATOMIC IONS
• Binary- Has only two elements, ex. NaCl
• Polyatomic- Multiple atoms with collective charge (e.g. Ammonium, Hydroxide,
Cyanide)
• Base unit has suffix of –ate, -ite has one less Oxygen, -per has one more, and hypohas two less than the base
• For some, it is the Hydrogen that changes, -bi means one more Hydrogen, and –di
means two more Hydrogen
• If there are multiple of the same polyatomic ion, the formula of the ion is put in
parentheses and a subscript indicating the number used is placed outside the
parentheses
WITH TRANSITION METALS
• Find which charge of the metal that needs to be used to make the compound
neutral
• Take the element name (cation) place it first, and then take the charge and
rewrite it a Roman numeral, then place the other element/polyatomic ion
name
• Example: Copper (II) Sulfate
NAMING IONIC COMPOUNDS
• If it is a nonmetal and metal that are bonding, the ending of the second
element is changed to –ide, the cation comes first.
• With transition metals, the transition metal cation’s Roman numeral is used
after the element name to indicate the charge, followed by the second
element with the suffix –ide, ex. MgO (Magnesium Oxide)
• With polyatomic ions, the cation is placed first and then polyatomic ion
name is placed (ex. Copper (II) Hydroxide)
Chemical
Formulas and
Names
By:
Emanuel Encarnacion
&
Michael Rudolph
Ionic Bonds/Compounds
Ionic: formed when neutral atom gains or loses an electron.
Cation: formed when a neutral atom loses an electron and becomes positively charged.
Anion: formed when a neutral atom gains an electron and becomes negatively charged.
Ionic Compound: a compound formed from a cation and an anion.
Typically metal + nonmetal reaction
Ex. Potassium Nitride
K^1+ N^3- K^1+K^1+K^1=K^3
K3N3 balanced
Covalent Bonds: Bonds between atoms when
they share electrons
Covalent
Bonds/
Compounds
Covalent Compounds: Formed when nonmetal
atoms share valence electrons
Polyatomic Ions
Polyatomic ions: a group of atoms that have a charge.
Polyatomic ions usually have a charge because the collection of atoms has either gained an extra
electron or else it has lost an electron.
Nomenclature: Choosing names for things,
especially in science
Ionic Compounds: Metals with only one charge
are written as their basic name; metals with
multiple charges are written with roman
numerals with specific charge
Nomenclature
Fe₃O₂ - iron(II) oxide
Covalent Compounds: First element has
element name with prefix of numer of atoms
present; second element has prefix of number
of atoms present and an ending of -ide
CHEMICAL FORMULAS
AND NAMES
BY: Kabir Bhakta & Kaushik Jampala 2B Chemistry
POLYATOMIC IONS
What is a polyatomic ion?
• A polyatomic ion is a charged group of covalently bonded atoms.
Therefore, they are either cations or anions created by an imbalance
between the amount of electrons and protons
• The prefix poly- means many, so a polyatomic ion is an ion with more than
one atom.
• Note that these are named differently than monoatomic ions which contain
only one atom.
POLYATOMIC IONS
Naming Polyatomic Ions
• The following are rules for identifying different polyatomic ions based on the
number of oxygens they have. They are identified using special prefixes and
suffixes.
•
•
•
•
The suffix-ite is added to the ion with one less oxygen atom.
The suffix –ate is added to the most common ion.
The prefix hypo- is placed in front of the ion with one less oxygen than the –ite ion
The prefix per- is placed in front of the ion with one more oxygen than the –ate
ion.
NAMING POLYATOMIC ION
COMPOUNDS
• The following are the rules for naming ionic compounds containing
polyatomic ions:
1. The name of the cation is written first.
2. This is followed by the name of the anion.
3. However, when there is multiple of a polyatomic ion present, parenthesis are
placed around the entire ion.
IONIC BONDS
• Ionic bonding is a type of chemical bond that involves two, oppositely
charged ions
• Ionic bonding occurs when one atom takes an electron from another atom
hence the phrase “taken, not shared”
• The metal loses electrons to become positively charged and the nonmetal
receives those electrons to become a negatively charged anion
IONIC COMPOUNDS
When naming ionic compounds follow these rules:
• An ionic compound is named first by its cation and then by its anion
• The cation keeps the same name as its element and the anion name’s
ending is replaced with –ide
• Ex: NaF = sodium fluoride
• If the cation is a transition metal, the charge is written in Roman numerals in
parentheses directly after the element name
• Ex: Cu(NO3)2 = copper (II) nitrate
• If either the cation or anion is a polyatomic ion, the polyatomic ion name is
used in the overall compound, the polyatomic ion always stays the same
COVALENT BONDS
• Covalent bonds form between two nonmetals
• Unlike ionic bonding, covalent bonding “shares” electron pairs between
atoms
• Covalent bonding is generally weaker than ionic bonding
• Ex: H2O = Oxygen needs two more valence electrons to fill its outer shell, and
hydrogen needs to get rid of one electron to satisfy the octet rule so both
hydrogen atoms will share their extra electron with oxygen creating a
balanced compound
COVALENT COMPOUNDS
When naming covalent compounds follow the rules below:
• The first element when naming will be the less electronegative element
• A prefix will be added to the first element unless its –mono
• The second element in the compound will be named using the appropriate
prefix and changing the ending to –ide
• Ex: SO2 = sulfur dioxide
Chemical Formulas and
Names
Ashton Mathew
Noah Kim
Ionic Bonds/Compounds
Ionic compound formula = Cation + Anion
Cation - positively charged ion
Anion - negatively charged ion
Cation naming
1. Group A – element name
2. Transition metals- element
name and the charge (with
Roman Numerals)
3. Polyatomic – ion name
Anion naming
1. Nonmetal - element name +
“-ide”
2. Polyatomic – ion name
Covalent Bonds/Compounds
Covalent Bonding = Results from the sharing of
electron pairs between two atoms
2 elements are used to form a covalent bond
First element + Second element = covalent
compound
First Element
1. Write the element name
2. Add a prefix if there are more
than one of the element (never
use mono)
Prefixes
1-mono
2-di
3-tri
4-tetra
5-penta
6-hexa
7-hepta
8-octa
9-nona
10-deca
Second Element
1. Write the element name with the
ending “-ide”
2. Add a prefix in the front if there
are more than one of the element
(mono can be used)
Polyatomic Ions
Polyatomic Ion: Molecular ion that is a
charged chemical species composed of
two or more atoms covalently bonded or
of a metal complex that can be considered
to be acting as a single unit.
Prefixes
“ate”: original ion
“ite”: subtract one oxygen
“per”: add one oxygen
“hypo”: adds two oxygen or subtracts two
oxygen based on ending “ite” or “ate”
List
Acetate: CH3O2(-)
Ammonium:NH4(+)
Bromate:BrO3(-)
Carbonate:CO3(2-)
Chlorate:CIO3(-)
Chromate:CrO4(2-)
Cyanide:CN(-)
Dichromate:Cr2O7(2-)
Iodate: IO3(-)
Hydroxide: OH(-)
Nitrate:NO3(-)
Oxalate:C2O4(2-)
Phosphate:PO4(3-)
Sulfate:SO4(2-)
MOLES AND CHEMICAL
FORMULAS
MOLES AND CHEMICAL
FORMULAS
By: Aftab Shaik, Aaron Purewal, Madar Marar
Mole Concept
•
A mole is a unit of measurement like a dozen is a unit of measurement that means
12 of a certain thing
•
In simple terms, it is the amount of substance that contains the same amount of
particles in 12 grams of carbon-12
•
According to the National Institute of Standards and Technology (the NIST) the
formal definition of 1 mole is a:
•
The mole is the amount of a substance of a system which contains as many elementary
entities as there are atoms in 0.012 of carbon-12; its symbol is "mol." When the mole is used,
the elementary entities must be specified and may be atoms, molecules, ions, electrons, other
particles, or specified groups of such particles.
Exactly how big is this number?
•
Avogadro’s number: the amount of particles in exactly one mole of a pure
substance
•
This number is represented as 6.022*10^23
•
This number is used as a conversion factor to convert between particles/atoms and
moles for the same chemical
•
Moles of a substance to atoms of that substance: multiply the mole amount by
Avogadro’s number. You will end up with atoms of that substance.
•
Atoms of a substance to moles of that substance: divide the atom amount by
Avogadro’s number. You will end up with moles of that substance.
Molar Mass
•
This is as the mass of 1 mol of the substance
•
•
Measured in grams per mol (g/mol)
This measurement can be found by taking the avg. atomic mass of all the
elements in the substance
•
You have to take the avg. atomic mass of each individual element and add all of these
together and you will result in the molar mass of the substance
Molar Mass as a Conversion Factor
•
You can use the molar mass of a substance to convert between moles and grams
of that substance
•
Moles of a substance to grams of that substance: Can be found by taking the mol
amount and multiplying the molar mass of that substance and you will end up with
grams of that substance
•
Grams of a substance to moles of that substance: Can be found by taking the
grams amount and dividing this amount by the molar mass of the substance and
you will end up with moles of that substance
Percent Composition
•
The percent composition is the percent by mass of each element in a compound
•
This can be found by taking the mass of a specific element and dividing that by the
total mass of the compound
•
•
This will result in a number less than 1 in decimal form
We then take this decimal and multiply it by 100% to get the percent composition
Empirical Formula
•
The empirical formula of a compound is the smallest whole number ratio between
a compound
•
•
Ionic compound’s ratios must always be reduced
The steps in order to find this are as follows:
1.
2.
3.
4.
5.
You will be given a measurement, either grams or percent composition of the
compound
Find the grams of each part of the compound
Divide this measurement by the molar mass of each individual chemical
This new measurement will be in moles and divide these numbers by the smallest
number of moles (out of all the individual chemicals) in the result
This will end with a small ratio that can be made into a formula
Molecular Formula
•
The molecular formula of a compound is the ratio between elements in one
molecule of compound
•
In order to find this, we must first find the empirical formula of the compound
•
Then, we will most likely be given the molar mass of the compound
•
Using the molar mass of the empirical formula, we must divide the given molar
mass by the molar mass we found with the empirical formula and we will find a
conversion factor
•
We take this number and multiply each of the element amounts of the empirical
formula and you will end up with the molecular formula
MOLES AND CHEMICAL
FORMULAS
BY: NEEL AND MOSTAFA
MOLE CONCEPT
• MOLE - AMOUNT OF SUBSTANCE THAT CONTAINS AS MANY PARTICLES AS THERE ARE ARE
ATOMS IN EXACTLY 12 GRAMS OF CARBON-12
• AVAGADRO’S NUMBER – NUMBER OF PARTICLES IN ONE MOLE OF A SUBSTANCE
• AVAGADRO’S CONSTANT = 6.022 X 10^23
PERCENT COMPOSITION
• MULTIPLY EACH AMOUNT OF ELEMENTS BY THEIR OWN MOLAR MASS
• DIVIDE EACH BY THE TOTAL MOLAR MASS OF THE COMPOUND
• MULTIPLY 100
MOLAR MASS
• MOLAR MASS = ATOMIC MASS OF AN ELEMENT
• MOLAR MASS OF COMPOUND = SUM OF ALL MOLAR MASS IN THE COMPOUND
EMPIRICAL FORMULA
• ASSUME 100 GRAMS OF ALL ELEMENTS
• DIVIDE BY MOLAR MASS OF THAT ELEMENT
• DIVIDE BY SMALLER VALUE
• ROUND ALL NUMBERS
• THESE NUMBERS WILL BE THE SUBSCRIPT OF THE ELEMENTS
MOLECULAR FORMULA
• DIVIDE GIVEN AMOUNT OF GRAMS OF COMPOUND BY MOLAR MASS OF COMPOUND
• ROUND THE PRODUCT
• MULTIPLY PRODUCT BY EACH SUBSCRIPT IN THE EMPIRICAL FORMULA
MOLES &
CHEMICAL
FORMULAS
Mole
Concept
• A unit of a substance that is equivalent to the
amount of atoms as there are in 12 grams of
carbon-12
– Represented by Avogadro’s Number, or NA
– 6.022*1023
• Avogadro’s number can be used as a conversion
factor when calculating the number of particles or
atoms within an amount of a substance
• 1 mole of a substance=6.022*1023 atoms of that
substance
• Atoms to moles: divide by Avogadro’s number
• Moles to atoms: multiply by Avogadro’s number
Mole
Concept
(cont.)
• Example: Determine the number of moles
in 6.755 atoms of silicon.
• Calculation: (6.755 atoms silicon)*(1 mol
silicon/6.022*1023 atoms silicon)=
1.122*10-23 mol silicon
• Example: Determine the number of atoms
in 0.00384 moles of francium.
• Calculation: (o.oo384 mol
francium)*(6.022*1023 atoms francium/1
mol francium)=2.05 atoms francium
Your Turn!
• Determine the number of moles in 3.95
atoms of potassium
– 6.56*10-24 mol K
• Determine the number of atoms in 2.56
moles of nitrogen
– 1.54*1024
Molar Mass
• The mass of one mole of a substance
• Measured in units of grams per mole, or g/mol
• Molar mass of an element is its average atomic mass
– Ex: magnesium’s molar mass is 22.990 g/mol
• The molar mass of a compound is the combined
average atomic masses of each of the elements
– Ex: LiF: 6.941+18.998= 25.939 g/mol
• You try! Determine the molar masses of the following
compounds:
– NaCl:
• 58.443 g/mol
– H2O:
• 18.015 g/mol
– MgCl2:
• 95.211 g/mol
Molar Mass
as a
Conversion
Factor
• Applicable when performing dimensional analysis
• Allows conversions between grams and moles of
the same element
• Grams to moles: divide by molar mass
• Moles to grams: multiply by molar mass
• Example: determine how many grams are present
in 2.14 moles of carbon dioxide
• Calculation: (2.14 mol CO2)*(44.009 g CO2/1 mol
CO2) =94.2 g CO2
• Example: determine how many moles are present
in 96.74 grams of potassium chloride
• Calculation: (96.74 g KCl)*(1 mol KCl/74.551g
KCl)=1.298 mol KCl
You Try!
• Determine the number of moles in
16.54 grams of H2O
– 0.9181 mol H2O
• Determine the number of grams in 5.20
moles of calcium chloride
– 577 g CaCL2
Percent
Composition
• Percentage by mass of each element in a compound
• Divide mass of one element by the compound’s molar
mass and multiply by 100
• Example: FrF
• (223/241.998)*100= 92.1% Fr
– Subtract 92.1 from 100 to determine percent
composition of fluoride
– 7.9%
• You try! Find the percent compositions of each of the
elements in CH2O
– C: 40.001%
– H: 6.714%
– O: 53.284%
• Percentages should add up to 100
Empirical
&
Molecular
Formulas
• Empirical Formula: The smallest whole
number ratio between elements in a
compound
– NaCl, H2O are empirical formulas
– Ionic compounds are always written in empirical
formulas
• Molecular formula: ratio between elements as
they actually exist in one molecule of a
covalent compound
– C8H18 is a molecular formula
• Empirical formulas can be derived from molecular
formulas
– C4H9 is the empirical formula for C8H18
Finding
Empirical
Formulas
from
Percent
Compositio
ns
Determine the empirical formula of a compound that
contains 52.11% carbon, 13.14% hydrogen, and 34.75%
oxygen.
1.
Assume there are 100 grams of the compound.
– 52.11g C, 13.14g H, 34.75g O
2.
Divide by respective molar masses to determine number
of moles of each element
– 52.11g C/12.011= 4.33852 mol C
– 13.14g H/1.008= 13.03571 mol H
– 34.75g O/15.999=2.17201 mol O
3.
Divide these values by the smallest value to determine a
whole number ratio. Round.
– 4.33852 mol C/2.17201=~2 mol C
– 13.03571 mol H/2.17201=~6 mol H
– 2.17201 mol O/2.17201= 2 mol O
Finding
Empirical
Formulas
from
Percent
Compositio
ns
4.
Write out empirical formula using the determined
number of moles of each element.
– C2H6O2
• Molecular formulas require a given molar mass to be
determined.
• Given: 186.204 g/mol is the molar mass. Find the
molecular formula of C2H6O2
1.
First, find the empirical formula mass
– 2(12.011)+6(1.008)+2(15.999)=62.068 g/mol
2.
Then , divide the given molar mass by the empirical
formula mass.
– 186.204/62.068= 3
3.
Multiply the subscripts by this value to determine the
molecular formula.
– C6H18O6
Your turn!
1. Determine the empirical formula of a
compound with 63.50% silver, 8.25%
nitrogen, and 28.25% oxygen.
– AgNO3
2. The given molar mass is 339.78 g/mol.
Find the molecular formula.
– Ag2N2O6
Moles and
Chemical
Formulas
By: Ihsan and Amirtha
Mole Concept
– Mole: The amount of a substance that contains as many particles as there are
atoms in 12 grams of carbon-12
– Avagadro’s Number: Used as a conversion factor to connect between particles
and goals for the same chemical
– 6.022 x 1023 𝑚𝑜𝑙1−
– Converting from moles to atoms: multiply the amount of moles by Avagadro’s
number, 6.022 x 1023 𝑚𝑜𝑙1−
– Converting from atoms to moles: divide the amount of atoms by Avagadro’s
number, 6.022 x 1023 𝑚𝑜𝑙1−
Examples
1) 8.08 x 1023 atoms Kr
8.08 x 1023 atoms Kr
? mol Kr
2) 1.04 mol Au
? atoms Au
1.04 mol Au
/
x
6.022 x 1023 atoms Kr
6.022 x 1023 atoms Au
=
1.34 mol Kr
=
6.26 x 1023 atoms Au
Molar Mass
– Molar Mass: The mass of one mole of a substance
– Units of grams per mole (g/mol)
– Molar Mass of an element is equal to the average atomic mass of the element
– Molar mass of a compound is equal to the sum of the masses of the elements
present
Examples
1) 𝑀𝑔𝐶𝑙2
2) 𝐶13 𝐻18 𝑂2
24.305+ (2 x 34.453) = 95.211 g/mol
(12.011 x 13) + ( 1.008 x 18) + (15.999 x 2) = 206.285 g/mol
Percent Composition
– Percent by mass for each element in a compound
– Mass of the element/ total mass of compound x 100= % of that element in the compound
1) Calculate the mass of element A
2) Calculate the mass of element B
3) Divide the mass of element A by the total mass of the compound and multiply by 100
4) Divide the mass of element B by the total mass of the compound and multiply by 100
Examples
1) 𝐻2 𝑂
2) 𝐹𝑒𝑆
H: (1.008 x 2) = 2.016 g
Fe: 55.845 g
18.015 g
87.911 g
O: ( 1 x 15.999) = 15.999 g
S: 32.066 g
(2.016/18.015) x 100 = 11.19 % H
( 55.845/87.911) x 100 = 63.52 % Fe
(15.999/18.015) x 100 = 88.81 % O
( 32.066/87.911) x 100 = 36.48 % S
Empirical/Molecular Formulas
– Molecular Formula: The ratio between elements represents the number of
atoms in one molecule of the molecular compound
– Empirical Formula: An empirical formula consists of the symbols for the
elements combined in a compound, with subscripts showing the smallest whole
number mole ratio of the different atoms in the compound
*Ionic compounds only have empirical formulas
*Covalent compounds may have both
Deriving Empirical and
Molecular Formulas
Empirical Formula
Molecular Formula
– If percent composition is given, first
convert into grams (assume 100 g)
– A molar mass is always given
– Divide by molar mass
– Take the smallest mole amount and
divide the other mole amounts by it
– Round to the closest whole number and
write out the empirical formula based on
the ratios
– Determine the molar mass of the
empirical formula derived
– If the 2 amounts are different, divide
the molar mass by the empirical
formula mass
– Round the quotient to nearest whole #
– Multiply the empirical formula by that
value to derive the molecular formula
Example: Empirical Formula
 Given: 63.52% Fe, 36.48% S
Fe- 63.52 g (molar mass 55.485 g/mol)
Assume 100g
S- 36.48 g (molar mass 32.066 g/mol)
 63.52/ 55.485 = 1.13743 mol Fe/ 1.13743 ≈ 1
 36.48/ 32.066 = 1.13765 mol S / 1.13765 ≈ 1
 Ratio- 1:1, so FeS
Empirical
Molecular
Given: 𝑃2 𝑂5
Empirical molar mass- 283.89 g/mol
Molecular formula-?
 Actual molar mass- 141.943 g/mol
 283.89/141.943= 2.00002818≈ 2
 So,(𝑃2 𝑂5 )2
𝑃4 𝑂10
Molecular Formula: 𝑃4 𝑂10
Moles & Chemical Formulas
By: Ryan Kwon &
Ashwin Thevarumuriyil
The Mole
• Mole: an amount of a substance
• Avogadro’s Number: the number of particles in one mole of a pure
substance
▫ 6.022 x 10^23 mol^-1
▫ Used as a conversion factor to convert particles (atoms to moles)
Molar Mass
• Molar Mass: The mass of one mole of a substance
▫ Measured in grams per mole
 g/mol
▫ Used as conversion factor to convert between grams and moles
• Molar mass of an element is equal to atomic mass
▫ Oxygen has molar mass of 15.999 g/mol
• Molar mass of compound equal to sum of individual molar masses
▫ CO^2
= (1 x 12.011) + (2 x 15.999) = 44.009 g/mol
Percent Composition
• Percent Composition: Percent of mass of each element in a
compound.
▫ (Part/whole) x 100 = %
▫ (Mass of element/molar mass) x 100 = %
 H202
 H: (2 x 1.008) = 2.016
 O: (2 x 15.999) = 31.998
▫ Molar Mass = 34. 014
▫ H: 5.93%
▫ O: 94.07
Empirical & Molecular Formulas
• Empirical Formula: Formula with the smallest whole number ratio between
elements in a compound
▫ Ionic compounds are always empirical
 MgCl2
• Molecular Formula: Formula with ratio between elements as they exist in one
molecule of a molecular compound.
▫ Can be empirical formula
 H20
How To Find Empirical Formula
• 63.52% Fe 36.48% S
1. Assume 100g for example
2. Divide by molar mass
▫
▫
63.52% Fe -> 63.52g Fe/55.845 = 1.13743 mol Fe
36.48% S -> 36.48g S/32.066 = 1.13765 mol S
3. Divide by smallest number of moles
▫
▫
1.13743/1.13743 = 1 mol Fe
1.13765/1.13743 = 1.00019 mol S
4. Round
Special Cases to Finding Empirical Formula
• After dividing by smallest number of moles, the quotient may end in one of
the following
▫ .25 (x4)
▫ .33 (x3)
▫ .5 (x2)
▫ .67 (x3)
▫ .75 (x4)
• If you end up with these numbers, just multiply by numbers next to them
until they become close enough to a whole number to round
REACTIONS
Balancing equations
__Al +__HCl ->__AlCl3 +__H2
Balancing equations
_2_Al +_6_HCl ->_2_AlCl3 +_3_H2
combustion
• When a chemical includes hydrogen and carbon as reactants and
always produces carbon dioxide and water as a product
_1_CH4+_2_O2->_1_CO2+_2_H2O
YOU WANT SPICY? HERE WE GO
Process into Equation
What happened between A and B? What about C and D? Omg really are they still
dating? Long story short??
RUMORS
TEXT TALK (aka equation); AB + C -> AC + B
..get ready to learn you some more! it ain't over yet yung'uns
Process into Words
What happened between A and B? What about C and D? Omg really are they
still dating? Tell me in full detail??
RUMORS
LONG STORY (aka Words); AyyBee and Cee yielded AyyCee and
Bee
Synthesis
When two or more substances combine to form a new
product.
You find the one.. They're cute, they give you butterflies,
and most of all they're good hearted. Y'all think it's real
and you'll last forever. Awh, good for you. #synthesis
A+B -> AB
Decomposition
When a compound breaks down to form two or
more elements, basically the opposite of synthesis
You remember that time you and that special
person thought y'all were gonna last forever?
Lolyeahmetoo #decomposition
AB->A+B
Double Replacement
When to compounds react to form two new
compounds by switching elements
You're with your bae, and your best friend is with his
bae, now your best friend has your ex bae, and you
have your best friend's ex bae. . #ol'switcharoo
AB+CD->AD+CB
Single Replacement
When an element reacts with a compound and
replaces an element within the compound
Lol nice bae you got there…Don't mind if I do…
AB+C->AC+B
REACTIONS
By: Dhruv Kota and Hasta Zain
HOW TO TRANSLATE EQUATIONS
TO WORDS?
• Translating chemical equations to words is done by using number of moles and type of
chemical.
• The format for writing a chemical equation is “__ moles of Chemical and __ moles of
chemical yields __ moles of chemical and __ moles of chemical.
• Example- 4NH3+ 5O2
4NO+6H2O
• Answer- four moles ammonia and five moles of oxygen gas yield four moles of nitrogen
monoxide and six moles of water.
TRANSLATING EQUATIONS TO
FORMULAS
• To translate a chemical equation to the formula, the same process as above must be used, but just
reversed.
• The worded equation must be used, and the correct amount of moles and chemical names are written
for the chemical reaction formula.
• Example- One mole of sodium chloride and one mole of sulfuric acid yield one mole of hydrochloric acid
and one mole of sodium hydrogen sulfate.
• Answer- 1NaCl + 1H2SO4
1HCl + 1NaHSO4
TYPES OF REACTIONS
• There are 5 types of reactions- Synthesis, Decomposition, single replacement,
double replacement and combustion.
• Synthesis: When the reactants all combine into one product.
• Decomposition: When the reactant splits into different products.
• Single replacement: When an element in the reactants take the place of
another element so the composition of the reactants end up switched with
one element in the product.
• Double replacement: When 2 compounds react, the positive ions in the
reactant react with the negative ions in the reactant switch places with
each other causing them to be switched in the product.
• Combustion: Reactions that occur when something burns-uses organic
molecules.
EXAMPLES OF TYPES OF REACTIONS
• Synthesis: Mg + O2 ---> MgO
• Decomposition: HgO ---> Hg + O2
• Combustion:
CH4 + O2 ---> CO2 + H2O
• Single replacement: CH4 + O2 ---> CO2 + H2O
• Double replacement: KOH + H2SO4 ---> K2SO4 + H2O
BALANCING EQUATIONS
• Chemical equations are written to express different chemical reactions.
• Most of them need to be balanced-they need to have the same amount of
each atoms on each side of the equation.
• How to balance chemical equations:
1. Write equation.
2. Figure out a coefficient to put next to each compound that would make
each side have equal amounts of each element.
• Ex:
• C3H8 + 5O2 --> 4H2O + 3CO2
STOICHIOMETRY
Stoichiometry
CHIRAYU T. AND KIRK A.
How to Convert

Converting grams, moles, and atoms

Basically dimensional analysis
Converting Grams

Grams to grams

Grams to moles
20g C = ?mol C

Grams to atoms
20g C = ?atoms C
Converting Moles

Moles to moles

Moles to grams
1mol Na = ?g Na

Moles to atoms
1mol Na = ?atoms Na
Converting Atoms

Atoms to atoms

Atoms to moles
5×1020 atoms K = ?mol K

Atoms to grams
20×1016 atoms Cl = ?g Cl
Converting Between Chemicals

Uses ratios between chemicals
From the reaction: B2H6 + O2
HBO2 + H2O
a. What mass of O2 will be needed to burn 36.1 g of B2H6?
b. How many moles of water are produced from 19.2 g of B2H6?
STOICHIOMETRY REVIEW
By: Aleena Hassan and Raisha
Ahmed
CONVERTING MOLES TO MOLES
Mole ratio (Equation)
Amount of GIVEN
substance (in mol)
mol UNKNOWN
———————
mol GIVEN
Given in the problem
Conversion factor
Amount of
UNKNOWN
substance (in mol)
Calculated
CONVERTING MOLES TO MOLES EXAMPLE
If you have 2.00 mol of N2 reacting with H2, how many moles of NH3 will be produced?
Equation: N2 + 3H2  2NH3
2.00 mol N2
Given in the problem
2 mol of NH3
———————
1 mol N2
Conversion factor
(mole ratio)
4.00 mol NH3
Calculated
CONVERTING MOLES TO GRAMS
Mole ratio (Equation)
Amount of GIVEN
substance (in mol)
Given in the problem
mol UNKNOWN
———————
mol GIVEN
Molar mass
(Periodic table)
molar mass of UNKNOWN
(in g/mol)
Conversion factors
Mass of
UNKNOWN
substance (in g)
Calculated
CONVERTING MOLES TO GRAMS EXAMPLE
If you have 3.00 mol of H2O react with CO2, how many grams of C6H1206 are produced?
Equation: 6CO2+6H20  C6H1206+6C02
Mole ratio (Equation)
3.00 mol H2O
Given in the problem
1 mol C6H1206
———————
6 mol H2O
Molar mass
(Periodic table)
180.18 g C6H1206
————————
1 mol C6H1206
Conversion factors
90.1 g C6H1206
Calculated
CONVERTING GRAMS TO MOLES
Inverted molar mass
(Periodic table)
Mass of GIVEN
substance (in g)
Given in the problem
1
—————————————
molar mass of GIVEN (in g/mol)
Conversion factors
Mole ratio
(Equation)
mol UNKNOWN
———————
mol GIVEN
Amount of
UNKNOWN
substance (in mol)
Calculated
CONVERTING GRAMS TO MOLES EXAMPLE
How many moles of mercury (II) oxide, HgO, are needed to produce 125 g of oxygen, O2?
Equation: 2Hg0 2Hg + O2
Inverted molar mass
(Periodic table)
125 g O2
Given in the problem
1 mol O2
—————
31.998 g O2
Mole ratio
(Equation)
2 mol HgO
———————
1 mol O2
Conversion factors
7.81 mol HgO
Calculated
CONVERTING GRAMS TO GRAMS
Inverted molar mass
(Periodic table)
Mass of GIVEN
substance (in g)
Given in the
problem
Mole ratio
(Equation)
1
—————————————
molar mass of GIVEN (in g/mol)
mol UNKNOWN
———————
mol GIVEN
Conversion factors
Molar mass
(Periodic
table)
molar mass of
UNKNOWN (in
g/mol)
Mass of
UNKNOWN
substance (in g)
Calculated
CONVERTING GRAMS TO GRAMS EXAMPLE
You are given this equation: 4Al +3O2  2Al2O3
If you were given 74.00 grams of Al, how many grams of Al2O3 would be produced?
Inverted molar mass
(Periodic table)
74.00 g Al
Given in the
problem
Mole ratio
(Equation)
1 mol Al
—————————————
26.982 g Al
2 mol Al2O3——
————
4 mol Al
Conversion factors
Molar mass
(Periodic
table)
101.96 g Al2O3—
—————
1 mol Al2O3
139.8 grams of
Al2O3
Calculated
CONVERTING MOLES TO ATOMS
Amount of element
(in mol)
Given in the problem
6.022*1023
Number of atoms
of element
Calculated
CONVERTING MOLES TO ATOMS EXAMPLE
If you have 2.00 mol of Helium, how many atoms do you have?
2.00 mol He
Given in the problem
6.022*1023
1.20*1024 atoms
He
Calculated
CONVERTING ATOMS TO MOLES
Number of atoms
of element
Given in the
problem
1
—————
6.022*1023
Amount of element
(in mol)
Calculated
CONVERTING ATOMS TO MOLES EXAMPLE
If you have 5 atoms of Sulfur, how many moles of Sulfur do you have?
5 S atoms
Given in the
problem
1 mol S
—————
6.022*1023 S atoms
8 mol S
Calculated
CONVERTING GRAMS TO ATOMS
Mass of element (in
grams)
Given in the problem
1
—————————
molar mass of element
1
—————
6.022*1023
Number of atoms
of element
Calculated
CONVERTING GRAMS TO ATOMS EXAMPLE
How many atoms of Sulfur are in 4.00 g of Sulfur?
4.00 g S
Given in the problem
1 mol S
—————————
32.065 g S
1 atom S
—————
6.022*1023 mol S
7.51*10-22 atoms
S
Calculated
CONVERTING ATOMS TO GRAMS
Number of atoms of
element
Given in the problem
1
—————
6.022*1023
Molar mass of element
Mass of element
(in grams)
Calculated
CONVERTING ATOMS TO GRAMS EXAMPLE
If you have 1 atom of Hydrogen, how many grams of Hydrogen do you have?
1 atom H
Given in the problem
1 mol H
—————
6.022*1023 atoms H
1.008 g H
——————
1 mol H
2gH
Calculated
DETERMINING LIMITING REACTANT
Limiting reactant: the reactant that limits the amounts of
other reactants that can combine and the amount of
product that can form in a chemical equation
To determine: Balance the equation, convert to moles, use
stoichiometry for each individual reactant to find the mass
of product produced. The reactant that produces a lesser
amount of product is the limiting reactant.
DETERMINING EXCESS REACTANT
Excess reactant: the substance that is not used up
completely in a reaction
To determine: Balance the equation, convert to moles, use
stoichiometry for each individual reactant to find the mass
of product produced. The reactant that produces a larger
amount of product is the excess reactant.
CALCULATING THEORETICAL YIELD
Theoretical yield: the maximum amount of product that
can be produced from a given amount of reactant
In most chemical reactions, the amount of the product
obtained is less than the theoretical yield
To calculate: find the moles in the limiting reactant, and
convert it from moles to grams
CALCULATING PERCENT YIELD
Percent yield: the ratio of the actual yield to the
theoretical yield , multiplied by 100
Actual yield: the measured amount of a product
obtained from a reaction
Example: If the actual yield is 20 and the theoretical
yield is 25, the percent yield is 80%