Download Unit 1 - Morgan Science

Midterm Review
Unit 1 Matter and Measure
Unit 2 Atomic Theory and Structure
Unit 3 Nuclear Chemistry
Unit 4 Periodic Table
Unit 5 Bonding
Unit 6 Naming and Moles
Unit 1 - Matter & Measure
Chapters 1-3
Matter

Anything that has mass and takes up
space, volume

Classified into two categories
◦ Substances (Pure)
◦ Mixtures
Atom

Simplest form of matter

Made up of Subatomic Particles

Different atoms have different properties
Pure Substances

Element
◦ simplest form of matter that has a unique set
of properties.
◦ Can’t be broken down by chemical means

Compounds
◦ substance of two or more elements
chemically combined in a fixed proportion
◦ Can be broken down by chemical means
Mixtures

Physical blend of two or more substances

Homogeneous
◦ Composition is uniform throughout
◦ Solution is a homogeneous mixture
◦ Aqueous Solution is something mixed in water

Heterogeneous
◦ Composition is not uniform throughout
2.3
Matter
Separating Mixtures


Differences in physical properties can be
used to separate mixtures
Examples:
◦ Filtration – Separates solids from liquids in
heterogeneous mixtures
◦ Distillation – Separates homogeneous liquid
mixtures based on different boiling points
◦ Evaporation – evaporate away liquid to leave solid
◦ Chromatography – separation of substances
based on polarity
Phases(States) of Matter

Solid (s)
◦ Definite shape and volume
◦ Particles are packed tightly together
 Liquid
(l)
◦ Definite volume, takes shape of container
◦ Particles can slide past each other
 Gas
(g)
◦ Takes shape and volume of container
◦ Particles are spread very far apart
Phase Changes
Solid  Liquid
 Liquid  Solid
 Liquid  Gas
 Gas  Liquid
 Solid  Gas
 Gas  Solid


Melting
Freezing
Vaporization
Condensation
Sublimation
Deposition
Temperature does NOT change during a
phase change
Aqueous Solutions
Dissolved in water
 (aq) used after chemical symbols

◦ NaCl(aq)
Temperature

Measure related to the heat of an object

Measured in °Celsius or Kelvin(no
degrees)

Conversion
K  C  273
Density

Amount of matter in a given amount of
space

Amount of mass in a given volume
m
D
V
Identifying Substances

Physical Property
◦ quality or condition of a substance that can be
observed or measured without changing the
substance’s composition
◦ Ex: Color, shape, size, mass

Physical Change
◦ some properties change, but the composition
remains the same
◦ Ex: melting, freezing, tearing
Identifying Substances (cont)

Chemical Change
◦ change that produces matter with a different
composition than the original matter
◦ Ex. burning, rusting, decomposing, exploding,
corroding

Chemical property
◦ property that can only be observed by changing
the composition of the substance.
◦ Ex: Reactivity with acids, reactivity with oxygen
Energy

Ability to do something

Exothermic
◦ Process when energy is released or given off
◦ Ex: Burning, freezing

Endothermic
◦ Process when energy is absorbed or taken in
◦ Ex: Melting
Scientific Method
Observation - make observations
 Hypothesis - proposed explanation
 Experiment - test
 Analyze Data - check to see if results
support hypothesis.
 Theory - well tested explanation
 Law - concise statement that summarizes
the results

Scientific Laws

Law of Conservation of Mass
◦ Mass can not be created or destroyed, only
changed into different forms

Law of Conservation of Energy
◦ Energy can not be created or destroyed, only
changed into different forms
Metric System Prefixes
Prefix
Power
Symbol
Giga
109
G
Mega
106
M
kilo
103
k
deci
10-1
d
centi
10-2
c
milli
10-3
m
micro
10-6
μ
nano
10-9
n
pico
10-12
p
Significant Figures

If the decimal point is present, start
counting digits from the Pacific (left) side,
starting with the first non-zero digit.
 123
0.00310 (3 sig. figs.)

If the decimal point is absent, start
counting digits from the Atlantic (right)
side, starting with the first non-zero digit.
32 1

31,400 (3 sig. figs.)
SigFigs for Math

Addition and Subtraction
◦ Answer has to have the same number of
decimal places as least number of decimal
places in what you are adding or subtracting

Multiplication and Division
◦ Answer has to have same number of Sigfigs as
least number of Sigfigs in what you are
multiplying or dividing
Percent Error
Main Menu
Atomic Theory and Structure
Chapters 4-5
Atomic Theories

Democritus ~ 400 BC
◦ believed that atoms were indivisible and
indestructible

Dalton ~ 1800’s
◦ Developed through experiments
◦ First Atomic Model
Dalton’s Atomic Model

All elements are composed of tiny
indivisible particles called atoms

Atoms of the same element are identical.
The atoms of any one element are
different from those of any other
element.
Dalton’s Atomic Model (cont)

Atoms of different elements can physically
mix together or can chemically combine
in simple whole-number ratios to form
compounds.

Chemical reactions occur when atoms
are separated, joined, or rearranged.
Atoms of one element, however, are
never changed into atoms of another
element as a result of a chemical reaction.
“Plum Pudding” Model
◦ Discovery of electron requires new model

Uniform positive sphere with negatively
charged electrons embedded within.
Rutherford Gold Foil Experiment
Shot alpha particles at gold foil
 Most went through foil with little or no
deflection.
 Some were deflected at large angle and
some straight back.

Rutherford Model

Conclusions from Gold Foil Experiment
◦ Atom is Mostly Empty Space
◦ Dense positive nucleus
◦ Electrons moving randomly around nucleus
Bohr Model

Dense positive nucleus

Electrons in specified circular paths, called
energy levels

These energy levels gave results in
agreement with experiments for the
hydrogen atom.
Bohr Model
Bohr Model

Each energy level can only hold up to a
certain number of electrons
Level 1  2 electrons
 Level 2  8 electrons
 Level 3  18 electrons
 Level 4  32 electrons

Wave Mechanical Model



More detailed view of the Bohr Model
Still maintains nucleus with protons and neutrons
Electrons are found in orbitals within the energy
levels


Regions of space where there is a high probability of
finding an electron
Modern Model
◦ AKA Quantum Mechanical Model, Electron Cloud Model
Subatomic Particles

Electron
◦ Discovered first
◦ Negative charge (-1)
◦ Approx mass ~ 0u

Proton
◦ Discovered second
◦ Positive charge (+1)
◦ Approx mass ~ 1u

Neutron
◦ Discovered last
◦ No charge (0)
◦ Approx mass ~ 1u
 Just slightly larger than a
proton
Atomic Structure

Atomic Number
◦ Number of protons in an element
◦ All atoms of the same element have the same
number of protons

Mass Number
◦ Number of protons and neutrons in an atom
Atomic Structure

# of Neutrons = Mass Number – Atomic
Number

Atoms of the same elements can have
different numbers of neutrons

Isotope – atoms of the same element
with different number of neutrons
Ion

Atom or group of atoms that have gained
or lost one or more electrons
◦ Have a charge

Example:
◦ H+, Ca2+, Cl-, OH-
Average Atomic Mass

Atomic Mass
◦ Weighted average based on the relative
abundance and mass number for all naturally
occurring isotopes

Example
◦ C-12 98.9%
12.011u
◦ C-13 1.1%
◦ 0.989*12 + 0.011*13 = 12.011u
Electron Configuration
The way in which electrons are arranged
in the atom
 Example: Na 2-8-1


Valence Electrons
◦ Electrons in the outermost energy level
Ground State vs. Excited State

Ground State
◦ When the electrons are in the lowest
available energy level
◦ Ex: Na 2-8-1

Excited State
◦ When one or more electrons are not in the
lowest available energy level
◦ Ex: Na 2-7-2 or 2-8-0-1 or 2-6-1-1-1
Energy Level Transitions

Gaining energy will move an electron
outward to a higher energy level
(Absorption)

When an electron falls inward to a lower
energy level, it releases a certain amount
of energy as light (Emission)
Line Spectra

Emission Spectra
◦ Shows only the light that is emitted from an
electron transition

Absorption Spectra
◦ Shows a continuous color with certain
wavelengths of light missing (absorbed)
Main Menu
Unit 3 - Nuclear Chemistry
Chapter 25
Radioisotopes

Nuclei of unstable isotopes are called
radioisotopes.

An unstable nucleus releases energy
and/or mass by emitting radiation during
the process of radioactive decay
Radiation
Type
Alpha
Mass
4
Charge
+2
Strength
Weakest
Symbol
α
Beta
Gamma
0
0
-1
0
middle
Strongest
β
γ
Nuclear Stability

For smaller atoms, a ratio of 1:1 neutrons
to protons helps to maintain stability
◦ C-12, N-14, O-16

For larger atoms, more neutrons than
protons are required to maintain stability
◦ Pb-207, Au-198, Ta-181
Nuclear Reactions

Unstable isotopes of one element are
transformed into stable isotopes of a
different element.

They are not affected by outside factors,
like temp and pressure. They can not be
sped up or slowed down.
Radioactive Decay

Radioisotopes will undergo decay
reactions to become more stable
Alpha Decay
 Beta Decay
 Positron Emission

Alpha Decay
220
87
Fr    At
4
2
216
85
Beta Decay
90
38
Sr    Y
0
1
90
39
Positron Emission
19
10
Ne e F
0
1
19
9
Transmutations

Any reaction where one element is
transformed into a different element

Natural
◦ Usually has one reactant
◦ Alpha and Beta Decay

Artificial
◦ Usually has more than one reactant
◦ Fission, Fusion
Example
9
4
Be  H  He  Li
1
1
4
2
6
X
3
Fission

Splitting of a larger atom into two or
more smaller pieces
◦ Nuclear Power Plants

One Example:
U  n Ba  Kr 3 n
235
92
1
0
141
56
92
36
1
0
Fusion

Joining of two or more smaller pieces to
make a larger piece
◦ Sun, Stars

Examples:
4 H  He 2 
2
2
3
1
1 H 1 H  2 He  0 n
2
1
1
1
4
2
0
1
H  H  He  n
3
1
4
2
1
0
Energy Production

Energy is produced by a small amount of
mass being converted to energy
◦ Fission and Fusion
◦ More energy is produced by fusion than any
other source

E=mc2
Fission vs. Fusion

Advantages of Fission
◦ Produces a lot of energy
◦ Can be a controlled reaction
◦ Material is somewhat abundant

Disadvantages of Fission
◦ Uses hazardous material
◦ Produces hazardous material
 Long Half Life
◦ Reaction can run out of control.
◦ Limited amount of fissionable material
Fission vs. Fusion

Advantages of Fusion
◦ Lighter weight material
◦ Easily available material
◦ Produces waste that is lighter and has shorter
half-life
◦ Produces more energy than fission

Disadvantages of Fusion
◦ Must be done at very high temperatures
 Only been able to attain 3,000,000K
◦ Have not been able to sustain stable reaction for
energy production
Radioisotopes

I-131

◦ Diagnosing and
treating thyroid
disorders

Co-60
◦ Treating cancer
C-14
◦ Dating once-living
organisms
◦ Compare to C-12

U-238
◦ Dating geologic
formations
◦ Compare to Pb-206
Half Life

Amount of time for half of a sample to
decay into a new element
Fraction
1
 
Re maining  2 
t
# HalfLives 
T
Mass Left
Original Mass
=
1
 
2
t
T
t
T
Example

How many half lives does it take for a
sample of C-14 to be 11430 yrs old?
t 11430 y

2
T 5715 y
More Practice

How much 226Ra will be left in a sample
that is 4797 years old, if it initially
contained 408g?
x
1
 
408g  2 
4797 y
1599 y
3
1 1
  
2 8
51g
And One More….

What is the half life of a sample that
started with 144g and has only 9g left
after 28days?
9g
1
 
144 g  2 
28d
x
1 1
  
16  2 
4
28d
4
x
7d
Main Menu
Unit 4 - Periodic Table
Chapter 6
History

Dmitri Mendeleev developed the first
Periodic Table
◦ based on increasing atomic mass.
◦ elements with similar chemical properties next to
each other
◦ out of order by atomic mass if chemical
properties lined up
◦ left open spaces for elements not yet discovered
◦ predicted properties of elements not yet
discovered
Periodic Table

The modern periodic table is arranged in
order of increasing atomic number.
Arrangement
Rows are called Periods
 Columns are called Groups

 Numbered 1-18
◦
◦
◦
◦
◦
◦
Group 1 - Alkali Metals
Group 2 - Alkaline earth metals
Group 17 – Halogens
Group 18 - Inert or Noble gases
Groups 3-11 – Transition Metals
Bottom 2 rows – Inner Transition
Phases at STP

Most elements are solids at STP

Hg and Br are liquids at STP

H, N, O, F, Cl and Noble Gases are all
gases at STP
Valence Electrons

Electrons in outermost occupied energy
level

Valence Electrons are responsible for
most chemical properties
◦ Elements in the same group have similar
properties because they have the same
number of valence electrons
Metals
Good conductors of heat and electrical
current
 High luster or sheen
 Many are ductile, meaning they can be
drawn into wires
 Most are malleable, meaning they can be
hammered into thin sheets

Metals

Metallic Character increases as you move
towards the lower left

Most Metallic Element is Francium, Fr
Nonmetals

Most are gases at room temperature,
some are solids, and one is liquid

Most are poor conductors

Most solids are brittle
Nonmetals

Non-Metallic Character increases as you
move towards upper right

Most nonmetallic element is Fluorine, F
Metalloids

B, Si, Ge, As, Sb, Te

Have properties of both metals and
nonmetals, based on conditions

Exceptions:
◦ Al and Po are metals
◦ At is a nonmetal
Group Characteristics

Alkali Metals (Group 1)
◦ H, Li, Na, K, Rb, Cs, Fr
◦ All have 1 valence electron, tend to form +1
ions
◦ Most reactive metals
◦ Not found in nature by themselves, always
combined with someone else
◦ Have properties of metals but are softer and
less dense
Group Characteristics (cont)

Alkaline Earth Metals (Group 2)
◦ Be, Mg, Ca, Sr, Ba, Ra
◦ All have 2 valence electrons, tend to form +2
ions
◦ Harder and more dense than alkali metals, but
also have higher melting and boiling points
◦ Highly reactive, but not as much as alkali
metals
◦ Not found by themselves in nature
Group Characteristics (cont)

Halogens (Group 17)
◦ F, Cl, Br, I, At
◦ All have 7 valence electrons, tend to form -1
ions
◦ Strongly non-metallic
◦ Most active nonmetals
◦ Have low melting and boiling points
◦ Combine readily with metals to form salts
Group Characteristics (cont)

Noble Gases (Group 18)
◦ He, Ne, Ar, Kr, Xe, Rn
◦ Colorless gases that are extremely nonreactive
◦ Full valence shell, non-reactive
◦ All are found in small amounts in our
atmosphere
Group Characteristics (cont)

Transition Metals (Groups 3-11)
◦ Most are excellent heat and electrical
conductors
◦ Most have high melting points and are hard,
except Hg
◦ Less active than group 1 and 2 metals
◦ Many combine with Oxygen to form oxides
(Chemical property)
◦ Many have more than one oxidation number
◦ Form compounds that are colorful
Diatomics

Eight elements are diatomic molecules
when alone in nature (exist as two atoms
bonded together)

H2, N2, O2, F2, Cl2, Br2, I2, At2
Diatomics

Hydrogen and the Magic 7
Trends

Atomic number
◦ increases across a period.
◦ increases down a group

Atomic mass
◦ generally increases across a period.
◦ increases down a group.
Properties

Atomic Radius – size of the atom

Ion
◦ Atom, or group of atoms, that has gained or
lost electrons
 Cation – positive ion
 Anion – negative ion

Ionic Radius – size of an ion
Ions

Positive element Ions
◦ have lost electrons
◦ radius becomes smaller
◦ Metals tend to lose electrons

Negative element Ions
◦ have gained electrons
◦ radius becomes larger
◦ Nonmetals tend to gain electrons
Trends

Atomic Radius
◦ decreases across a period
◦ increases down a group

Ionic Radius
◦ decreases across a period for positive ions
◦ decreases across a period for negative ions
◦ increases down a group
Properties

Ionization Energy
◦ Amount of energy required to remove an
electron from an atom

Electronegativity
◦ Ability of an atom to attract an electron from
another atom when in a compound.
 Noble gases are usually omitted since they don’t
form compounds
Trends

Ionization Energy
◦ tends to increase across a period
◦ tends to decrease down a group

Electronegativity
◦ tends to increase across a period
◦ tends to decrease down a group
Trends Summary
Property
Atomic Number
Atomic Mass
Atomic Radius
Ionic Radius
Ionization Energy
Electronegativity
Period (LR)
Group (TB)
Main Menu
Unit 5 - Bonding
Chapters 7-8
Octet Rule

Atoms tend to lose or gain electrons to
achieve a full valence shell (8)
◦ Exception: First Energy Level is full with 2
electrons
Electron Dot Structures

Diagrams that show valence electrons,
usually as dots
◦ AKA Lewis Electron Dot Diagrams

Rules
◦
◦
◦
◦
Start on any side
First two get paired together
Next three are separated
Fill in as needed
O
Ions

Atoms that have gained or lost electrons,
and now have a charge

Must show charge
+
Na
-
F
-2
O
Compounds

Two Main Types of Compounds
◦ Ionic
◦ Molecular (Covalent)

Based on type of bonding involved
Bonding

Bond
◦ Shared or exchanged electrons that hold two
atoms together

Three Main Types
◦ Covalent
◦ Ionic
◦ Metallic
Covalent Bonds

Electrons are shared between two atoms
to hold them together
◦ Each atom will try to achieve a full valence
shell
◦ 2 nonmetals

Two types of covalent bonds
◦ Non-Polar Covalent – Shared equally
◦ Polar Covalent – Shared unequally
Covalent Bonding
HH
O O
N N
OH
H
More Examples
H Cl
H
H C H
H
H N H
H
O C O
Determining Bond Type

Based on electronegativity DIFFERENCE
between two bonding atoms

Nonpolar Covalent Bond
◦ 2 same Nonmetals

(no difference in electronegativity)
Polar Covalent Bond
◦ 2 different Nonmetals (small difference in electronegativity)
Determining Bond Polarity

The larger the difference in
electronegativity, the more polar the
bond.
Bonding

Ionic Bond
◦ Electrons are transferred from one atom to
another (one gives, one takes)
◦ Metal and nonmetal, NaCl
 Large electronegativity difference
Properties

Ionic Compounds
◦ Most ionic compounds are hard, crystalline
solids at room temperature
◦ High melting points
◦ Mostly soluble in water
◦ Can conduct an electric current when melted
or dissolved in water(aq).
Properties

Covalent Compounds
◦ Most molecular compounds tend to have
relatively lower melting and boiling points
than ionic compounds.
Ionic Compounds

Ionic compounds are electrically neutral,
even though they are composed of
charged ions
◦ Total positive charge equals total negative
charge
Determining Formulas

Must be electrically neutral
◦ Total positive charge must equal total negative
charge

Use oxidation numbers from Periodic
Table
◦ Group 1  +1
◦ Group 13  +3
◦ Group 16  -2
Group 2  +2
Group 15  -3
Group 17  -1
Determining Formulas

Determine number of each ion to balance
out charge
◦ Use as subscript for element symbol
◦ Ex: CaCl2, Na3PO4, Mg(NO3)2
◦ Write Positive Ion First

Formula must be smallest wholenumber ratio
Short-cut (criss-cross method)

Magnesium and Phosphate
Mg+2
PO4-3
Mg3(PO4)2
Short-cut (criss-cross method)

Magnesium and Carbonate
Mg+2
CO3-2
Mg2(CO3)2
MgCO3
Must Simplify
Polyatomics

Compounds with polyatomic ions contain
BOTH ionic and covalent bonds
◦ Example: NaNO3
Allotropes

Two or more different molecular forms of
the same element in the same physical
state
◦ Different properties because they have
different molecular structures
◦ O2 vs O3
◦ Diamond and Graphite (carbon)
Metallic Bonding

Bonding within metallic samples is due to
highly mobile valence electrons
◦ Free flowing valence electrons
◦ “Sea of Electrons”
Network Solids
All atoms in a network solid are
covalently bonded together
 Network solids have very high melting
and boiling points, since melting requires
the breaking of many bonds throughout
the compound.
 Some of the strongest materials known to
man are network solids.

Bond Energy

When two atoms form a bond, energy is
released
◦ Example: Cl + Cl  Cl2 + energy

Energy needs to be added to break a
bond
◦ Example: Cl2 + energy  Cl + Cl
Structural Formulas

Shared electrons are written as a line,
unshared electrons are not written
◦ Each line represents 2 electrons
O H
H
H Cl
H
O O
N
H
H
Molecular Polarity

Polar Molecule
◦ one end of a molecule is slightly negative(δ-)
and the other end is slightly positive(δ+).
◦ Asymmetrical charge distribution

Nonpolar Molecule
◦ Can not be separated into different ends
◦ Symmetrical charge distribution
Polar Molecule

H2O
◦ Polar Covalent Bond
◦ Electrons shared Unequally
δ-
O Hδ
Hδ
+
+
More Examples

HCl
δ+

H
Cl
δ-
δ-
NH3
δ+
H
H
N
H
δ+
H
Hδ
+
Another Example

CH4
δ+
H
δ+
H
C
H
δ
δ-
+

Nonpolar Molecule
H
δ+
Polarity

Ionic Compounds are Ionic

Nonpolar Covalent Bonds always indicate
Nonpolar Molecules

Polar Covalent Bonds
◦ Determine Symmetry
“Like Dissolves Like”

Polar and Ionic substances will dissolve in
other Polar Substances

Nonpolar substance will dissolve in other
nonpolar substances
Intermolecular Forces

Intermolecular Forces of Attraction
◦ attraction between two molecules or ions
that hold them together (not a bond)
◦ Determines melting and boiling points of
compounds
 Stronger intermolecular forces, higher melting
and boiling points
Hydrogen Bonding
Hydrogen bonded to N, O, or F, is
attracted to the N, O, or F of another
molecule.
 Not actual bond, just attraction

Hydrogen “Bond”
H F
H F
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Unit 6 – Naming and Moles
Chapter 9-10
Naming Ions

Positive Ions, cations, simply retain their
name.
◦ Na+  Sodium Ion
◦ Mg2+  Magnesium Ion

Negative Ions, anions, change ending of
element to –ide
◦ Cl-  Chloride Ion
◦ Br-  Bromide Ion
Naming Systems

Binary Covalent System (Prefixes)
◦ 2 nonmetals/metalloids

Stock System (Roman Numerals)
◦ Metal has more than one positive oxidation
number

Ionic System
◦ Everything else (for now)
Naming Ionic Compounds

Name positive ion first, then negative ion.
◦ NaCl  Sodium chloride
◦ Mg(OH)2  Magnesium hydroxide
Stock System

We must indicate which charge the metal
ion has, using Roman Numerals
◦
◦
◦
◦
CuCl
CuCl2
Cu+1
Cu+2
Copper(I)Chloride Copper(II)Chloride
Roman Numeral is the charge of the metal
ion
Roman Numerals
Cation
Charge
+1
+2
+3
+4
Roman
Numeral
I
II
III
IV
+5
+6
+7
+8
V
VI
VII
VIII
Binary Covalent Compounds

Use when compound is 2 nonmetals
◦ Including metalloids

Use a prefix system to indicate the number
of atoms for each element

Second element ends in –ide

N2Cl3
◦ Dinitrogen Trichloride
Prefixes
Number of atoms
1
2
Prefix
monodi-
3
4
5
6
tritetrapentahexa-
7
8
heptaocta-
Exceptions

When there is only one atom of the first
element, do not use mono- prefix.
◦ CO2
 Carbon dioxide

When an element starts with a vowel, drop
any o or a at the end of a prefix
◦ CO
 Carbon Monoxide
◦ P2O5
 Diphosphorus Pentoxide
Mole

6.02 x 1023
◦ Avogadro’s Number
Number of representative particles in a
mole
 1 mol He = 6.02 x 1023 atoms
 1 mol H2 = 6.02 x 1023 molecules

Gram Formula Mass
Mass of the formula in g/mol
 Simply add the atomic masses of each
element in the formula together
 H2O = 1 + 1 + 16 = 18 g/mol

◦ Also known as gram atomic mass, gram
molecular mass, molar mass
Practice

KNO3 = 39 + 14 + 16(3) = 101 g/mol

C6H14 = 12(6) + 1(14) = 86 g/mol

CuSO4 = 63.5 + 32 + 16(4)
= 159.5 g/mol
Mole - Mass Conversion

Example: 96 g of Oxygen gas = ? mol
96 g
# mol 
 3mol
32 g / mol
Molar Volume
At STP, 1 mol of any gas occupies 22.4L of
space
Examples:
 2 mol of He occupies how much space at
STP?
44.8L
 11.2L will have how much H2 gas at STP?

0.5 mol
Mole Road Map
Percent Composition
Part
*100%
Whole
Another Example

NH3
N
H
14
*100
17
3
*100
17
82.4% N
17.6% H
Hydrates

Compounds that have a specific number
of water molecules attached
CuSO4·5H2O
◦ Dot means plus (+)
◦ gfm = 159.5 + 5(18) = 249.5g/mol
Empirical Formula

Simplest Whole-Number ratio of atoms in
a compound
◦ Molecular Formula is a multiple of the
Empirical Formula
Examples
 CO2
 P4O10
 C6H12O6

P2O5
CH2O
Empirical Formula

A molecular formula has an empirical
formula of CH2 and a molecular mass of
28 g/mol.
C 2H 4

A molecular formula has an empirical
formula of CH2 and a molecular mass of
42 g/mol.
C 3H 6
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