Section 6.1
Atoms and Moles
B. Atomic Masses: Counting Atoms by Weighing
• Atoms have very tiny masses so scientists made a unit
to avoid using very small numbers.
1 atomic mass unit (amu) = 1.66 10-24 g
• The average atomic mass for an element is the weighted
average of the masses of all the isotopes of an element.
Section 6.1
Atoms and Moles
C. The Mole
• One mole of anything contains 6.022 x 1023 units of that
substance.
– Avogadro’s number is 6.022 x 1023.
• The mole is defined as the number of atoms in
exactly 12g of carbon-12.
Section 6.1
Atoms and Moles
Do standard 3e (front left only) for review
Section 6.1
Atoms and Moles
Section 6.1
Atoms and Moles
A. Molar Mass
• A compound is a collection of atoms bonded together.
•
The molar mass of a compound is obtained by summing the masses of the
component atoms.
Section 6.1
Atoms and Moles
B. Percent Composition of Compounds
• Percent composition consists of the mass percent of
each element in a compound:
Mass percent =
mass of a given element in 1 mol of compound
100%
mass of 1 mol of compound
Section 6.1
Atoms and Moles
B. Calculation of Empirical Formulas
Section 6.1
Atoms and Moles
Find the Empirical formula
• 25.95% Nitrogen
• 74.06% Oxygen
Section 6.1
Atoms and Moles
B. Mole-mole Relationships
• A balanced equation can predict the moles of product that a
given number of moles of reactants will yield.
•
The mole ratio allows us to convert from moles of one substance in a balanced
equation to moles of a second substance in the equation.
Section 6.1
Atoms and Moles
9.2 B. Mass Calculations Using Scientific Notation
• Stoichiometry is the process of using a balanced
chemical equation to determine the relative masses of
reactants and products involved in a reaction.
– Scientific notation can be used for the masses of any
substance in a chemical equation.
• To solve: gA -> mol A -> mol B -> g B
• To convert between moles and grams we use the molar
masses of the substance. To convert between moles we
use the mole ration from the balanced equation.
Skip to mole
What is
temperaure?
Temperature

7a
What is
temperature?
Temperature is a measure of the average
kinetic energy of the molecules (/atoms)
of a substance.
◦ In a hot sample, the molecules are moving
much faster than in a cold sample.
Heat

What happens when
you heat up a
substance?
Heat is energy transferred from
molecules at a higher temperature to
molecules at a lower temperature.
Heat Flow Problems

To calculate heat, use the following formula:
Always given to you!
◦ Energy = mass * specific heat * temperature change
Latent Heat

7d cont
Latent (hidden) heat is energy added to a
substance that doesn’t change the
temperature.
◦ Instead, it is used to change the phase
solid  liquid
liquid  gas
What is latent heat and
what is it used for?
Heat added here doesn’t
change the temperature of the
substance, it is causing the
phase change.
ΔH – Heat of Reaction

7b cont
Endothermic
◦ The products have more energy than the
reactants.
◦ Heat must be put into the reaction, so ΔH is
positive.

Exothermic
◦ The products have less energy than the
reactants.
◦ Heat must be released from the reaction, so
ΔH is negative.
Summary
What is temperature
 What happens when you heat a substance?
 Describe endo/exothermic reactions.
 What is ΔH for endo/exothermic reactions?
 Where does the energy go during phase
changes?
 Be able to calculate heat problems.
 What is latent heat used for?

B. COUNTING SIGNIFICANT FIGURES
All digits in a number are significant except:
1. 0s to the left of ALL the other digits.
2. 0s to the right of ALL the other digits when there
is no decimal place.
-Why aren’t these 0s significant? Because they
weren’t measured, they are just there to tell you
what place the first number is in.
Examples
a. 1457
b. 0.0025
c. 1.008
4 significant figures
2 significant figures
4 significant figures
d. 100
e. 100.
f. 120.0
1 significant figure
3 significant figures
4 significant figures
Section 5.1
Scientific Notation and Units
A. Scientific Notation
• Representing Large Numbers
• Representing Small Numbers
0.000167 To obtain a number between 1 and 10 we must
move the decimal point.
0.000167 = 1.67 10-4
Section 5.1
Scientific Notation and Units
B. Units
• Units provide a scale on which to represent the results of a
measurement.
C. Density
 Density is the amount of matter (atoms) present in a
given volume of substance.
The Scientific Method
1.
2.
3.
4.
5.
6.
7.
8.
Observation
Ask a Question
Form a Hypothesis
Set up an Experiment
Record Data
Draw a Conclusion
Repeat…..Repeat…..Repeat
Theory….maybe
Hypothesis vs. Theory
 Hypothesis – A proposed scientific explanation
for a set of observations
 Theory – A well-tested explanation that unifies
a broad range of observations
 Plate Tectonics
How does a hypothesis become a theory??
Theory vs. Law
Theories Do Not Become Laws
• A natural law is a summary of behavior, it tells what
happens.
• A theory is our attempt to explain why it happens.
PHYSICAL PROPERTY
• A physical property is any
characteristic of a material that
can be observed without
changing the material, such as
color, length, or shape.
• Substances have physical properties
that can be described and physical
changes that can be observed.
CHEMICAL PROPERTY
A chemical property is the
ability or inability of a substance
to combine with or change into
one or more new substances.
• In a chemical change, the
properties that give a substance
its identity change.
• A physical change is any change in
the size, shape, or state of matter in
which the identity of the substance is
not changed.
Dissolving is mixing a
substance into another substance to
form a solution.
Mixing is a physical change
in which neither substance dissolves
into the other.
• Chemical changes change one
substance into another substance.
• Usually chemical changes cannot be
easily reversed.
Forming New Substances
• All chemical changes produce
substances that are different from the
starting substances.
The Atoms, molecules or bonds
have been rearranged!!!
Review, don’t copy
Elements
• An element is a pure substance
made from atoms that all have the
same number of protons.
• Atoms of a particular element always
have the same number of protons.
• The number of protons in an atom of
an element is the element’s atomic
number.
Compound
• A compound is a distinct, pure substance made
of 2 or more elements in the same, fixed ratio.
• Compounds are not a mixture of elements!
They have their own properties and can only be
broken down into elements by chemical
processes. (Think water!)
• Compounds are written with a chemical formula
showing the type and number of atoms present.
E.g. H2O, C6H12O6.
• The subscripts show the # of each element in a
molecule, or the smallest ratio of elements in the
crystal lattice of an ionic compound.
Atomic Number (Z)
• The number of protons.
• The atomic number is also equal to
the number of electrons in a neutral
atom (not an ion).
• The elements on periodic table are
ordered by atomic number.
Mass number (A)
• The sum of the number of protons and
neutrons in the nucleus is called the
mass number.
– Protons & neutrons each count as 1
– Electrons have so little mass that their mass
is negligible and is ignored.
Average Atomic Mass
• Is the weighted average of the masses of the
naturally occurring isotopes of the element. It is
measured in atomic mass units (amu).
• Also called average atomic “weight,” sometimes
(incorrectly) shortened to “atomic mass”.
• For Si…
• 28.09 amu is the _______
• and 28.09 g is the _______
Isotopes
Atoms of the same element always
have the same number of protons,
but they may have different numbers
of neutrons.
Isotopes
Ex. Isotope Of Neon
Ex. Carbon Isotopes
•The average atomic mass of an element is
the weighted average mass of the mixture of
an element’s isotopes.
4.2 Discovering Parts of the Atom
Thomson’s Experiments (cont.)
• Opposite charges attract each other.
• Thomson concluded the cathode ray
must have a negative charge and named
the particles electrons.
4.2 Discovering Parts of the Atom
Discovering the Nucleus
• In Rutherford’s gold foil experiment, particles were shot
through a thin sheet of gold into a detector behind the foil.
Copy this slide
Rutherford’s Conclusions
• Ernest Rutherford showed
that atoms have internal
structure.
– A nucleus that is very small
with a positive chare and a
large mass.
– Since the atom is mostly
empty space he proposed
that electrons move around
the nucleus.
Section 3.4
Using the Periodic Table
Section 3.4
Using the Periodic Table
B. Natural States of the Elements
• Diatomic Molecules
Nitrogen gas contains
N2 molecules.
Oxygen gas contains
O2 molecules.
Section 3.4
Using the Periodic Table
B. Natural States of the Elements
• Diatomic Molecules
Review
 Valence electrons are electrons that are
found in the outer shell of the atom.
 Remember that the Group Number relates to
the number of Valence electrons
 All atoms want a full valence shell (8
electrons). This is called the octet rule.
 Electronegativity-ability to attract e-(how
much an atom wants an e-)
Section 3.4
Using the Periodic Table
Electronegativity (RECALL)
• Electronegativity – the relative ability of an atom in a
molecule to attract shared electrons to itself
Patterns (cont.)
Ion Charges and the Periodic Table
The Sodium and Chlorine Ion are held
together by an electrostatic force called
an Ionic Bond.

A covalent bond is a
shared pair of
electrons between two
atoms. (nonmetals)

Atoms with similar
electronegativities
share the electrons
fairly equally.

2+ covalently- bonded
atoms make a molecule

When atoms of different electronegativities
bond, they don’t share electrons equally.

This unequal sharing of electrons results in a
polar covalent bond.
How are polar covalent
bonds different from
regular (nonpolar)
covalent bonds?

A dipole moment results when a polar
molecule has a center for positive charge
separate from a center for negative charge

Water molecule dipole moment

The polarity of water affects its properties
– Permits ionic compounds to dissolve in it
– Causes water to remain liquid at higher temperature

In water, Oxygen pulls
on the electrons more
than the Hydrogens.

This makes the Oxygen
end slightly negative
and the Hydrogen end
slightly positive.

Water molecules
attract to each other
with hydrogen bonds.
Describe the polar
covalent bond and it’s
effects between water
molecules.
Guidelines for Drawing Lewis Structures
1. Determine the total number of valence electrons;
(for ions adjust for charge).
2. Arrange atoms in a skeleton structure and connect
them with single bonds.
3. Complete octets of the outer atoms
4. If not all of the valence electrons have been used
place any extra electrons on the central atom.
5. If the central atom does not have an octet, use lone
pairs from terminal atoms to form multiple bonds.
-Only C, N, O and sometimes S or P (in combination
with C or O) form multiple bonds.
Note: If more than one acceptable Lewis structure
can be drawn they are called resonance structures.
•
Writing Lewis Structures
Section 12.4
Structure of Molecules
[NO2]-
-
B. The VSEPR Model
Polar vs. Non-Polar Molecules
•If a molecule has a center of positive
charge that is at a different location than the
center of negative charge then the molecule
has a dipole and is a polar molecule.
•H2
•HCl
•H2O
Section 11.2
The Hydrogen Atom
A. The Energy Levels of Hydrogen
• Quantized Energy Levels
– Since only certain energy changes occur the H atom
must contain discrete energy levels.
Section 11.3
Atomic Orbitals
A. The Hydrogen Orbitals
Hydrogen Energy Levels
• Each principal energy level is divided into sublevels.
– Labeled with numbers and letters
– Indicate the shape of the orbital
4s
4p
3s
4d
3p
2s
3d
2p
1s
4f
Section 11.3
Atomic Orbitals
Sublevels & Orbitals
• The s sublevel has 1 orbital which is spherically symmetric.
• The p sublevel has 3 orbitals, shaped like dumbbells (or
double tear drops).
Section 11.3
Atomic Orbitals
B. The Wave Mechanical Model: Further Development
• Pauli Exclusion Principle - an atomic orbital can hold a
maximum of 2 electrons and those 2 electrons must have
opposite spins.
-Another way to say it is that no 2 electrons can have
the same 4 “quantum numbers”: can’t have the same
principal energy level (n), same sublevel (s,p,d,f), same
orbital (e.g. px, py, pz), AND the same spin (+1/2 or -1/2).
Section 11.4
Electron Configurations and Atomic Properties
A. Electron Arrangements in the First 18 Atoms on the
Periodic Table
Section 11.4
Electron Configurations and Atomic Properties
B. Electron Configurations and the Periodic Table
• Orbital filling and the periodic table
Section 11.4
Electron Configurations and Atomic Properties
A. Naming Compounds That Contain a Metal and a
Nonmetal
Type 2 Binary Ionic compounds
•
Since the metal ion can have more than one charge, a Roman numeral is
used to specify the charge.
Ex:
gold (I) chloride & gold (III) chloride
Section 11.4
Electron Configurations and Atomic Properties
B. Naming Binary Compounds That Contain Only
Nonmetals
Type 3 Compounds
Section 11.4
Electron Configurations and Atomic Properties
Describe the metallic
bond.
Metallic Bond – 2a.1
• Metal atoms don’t hold on to electrons well in metallic
bonds.
• This means that electrons are delocalized - free to roam
from metal atom to metal atom
• These are called metallic bonds and give metals the
ability to conduct electricity (electrons)
Section 11.4
Electron Configurations and Atomic Properties
Section 11.4
Electron Configurations and Atomic Properties
Section 11.4
Electron Configurations and Atomic Properties
Intermolecular Forces – 2d
 In any substance at any temperature, the forces holding
the material together are constantly working against the
internal energy of particle motion.
 In a solid, there is not enough energy to overcome the
Intermolecular forces
 In a liquid, there is more energy, but only enough to
allow the liquid to take the shape of its container
 In a gas, the energy has completely overcome the
Intermolecular forces
Section 11.4
Electron Configurations and Atomic Properties
• As temperature (molecular motion) increases, molecules
can break free of their IM forces
– Melting
– boiling
Explain how
Intermolecular forces
and molecular motion
interact to create
solids, liquids, and
gases.