Chemistry Fall 2013 Review
 Fold the sheet of paper into 4 squares.
 Give each box a chapter name:
Box 1: Chapter 1 Scientific method
 Box 2 : Chapter 3- Properties of matter
 Box 4: Chapter 3_ Mixture and compoiunds
 Box 5 : Chapter 4Structure of the atom
 Chapter 6: Chapter 4 Isotopes
 Box 7: Chapter 5 Atomic emission/ flame test
 Box 8: Chapter 6The periodic table and
periodic law

Chapter 1 and 2
 Density= mass / volume
 Temperature is a measurement of the average kinetic
energy of the particles that make up an object.
 Increase in K.E. can result in melting.
( Review notes and powerpoint to add any additional
information here)
Chapter 3
 3 states of matter:
 Solid: (s)definite shape and definite volume
Has least random molecular motion
 Particles in a rigid and fixed pattern (tightly packed)
 Strongest attractive force


Liquid: (aq) or (l) no definite shape, definite volume


Stronger attractive force than gas, but weaker attractive force
than a solid in the same conditions
Gas: no definite shape and no definite volume
Has greatest random molecular motion and energy
 Particles farthest apart
 Weakest attractive force

Chapter 3
 Physical change vs Chemical Change
 Physical change:


Examples: melting, evaporation, crystallization, dissolving,
bending
Chemical change:

Examples: igniting, burning, corrosion, rusting, decomposing,
fermenting, exploding.
Chapter 3
 Endothermic vs Exothermic
 Endothermic: absorbs heat from surroundings.


Example: melting an ice cube ( solid to liquid), evaporation,
sublimation ( Dry ice)
Exothermic: releases heat to surroundings.

Example:
 Liquid to solid ( Freezing)
 Gas to liquid ( Condensation)
Chapter 3
 Two types of matter:
1. Pure substance : uniform and unchanging composition.
2 types: Elements and compounds
- compounds: made up of 2 or more different elements that
are chemically combined in a definite ratio.
example of compound: NaCl
example of element: Na
2. Mixture = Consists of 2 or more substances that can be
physically separated.
2 types: -heterogenous
-homogenous
Chapter 3
 Separation techniques:
 Distillation- separates liquids based on their different boiling
points.


Example: alcohol and water
Filtration – a technique that uses a porous barrier to separate
a solid from a liquid.( when the solid does not dissolve in
water)

Example: sand and water
Chapter 4
 Atom- smallest particle of an element
 Experiments about the history of the atom, led to




the conclusion that an atom is mainly empty space,
and the nucleus has a positive charge.
The subatomic particles of an atom include
protons, neutrons and electrons.
The nucleus of the atom is made up of protons and
neutrons.
The electrons of an atom travel outside the nucleus
in different orbitals.
Define Heisenberg’s uncertainty principle
Particle
Electron
Symbol
e-
Location
Relative Charge
Relative Mass
Outside nucleus
-1
1/1840
traveling in orbitals
( relative mass of
zero)
Much smaller than
protons and neurtons
Proton
Neutron
p
n
Inside nucleus
Inside nucleus
+1
1
0
~1
Chapter 4
 Mass number= # of protons + # of neutrons
 Number of neutrons= mass number – atomic
number
 In a neutral atom, the # of protons equals the # of
electrons.
Chapter 4
 Isotopes:

Atoms that have the same number of protons but a
different number of neutrons.

Example: Rb-85 and Rb 87 ( the difference between them
is there mass number)
 Isotope Abundance:

The atomic mass is a weighted average of the masses
of all of its naturally occurring isotopes. Most of the
time it can also tells you which isotope is most
abundant.
Chapter 5
 The energy level of electrons in shells farther away
from the nucleus have a greater amount of energy
than those closer to the nucleus
 The color we see from fireworks and flame tests is
due to electrons releasing energy when they drop to
their ground state (lower energy).
Chapter 5
 Electron configuration example:
1s22s22p63s23p64s23d8
 Orbital diagram example:
Chapter 6
 Metals vs Non-metals


Metals: tend to lose electrons easily when bonding and form
positive ions (cations). Good conductors of heat and electricity,
malleable and ductile.
Non-Metals: tend to gain electrons and form negative ions
(anions). Poor conductors of electricity.
Chapter 6
 Noble gases- group 8 elements, all have the most
stable outer electron configuration
 Groups/families= vertical columns on periodic table
 All elements in the same group share similar properties (like a
family)
 All have the same number of valence electrons
 Periods= rows on periodic table

Represent energy levels
Chapter 7
 Ionic bonding involves the transfer of electrons from
the metal atom to the non-metal atom.
 Electrons are the subatomic particles that are
involved with chemical bonding
Chapter 7
 List examples of naming ionic compounds and
writing formulas:
 Lewis dot diagrams with ionic compound example: